Class 10 Physical Science

WBBSE Chapter 2 Behaviour Of Gases Pressure Exerted By A Gas And Its Volume

Pressure: Gases are made up of a large number of tiny particles (called molecules) which are moving randomly in all possible directions at all possible speeds within the available space of the container.

As a result, gas molecules collide with the inner walls of the container. After a collision, gas molecules bounce off and their direction of motion changes.

This creates a change in momentum of the gas molecules and produces a force \(\left(\mathrm{F}=\frac{m v-m u}{t}\right)\) normally on the wall.

Wbbse Class 10 Physical Science Notes

WBBSE Notes For Class 10 Physical Science And Environment

The pressure of a gas is defined as how much force is exerted normally per unit area on the wall of the container.

If a force F is applied on a surface of area A then the pressure of the gas, p=F/A. Volume Gases have neither a fixed shape nor a fixed volume.

Gases occupy the available volume of the container in which it is kept, whether the container is big or small.

So, the volume of gas = the volume of the container available for the free movement of gas molecules.

Units of pressure and volume: In the CGS and SI systems, units of force are respectively dyn and N and units of area are respectively cm² and m².

So, the unit of pressure in the CGS unit is dyn/cm² and in the Sl system, it is N/m²or pascal (Pa). (1 Pa = 10 dyn/cm²).

A commonly used unit of pressure is ‘atm’ (atmospheric pressure).

Wbbse Class 10 Physical Science Notes

1 atm =1.01325 × 10⁵ Pa = 101-325 KPa

= 760 mm of Hg = 760 torr = 76 cm of Hg= 1.01325 bar.

And, in CGS and SI systems, units of volume (V) are respectively cm³ and m³. Another commonly used unit of volume is the litre (symbol L) or dm³.

1 m³ = 1000 L = 1000 dm³

As 1L = 1 dm³ and IL 1000 mL = 1000 cm³ so,

1 m³ 106 cm³ 10⁶ mL 10⁹ mm³.

Wbbse Class 10 Physical Science Notes

WBBSE Chapter 2 Behaviour Of Gases Measurement Of Pressure Exerted By A Gas

You know that a barometer is used to measure atmospheric pressure. But to measure the pressure of an enclosed gas, we use an instrument known as a manometer or pressure gauge.

Manometer is a U-shaped tube whose one limb is short and another is long. Little mercury is poured into the U-tube.

Under normal conditions, when the stopcock is closed, in two limbs mercury level remains at the same height.

Now a container containing the gas is connected to the short limb. As a result, the gas inside the container exerts pressure on the mercury in the short limb.

Here two different cases may arise:

Wbbse Class 10 Physical Science Notes

Case 1: When the pressure of the gas is more than the atmospheric pressure then the mercury level in the long limb will go higher than in the short limb.

The difference in height of mercury level in two limbs is measured. Suppose this height is ‘h,’ when atmospheric pressure obtained from the barometer is ‘po’.

Since at the same horizontal levels inside the liquid (at points A and B) pressures are the same PA (pressure at point A) = Pg (pressure at point B)

⇒ \(p_{\text {gas }} \text { (pressure of gas) }=p_0+\rho g h_1\)

⇒ \(\text { [where } p_{\mathrm{B}}=p_0+\rho g h_1 ; \rho \rightarrow \text { density of mercury, } g \rightarrow \text { acceleration due to gravity] }\)

Case 2: When the pressure of a gas is less than atmospheric pressure than the mercury level in the short limb.

will go higher than in the long limbs. Suppose ‘h’ be the difference in height of mercury levels in two limbs.

Since at same horizontal levels inside liquid at points C and D, pressures are same so that Pc (pressure at point C) = PD (pressure at point D)

⇒ \(p_{\text {gas }}+\rho g h_2 \quad=p_0\)

⇒ \(p_{\text {gas }} \text { (pressure of gas) }=p_0-\rho g h_2\)

WBBSE Chapter 2 Behaviour Of Gases Boyle’s Law

Basically, the physical behaviour of a given amount of gas (mass or a number of moles fixed) depends on three factors, namely, pressure (p), volume (V) and temperature (T).

Keeping one-factor constant out of these three, we will see how the remaining two factors depend upon one another in studying some important generalisations of Gas Laws.

In 1662 Irish scientist Robert Boyle extensively studied the relationship between the pressure of a gas and the volume it occupies. He made experiments with an air pump designed by his assistant Robert Hooke.

Boyle observed that whenever the pressure of a gas is changed, its volume also changes. As pressure is increased twice, volume decreases to half and as pressure increases to 4 times, volume decreases one-fourth and vice-versa.

In his experiment, he took a fixed mass of gas at a constant temperature. Based on these observations, he discovered a gas law that is known as Boyle’s law.

Wbbse Class 10 Physical Science Notes

Boyle’s law: At a constant temperature, for a given mass of gas, the pressure exerted by the gas is inversely proportional to the volume it occupies.

Here ‘inversely proportional relation’ means that as the pressure goes up, its volume goes down and as the volume goes up, its pressure goes down.

The mathematical expression of Boyle’s law:

For a fixed mass of gas at a constant temperature if V be the volume and p is the corresponding pressure then according to Boyle’s law

⇒ \(V \propto \frac{1}{p} \text { (at constant temperature) or, } V=\frac{k}{p} \text { or, } p V=k\)

where k is a constant of proportionality whose value depends on the temperature and mass of the gas concerned.

The relation “pV=K” explains that the values of p and V can change in such a way that the product p and V always remain the same, provided temperature and mass are kept constant.

If V1 be the initial volume occupied by a gas at initial pressure p1 then according to Boyle’s law: P₁V₁ = k. Let’s assume that the final volume of the same quantity of gas is V₂ at final pressure P₂, then P₂V₂ = k. Now, relate these initial and final conditions :

⇒ \(p_1 V_1=k \text { (constant) }=p_2 V_2 \Rightarrow p_1 V_1=p_2 V_2\)

Wbbse Class 10 Physical Science Notes

1. p vs V graph mathematical form of Boyle’s law is \(p \times \frac{1}{V} \Rightarrow p V=K\) this relation is similar to \(y \propto \frac{1}{x} \Rightarrow y x=c\)

Along the Y-axis, then the nature of the graph is like a rectangular hyperbola.

The graph shows that if V is very high then p is very low and if V is very low then p is very high such that pV = constant provided temperature remains constant at all points.

This type of graph at constant temperature is called an isotherm. (iso = same, therm temperature).

2. pV vs p graph:

The mathematical form of Boyle’s law is pV = K. If p is changed then the corresponding value of V also changes but the value of product pV remains constant.

So if p is plotted along X-axis and pV along Y-axis, then a straight line parallel to the pressure axis is obtained.

Simple problems related to Boyle’s law:

The volume of air bubbles increases as they ascend in water. As the bubbles are rising up, the surrounding liquid exerts lesser pressure on the bubbles.

According to Boyle’s law, with a decrease in pressure, their volume will increase. Here the temperature is considered constant.

When a balloon is filled with gas both its volume and pressure increase. But in that case, Boyle’s law is not violated, since the mass of the gas filled in the balloon (does not remain the same) increases.

At higher altitudes, the atmospheric pressure is low and the air is less dense. As a result, for breathing less oxygen becomes available.

So mountaineers have to carry oxygen cylinders for breathing while climbing the mountain.

WBBSE Chapter 2 Behaviour Of Gases Simple Numerical Problems

Question 1. At room temperature and 740 mm Hg pressure, the volume of a certain mass of a gas is 2L. At the same temperature, what will be its volume at 760 mm Hg?
Answer:

Given:

At constant temperature, according to Boyle’s law: P₁V₁ = P₂V₂ So that 740 x 2 = 760 x V₂ ⇒ V₂ =

⇒ \(\frac{37}{19}=1.94 \mathrm{~L}\)

Wbbse Class 10 Physical Science Solutions

Question 2. A syringe has a volume of 10 cm³ at a pressure of 1.0 atm. Its plunger is pushed down to change the volume to 2.9 cm3 at the same temperature. What will be the final pressure?
Answer:  Given:

At constant temperature, according to Boyle’s law: P1V1 = P2V2

So that, 1.0 x 10 = P2 x 2.9 \(\Rightarrow p_2=\frac{10}{2 \cdot 9}=\frac{100}{29}=3 \cdot 44 \text { atm }\)

Question 3. A balloon contains 7-20 L of He. Its pressure is reduced to 2-00 atm and the balloon expands to occupy a volume of 25-2 L.

What was the initial pressure exerted on the balloon? (Assume that temperature remains constant throughout)

Answer: Given: 

At constant temperature, according to Boyle’s law: P₁V₁ = P₂V₂

So that p₁ × 7.20 = 2.00 x 25.2  \(\Rightarrow \quad p_1=\frac{2.00 \times 25 \cdot 2}{7 \cdot 20}\)

=7. 00 am

WBBSE Chapter 2 Behaviour Of Gases Charles’ Law

In 1787 French scientist Jacques Charles investigated the effect of a change in temperature on the volume of a fixed mass of gas keeping pressure constant, According to him, as the temperature of the gas increases its volume increases and as temperature decreases volume decreases. Why?

Because as you heat up a gas, the gas molecules move faster K.E. increases and they take up more space (volume).

This is possible only if pressure is kept constant. Similarly, we can explain the vice-versa case. Based on such observations, he discovered a gas law that is known as Charles’ law.

Charles’ law: At constant pressure, the volume of a fixed mass of gas increases or decreases by \(\frac{1}{273}\) times the volume of the gas at 0°C for every 1°C rise or fall in temperature.

Here \(\frac{1}{273}\) is called the coefficient of volume expansion. Its value is the same for all gases.

Explanation of Charles’ law: Suppose at 0°C temperature, the volume of a given mass of gas at constant pressure be Vo.

Keeping pressure fixed, if the temperature is increased by 1°C, then volume increases by \(\frac{V_0}{273}\)

Therefore, the volume of the gas at 1°C is V₁= Original volume + Increase in volume + increase in volume = \(V_0+\frac{V_0}{273}\)=\(V_0\left(1+\frac{1}{273}\right)\)

At the same pressure, the volume of the gas for a 2°C rise in temperature is \(V_2=V_0\left(1+\frac{2}{273}\right)\)

Similarly, the volume of the gas at t°C is \(V_t\) = \(V_0\left(1+\frac{t}{273}\right)\)

Again, at the same pressure, for lowering in temperature by 1°C (i.e. at 1°C temperature), the volume of the gas decreases to \(V_{-1}=V_0\left(1-\frac{1}{273}\right)\)

Hence, the volume of the gas at t°C is \(V_{-1}=V_0\left(1-\frac{t}{273}\right)\)

Graphical representation of Charles’ law (V vs t (At constant pressure) graph): If a graph is plotted for Charles’ law, take the V vs t graph.

Along the y-axis and t (in °C) along the x-axis, we get a straight line. Here we see that volume is directly proportional to temperature, i.e, as temperature increases, volume increases and vice-versa.

Now, if we extend the graph backwards, the line meets the temperature axis at -273°C.

Theoretically, it means that at -273°C temperature, all gases occupy zero volume, according to Charles.

Scientist Kelvin named this temperature absolute zero temperature. Because -273°C is the lowest possible temperature in the universe.

The accurate value of this temperature is -273-15°C. In reality, before reaching -273°C temperature, all gases convert into liquids.

(Charles’ law is not applicable for liquid). This type of graph at constant pressure is called isobar (iso = same, bar = pressure).

WBBSE Chapter 2 Behaviour Of Gases  Absolute Scale Of Temperature

Kelvin observed that at constant pressure, the V-t graph for different gases is a straight line and all lines meet the temperature axis at 273°C.

Thus, -273°C is such a temperature which does not depend on the properties of the gas.

Using this concept, Kelvin developed a new scale of temperature which has its zero point (0) at -273°C, and the scale is called the absolute scale of temperature or Kelvin scale.

In this scale, the temperature reading is denoted by K unit, the temperature by T (absolute temperature) and each degree is taken equal to a degree in Celsius scale.

By definition: OK = -273°C and 1°C = 1K. Adding 273 with the Celsius scale reading, we can get the absolute scale reading.

A temperature t°C in Celsius scale will be in absolute scale: TK = (t + 273)°C.

The freezing point of water (0°C) in absolute scale is 273 K and the steam point of water (100°C) is 373 K. The value of absolute zero (OK) Celcius Scale in Fahrenheit scale is – 459-4°F.

Note: 

1. Absolute temperature is independent of the properties of the gas.
2. In this scale, a temperature may be zero or positive, no negative absolute temperature is possible.

WBBSE Chapter 2 Behaviour Of Gases Representation Of Charles’ Law In Terms Of Absolute Temperature

Suppose at constant pressure a fixed mass of gas has volume V₀ at 0°C, V₁ at t₁ °C and V₂ at t₂ °C.

According to Charles’ law: \(V_1=V_0\left(1+\frac{t_1}{273}\right)\)=\(V_0\left(\frac{273+t_1}{273}\right)=V_0 \cdot \frac{T_1}{273}\)

[Where T₁ is absolute value of t°C i.e. T₁ K = (t + 273)°C]

Similarly, \(V_2=V_0 \cdot \frac{T_2}{273} .\)

2. where T₂ K=(t₂+273)°C

Dividing 1 or 2: \( \) or,\(\frac{V_1}{V_2}=\frac{T_1}{T_2} \text { or, } \frac{V_1}{T_1}=\frac{V_2}{T_2}\)

Thus, the ratio of volume and absolute temperature for a given mass of gas is constant, provided pressure is kept constant.

So, Charles’ law can be used to relate initial and final conditions.

= constant or, \(\frac{V}{T}\) at constant pressure.

Hence, we arrived at the alternate statement of Charles’ law At constant pressure, for a fixed mass of gas, the volume of gas is directly proportional to its absolute temperature.

V vs T (in K) graph:  As the temperature at const. press p1 increases the volume of gas increases and vice-versa.

So the V-T graph is a straight line. How far can this volume decrease?

When the gas reached absolute zero temperature (-273°C or OK), its volume becomes zero (but in reality, this is not possible).

Simple non-numerical problem related to Charles’ law: In Diwali, when you light up a sky lantern or air balloon, the volume of air inside it increases.

(As temperature increases, the volume also increases). As a result, the density decreases. This makes the balloon lighter than its surrounding air.

Hence, a buoyant force acts on the balloon in an upward direction which helps the balloon to rise into the air.

WBBSE Chapter 2 Behaviour Of Gases Simple Numerical Problems

Remember In Charles’ law, whenever we talk about temperature we always use the absolute value. But units of pressure and volume can be different.

Wbbse Class 10 Physical Science Solutions

Question 1. A balloon is filled with 25 mL H₂ gas at 15°C. The balloon is taken to a high altitude where pressure remains constant, the temperature becomes 35°C. Find the volume of the balloon.
Answer:

∴ Pressure remains constant. According to Charles’ law:

⇒ \(\frac{V_1}{T_1}=\frac{V_2}{T_2}\)

Given: 

So that \(V_2=\frac{V_1 \times T_2}{T_1}\)=\(\frac{25 \times 308}{288} \mathrm{~mL}=26.74 \mathrm{~mL}\)

Question 2. To what temperature must a gas at 27°C be cooled in order to reduce its volume to one-third of the original volume keeping the pressure constant?
Answer: 

Given: 

At constant pressure, according to Charles’ law: \(\frac{V_1}{T_1}=\frac{V_2}{T_2}\)

So that

⇒T₂ =100 K=(100-273)°C=-173°C.

Question 3. In a glass, vessel air is kept at a temperature of 67°C. Keeping pressure the same, the temperature is increased. At what temperature, one-third air will go out of the vessel?
Answer:

Given:

At constant pressure, according to Charles’ law: \(\frac{V_1}{T_1}=\frac{V_2}{T_2}\)

So that,

⇒ \(T_2=\frac{4 \times 340}{3}=\frac{1360}{3}\)= 453-K (453.3-273)°C 180.3°C

WBBSE Chapter 2 Behaviour Of Gases The Combined Form Of Boyle’s And Charles’ Laws

Suppose for a fixed mass of gas the pressure, volume and absolute temperature are p, V and T respectively.

According to Boyle’s law: \( V \propto \frac{1}{p}\)

According to Charles’ law: \(V \propto T\)  Combining these two relations:

⇒ \(V \propto \frac{T}{p}\) when both T and p vary

⇒ \(V=\frac{k T}{p} \Rightarrow \mathrm{PV}=k T\)…..1. So that, \(\frac{p V}{T}=k\)(constant)….2

Where k is a constant whose value depends on the units of p, V, T and the mass of the gas. The equation is obtained by combining Boyle’s and Charles’ laws into a single expression.

The equation pV=kT is called the mathematical form of combined gas law. This equation is also called the equation of state of an ideal gas because the physical property of an ideal gas depends on its pressure, volume and temperature.

The equation shows that for a given quantity gas how p, V and T affect each other when all the three parameters vary for initial and final conditions.

Thus from the equation, we can write for the initial condition \(\frac{p_1 V_1}{T_1}=k\) and for the final condition \(\frac{p_2 V_2}{T_2}=k.\)

WBBSE Chapter 2 Behaviour Of Gases Simple Numerical Problems

Question 1. At 27°C a sample of N2 gas is placed in a flexible 9-0 L container at a pressure of 1.5 atm. The container is compressed to a volume of 3.0 L and the gas is heated up to 327°C. Then what will be the new pressure inside the container?
Answer:

Given:

From combined gas law: \(\frac{p_1 V_1}{T_1}=\frac{p_2 V_2}{T_2}\)

So that,

Wbbse Class 10 Physical Science Solutions

Question 2. A definite amount of CO2 gas occupies a volume of 512 cm³ at STP. Find its volume at 17°C and 750 mm of Hg.
Answer:

P₁ = 760 mm of Hg; T₁ = 273 +0 = 273K; V₁=1.5dm³

P₂ = 750 mm of Hg; T₂ = 273 +17= 290K; V₂ =?

Gas equation is ,\(\frac{p_1 V_1}{T_1}=\frac{p_2 V_2}{T_2}\)

So that \(\frac{760 \times 512}{273}=\frac{750 \times V_2}{290}\)\(\Rightarrow V_2=\frac{760 \times 512 \times 290}{273 \times 750} \mathrm{~cm}^3=551.13 \mathrm{~cm}^3\)

Wbbse Physical Science Class 10

Question 3. In the laboratory 1.5 dm3 of dry H₂ gas was prepared at a pressure of 750 mm Hg and temperature of 27°C. Find the volume of the gas at 17°C and 760 mm of Hg.
Answer:

According to the gas equation \(\frac{p_1 V_1}{T_1}=\frac{p_2 V_2}{T_2}\)

P₁ = 750 mm of Hg; T₁ = 273 +27= 300K; V₁ = 1.5 dm³

P₂ = 760 mm of Hg; T₂ = 273 +17= 290K; V₂ =?

So that \(\frac{760 \times 512}{273}=\frac{750 \times V_2}{290}\) \(\Rightarrow V_2=\frac{760 \times 512 \times 290}{273 \times 750} \mathrm{~cm}^3=551.13 \mathrm{~cm}^3\)

WBBSE Chapter 2 Behaviour Of Gases Ideal gas

The gases which obey the equation of state of an ideal gas “pV = kr” at all pressures and temperatures are called ideal gases.

But in real situations, no ideal gas can exist. All gases in the universe like O₂, N₂, H₂, CO₂, etc. are non-ideal or real gases.

Some real gases behave like ideal gases under certain conditions. Let’s discuss the conditions: We know that gases are affected by pressure and temperature.

From the equation \(p V=k T, V=\frac{k T}{p}\). At constant temperature,

⇒ \(V \propto \frac{1}{p}\)(Boyle’s law) So, keeping the temperature constant, at very high pressure (suppose p→ ∞), V→ O.

In reality, at very high pressure, gases change into liquids. So, the relation pV = kT is not applicable for real gases at very high pressure.

On the other hand, from pV = kT, we can say that at constant pressure, V∞ T (Charles’ law).

In reality, as temperature decreases volume also decreases, but before reaching an OK temperature, gases convert into liquids.

So, the relation pV = kT does not hold well for real gases at very low temperatures.

For example, at high temperatures and low pressure, molecules of real gases move far away from each other (at the maximum possible distance).

In contrast, at low temperatures and high pressure, molecules of real gases come close to each other.

Therefore we can say that at very high temperatures and very low pressures, a real gas behaves like an ideal gas.

Hence, the gases, which do not obey the equation “pV = kT” except at very high temperatures and very low pressures are called real gases.

A natural question arises: Why do we study ideal gas? Although the concept of ideality is an abstraction, to estimate p, V and T of any real gas, the relation pV=KT is undoubtedly a very useful idea under certain conditions.

WBBSE Chapter 2 Behaviour Of Gases Avogadro’s law

At constant temperature and pressure, the volume occupied by 1 mol of any gas is called its molar volume (1 mol = 6.02 x10²³ no. of molecules).

If the mol of a gas has a volume of V then the molar volume will be\(\frac{v}{n}\). Due to changes in temperature and pressure, the volume of gas changes so molar volume also changes.

Experimental results show that molar volumes \(\left(\frac{v}{n}\right)\) for different real gases are approximately the same, i.e., \(\left(\frac{v}{n}\right)\) are independent of the nature of the gas, provided temperature and pressure remain the same.

At standard temperature and pressure (STP), the molar volume of all real gases is found to be nearly equal and its limiting value is (22.4 L) or 22400 mL.

This volume (22-4 L) is called the molar volume at STP. Thus, at STP, 1 mol of H₂, N₂, CO₂, O₂ … occupies a volume equal to 22.4 L.

Based on these experimental results, in 1811 Italian chemist Amedio Avogadro discovered a law relating the volume of gas to the number of mol or molecules which is known as Avogadro’s law or Avogadro’s hypothesis.

Avogadro’s law: Under similar conditions of temperature and pressure, an equal volume of all gases contains an equal number of molecules.

Explanation of Avogadro’s law: Suppose we have a container of volume 100 mL filled with H₂ gas and the no. of molecules in it is x at a given T and p.

If H₂ is replaced by another gas say N₂ then under the same T and p, the no. of molecules will be x. Under similar conditions, 100 mL of all other gases like O₂, NH₃, CH₄, SO₂…. will have x no. of molecules.

So 200 mL of CO contains 2x molecules, 150 mL NO₂ contains \(\frac{3}{2} x\) molecules, 300 mL NH₃ contains 3x molecules, 50 mL H₂ contains molecules, etc.

According to Avogadro’s law, we can say that under similar conditions of temperature and pressure, the volume of a gas is directly proportional to the no. of molecules present in it.

Mathematically, V∞on where V = volume, n = no. of mol. Therefore \(\frac{v}{n}=k\)(constant).

During Avogadro’s time, the concept of atoms and molecules was purely hypothetical. Avogadro first explained the distinction between these two in his molecular theory.

According to him, the smallest particles of gases which can exist independently are molecules, not atoms.

In this context, it is important to mention that in Avogadro’s law ‘volume’ means the volume occupied by the gas- not the volume of gas molecules.

For example, at STP volume of 1 mol or 6-023 x 1023 molecules of a gas is very much negligible as compared to 22.4 L.

Now let us take the reaction: H₂+ Cl₂ → HCI.

Here, 1 volume of H₂ combines with 1 volume of Cl₂ to form 2 volumes of HCI. According to Avogadro’s law,

1 molecule of H₂ + 1 molecule of Cl₂ → 2 molecules of HCI

⇒ \(\frac{1}{2}\) molecule of H₂ +\(\frac{1}{2}\)  molecule of Cl₂ → 1 molecule of HCI

An atom is an individual. A group of atoms are called a molecule. That is,

H-atom + H-atom → 1 H₂ molecule⇒ 1 H-atom → ½H₂ molecule, according to Dalton’s atomic theory, this can exist and this concept was valid.

Gay-Lussac’s law of combining volumes: When gases react together to form gaseous products under similar conditions of temperature and pressure, then the ratio between volumes of reactant gases and product will always be in simple whole numbers.

Explanation of Gay-Lussac’s law using Avogadro’s law: This law is applicable only to the reaction of gases.

Let at the same temperature and pressure molecules of A gas combine with ‘b’ molecules of B gas to form ‘c’ molecules of C gas.

where a, b, and c are simple whole numbers. According to Avogadro’s law, at the same temperature and pressure equal volumes (V) of all gases contain an equal number of molecules (n).

Then, volume of ‘a’ molecules of A gas = \(\frac{\mathrm{V} \times a}{n}\)

the volume of ‘b’ molecules of B gas = \(=\frac{V \times b}{n}\)

and volume of ‘c’ molecules of C gas =\(\frac{\mathrm{V} \times c}{n}\)

The ratio of the volumes is at the same temperature and pressure =

⇒ \(\frac{\vee \times a}{n}: \frac{\vee \times b}{n}: \frac{\vee \times c}{n}\) = a: b: c.

The ratio is simple [as a, b, and c are whole simple numbers. This is Gay-Lussac’s law.]

Moist air is less dense than dry air: Dry air contains 78% N2, 21% O₂ and the remaining 1% trace gases, whereas moist air contains N₂, O₂ and H₂O (water vapour).

Molar mass of N₂, O₂ and H₂O are respectively 28 g.mol^-1, 32 g.mol^-1 and 18 g.mol^-1.

That is H₂O is relatively less dense than N₂ and O₂. According to Avogadro’s law, we know that equal volumes of all gases under the same conditions of temperature and pressure contain equal no. of molecules.

This can be understood easily with the help of imagining a container of volume V containing dry air at a certain T and p.

Imagine there are 6 molecules of N2 and 9 molecules of O₂, then the total gram molecular mass would be = 28 × 6 +32 × 9456 g.

If water vapor molecules are allowed to introduce in the same container then heavier N₂ or O₂ molecules must leave so that the total no. of molecules remains the same.

There are 5 molecules of N₂ 7 molecules of O₂ and 3 molecules of H₂O, then the total gram molecular mass of moist air becomes 28 × 5 + 32 × 7 + 18 × 3 = 418 g.

WBBSE Chapter 2 Behaviour Of Gases Combination Of Boyle’s, Charles’ And Avogadro’s Law Ideal Gas Equation

Suppose at pressure p and absolute temperature 7 the volume of nmol of gas is V. (when n and T remain constant)

According to Boyle’s law:  \(V \propto \frac{1}{p}\) (when n and p remain constant)

According to Charles’ law: \(V \propto T\) (when T and p remain constant)

According to Avogadro’s law: \(V \propto n\) (when T and p remain constant)

Combines these three relations when p, T and n are variables \(V \propto \frac{n T}{p} \Rightarrow V=\frac{n R T}{p}\)  where ‘R’ is a constant.

Putting n = 1 in the equation, we get R.

So for 1 mol of any gas, the value of R is equal and the value does not depend on the nature of the gas.

By this logic, ‘R’ is called the molar gas constant or universal gas constant.

In the equation, there is no such parameter which depends on the nature of the gas. By this logic, equation P is called the equation of state for n mol of an ideal gas or simply the ideal gas equation.

p-T graph: From the equation \(p=\frac{n R T}{V}\) we can write p= Keeping mass no. of mol (n) and volume (v) V unchanged we see that p∞ Tp = kT (k → being a proportionality constant)

⇒ \(\Rightarrow \frac{p}{T}=k \Rightarrow \frac{p_1}{T_1}=\frac{p_2}{T_2} .\)

If a graph is plotted taking absolute temp (T) along the x-axis and pressure (p) along the y-axis, we will get a straight line Extending the line backwards, it meets the temperature axis at OK (or -273°C) temperature.

Theoretically, at temp OK, the pressure of gas becomes equal to zero (but in reality, this is not possible).

Representation of equation pV = nRT as pV= \(\left(\frac{\mathbf{W}}{\mathbf{M}}\right)\) Suppose the mass of n mol of an ideal gas is W g. If its molar mass is M g.mol^-1, then no. of moles, \(n=\frac{W}{M}=\frac{\text { Given mass }}{\text { Molar mass (in gram) }}\) So, from the ideal gas equation, we can write, \(p V=n R T \Rightarrow p V=\left(\frac{W}{M}\right)\)

Note: M is the molar mass (unit: g.mol^-1)-not the molecular weight (which is a dimensional quantity).

Unit of ‘R’ from dimensional analysis of ideal gas equation: From ideal gas equation pV=nRT we get \(\mathrm{R}=\frac{p \mathrm{~V}}{n T}\) Remember, ‘T’ is always expressed in Kelvin (K) so for simplicity we can write \(R=\frac{p V}{n T(i n K)}\)

⇒ \([\mathrm{R}]=\frac{[p][\mathrm{V}]}{[n][\mathrm{K}]}=\frac{[\mathrm{F} / \mathrm{A}][\mathrm{V}]}{[n][\mathrm{K}]}=\frac{\left[\frac{\mathrm{MLT}^{-2}}{\mathrm{~L}^2}\right]\left[\mathrm{L}^3\right]}{[n][\mathrm{K}]}=\frac{\left[\mathrm{ML}^2 \mathrm{~T}^{-2}\right]}{[n][\mathrm{K}]}=\frac{[\mathrm{W}]}{[n][\mathrm{K}]}\)

Here unit of [W]= work done or energy is J.mol^-1. K^-1 and CGS unit erg. Mol^-1. K^-1. different in different systems.

So SI unit of ‘R’ Similar unit is cal.mol-1, K1. These units are mainly used for energy calculations. The most commonly used unit of ‘R’ is atm. litre/mol. K.

We know that 1 mol of an ideal gas occupies a volume of 22.4 L at STP (at temperature T = 273 K and pressure = 1 atm). Then the value of ‘R’ is found to be

⇒ \(R=\frac{1 \mathrm{~atm} \times 22.4 \mathrm{~L}}{1 \mathrm{~mol} \times 273 \mathrm{~K}}=0.082 \mathrm{~atm} \text {.litre } / \mathrm{mol} . \mathrm{K} \text {. }\)

Another value: R = 8.31 J.mol^-1.K^-1 8.31 x 107 erg. mol^-1.K^-12 cal. mol^-1, K^-1

WBBSE Chapter 2 Behaviour Of Gases Simple Numerical Problems

Wbbse Physical Science Class 10

Question 1. Find the volume of 32 g of methane at 27°C and 400 mm Hg.
Answer: 

Given: \(n=\frac{W}{M}=\frac{\text { Given mass }}{\text { Molar mass }}=\frac{32 g}{(12+4) g}=2\)

T = 27°C = (27 + 273) K = 300 K.

p = 400 mm Hg= \(\frac{400}{760} \mathrm{~atm}\)

V =?

We know that pV = nRT

So that \(\frac{400}{760}\) V =2 x 0.082 x 300

Wbbse Physical Science Class 10

Question 2. Find the pressure (in mm Hg unit) exerted by 14 g of N2 at 127°C kept in a container of volume 20L.

 Answer:

Given: \(n=\frac{W}{M}=\frac{14}{14 \times 2}=0.5\)

T= 127°C=(127+273) K = 400 K

V=30L

p=?

We know that Pv=Nrt

So that p×30=0.5×0.082×400

WBBSE Chapter 2 Behaviour Of Gases The Behaviour Of Ideal Gas The Molecular Level

We know that no real gas is ideal, but some real gases behave like ideal gas under certain conditions.

Although they have some common properties like high compressibility, expansibility, and diffusion. In the 18th century Boltzmann, Clausius, and Maxwell put forward a model applicable to all gases.

This is known as the kinetic theory of gases and is applicable to ideal gases only. Some important postulates (or assumptions) of the kinetic theory of gases are:

Gases are made up of very tiny particles called molecules (or atoms). Molecules of the same gas are identical in all respects (mass, shape, size).

The gas molecules are moving randomly in all possible directions with all possible speeds ranging from zero to infinity.

During such chaotic motion, molecules collide against each other and with the wall of the container. This gives rise to gaseous pressure exerted on the wall. 

All collisions of gas molecules among themselves and with the wall of the container are elastic in nature (which means no loss in momentum and K.E.).

The K.E. of a gas depends only upon (absolute) temperature -not upon the nature of the gas. It has been established that the temperature of a gas is directly proportional to the average K.E. of its molecules.

Note that, it is not possible to determine the speed or K.E. of a single molecule at any particular instant. 

WBBSE Chapter 3 Stoichiometric Equations Chemical Reactions

Stoichiometry is a kind of calculation that chemists use. We know that a chemical equation is a complete description of a chemical reaction by using symbols and formulae of all reactants and products.

Reactant 1 + Reactant 2→Products.

Stoichiometry helps us to find out how much reactance is used up or what product is produced in a chemical reaction.

It deals with the calculation of various quantities (no. of mol/mass/volume) of reactants and products. We need two things to do in stoichiometry,

A balanced chemical equation is a concise description in which the no. of atoms of each reactant/product remains the same on either side,

Wbbse Class 10 Physical Science Notes

Some measured values of the chemicals involved in the reaction. These values are used to find out whatever (no. of mol/mass/volume) is unknown.

WBBSE Notes For Class 10 Physical Science And Environment

(Is. 1.2. Conservation of mass in chemical reactions Stoichiometry works on the basis of the law of conservation of mass. Scientist Lavoisier proposed this law in the year 1774.

It states that in chemical reactions matter (or mass) is neither created nor destroyed. That is the total mass of reactants before the reaction = the total mass of products after the reaction.

If you have a certain no. of H atoms at the start of a reaction, you will have the same no. of H-atom at the end of the reaction.

As an example, we can consider the following reaction:

Wbbse Class 10 Physical Science Notes

However, for accurate measurement, it is seen that every chemical reaction is associated with the emission or absorption of energy (heat).

This was successfully explained by the great physicist Albert Einstein. He explained that mass and energy are interconvertible according to the relation E = mc²,

Wbbse Class 10 Physical Science Notes

Where the amount of mass disappears to produce an equivalent amount of energy and is the speed of light in a vacuum.

For example, in a reaction 1000J of energy would be released due to a loss in mass \(m=\frac{E}{c^2}=\frac{1000 \mathrm{~J}}{\left(3 \times 10^8 \mathrm{~m} / \mathrm{s}\right)^2}=1 \cdot 11 \times 10^{-11} \mathrm{~g}\).

This loss in mass is so small that it is negligible. Hence, no detectable change in mass occurs in ordinary chemical reactions.

But in nuclear fission and fusion reactions, the change in mass is released as. nuclear energy.

WBBSE Chapter 3 Stoichiometric Equations Weight Versus Weight Calculations

In stoichiometry, if we have a balanced chemical equation and some measured values then we can use this information to calculate whatever is unknown.

Simple problems based upon stoichiometric equations are of three types: 

1. Problems on the mass-mass relationship.
2. Problems with mass-volume relationship and
3. Problems related to vapor density.

Let us illustrate these three types of problems one by one through various examples.

Problems on mass-mass relationship: If the mass of the reactant is given then the mass of the product can be easily calculated and vice-versa is also true.

Wbbse Class 10 Physical Science Notes

To do so, we are to follow three steps: (1) First, we are to convert from mass (in grams) to no. of mol using the formula “no of mol \((n)=\frac{\text { mass given }}{\text { molar mass }}\)

Follow the examples: No. of mol in 48 g of CH₄ (C = 12, H = 1),

⇒ \(n=\frac{48}{12+1 \times 4}=3\);

No. of mol in 18g of H₂O is \(n=\frac{18}{1 \times 2+16}=1\)  and so on.

(2) Then from the balanced equation tries to bring out the relation among known and unknown chemicals in terms of no. of mol.

Because chemical equation gives us a molar relationship, not a mass relationship.

Finally, convert back from no. of mol to grams (mass) using the formula Required mass = (no. of mol) x (molar mass).

For example, mass of 2 mol of H₂SO₄ = 2 x (1 x 2 + 32 + 16 x 4) = 2 x 98 = 196 g.

Per centage composition of an element in a compound (by weight) :

= \(\frac{\text { mass of that element in the compound }}{\text { total mass of compound }}\) x 100

Example: 

Wbbse Class 10 Physical Science Solutions

Question 1. If 28 g of N₂ reacts with an excess of H₂ then how many grams of NH₃ will be produced?
Answer: 

Question 2. How many grams of KCIO₃ would be required to be heated to produce 48g of O₂? (K = 39, Cl = 35-5, O = 16)
Answer:  

Wbbse Class 10 Physical Science Solutions

Question 3. If 131 g of Zn reacts with dil. H₂SO₄, then find the amount of H₂gas liberated in the reaction. (Zn = 65-5, H = 1, O = 16)
Answer: 

Question 4. How many grams of CaCO₃ will be required to produce 2-2 g of CO₂ by the action of dil. HCI?
Answer: Question 5. Due to the thermal dissociation of 10 g of limestone (CaC0₃), how many grams of quick lime (CaO) and C0₂ are formed? Also, find the percentage composition of Ca in the given limestone.
Answer: Wbbse Class 10 Physical Science Solutions

Question 6. If 48 g of methane is burnt in excess air, find how many grams of C0₂ will be produced. Also, find the no. of molecules present in that amount of C0₂
Answer: 

Question 7. To obtain 85 g of NH₃, how much amount of NH₄CI is to make a reaction with an excess amount of CaO?
Answer: 

Wbbse Class 10 Physical Science Solutions

Question 8. How many grams of iron are needed to get 1 mol of hydrogen by running steam on red hot iron? (Fe = 56)
Answer: 

Wbbse Class 10 Physical Science Solutions

Problems on mass-volume relationship: If the mass of one reactant is known then the volume of a gaseous product can be calculated or if the volume of a gaseous chemical is known then the mass can be calculated, under similar conditions of temperature and pressure.

In general, this condition would be STP. This is applicable only to gases, not solids or liquids. Like mass

The mass relationship here also we are to follow three steps:

1. Convert volume (at STP) into no. of mol using the formula “no. of mol \((n)=\frac{\text { volume given (at STP) }}{\text { molar volume (at STP) }}=\frac{\text { Volume given }}{22.4 \mathrm{~L}}\)

For example no. of mol of 11-2 L of H₂O   vapor (at STP) \(n=\frac{11 \cdot 2}{22 \cdot 4}=0 \cdot 5\)

2. Then, from the balanced equation bring out the relation between known and unknown chemicals,

3. Finally, convert back from no. of mol to volume in liters at STP using the formula…

For example Volume of 2 mol of C0₂ at STP = (no. of mol) (molar volume at STP) = 2 x 22-4 L = 44.8 L.

WBBSE Chapter 3 Stoichiometric Equations Simple Numerical Problems

Question 1. When 367-5 g KCI0₃ is heated, how many liters of 0₂ gas will be produced at STP
Answer: 

Question 2. On heating some quantity of CaC0₃, the volume of C0₂ at STP is measured to be 11-2 L. Find the mass of CaC0₃.
Answer: 

Wbbse Physical Science Class 10

Question 3. Find the volume of CO₂ and H₂O vapor produced when 64g of CH₄ is combusted in presence of excess air.
Answer: 

Question 4. Calculate the mass and volume (at STP) Of O₂  involved, when 34 g of pure hydrogen peroxide decomposes.
Answer: 

Wbbse Physical Science Class 10

Molecular weight and vapor density of gas under a given condition: Gases or vapors have a very low density in comparison to the density of solids and liquids. Usually, the density of a gas is measured in comparison with the density of H₂ gas, under similar conditions of T and p.

Definition: The ratio of the density of a gas (or vapor) with the density of an H₂ gas under similar conditions of temperature and pressure is called vapor density or relative density of that gas (or vapor).

By definition,

As molar mass for any gas at STP = 22.4 L, so at STP For 1 mol gas

Remember:

1. Vapor density is a unitless, dimensionless quantity,
2. Vapor density is applicable only for gaseous states of matter,
3. M = 2D formula is applicable for an equal volume of both gases at STP.

Question 5.  The vapor density of a gas is 11.2. Calculate the volume occupied by 11.2g of the gas at STP.
Answer:

Wbbse Physical Science Class 10

Since M=2D Since M =2×11.2=22.4g

Question 6. Calculate the volume occupied by 40 g Ch4 at STP if its V.D. is 8.
Answer: 

Wbbse Physical Science Class 10

Question 7. Calculate the V.D and Molecular mass of CO₂  if 200 ml of the gas at STP weighs 0.40g. (1l of H₂  at STP weighs 0.09 g)
Answer: 

Chapter 1 Concerns About Our Environment Physical Science And Environment

Physical science explains the naturally occurring phenomena in our everyday life. With the advancement of industrialization, at present, the entire environment is under greater threat.

We are to know the reasons for the threats, as well as to be conscious of the remedial measures for sustaining life on the earth.

Wbbse Class 10 Physical Science Notes

Chapter 1 Concerns About Our Environment Structure Of The Atmosphere

The Earth is surrounded by a blanket of air called the atmosphere. It keeps the average temperature of the earth nearly constant.

The earth’s atmosphere extends up to a height of about 1600 kilometers from the surface of the earth.

WBBSE Notes For Class 10 Physical Science And Environment 

The atmosphere is held near the surface of the earth by the force of gravity and it is made up of a mixture of gas molecules like N₂, O₂, CO₂, Ar, O₃, water vapors, and dust particles.

The density (mass per unit volume) and also temperature is not the same throughout the atmosphere at all heights.

Earth’s atmosphere can be divided into four major layers separated from one another due to changes in air temperature.

The layers are:

1. Troposphere,
2. Stratosphere,
3. Mesosphere and
4. Thermosphere.

Wbbse Class 10 Physical Science Notes

1. Troposphere: This is the layer where we all live in. It extends from the ground roughly 12 kilometers above the earth’s surface. About 80% of the total mass of the atmosphere is contained in this layer.

Almost all dust particles and water vapors are present in this layer, so all the weather phenomena like storms, clouds, rains, air currents, etc occur in this layer.

Components like O₂, CO₂, and N₂ useful for sustaining life are available here.

The temperature in the troposphere is constantly changing and it falls with increasing altitude at an average rate of 6-5°C per kilometer of height above the earth’s surface.

The upper limit of the troposphere at which the temperature stops decreasing is called tropopause. The temperature at this level may be as low as – 56°C.

2. Stratosphere: It is the second layer of the atmosphere. Above the tropopause, it extends up to 45 km above the earth’s surface.

The tropopause separates the troposphere from the stratosphere.

There is very little air (N₂, O₂) but practically no moisture in this layer, so no cloud is formed here and disturbances like storms, and tornadoes do not take place here.

Due to this fact, the layer is calm and quiet. Jet planes fly through this layer as the stratosphere is free from all weather disturbances. Due to a lack of O₂, normal breathing is difficult in the upper levels of the stratosphere.

The farther side of the stratosphere is called the ozonosphere and this part contains most of the ozone (O₃) gas molecules of the atmosphere.

The O₃ molecules absorb most of the harmful UV rays of the sun and protect life on the earth. Thus, the ozonosphere acts as a filter.

The absorption of UV rays warms the stratosphere from – 56°C to 3°C. Here also temperature increases with the rise in altitude.

The end of the stratosphere where the temperature remains almost constant is called stratopause.

O₃ molecules keep the upper layer of the stratosphere warm and it contains too little O₂. While in the lower part of the stratosphere, no UV light-all is absorbed by O₃ molecules in the upper layer.

Wbbse Class 10 Physical Science Notes

3. Mesosphere:  Above the stratosphere, the layer of atmosphere lying between 45-85 kilometers from the earth’s surface, is called the mesosphere.

In this layer, very little quantity of N₂, O₂, O₂, and NO₂ is present. In the mesosphere, the temperature drops again with altitude, reaching a minimum of 92°C.

It is the lowest temperature in the atmosphere. Thus, the mesosphere is the coldest layer. Because of the very low temperatures, ice crystals can be formed here.

It protects the earth from being struck by meteors. The upper end of the mesosphere where the temperature remains almost constant is called mesopause.

It acts as the boundary between the mesosphere and the thermosphere.

4. Thermosphere: The layer of the atmosphere above the mesosphere extending about 85-500 kilometers above the earth’s surface is called the thermosphere.

It is the layer where the air is very very thin. There is nothing protecting this layer from the heat of the sun, that’s why its temperature increases with altitude up to about 1200°C.

Wbbse Class 10 Physical Science Notes

It has two different parts:

1. ionosphere
2. Exosphere.

1. ionosphere: The lower part of the thermosphere consists of ions of lighter atoms like NO+, O+, electrons, etc.

Because of the presence of ions, this particular zone becomes highly electrically conducting, making radio communications possible around the world. The northern lights or aurora borealis are seen here.

2. Exosphere: It is the upper portion of the thermosphere, extending somewhere about 500- 1000 kilometers above the earth, having traces of lighter gas molecules like H₂, He, etc.

The highest temperature of this layer is about 1200-1700°C. It is colder at night and much hotter during the day. Artificial satellites and space stations operate from this portion.

Beyond the exosphere, there is practically no air. It is almost a vacuum.

Variation of pressure with altitude: Normally, the pressure of the atmosphere is maximum at the lowest layer and it decreases as the height above the earth’s surface increases.

At the ground, the air pressure is about 1.013 atm (where 1 atm = 76 cm of Hg). The variation of pressure with altitude.

Convection currents and storms: Consider the first 50 kilometers where 99% of the atmosphere exists. Sunlight passes through the atmosphere and warms the ground heating up the air adjacent to the earth’s surface.

It expands and becomes less dense. While the cold air from above sinks, it pushes up the warm air to take the vacant space.

This action sets up the flow of air which is called convection currents. Sea breeze, and land breeze flow due to the convection of air.

In liquids and gases, heat is transmitted by convection as their molecules are quite free to move about.

On the other hand, molecules in a solid, being closely packed, cannot move like molecules of liquids and gases. So, convection currents does not occur in solids.

All kinds of wind are caused due to uneven heating of the earth. When the air moves at a high speed over a surface, it results in decreasing pressure.

Thus, strong winds start blowing from the high-pressure area to the low-pressure area. This difference in air pressure results in the formation of a storm. Higher wind speed creates cyclones/tornados/typhoons/hurricanes.

Chapter 1 Concerns About Our Environment The Ozone Layer

Ozone (O₃) is an allotrope of oxygen. Its molecules are made up of 3 oxygen atoms bonded together. The ozone layer in the stratosphere absorbs most of the harmful UV rays coming from the sun and protects life on Earth.

Formation of the ozone layer: Ozone (O₃) is formed in the upper part of the stratosphere by a photochemical reaction of O₂.

UV-rays of small wavelength 240 nm on being absorbed by O₂ molecules in the stratosphere splits into 2 fast

Moving O-atoms. Each O-atom collides with air molecules (N₂, O₂) around them. This slows down the free O-atoms, and abling them to make bonds with O, molecules. In this way, O₃ molecules are formed.

Depletion of the ozone layer: O₃ molecules break down again by absorbing UV-radiation of high wavelength 300 nm and O₂ is back again.

The density of the ozone layer is measured by the Dobson unit (DU). 1 DU refers to the 0.01 mm thick ozone layer at STP.

In this way, O₃ is continuously formed and broken by UV light in the ozone-oxygen cycle. The thinning of the ozone layer in the upper atmosphere is called ozone layer depletion.

The main cause of it is the release of chlorofluorocarbons (CFCs mainly CF₂, CI₂, and CFCI₃) used in refrigeration, air conditioning, fire extinguishers, foam production, and aerosol sprays and also by nitric oxide (NO), nitrogen dioxide (NO₂).

The thickness of the ozone layer is thus reduced. The thinning looks like a hole, called the ozone hole. This was observed by scientists in the 1980s over Antarctica.

In the Montreal Protocol (an international treaty), on 1 January 1989, thirty nations worldwide agreed to reduce the usage of CFCs and encouraged other countries to do so.

Role Of Cfcs In The Destruction Of Ozone Layer:

CFCs are comparatively inactive in the lower atmosphere, but highly reactive in the stratosphere.

When CFCs are hit by photons, they break up into ‘free chlorine’ (CI) atoms which grab one of the atoms of the O₃ molecule.

As a result, O, molecules are broken down very rapidly (1 Cl-atom can break nearly 1000 no of O₃ molecules). This led to the depletion of the ozone layer.

Role Of No, No, In The Destruction Of Ozone Layer:

⇒ \(\mathrm{NO}+\mathrm{O}_3 \rightarrow \mathrm{NO}_2+\mathrm{O}_2 ; \mathrm{O}_2 \stackrel{\text { UV-rays }}{\longrightarrow} \mathrm{O}+\mathrm{O}\)

Sources of NO, NO₂ are supersonic jet planes, and nitrogenous fertilizers. Also, NO emitted from transportation oxidizes into NO₂. However, NO reacts with O₃ molecules in the stratosphere to form O₂ and NO₂.

The O₂ molecules break down to atomic O by absorbing UV radiation. NO₂ reacts with atomic O and forms NO back.

So, the amount of NO remains the same. As a result, O molecules get decomposed. In this way, the concentration of ozone in the stratosphere is destroyed.

Effects of ozone depletion on living bodies: 4th Proof User 5 (ASHIS) For humans Direct exposure to UV radiation can lead to serious health problems like sunburn, skin cancer, skin aging, cataracts, weak immune system, etc.

For animals: Skin cancer, eye problems, etc.

For plants, UV rays can affect the photosynthesis process, growth rate, flowering, etc. As a whole, it can destroy the total life cycle.

For marine life: Zooplanktons and phytoplanktons (the base of the food web) are also affected by strong UV rays.

Chapter 1 Concerns About Our Environment Greenhouse Effect And Global Warming

The earth is getting an enormous amount of heat radiation from the sun in the form of UV, visible, and IR rays.

While passing through the earth’s atmosphere, the ozone layer blocks some parts of UV rays. The remaining rays reach the earth’s surface.

A part is absorbed by the earth and especially long wavelength IR-rays are reflected back. But some gases in the atmosphere, trap the heat radiation and obstruct to escape to space.

This leads to the warming up of the earth’s surface and also the atmosphere. This effect is known as the greenhouse effect.

The increase in the average temperature of the earth due to the greenhouse effect is known as global warming.

The gases involved in warming up the earth are known as greenhouse gases.

Question 1. Why it is called the greenhouse effect?
Answer:

Greenhouse effect

A greenhouse is a glass building in which plants that require protection from cold weather are kept. The glass cover traps all the heat radiation but cannot allow longwave IR rays.

So the inside environment keeps warm. For our planet Earth, the atmosphere acts as the glass cover and traps the heat.

Utility of greenhouse gases: If there were no greenhouse gases to trap the sun’s radiation, the temperature of the earth’s surface would have been -20°C~-30°C. It would have been impossible for the survival of living beings.

Greenhouse gases: The gas molecules of CO₂, H₂O vapor, CH₄, N₂O, CFCs, etc. absorb the sun’s IR radiation.

Among these, CO₂ is the main greenhouse gas. They trap too much heat and the earth becomes hotter and hotter.

Remember: N₂ and O₂ which form most of the atmosphere do not absorb IR radiation. So N₂ and O₂ are not greenhouse gases.

Causes of global warming: Both natural and human activities are responsible for the emission of greenhouse gases such as volcanic eruption, plants, forest fires, deforestation, smoke from vehicles & industries, etc.

Effects of global warming:

With the increase in the earth’s temperature, the glaciers melt faster than they accumulate new snow. As a result, the water level in oceans and seas rises up.

It causes flooding of lowlands, agricultural soil contamination with salt, soil erosion, etc.

From rising water levels more water vapors are evaporated-this leads to the formation of hurricanes/cyclones.

Because of global warming, birds, polar bears, and different insects migrate to move from one area to another.

Global warming helps mosquitoes to grow vigorously in tropical regions, and it causes the spread of diseases like dengue, chikungunya, malaria, etc.

How can we contribute to reducing global warming?:

1. Planting more and more trees, to decrease the emission of CO₂,
2. Using renewable energies like solar energy, wind energy, hydropower, etc.
3. Reducing the use of products that generate greenhouse gases.

Chapter 1 Concerns About Our Environment Calorific Value Of Fuels

Fuels are substances that can burn or which can undergo combustion in excess air (oxygen) such as coal, cow-dung cakes, matchstick, petrol, diesel, kerosene, alcohol, LPG (Liquefied Petroleum Gas), CNG (Compressed Natural Gas), etc.

So, fuels can be solid, liquid, or gas. All the fuels produce heat on burning. But, different fuels produce different amounts of heat on burning.

Such a difference in producing heat energy i.e. fuel efficiency is expressed in terms of its calorific value. The quality of fuel depends on its calorific value.

The total amount of heat produced in kilojoule by complete burning or complete combustion of one kilogram (or one gram) of solid or liquid fuel in excess air (O₂) is called its calorific value.

It is expressed in the unit of ‘KJ per Kg’ or ‘KJ per g’. In the case of gaseous fuel, it is the total amount of heat produced by the complete combustion of 1 m3 of its volume at STP the calorific value of some common fuels.

7 Higher calorific value of fuel means that the more heat it produces on burning, the better fuel it is.

The necessity of conservation of fossil fuels: Most of our energy requirements are met by the burning of fossil fuels like coal, petroleum, and natural gas.

Today we use them carelessly, without thinking that fossil fuels are non-renewable sources of energy that do not develop in a day or two but rather will take millions of years to form.

If we continue to burn these fuels at this rate, it leads to an increase in CO₂ in the environment. This contributes to greenhouse gases and ultimately global warming.

So, the lesser would be the consumption, the lesser would be the pressure on the next generations.

Concept of sustainable development: Sustainable development refers to the development that will allow our future generations to lead a quality lifestyle same as what is being enjoyed by the current generation by utilizing all natural resources. We can use fossil fuels in fewer amounts only when required.

Harnessing alternate sources of energy for sustainable growth and development: Fossil fuels are non-renewable sources of energy. They are used in almost 90% of areas.

These, on combustion, produce harmful substances like C, CO, oxides of nitrogen (NO), oxides of sulfur (SO), soot, etc. which pollute the air around us.

Their stocks are also very limited. So, misuse of these fuels should be avoided.

Rather, we are looking for alternate or renewable or non-conventional sources of energy such as solar energy, wind energy, tidal energy, geothermal energy, biomass energy-biofuel, biogas, etc.

Solar energy: The Sun is the main source of all energy. The energy obtained from the sun is called solar energy.

Today scientists are able to make solar cookers, solar water heaters, and many other devices that work on solar energy.

A single solar cell can produce electricity on a very small scale. The solar cell is used in watches, transistors, calculators, and for domestic purposes in remote village areas. Solar batteries are used in space flight.

To obtain more electricity and more electric power, a large number of solar cells are joined together to form a solar panel.

These are used in artificial satellites, space stations, water pumps, street lighting, traffic signals, etc.

Wind energy: The energy obtained from very fast-moving wind is called wind energy. It is a natural, conventional, renewable source of energy.

It is available in high-wind regions without any cost. It also causes no pollution.

Nowadays high-speed wind energy is used in a wind turbine to generate electricity. The basic principle is here the same as any power plant. Here, the turbine is connected to the windmill.

As wind strikes the blades of the windmill, it starts rotating. Hence, the turbine and the generator start rotating and electricity is generated. Also, windmills are used to grind grains.

The working of a windmill depends on the speed of the wind.

Tidal energy: The energy obtained from rising tides in coastal areas is called tidal energy. It’s a renewable source of energy.

In general, it is not enough to generate electricity on a large scale with tidal energy.

But, it is harnessed for generating electricity by constructing a dam across a comparatively narrower opening to the sea.

Geothermal energy: It is the energy obtained from the heat of the core of the earth. In a few places, rocks under the earth’s surface remain very hot, and the heat turns underground water into steam which is compressed at higher pressure between the rocks.

Drilling holes through pipes into such places, steam is ejected through the pipes, which in turn, rotates the turbine to generate electricity.

In our country, Madhya Pradesh, there is one such power station.

Biomass energy: A mixture of waste materials like wood filings, crop residues, garbage, excreta of animals, sewage, dead parts of plants, trees, animals, etc is known as biomass.

It mainly contains carbon-rich compounds. The chemical energy stored in biomass is known as bioenergy or biomass energy. Biomass can be used as a source of energy.

For example-

Dry biomass, like cow dung cake, is burnt to produce heat energy, which is used for cooking purposes.

In biogas plants, cow dung and sewage are used to produce biogas. The residue is used as organic manure in the fields.

Biogas: It is prepared in biogas plants. When the slurry of cow dung and water (or a mixture of crop residue, sewage, or other waste materials) is allowed to make anaerobic fermentation in the absence of oxygen but in the presence of water,

They produce a mixture of methane, carbon dioxide, hydrogen, and traces of hydrogen sulfide. This mixture is called biogas.

The fermentation is carried out by anaerobic bacteria known as methanogenic bacteria which decompose the carbon compounds present in biomass (organic wastes) into methane gas and carbon dioxide.

Biogas is a renewable source of energy. It can never replace natural gas.

Uses: Biogas mainly contains methane (about 75%) which is an excellent fuel for cooking. It causes no environmental pollution. It can be used for street lighting and for running engines.

The waste products such as nitrogen and phosphorus left from biogas plants are used as fertilizers.

Advantages of using biogas as fuel:

On burning, biogas does not produce any smoke or ash, which can be used for domestic cooking. The calorific value of biogas is very high (about 35000-40000 KJ/Kg).

Electricity from waste:

Tons of organic waste, foods, and unsold vegetables in everyday market pollute our environment and are used to turn into electricity in biogas plants.

These wastes are converted into a slurry and put into a large container (pit).

Using an anaerobic digester, the waste is converted into biofuel and put into a biogas generator which converts it into electricity. This type of project can make life more sustainable.

Biofuel: It is a form of biomass energy produced from biomass. Its examples are bio-ethanol (liquid) and bio-diesel (gaseous like gobar gas).

Bio-Ethanol: It is an alternative to petrol in road vehicles. It is produced by the sugar fermentation process of sugarcane, corn, rice straws, wheat, etc.

Bioethanol is added to petrol to minimize environmental pollution. It is carbon neutral and produces fewer greenhouse gases.

Biodiesel: It is an alternative to diesel. It is manufactured by the transesterification process from vegetable oil, animal fats, or unused oils from households. Bio-diesel can be used as a blend with regular diesel.

Coal-bed methane or CBM: In coal mines, plenty of methane gas (>95%) remains adsorbed in hard layers of coal. It is called coal-bed methane (CBM).

It is formed during the process of coalification by the transformation of plants into coal.

The coal-bed methane is a form of natural gas extracted from the coal bed by drilling steel pipes 100 -1500 m deep below ground into the layers of coal.

It is now being considered an alternative source for meeting household energy requirements.

Methane hydrate: It is a solid compound of methane. It is mainly found in polar regions under the ice and also under the sea level.

In methane hydrate, the density of methane is very high (about 160 m3 of methane remains in a block of methane hydrate of volume 1 m3).

Methane hydrate can rapidly dissociate with an increase in temperature and a decrease in pressure.

Other names of methane hydrate are methane clathrate, hydro methane, fire ice, etc. The name is ‘fire-ice’ because methane in it starts burning in contact with fire It appears as if ice is catching fire.

Due to the abundance of methane, the compound methane hydrate can be considered a source of energy. Scientists predict, it to be one of the largest sources of energy in the future.

If by adopting effective technology we are able to extract methane hydrate that could change the future energy needs by the replacement of coal, oil, and natural gas.

WBBSE Chapter 8 Organic Chemistry Organic Compounds Are Compounds Of Carbon

In the past when there was little knowledge about chemistry, scientists classified compounds on the basis of their sources.

They thought that substances like starch, sugar, proteins, lipids, acetic acid, etc. could only be obtained from plants and animals (living organisms) and called them organic compounds.

On the other hand, the compounds like salts (NaCI), sulfates (CaSO₄), nitrates (KNO₃), carbonates (CaCO₃), etc.

Are obtained from minerals and non-living organisms and they are called inorganic compounds. Lavoisier also accepted this type of idea.

Later, in 1809, Berzelius explained in his vital force theory that organic compounds are only synthesized by living organisms by some unknown vital force and they cannot be prepared in the laboratory.

This old belief has been changed in 1828 when a German scientist Friedrich Wohler was successful to prepare the first organic compound urea in the laboratory.

WBBSE Notes For Class 10 Physical Science And Environment

He took an inorganic salt ammonium cyanate which after heating converted to urea. For this conversion, no vital force was required.

Already Lavoisier proved in 1784 that the main constituent of organic compounds is carbon.

In 1845, Kolbe prepared acetic acid and in 1856, Berthelot prepared methane (bio-gas) in the laboratory.

After that, the old concept of chemistry has been changed. Scientists accepted that all organic compounds essentially contain carbon and “organic chemistry is essentially the chemistry of carbon compounds.”

Exceptions: Oxides of carbon (CO, CO₂), metallic carbonates (CaCO₃), carbides (AI₄CI₃), metal cyanides (NaCN), etc. They are inorganic compounds.

Organic compounds: The compounds which contain carbon atoms excluding oxides of carbon, metallic carbonates, bicarbonates, carbides, and cyanides are called organic compounds.

They show some special properties like isomerism and catenation which the inorganic compounds do not show.

Difference between organic and inorganic compounds:

Organic compounds must contain carbon, but inorganic compounds may or may not contain carbon.

Organic compounds do not form ionic bonds.

They contain covalent bonds whereas inorganic compounds (like table salt) mostly contain ionic bonds.

Due to the difference in chemical bonding, certain special/different qualities are seen in organic compounds such as-

Most of organic compounds are volatile in nature,

The m.p. and b.p. are relatively lower than those of inorganic compounds,

Organic compounds are usually insoluble in water but soluble in organic solvents like benzene, alcohol, etc.

They are non-electrolytes as they are not ionized in an aqueous or fused state,

They have the property of catenation and isomerism,

Most of them form covalent bonds, due to which they react slowly in comparison to inorganic compounds.

WBBSE Chapter 8 Organic Chemistry Tetravalency And Catenation Property Of Carbon

Tetravalency of carbon: The electronic configuration of ₆C is = 2 (K) + 4 (L).

It has 4 valence e¯s in the outermost shell. How does C achieve its stable state (duplet/octet state)? Is it either by losing 4e¯s or gaining 4e¯s?

But it doesn’t do any one of this. Because to lose 4e¯s, it would need extra energy and to gain 4e¯s it would be difficult for 6ps to hold 10e¯s.

So for the C atom, transfer of e¯s is not possible. C atom likes to share its e^–s. That’s why, C forms 4 bonds.

C is tetravalent, has valency = 4 i.e. its combining capacity is relatively high. It can combine with 4 other atoms having valency 1 or 2.

In the simplest organic compound methane molecule, a 1C atom is surrounded by 4 bonds.
Scientist Kekule thought that the structure is planer— something looks like where the bond angle (angle between 2 bonds) should be 90°.

But this general model couldn’t explain all properties of organic compounds.

Tetrahedral model: In 1874, Van’t Hoff and Le Bel explained the properties of some organic compounds while trying to rationalize experimental data.

According to this explanation, the structure in methane is tetrahedral (a 3D structure). In case of which all the 4 bonds are not visible.

3 bonds out of 4 are above the plane (towards the observer) and 1 bond is below the plane (away from the observer).

Due to the tetrahedral structure, the bond angle is not 90° it is 109.28′. Later, by X-ray analysis, the model has been proved correct.

Catenation property of carbon: The self-linking property or tendency of a C-atom to join with another C-atom with a single covalent bond to form a long chain-like structure is called catenation.

Different types of chains can be formed like straight/open chains, branched chains, cyclic closed chains,s or a combination of all these.

For example:

C is tetravalent in nature for which other atoms like another C/H/O/N/F/CI/Br/l/S and so on can join with it and thus different organic compounds are formed.

For example:

Remember:

Most of the time, C and H atoms are joined. C can form bonds in different forms. It could be a single bond/double bond/triple bond.,

(Each covalent bond consists of 2 e¯s).

For example:

(Here the 3D tetrahedral model is projected in the 2D plane).

Structure: Trigonal; Bond angle = 120° (All 3 bonds are visible).

Structure: Coplaner; ingle triple bond; Bond angle = 180° (linear)

WBBSE Chapter 8 Organic Chemistry Structures Of C₂H₆, C₂H₄, C₂H₂

Millions of organic compounds are there. About 10 million C-compounds have already been discovered and still new compounds are being found and synthesized.

Hydrocarbons are the simplest organic compounds.

They contain only C and H atoms—no other elements. The open-chain hydrocarbons are divided into three groups —

1. alkanes,
2. alkenes and
3. alkynes.

(Here ‘o’, ‘e‘, and V are in alphabetical order).

Hydrocarbons are of two types—

Saturated (means less reactive) and unsaturated (very reactive).

The saturated hydrocarbons are those in which the C-atoms are connected by single bonds only (like C – C) while in unsaturated hydrocarbons the C-atoms are connected by at least one double bond (C = C) or triple bond (C = C).

Saturated hydrocarbons are also called alkanes, but unsaturated hydrocarbons are of two types alkenes and alkynes.

Structural and molecular Formula:

In alkanes, no. of H-atoms = 2 x no.of C-atoms +2; in alkenes,  no.of H-atoms = twice the no. of C-atoms, and in alkynes, no. of H- atoms = 2 x no. of C-atoms – 2. For example hex and (n = 6) = I
C6H14 Pent ene (n = 5) = C₅H₁₀, hex Yne (n = 6) = C₆H₁₀ etc.

Note:

(1) Methane is the simplest alkane, similarly ethene (or ethylene) is the simplest alkene, and ethyne (or acetylene) is the simplest alkyne.

(2) As a whole, the alkanes, alkenes, and alkynes are collectively known as aliphatic hydrocarbons.

WBBSE Chapter 8 Organic Chemistry Functional Groups

We see a wide variety of properties in C-compounds. Because elements other than C and H like 0, N, S, Cl, Br, F … etc. are present in them.

Definition: Functional groups are the groups of atoms that determine all characteristic properties (both physical and chemical) of the compounds in which they are present.

The presence of functional groups determines the properties of C-chains (irrespective of straight or branched).

There are many different functional groups like — OH (alcohol), — CHO (aldehyde), — COOH (carboxylic acid), — NH₂ (amine), etc.

Note: C = C double bond in alkenes and C = C triple bond in alkynes are usually considered functional groups because of which the alkenes and alkynes participate in an addition reaction.

WBBSE Chapter 8 Organic Chemistry Isomerism

The term Isomerism’ originated from two Greek terms Isos→ meaning same or equal and meros → meaning parts. In organic chemistry, isomerism is a very important concept.

There are many C-compounds that have the same molecular formula but different physical and/or chemical properties they are called isomers of each other.

Isomers are made up of the same type of atoms but the atoms are arranged differently.

Basically, two different types of isomerism are seen —

1. constitutional isomerism and
2. stereoisomerism.

Among these, constitutional isomerism is of different types. Here we will read about two types of constitutional isomerism –

Functional group isomerism and

Positional isomerism.

Functional group isomerism: Same molecular formula but different functional groups having different structures.

Common examples of functional group isomers:

We already got the idea that if functional groups are part of a compound, the entire compound will take the properties of functional groups.

For this reason, functional group isomers show properties different from each other. it is cited here-

Positional isomerism: In positional isomerism, the difference in structure occurs due to the position of the functional groups must having the same parent chain. Example:

Both have a parent chain of 3 C-atoms.

Molecular Formula: C₃H_gO In the 1st compound, the functional group (- OH) is attached to the 1st C-atom, whereas in the 2nd compound, with the 2nd C-atom.

So the position of the functional group is different.

Note: Positional isomerism also occurs due to unsaturation (position of double/triple bond in the compound).

WBBSE Chapter 8 Organic Chemistry Homologous Series Homo→Means Same Functional Groups

It is the series of compounds having the same general formula, the same type of molecular structure, and similar chemical properties which appear in immediate succession of molecular mass/Compounds in homologous series (alkanes/alkenes/alkynes) are called homologs.

Characteristic properties of homologous series:

They basically differ in no. of C and H atoms and also molecular mass.

Successive compounds in this series differ by – CH₂.

Physical properties like m.p., b.p., and solubility of the compounds that exist in the same series vary in succession.

But the chemical properties are definitely similar as they are dependent on the functional groups present.

The compounds can be artificially prepared by the same method.

For examples: 

WBBSE Chapter 8 Organic Chemistry IUPAC Nomenclature Of Simple Organic Compounds

Need for IUPAC nomenclature: Carbon has millions and millions of compounds.

Different common names are used for naming organic compounds over the years in different countries.

Common names are like a nickname (what your family members or friends call you). They are not systematic or random.

In 1967, a group of scientists decided to follow a very systematic scheme for naming the organic compounds which are known as IUPAC nomenclature

(Full form: International Union of Pure and Applied Chemistry). IUPAC naming scheme is based on the number of C-atoms in the longest chain/position of functional groups/nature of bonds/…. etc.

Basic IUPAC naming rules:

Root word: First count the number of C-atoms in the longest continuous C-chain. If the compound contains lC-atom, the name starts with meth, for 2C-atoms → eth, 3C -atoms → prop,…. etc. This gives the 1st part of the IUPAC name.

Suffix: Used after root word,

Primary suffix? Check the type of bond between the C-atoms. This gives the 2nd part of the IUPAC name.

Secondary suffix: The name of functional groups comes after the primary suffix.

While combining the IUPAC suffix remove ‘e’ from the alkane,

Make numbering in C-chain forward and backward and chose the lowest possible number for indicating the position of functional groups, Between the number and the letter always use a dash (-). (d) No. 1 in between two parts of the IUPAC name is not necessary.

let us take a few examples:

WBBSE Chapter 8 Organic Chemistry Industrial Source And Major Uses Of CH₄, C₂H₄, C₂H₂, LPG And CNG

[Methane]→

1. Colourless  Tasteless Odourless Non-toxic
2. Highly flammable
3. Abundantly available in marshy areas like wetlands, and swamps where organic substances are decomposed in the absence of O₂ by micro-organisms
4. Naturally present in natural gas (about 90%)
5. second most abundant greenhouse gas after CO₂.

Industrial source of CH₄:

 By destructive distillation of bituminous coal (continuous heating in the absence of air), a gaseous mixture (called coal gas) comes out.

1. CH₄ occurs in about 30 – 50% of coal gas.
2. Biogas or gobar gas contains about 60% CH₄.
3. CH₄ is a major component ( ~90%) of natural gas.

Uses of CH₄:

1. CH₄ is primarily used as a fuel for domestic uses, industries, automobiles, etc.
2. Refined liquid CH₄ is used as rocket fuel.
3. For the production of methyl alcohol, formal dehyde, and methanoic acid by slow combustion (catalytic oxidation) of methane.
4. In manufacturing carbon black, printing ink, etc.

[Eythlene]→ Colourless

1. Sweat taste and odor
2. Natural source of ethylene is petroleum or crude oil.
3. Naturally occurring hormone for plants.

Industrial source of C₂H₄:

By steam cracking of petroleum hydrocarbons at very high temperatures in the presence of a catalyst.

Fractional distillation of petroleum gives a mixture of gases from which C₂H₄ can be separated.

Uses of C₂H₄:

1. Widely used in the production of polymers like polythene, synthetic rubber, etc.
2. Used as a refrigerant, especially in LNG liquefaction.
3. Artificial ripening of fruits.
4. Oxy-ethylene flame in metal cutting, and welding.

[Eythlene or Acctylene] ⇒

1. Colourless, odorless when pure
2. Industrial acetylene is
3. Extremely explosive
4. Produces oxy-acetylene flame (~ 3100°C)
5. Non-polluting

Industrial source of C₂H₂:

1. By fractional distillation of crude oil.
2. The thermal cracking process by raising the temperature of some suitable hydrocarbons.
3. Hydrolysis of calcium carbide.

Uses of C₂H₂:

1. For welding and cutting.
2. Used in carbide lamp to get light.
3. Production of several inorganic compounds by chemical synthesis of C₂H₂.
4. Used in the production of different variants of plastic like PVC etc.
5. (Stands for Liquefied Petroleum Gas) A mixture of (45%) propane (C₃H₈) and (55%) butane (C₄H₁₀)
6. Colourless
7. Odourless but ethyl mercaptan added to it gives a strong smell for detecting any leakage
8. It’s a liquid under high pressure and very low temperature and turns back into gaseous vapor on releasing pressure
9. Heavier than air.

Source: LPG is extracted from petroleum/crude oil.

Uses:

1.  LPG has a higher calorific value (~ 94 MJ/m³) than other conventional fuels and is widely used as cooking gas.
2. Also in the petroleum Industry.

CNG⇒ Compressed Natural Gas

Compressed at very high pressure in the cylinders

To avoid transportation problems, it is delivered with the help of pipelines

Composition: Natural gas-basically methane (~ 85 – 90%) with (15 – 10%) ethane, propane, and butane.

Source: Natural gas.

Uses: CNG forms less CO₂ and footprints. That’s why it is used as an alternative to petrol/ diesel and can.be used in transportation by bus/taxi.

Advantages of using LPG and CNG: Higher calorific value than conventional fuels like petrol/diesel. Produces no smoke, no soot, and almost no unburnt C-particles.

WBBSE Chapter 8 Organic Chemistry Reactions of CH₄, C₂H₄, C₂H₂

Some reactions of methane (CH₄):

Combustion: Methane is highly flammable but it is not a supporter of combustion.

In a free supply of air (presence of O₂), CH₄ burns with a non-sooty blue flame and produces CO₂ (g) and water vapor.

CH₄ + 2O₂ CO₂ + 2H₂O + Heat

The reaction is exothermic (produces a large amount of heat). No C-particles are produced (i.e. pollution-free). That’s why, CH₄ can be used as a good fuel.

Substitution reaction with chlorine: CH₄ is a saturated hydrocarbon for which it is less reactive. It shows only substitution reactions.

In diffused sunlight (not in direct sunlight), CH₄ reacts with Cl₂ slowly in which 4 H-atoms of CH₄ are^ubstituted by Cl-atom and produce different chloro compounds step by step. \(\underset{\text { sunlight }}{\stackrel{\text { Diffused }}{\longrightarrow}}\)

 Some reactions of ethylene or ethene (C₂H₄): It is an unsaturated hydrocarbon (i.e. very much reactive) because of a double bond in it. It takes part in additional reactions.

(1) Addition of H₂ or hydrogeneration: At 200°C, in presence of Ni-catalyst, C₂H₄ reacts with H₂ to produce ethane (unsaturated hydrocarbon)

In this reaction, the double (=) bond of C₂H₄ is broken into a single (-) bond i.e. unsaturated compound is converted into the saturated compound.

Addition reaction with bromine: In the presence of a non-polar solvent like carbon tetrachloride (CCI₄), Br₂ combines with ethylene to produce ethylene di bromide.

In this reaction, the red color of the bromine solution is decolorized. This reaction can be used as a test for the unsaturated nature of ethene.

Some reactions of acetylene or ethyne (C₂H₂): It is an unsaturated compound having a triple bond (=). It is also very reactive and does additional reactions.

Addition of H₂/hydrogenation: In the presence of Ni-catalyst, at 200°C temperature, C₂H₂ is added with 2 mol H-, and produces ethane (a saturated compound).

Addition of Br₂in the presence of CCI₄solvent: This reaction is called

Polymerization of ethylene: Ethene or ethylene is an alkene where 2 C-atoms are joined by a double In this reaction, the red color of the bromine solution is decolorized.

This reaction can be used as a surer test for unsaturated compounds like ethene and ethyne bonds (=).

Polythene or polyethylene is formed when this double bond breaks at a temperature of about 200°C and very high pressure of around 2000 atm in the presence of an organic peroxide (not H₂O₂) catalyst.

When the double bond breaks, adjacent other molecules link together to form a very large (solid) molecule (macro-molecule) having many small repeating units with very high molar mass.

The repeating unit is called the monomer unit and the long molecular chain is called the polymer.

In this reaction, ethylene is the monomer and polyethylene is the polymer.

Homopolymer is made up of only one type of monomer unit, while copolymer is made from two or more types of monomer units.

The process of formation of polymers is called polymerization.

General Difference Between Ethylene and polythylene:

WBBSE Chapter 8 Organic Chemistry Synthetic organic polymers

On the basis of origin, polymers are of two types:

Natural polymers like starch, cellulose, DNA, protein, silk, etc. which are obtained from natural sources,

Synthetic polymers are artificially prepared in the laboratory. Examples of synthetic polymers are plastic, nylon, Teflon, PVC, polystyrene, etc. In our daily life, we use many types of polymers.

WBBSE Chapter 8 Organic Chemistry Biodegradable Polymers

Environmental impact of using synthetic polymers: The biggest problem in using synthetic polymers is that they are not degradable biologically through micro-organisms so the soft drink bottles or the plastic wrappers what we throw away are not biodegradable.

They are not destroyed for hundreds of years. In addition to this, when the rejected plastic materials are burnt toxic fumes directly mix with the atmosphere, the CO₂(g) also contributes to global warming.

As a whole, they cause a devastating impact on the environment.

Biodegradable polymers: Materials that are designed such that they get decomposed wholly when exposed to micro-organisms (fungi and bacteria) through aerobic and/or anaerobic processes.

Many natural biodegradable polymers are formed naturally by plants.

Examples: Cellulose, starch, silk, wool, and natural rubber.

Examples of synthetic biodegradable polymers: are bioplastic or green plastic (plant-derived materials), polylactic acid, bio pol resin (PHB + PHV), polyglycolide (polyglycolic and lactic acid,

PHB- Polyhydroxy butyrate; PHV – polyhydroxy venerate).

Long polymer chains are broken down by bacterial action when disposed of without producing toxic substances in the environment.

This is the main advantage of using biodegradable polymers. There are many applications of biodegradable polymers.

Examples: Polyglycolic and lactic acid→heart repair, Dextron→post-operative stitches, Polylactic acid, and lactic acid→drug delivery. –

To avoid non-biodegradable polymers, we can use jute and paper in packing. The use of jute and paper is safe and eco-friendly.

WBBSE Chapter 8 Organic Chemistry Uses And Properties Of Ethyl Alcohol And Acetic Acid

Properties and uses of ethyl alcohol or ethanol (H₃C – H₂C – OH):

Physical properties:

1. Colourless
2. Pleasant sweety
3. Volatile liquid
4. Highly soluble in water

Chemical properties: Ethanol is slightly acidic in nature. It has a tendency to lose H⁺

Reaction with sodium (a very reactive metal): Ethanol reacts with Na metal vigorously and produces sodium ethoxide (a Na—salt) and releases H₂(g).

In this reaction, Na replaces H from ethanol (substitution) to produce C₂H₅ONa.

Dehydration reaction (removal of water): At about 170°C temperature, ethanol is dehydrated by the cone. H₂SO₄(a dehydrating substance).

During this reaction, ethanol is converted into ethene.

Uses: Ethanol is used

1. As a solvent for gum,
2. As antifreeze liquid (due to very low freezing temperature),
3. In thermometers,
4. In car radiators,
5. Mostly for making wine/ beer/whisky.

Properties and uses of acetic acid or ethanoic acid (CH₃ – II – OH):

Physical properties: Colourless liquid Characteristic unpleasant pungent odor Soluble in water 5 – 8% solution of acetic acid is called vinegar.

Chemical properties: Less acidic in nature than mineral acids. It has ‘carboxylic acid’ as a functional group.

Reaction with sodium bicarbonate (a weak base):

Reaction with sodium hydroxide (strong base): Ethanoic acid reacts with NaOH to form salt and water (neutralization reaction).

Reaction with ethanol (esterification reaction): The reaction of an alcohol with carboxylic acid is called esterification which produces a new compound called ester.

In this reaction, OH from ethanoic acid and H from ethanol are removed to form water, and ethyl acetate an ester, having sweet fruity odour is formed.

This reaction establishes the identification of carboxylic group.

Uses: Used

1. In manufacturing vinegar
2. Esters
3. Many polymeric materials like cellulose acetate (used in photographic film)
4. As solvent.

Important uses of synthetic polymers:

WBBSE Chapter 8 Organic Chemistry Harmful Effects Of Methanol And Ethanol

Methanol (CH₃– OH) is highly poisonous/highjy toxic : It makes alcohol unfit for drinking. If methanol is taken with wine, it oxidises to form toxic formic acid which can cause blind ness, madness or even death.

Ethanol: (CH₃ – CH₂ – OH) is mainly used in making wine/beer/whisky. In low proportion, it affects cerebral cortex temporarily.

But for excessive intake of ethanol affects the nervous system and harmful for liver/kidney.

WBBSE Chapter 8 Organic Chemistry Denatured Spirit Or Methylated Spirit

In industry, ethanol is used as a solvent. 95% ethyl alcohol is called rectified spirit.

Rectified spirit is denatured by adding toxic substances like methyl alcohol (nearly poisonous), pyride (highly poisonous) and naptha.

This mixture is commercially known as denatured spirit. It is sold excise duty free for industrial use.

In market two types of spirits are available:

Industrial denatured spirit which contains 95% rectified spirit and 5% methyl alcohol,

Mineralised denatured spirit which contains 90% rectified spirit, 9% methyl alcohol and 1% pyride and naptha.;

Uses: Denatured spirit is used as a solvent for varnish.

WBBSE Chapter 8 Metallurgy Uses Of Fe, Cu, Zn, And Al And Their Alloys

Advantages of using alloys:

Alloys are harder than their constituent metals but less ductile and malleable. Example: Aluminium is lightweight, but its alloy duralumin (Al + Cu + Mg + Mn) is light but strong. Copper is soft, but it is alloy brass and bell metal are strong.

WBBSE Notes For Class 10 Physical Science And Environment

Alloys are resistant to corrosion: For example, Iron rusts, but stainless steel (Fe + C) does not. The melting point of an alloy may be lower than any of its original constituents.

Example: For soldering, an alloy of Pb and Sn is used whose m.p. is lower than that of Pb or Sn.  Using alloys the electrical conductivity can be changed.

Example: The alloy of Ni + Cr + Fe, called nichrome, resistivity is high.

In making 24-carat pure gold jewellery, 22 parts of gold are mixed with 2 parts of either Cu or Ag. So the jewellery we use is alloys.

Some important alloys:

WBBSE Chapter 8 Metallurgy Ores And Minerals

Most of the metals exist in the earth’s crust in combined states with some impurities such as sand/stones/rocks/limestone.

These naturally occurring substances in the earth’s crust which contain metals are called minerals. For example, iron exists in the earth’s crust as sulphides/carbonates/oxides/nitrates, with some impurities like C, Si, S, Mn etc.

Metallurgy refers to the process of extraction of pure metals from some minerals (ores). But from all minerals, metals cannot be extracted.

Minerals from which pure metals can be extracted cheaply and easily/conveniently are called ores. For example,

Al is most abundant in the earth’s surface /clay. But Al is not extracted from clay. Because the process is very costly. Rather, Al is extracted from bauxite.

So that bauxite is the ore of Al.

Haematite (Fe₂O₃) and iron pyrites (FeS₂) both contain a high percentage of Fe. However the removal of S from FeS₂ is very difficult and costly also. So FeS₂ is not considered as the ore of Fe.

Conclusion: All minerals are not ores, but all ores are minerals.

Major ores of Fe, Cu, Zn and Al :

WBBSE Chapter 8 Metallurgy Electronic Theory And Redox Processes

According to electronic theory, oxidation is loss of e~ and reduction is gain in e¯.

Examples of oxidation: \(\mathrm{Na}-e^{-} \rightarrow \mathrm{Na}^{+} ; \mathrm{Ca}-2 e^{-} \rightarrow \mathrm{Ca}^{2+}\)

Examples of reduction: \(\mathrm{Al}^{3+}+3 \mathrm{e}^{-} \rightarrow \mathrm{Al} ; \quad \mathrm{Fe}^{3+}+3 e^{-} \rightarrow \mathrm{Fe} ; \quad \mathrm{Zn}^{2+}+2 e^{-} \rightarrow \mathrm{Zn}\)

In metallurgy reduction/dressing/removal of oxygen of metal oxide is a very important step.

In order to free metal from its oxide, the reduction is done by excluding the non-metallic part Usually, the reduction is done in two ways:

Oxides of Cu, Pb, and Fe can be reduced by using chemical-reducing agents like C (coke), CO, H₂, and NH₃.

ZnO can only be reduced using coke (C) heated at high temperatures.

⇒ \(\mathrm{ZnO}+\mathrm{C} \rightarrow \mathrm{Zn}+\mathrm{CO}\) (carbon reduction of ZnO) Reduction

Note: During electrolytic extraction, reduction takes place at the cathode. General equation:

Mⁿ⁺+ ne¯→ M (where M = metal, n = number of electrons)

WBBSE Chapter 8 Metallurgy Thermite Reaction

Al has a high affinity to react with O₂ at high temperatures and the reaction is exothermic in nature. Using this principle, Fe is extracted from its oxide using the Goldschmidt thermite process.

Thermite mixture is basically a mixture of fine aluminium oxide (Al₂O₃) and ferric oxide (Fe₂ O₃) in a ratio of 1: 3 by mass.

The mixture also consists of a very minimal amount of igniting material like barium peroxide (BaO₂).

Taking the whole mixture in a crucible, ignite it with an Mg-ribbon. As soon as Mg- ribbon is burnt, a high temperature is created. As a result of which, both Al₂O₃ and Fe₂O₃ react with each other.

⇒ \(2 \mathrm{Al}(s)+\mathrm{Fe}_2 \mathrm{O}_3(s) \rightarrow 2 \mathrm{Fe}(s)+\mathrm{Al}_2 \mathrm{O}_3(s)+\text { Heat }\)

The reduction of Fe₂O₃ by Al is highly exothermic (2500 – 3000°C), which takes place within a maximum time of 30 seconds.

This produces molten Al₂O₃ and molten Fe. Molten Fe goes at the bottom since it is heavier than molten Al.

Application: Molten Fe is used for welding ferrous metals like the joining of rails, pipes, and broken parts of large gears.

Activity series of metals: It is a list of metals in ascending order of their activity. From top to bottom, the tendency of gaining e“s decreases i.e. reducing ing property gradually decreases.

Hence, the most reactive metal (K) exists at the top and the least reactive metal (Au) at the bottom. Any metal in this series can displace any other metal below it from its salt solution.

Example: Fe + ZnSO₄ does not take place, but \(\mathrm{Zn}+\mathrm{FeSO}_4 \rightarrow \mathrm{ZnSO}_4+\mathrm{Fe}\) Fe can take place as the position of Zn is above Fe in the series.]

Reactive metals like Au and Pt exist in a free state in nature. So, for these metals reduction is not needed.

Oxides/sulphides of Ag, and Hg are less stable. They can be reduced by heat. ng.

Oxides/sulphides of Zn and Fe can only be reduced by using Coke or any other suitable substances.

Oxides of most reactive metals (K, Na, Ca, Mg, Al) are stable and ionic compounds. They can not be reduced using coke or any other reducing substance.

Electrolysis is the only process (applying an electric field or potential difference) -in which reduction takes place at the cathode.

Activity: Carry out the following activities and note down your observations :

Add Cu-wire in ferrous sulphate solution (FeS0₄) → No reaction.

Add iron nails in copper sulphate solution (CuSO₄) → The Blue colour of CuSO₄ fades away and brown-coloured FeSO₄ deposits on Fe-nails. (Explain the reason)

WBBSE Chapter 8 Metallurgy Metal Corrosion

Usually, metals are lustrous/shining. But with time, metals start appearing dull, less shining. This is because of metal corrosion.

You might have noticed the difference in the appearance of a gold ring, silver spoon, copper coin, iron nail, and aluminium toy left in the air for a long time.

Silver spoon will appear less lustrous having a black coloured layer [Ag + H₂S (from the air) → Silver spoon Copper coin Ag₂S (black silver sulphide)]

On the copper coin, a green-coloured layer will form, and the iron nail will appear in the worst condition having a reddish-brown layer (rust),

The gold ring will remain in the best condition (as gold is the least reactive), and the aluminium toy will also be preserved very well.

Metals are considered to be corroded when they react with air/water. Corrosion happens very slowly.

Corrosion Definition: Corrosion is described as the formation of compounds on the surface of a metal when it is exposed to air/water or moisture/acid or electrolyte like salt water.

Rusting (Corrosion of iron): If impure iron is left in moist air, it gets corroded with a reddish brown coating on its surface.

This coating is called rust which is a hydrated ferric oxide (Fe₂O₃ . x H₂O where x represents an unknown quantity of water used in rusting and it is a variable) and the entire process is called rusting.

Rust is porous and can be removed. Iron corrodes very easily as compared to other metals like Cu/Ag/AI.

Because Fe is more reactive than Cu/Ag/AI. In this process of rusting, the upper surface of Fe is eaten up gradually.

Rusting weakens Fe objects and cuts short their lives. So, it’s an economic loss also. Pure Fe does not rust.

Mechanism of rusting of iron: For rusting to occur -Fe, water and O₂ are required. Here we will study the roles of water and O₂ in rusting.

The overall change “Fe→ Fe²⁺(unstable state) → Fe³⁺ (very stable state) → Fe₂O₃ • x H₂O¯ is a redox reaction i.e. an irreversible process once the rust is formed, it is not possible to get back the original Fe from rust.

When Fe-surface is exposed to water droplets, the Fe^– at- oms at the centre of water droplets give up 2 e¯sto form Fe²⁺ Fe (s) → Fe²⁺ (aq.) + 2e^– (oxidation takes place – here Fe is oxidised)

The e^–s move from the centre of the water droplet towards the edge all around. At the edge of the water droplet, there is an abundance of O₂.

O₂ is reduced to OH^– ions in the presence of water.

⇒ \(\mathrm{O}_2(g)+2 \mathrm{H}_2 \mathrm{O}(\mathrm{I})+4 e^{-} \rightarrow 4 \mathrm{OH}^{-} \text {(aq.) (reduction) }\)

We know that the electrode anode always donates e^– s and the cathode gains e^– s. 

So, on Fe^– surface plenty of small local chemical cells are formed, where each Fe atom at the centre of the water droplet acts as an anode,

The edge of the water droplet where C (impurity)-atoms are present acts as the cathode and the water molecule as an electrolyte.

Fe²⁺ ions react with OH^– ions to form solid ferrous hydroxide [Fe (OH)₂

⇒ \(\begin{aligned}
& 2\left[\mathrm{Fe} \rightarrow \mathrm{Fe}^{2+}+2 \mathrm{e}^{-}\right] \\
& \frac{\mathrm{O}_2+2 \mathrm{H}_2 \mathrm{O}+4 \mathrm{e}^{-} \rightarrow 4 \mathrm{OH}^{-}}{2 \mathrm{Fe}+2 \mathrm{H}_2 \mathrm{O}+\mathrm{O}_2 \rightarrow 2 \mathrm{Fe}^{2+}+4 \mathrm{OH}^{-}} \rightarrow 2\left[\mathrm{Fe}^{2+}+2 \mathrm{OH}^{-}\right] \rightarrow 2 \mathrm{Fe}(\mathrm{OH})_2 \text { (s) (unstable) } \\
&
\end{aligned}\)

Fe (OH)₂ is further oxidised by O₂ and forms ferric hydroxide [Fe (OH)₃].

⇒ \(4 \mathrm{Fe}(\mathrm{OH})_2(s)+2 \mathrm{H}_2 \mathrm{O}(\mathrm{I})+\mathrm{O}_2(g) \rightarrow 4 \mathrm{Fe}(\mathrm{OH})_3(s) \text { (a stable compound) }\)

Fe (OH)₃ decomposes into hydrated ferric oxide.

Fe (OH)₃ (s)→Fe₂O₃ • x H₂O (s)

Overall electro-chemical rusting reaction: 4Fe (s) + 2 x H₂O (l) + 3O₂ (g) 2Fe₂O₃ • x H₂O(s)

Note: Rusting is faster in the presence of chloride (Cl^–) ions. It is a serious problem in sea-going ships or submerged parts of pipelines.

Prevention of rusting:

Barrier protection: The simplest way of preventing rusting is by painting/oiling/greasing iron objects because of which iron will not get exposed to moisture.

Metallic coating: Rusting can be prevented by making a coating of another more reactive metal.

For example Galvanization

(Coating of Zn on Fe): Zn is more reactive than Fe. So, Zn reacts more rapidly than Fe forming a thin layer of ZnO on the surface of the Fe-object.

It does not allow more O₂ to react with the inner layer of Fe. Thus, galvanization is a better way to prevent Fe from rusting,

Tin plating, Chromium plating (plating of chromium metal Nickel plating This is done by electrolysis.

Alloying: Fe rusts very easily. But when Fe is mixed with Cr and Ni (an alloy), it becomes strong as well as does not rust. (Cr, Ni give a shiny look to Fe).

Cathodic protection: In an electro-chemical reaction, Fe-atoms (anode) at the surface give up e“s which flow towards the cathode because of this the anode gets corroded slowly.

In the cathodic protection method, the flow of e”s can be stopped by the use of (a) a sacrificial anode.

A more active metal like Mg-block is connected with Fe so that Mg acts as an anode instead of Fe. So, oxidation takes place in Mg.

[Mg (s)→ Mg²⁺ (aq.) + 2e ]. Instead of Fe- Mg gets corroded. This process is used for the protection of underground pipes, tanks etc.

When the Mg-block is corroded completely, it has to be replaced by a fresh one.

Sherardizing is the process of formation-corrosion-resistant Zn-layer by vapour galvanizing on the surface of iron or steel.

The Delhi Pillar of the Gupta Age is a unique metallurgical marvel of high-quality of steel production in ancient India. It was made from 98% corrosion-free wrought iron.

Corrosion of other metals and its health implications:

Al is a good conductor of heat and electricity. It is not affected in dry air. But in contact with moist air, Al reacts with the O₂ from air forming aluminium oxide (Al₂O₃) which is very much unreactive.

Once the Al₂O₃ layer is formed, it binds very tightly with the Al-surface. As a result of which, the thermal and electrical conductivity of Al decreases. Although, the Al₂O₃ layer keeps Al unaffected in moist air.

We have learnt that metals get corroded in moist air for a long time. On Cu or alloys of Cu, left in the open air for many years, green-coloured patches are formed. CO₂ and water vapour reacting with Cu form basic copper carbonate [Cu + O₂ + CO₂ + H₂O →Cu (OH)₂ • CuCO₃].

Al, Zn react with weak organic acidic substances like vinegar or lemon juice, and they react to form soluble metallic compounds (harmful to our health).

Cu also forms soluble metallic compounds by reacting with acidic pickles or fruits. For this reason, acidic foods should not be kept or processed in Al, Zn or Cu containers.

Also, tarnished metallic utensils should be cleaned well before use to avoid metal poisoning.

WBBSE Chapter 8 Inorganic Chemistry In The Laboratory And In Industry Laboratory Preparation Of Ammonia(NH₃)

Principle: Ammonium salt + Alkali → Salt + Ammonia Gas+ H₂O.

When ammonium salts are heated with less volatile strong alkalis, then comparatively more volatile ammonia gas is produced. For example:

⇒ \(\left(\mathrm{NH}_4\right)_2 \mathrm{SO}_4(s)+\mathrm{CaO}(s) \stackrel{\Delta}{\longrightarrow} \mathrm{CaSO}_4(s)+2 \mathrm{NH}_3(g)+\mathrm{H}_2 \mathrm{O}(I)\)

⇒ \(\mathrm{NH}_4 \mathrm{Cl}(s)+\mathrm{NaOH}(\text { aq. }) \stackrel{\Delta}{\longrightarrow} \mathrm{NaCl}(\mathrm{s})+\mathrm{NH}_3(\mathrm{~g})+\mathrm{H}_2 \mathrm{O}(I)\)

⇒ \(\left(\mathrm{NH}_4\right)_2 \mathrm{SO}_4(s)+2 \mathrm{KOH}(\text { aq. }) \stackrel{\Delta}{\longrightarrow} \mathrm{K}_2 \mathrm{SO}_4(s)+2 \mathrm{NH}_3(g)+2 \mathrm{H}_2 \mathrm{O}(I)\)

Caustic alkalis like NaOH/KOH are not used because they are deliquescent i.e. from an atmosphere they absorb moisture and dissolve.

In the laboratory, slaked lime [Ca (OH)₂] is used as an alkali.

Chemicals required: Dry ammonium chloride (NH₄CI) and dry Ca (OH)₂in 1: 3 mass ratio (Both are finely ground).

Condition:

1. Direct contact of the reactants and
2. Application of heat.

Chemical equation: \(2 \mathrm{NH}_4 \mathrm{Cl}(s)+\mathrm{Ca}\begin{aligned}
& (\mathrm{OH})_2(s)-\stackrel{\Delta}{\longrightarrow} \mathrm{CaCl}_2(s)+2 \mathrm{NH}_3(g)+ \\& 2 \mathrm{H}_2 \mathrm{O}(g)\end{aligned}\)

Laboratory arrangement: The round bottom flask is kept in a tilting position.

Drying agent: Drying is essential because water vapor is one of the products.

Moist NH₃(g) is dried by passing it through quick lime (CaO). Ammonia gas is basic in nature. Quicklime is also a base. So they never react to one another.

WBBSE Notes For Class 10 Physical Science And Environment

Gas collection: Dry NH₃(g) is collected by downward displacement of air in the inverted gas jar. Because-

1. NH₃(g) is lighter than air and
2. NH₃ (g) is highly soluble in water. A moist red litmus paper held near the mouth of the gas jar turns blue in the presence of NH₃ (g).

Precautions:

The reagents are finely ground in dry condition, otherwise, NH₄CI may sublime on heating,

The round bottom flask is kept tilted so that water may not trickle back and crack the hot flask,

In the laboratory process, we should not use ammonium nitrate (NH₄NO₂).

Because NH₄NO₂ is explosive in nature, it decomposes on heating forming nitrous oxide (N₂O) gas and water vapor.

⇒ \(\mathrm{NH}_4 \mathrm{NO}_3(s) \stackrel{\Delta}{\longrightarrow} \mathrm{N}_2 \mathrm{O}(g)+2 \mathrm{H}_2 \mathrm{O}(g)\)

In drying NH₃(g), common drying agents are like a cones. H₂SO₄, P₂ O_5, and anhydrous CaCl₂ are not used.

Because NH₃ reacts  with these substances: 2NH₃ + H₂SO₄ →(NH₄)₂SO₄; 6NH₃ + P₂O₅ + 3H₂O→ 2 (NH₄)₃ PO₄ (P₂O₅ is acidic in nature); x NH₃ + (anhydrous) CaCI₂ → CaCI₂ . × NH₃ (addition compound)

WBBSE Chapter 8 Inorganic Chemistry In The Laboratory And In Industry Properties Of Ammonia

Physical properties: Ammonia is a colorless gas having a characteristic pungent choking smell.

The pungent smell in the washroom is due to ammonia.

Density: Under standard conditions, the specific density of ammonia is 8-5 whereas the specific density of air is 14-4. This means that NH₃ (g) is lighter than air.

Solubility: At ordinary temperature, NH₃ (g) is highly soluble in water (in 1 volume of water about O₂ volume of NH₃ (g) is soluble).

Aqueous ammonia: Ammonia, being very much soluble, is dissolved in water and forms ammonium hydroxide (NH₄OH) which is a weak base due to the presence of OH¯ ions.

⇒ \(\mathrm{NH}_3+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{NH}_4 \mathrm{OH} ; \mathrm{NH}_4 \mathrm{OH} \rightleftharpoons \mathrm{NH}_4^{+}+\mathrm{OH}^{-} \text {(Arhenious theory) }\)

This is called aqueous ammonia (or in the aqueous form of ammonia) called ammonium hydroxide. Under high pressure, the volume of NH₃ (g) is compressed, and when the temperature is decreased, the gas gets liquified.

By cooling NH₃(g) to a temperature – 33-4°C under normal pressure, it changes into a colorless liquid called anhydrous or liquid ammonia /moisture-free ammonia.

Liquor ammonia: It is the 35% concentrated/saturated solution of ammonia in water. Its specific gravity is 0.88. That’s why liquor ammonia is treated as a very strong solution of NH₃ (g) in water.

Remember: Liquor ammonia is not liquid ammonia. Liquor ammonia contains NH₄OH (NH₄⁺and OH¯ions) whereas liquid ammonia contains NH₃

Chemical properties:

Reactions due to basic nature: Ammonia is a weak base, so reacting with acids (HCl, H₂S0₄) it forms ammonium salts.

For example:

⇒ \(\mathrm{NH}_3 \text { (aq.) }+\mathrm{HCl} \text { (aq:) } \rightarrow \mathrm{NH}_4 \mathrm{Cl} \text { (aq.) (ammonium chloride salt) }\)

The same reaction happens in the gaseous state also.

⇒ \(\mathrm{NH}_3(g)+\mathrm{HCl}(g) \rightarrow \mathrm{NH}_4 \mathrm{Cl}(s) \text { (dense white fumes) }\)

When a glass rod is dipped in a cone. HCl is brought in contact with NH₃ (g), and it produces dense white fumes (fine dust of solid NH₄CI).

This is a reaction between two gases that produces a solid compound. This reaction is also used for the detection of ammonia gas.

⇒ \(2 \mathrm{NH}_3 \text { (aq.) }+\mathrm{H}_2 \mathrm{SO}_4 \text { (aq.) } \rightarrow\left(\mathrm{NH}_4\right)_2 \mathrm{SO}_4 \text { (ammonium sulphate salt) }\)

NH₄OH (aq.) turns red litmus blue and colorless phenolphthalein solution pink.

Reducing property: NH₃ is a good reducing agent. This means NH₃ reduces i.e. removes O₂ from another substance. Example:

When NH₃ (g) is passed over red hot copper (II) oxide (black), then NH₃ reduces CuO to form metal copper [red).

At the same NH₃ gets oxidized in N₂ (g). This reaction proves that NH₃ contains nitrogen.

Catalytic oxidation of NH₃:

⇒ \(4 \mathrm{NH}_3(g)+5 \mathrm{O}_2(g) \frac{\mathrm{Pt}}{800^{\circ} \mathrm{C}} 4 \mathrm{NO}(g)\)+ 6 \(\mathrm{H}_2 \mathrm{O}(g)+\text { Heat (Exothermic reaction). }\)

NH₃ (g) reacts with O₂ in the presence of Pt- catalyst heated at 800°C and colorless nitric oxide gas (NO) is produced (oxidation reaction).

⇒ \(4 \mathrm{NH}_3+3 \mathrm{O}_2 \rightarrow 2 \mathrm{~N}_2+6 \mathrm{H}_2 \mathrm{O}(\mathrm{g})\)

The oxidizing tendency is so high that NO (g) further oxidizes into reddish-brown nitrogen dioxide gas (NO₂).

⇒ \(\mathrm{NO}(g)+\mathrm{O}_2(g) \rightarrow \mathrm{NO}_2(g) \text { (reddish brown gas) }\)

4. Precipitation of metal hydroxides from aq. solution of metal ions (Fe³⁺, Al³⁺) :

Analytical chemistry: Usually, all cations like Na⁺, K⁺, Ca²⁺, Zn²⁺, Mg²⁺, Al³⁺, NH₄⁺ are colourless and all anions like Cl^–, Br^–, CO₃^2-, NO₃^1-, SO₄^z-

Are colorless except for some cations: Fe²⁺ (-ous) → light green, Fe³⁺ (-ic) → yellowish/brown, Cu²⁺ (-ous) → blue and some anions MNO₄^1- (permanganate) →pink, Cr₂O₇^2- (dichromate) → orange.

That’s why Zn (NO₃)₂ → colourless, FeSO₄ → light green, FeCI₃ →yellow, AICI₃ → colourless, CuSO₄ → blue, …. etc.

Here we will study what happens due to the addition of NH₄OH/ NaOH solution into the aq. solution of Fe³⁺ (ferric), and Al³⁺ salts—especially to note down the color of the salt and its solution.

⇒ \(\begin{array}{ll}
\mathrm{FeCl}_3 \text { (aq.) }+3 \mathrm{NH}_4 \mathrm{OH}(\text { aq.) } \longrightarrow & \mathrm{Fe}\left(\mathrm{OH}_3\right)(s) \downarrow+3 \mathrm{NH}_4 \mathrm{Cl} \text { (aq.) } \\
\text { yellow } & \text { Brown ppt. } \\
\mathrm{AlCl}_3 \text { (aq.) }+3 \mathrm{NH}_4 \mathrm{OH} \longrightarrow & \mathrm{Al}(\mathrm{OH})_3(s) \downarrow+3 \mathrm{NH}_4 \mathrm{Cl} \text { (aq.) } \\
\text { Colourless } & \text { Gelatinous } \\
\text { white ppt. }
\end{array}\)

Significance of such reactions: Observing the color of metal hydroxide precipitate, we can easily detect /identify the presence of metallic cation (Fe³⁺, Al³⁺).

Physical observation: In aq. CuSO₄, aq. NH₃ is added :

⇒ \(\mathrm{CuSO}_4 \text { (aq.) }+2 \mathrm{NH}_4 \mathrm{OH} \text { (aq.) } \longrightarrow \mathrm{Cu}(\mathrm{OH})_2(\mathrm{~s}) \downarrow+\left(\mathrm{NH}_4\right)_2 \mathrm{SO}_4 \text { (aq.) }\)

But this blue ppt. of Cu(OH)₂ dissolves in excess NH₄OH forming a deep blue soluble com plex salt-tetramine copper (II) sulfate.

The alkaline solution of potassium mercuric iodide (K₂Hgl₄) is called Nessler’s reagent. It is used to identify/test the presence of ammonia.

Ammonia in contact with Nessler’s reagent produces yellowish/brown ppt. A bottle of liquor ammonia is cooled before opening.

We know that a large quantity of NH₃ vapor is kept under high pressure inside the bottle. If the bottle is extremely suddenly opened due to the release of pressure, NH₃ vapor may spurt danger to the eyes, out and come in contact with the eyes.

This is very much harmful to the eyes.

Use of liquid ammonia as a coolant: It is used as a refrigerant in ice plants.

Liquid NH₃ is capable of absorbing heat as its specific heat capacity is very high and cools another substance.

In the refrigerator, tetra fluoro ethane is used. Why not liquid NH₃? Because- NH₃ reacts with Cu (refrigerator pipe metal) a high quantity of NH₃ is poisonous.

Ammonia gas is non-poisonous. If inhaled, affects the respiratory system and brings tears to the eyes.

NH₃ concentration (300 ppm) is immediately dangerous to life. Workers should use personal protective equipment:

In case of accidental leakage of NH₃ from cold storage, the first and foremost duty is to wash out the affected body parts with water as NH₃(g) is highly soluble in water.

In the factory, after NH₃ (g) pipeline leakage inhalation of NH₃ can cause severe irritation of the nose and throat, coughing, and shortness of breath/difficulty in breathing resulting in respiratory failure. People fell unconscious and they need to be immediately hospitalized.

WBBSE Chapter 8 Inorganic Chemistry In The Laboratory And In Industry Major Industrial Uses And Industrial Manufacture Of NH₃ And Urea

Industrial uses of ammonia: Ammonia is used in the manufacture of

Nitrogenous fertilizers: examples – ammonium sulfate [(NH₄)₂SO₄], ammonium nitrate (NH₄NO₃), ammonium phosphate [(NH₄)₃P0₄], urea [CO (NH₂)₂], etc.

Nitric acid by Ostwald’s process, sodium carbonate (Na₂CO₃) by Solvay’s process

Explosives: examples-TNT (tri nitro toluene), RDX, etc.

Polymers: Examples – nylon, rayon, plastics, dyes, etc.

Ammonium chloride is used in dry cells, ammonium carbonate is used as a smelling salt for reviving a fainted person

As a refrigerant in ice-plants

Many pharmaceutical products, and household cleaning products – removing grease/perspiration strains, cleaning tiles/windows, etc.

Industrial manufacture of NH₃ (Haber’s process): (For industrial production on a large scale) From the synthesis of dry and pure N₂ (g) and H₂ (g) in the ratio 1 : 3 by volume under some favorable conditions NH₃ (g) is produced.

Favorable conditions:

High temperature: 450 – 500°C
High pressure: 200 – 900 atm
Catalyst: Finely divided Fe
Promoter: Molybdenum (Mo) or Al₂O₃

Relevant Reaction:

In this reaction, both the reactants and product are gaseous; the reaction is reversible and exothermic; volume decreases in the forward direction.

Collection Of NH₃: The gaseous mixture contains produced NH₃ along with unrelated N₂ and H₂ Then by condensation through a cooling chamber, NH₃ (g) is liquefied easily as compared
to N₂ and H₂ molecules and dissolved in water (because NH₃ is highly soluble in water).

By recirculating unreacted N₂ and H₂, an eventual yield of NH₃ (~ 98%) can be obtained.

Uses of urea [CO (NH₂)₂]: Urea is used both as a nitrogenous fertilizer (about 90%) and animal feed.

In the manufacture of urea-formaldehyde resin/plastic used for lamination/fabrication purposes.

In the production of urea stibine, a medicine for leishmaniasis (kala-azar), and barbiturates, used as sleeping pills.

Manufacture of urea: Raw materials are:  CaCO₃ (indirectly used). \(\mathrm{CaCO}_3 \stackrel{\Delta}{\longrightarrow} \mathrm{CaO}+\mathrm{CO}_2\)

This CO₃ is used in urea production, NH₃ (directly used through Haber’s process).

Condition: Liquid NH₃ and liquid CO₂ react at 200°C temperature Under 200 atm pressure.

Chemical reaction:

WBBSE Chapter 8 Inorganic Chemistry In The Laboratory And In Industry Laboratory Preparation Of Hydrogen Sulphide(H₂S)

Principle: A strong acid can substitute the weak acid part from its salt. Example- (where H from strong acid HCI substitutes Na of Na₂S)

Chemicals required : A few pieces of solid ferrous sulfide (FeS) [dark brown in color] and

H₂SO₄(or dil. HCI).

Condition: Direct contact of FeS and dil. H₂SO₄ at normal temperature.

Chemical equation: FeS (s)+ H₂SO₄(aq.) →FeSO₄+ H₂S (g)

(where H₂SO₄→strong acid, H₂S→ weak acid, FeS→ salt of weak acid)

Laboratory arrangement: H₂S (g) is collected by the upward displacement of air, as it is heavier than air.

How Laboratory make H₂S(g) free from water vapor? this is done by passing H₂S (g) through acidic Phosphorous pentoxide (P₂O_s) (because H₂S is acidic and reducing in nature).

Why drying agents like

conc. H₂SO₄ Quicklime(CaO) or  Anhydrous CaCl₄are not used? Conc. H₂SO₄ is a strong oxidises H₂S to sulfur(s)

⇒ \(\mathrm{H}_2 \mathrm{SO}_4 \text { (aq.) }+\mathrm{H}_2 \mathrm{~S}(g) \longrightarrow 2 \mathrm{H}_2 \mathrm{O}(I)+\mathrm{SO}_2(g)+\mathrm{S}(s)\)

CaO is a base. \(\mathrm{CaO}(s)+\mathrm{H}_2 \mathrm{~S}(g) \longrightarrow \mathrm{CaS}(s)+\mathrm{H}_2 \mathrm{O}\)

(anhydrous s) \(\mathrm{CaCl}_2+\mathrm{H}_2 \mathrm{~S}(g) \rightarrow \mathrm{CaS}(s)+2 \mathrm{HCl}(g)\)

In the laboratory preparation of H₂S(g) from FeS (s)why conc. H₂SO₄ conc. HNO₃ conc. H₂SO₄ is not taken?

Reaction: FeS + H₂SO₄ → FeSO₄ + HS

A, B, and C are respectively bottom, middle, and top chambers or globes of Kipp’s apparatus. Chamber C connects chamber A passing through middle chamber B. A tap is fitted to the middle chamber B.

Take solid FeS in B. Add dil. H₂SO₄ through C, until acid touches FeS in B. At once, the reaction starts and H₂S (g) is produced. By opening the tap, H₂S (g) can be collected.

When the tap remains closed (air-tight), the high pressure of gas in B pushes H₂S0₄ down to bottom chamber A and up into the top chamber C.

This separates FeS and H₂SO₄. Immediately, further production of gas is stopped.

As per need, when a tap is opened, pressure in the middle chamber decreases, which allows H₂SO₄ to enter into B from A making contact with FeS. Again H₂S (g) starts producing.

WBBSE Chapter 8 Inorganic Chemistry In The Laboratory And In Industry Properties Of Hydrogen Sulphide

Physical properties:  Foul odor of rotten egg, colorless, poisonous gas.

G Density: Heavier than air. H₂S has a vapor density of 1-2, it is large as compared to air which is 1 Solubility Moderately soluble in cold water. But insoluble in hot water.

Chemical properties:

Acidic property: Aq. solution of H₂S (g) behaves like a mild acid. It is a di-basic acid.

⇒ \(\mathrm{H}_2 \mathrm{~S}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{H}_3 \mathrm{O}^{+}+\mathrm{HS}^{-} \text {i.e. } \mathrm{H}_2 \mathrm{~S} \text { (aq.) } \rightleftharpoons \mathrm{H}^{+}+\mathrm{HS}^{-}\)(partial ionization)

⇒ \(\mathrm{HS}^{-}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{S}^{2-}+\mathrm{H}_3 \mathrm{O}^{+} \text {i.e. } \mathrm{H}_2 \mathrm{~S} \text { (aq.) } \rightleftharpoons 2 \mathrm{H}^{+}+\mathrm{S}^{2-}\)(complete ionization)

Reaction with alkali [ex. NaOH (aq.)]: Since aq. H₂S behaves as a di-basic acid, it reacts with NaOH solution to form both acid salt and normal salt.

⇒ \(\begin{array}{ll}
\mathrm{NaOH}+\mathrm{H}_2 \mathrm{~S} \longrightarrow & \mathrm{NaHS}+\mathrm{H}_2 \mathrm{O} \text { (partial ionization) } \\
& \text { acid salt (sodium hydrogen sulfide) } \\
\mathrm{NaHS}+\mathrm{NaOH} \longrightarrow & \mathrm{Na}_2 \mathrm{~S}+\mathrm{H}_2 \mathrm{O} \text { (complete ionization) } \\
\text { normal salt (sodium sulfide) }
\end{array}\)

Combustibility of H₂S (g): H₂S (g) burns in air (O₂) with blue flame but does not support in burning.

In the excessive supply of O₂, sulfur dioxide gas (another toxic gas) is produced, and in low O₂-level, sulfur deposits.

(excessive supply of O₂) 2H₂S + 3O₂ 2H₂O + 2SO₂ (low supply of O₂) 2H₂S + O₂→2H₂O+ 2S

As a strong reducing agent (reaction with acidified potassium dichromate solution): H₂S (g) is passed through acidified K₂Cr₂O₇ solution (orange).

Result: Orange Coloured Solution Solution is Converted Into a Green Colour

This Colour Change occurs only because of the reducing property of H₂S.

(4) Precipitation of mental sulfides (Cus, Pbs, Ag₂S- black) From Aqueous solution of copper sulfate CuSO₄ (blue), lead nitrate pb (NO₃)₂ (Colourless): H₂S (g) is passed through the Queous solutions separately in each case, Black ppt. of mental sulfide is obtained.

⇒ \(\begin{array}{ll}
\mathrm{CuSO}_4 \text { (aq.) }+\mathrm{H}_2 \mathrm{~S}(g) \longrightarrow & \mathrm{CuS}(s) \downarrow+\mathrm{H}_2 \mathrm{SO}_4 \text { (aq.) } \\
\text { (Blue) } & \text { (Black) } \\
\mathrm{Pb}\left(\mathrm{NO}_3\right)_2 \text { (aq.) }+\mathrm{H}_2 \mathrm{~S}(g) \longrightarrow & \mathrm{PbS}(s) \downarrow+2 \mathrm{HNO}_3 \text { (aq.) } \\
\text { (Colourless) } & \text { (Black) } \\
2 \mathrm{AgNO}_3 \text { (aq.) }+\mathrm{H}_2 \mathrm{~S}(g) \longrightarrow & \mathrm{Ag}_2 \mathrm{~S}(s) \downarrow+2 \mathrm{HNO}_3 \text { (aq.) } \\
\text { (White) } & \text { (Black) }
\end{array}\)

 Identification test for H₂S: Dip a filter paper into lead acetate solution, which in contact with H₂S (g) immediately turns black due to the formation of insoluble black lead (II) sulfide.

This method is important to determine the presence and amount of H₂S (g).

Reaction with alkaline sodium nitroprusside solution: When H₂S {g) is passed through freshly prepared alkaline sodium nitroprusside solution (colorless), the solution becomes violet.

Toxicity of H₂S: It is a highly toxic gas.

Rapidly affects the central nervous system and respiratory system. Inhaling high concentrations (over 500 -1000 ppm) can cause instant death.

There is no proven antidote for H₂S poisoning. Wear a nose mask and avoid inhaling H₂S.

WBBSE Chapter 8 Inorganic Chemistry In The Laboratory And In Industry Laboratory Preparation And Major Uses Of Nitrogen (N₂)

Laboratory preparation: Chemical required: A concentrated aqueous solution of ammonium chloride (NH₄CI) and sodium nitrite (NaNO₂) in 1:1 molar proportion.

Condition:  In the laboratory N₂(g) is prepared by (gently) heating the mixture of NH₄CI and NaNO₂.

Chemical reaction: The reactants initially undergo double decomposition to form ammonium nitrite (NH₄NO₂) and sodium chloride (NaCI). NH₄NO₂ (aq.) + NaCI (aq.)

⇒ \(\mathrm{NH}_4 \mathrm{Cl} \text { (aq.) }+\mathrm{NaNO}_2 \text { (aq.) } \stackrel{\Delta}{\longrightarrow}\)

⇒ \(\mathrm{NH}_4 \mathrm{NO}_2 \text { (aq.) }+\mathrm{NaCl} \text { (aq.) }\)

Thereafter the NH₄NO₂ formed here decomposes to produce N₂ (g).

⇒ \(\mathrm{NH}_4 \mathrm{NO}_2 \text { (aq.) } \stackrel{\Delta}{\longrightarrow} \mathrm{N}_2(g)+2 \mathrm{H}_2 \mathrm{O}(g)\)

Gas collection: N₂(g) is collected by the downward displacement of water.

Drying of laboratory-prepared N₂ (g): By passing through

Cone. H₂S0₄, existing moisture can be removed

Cone. NaOH/KOH, Cl₂ (g) can be removed

Finally passing over red hot Cu, oxides of nitrogen are removed (reduced) to form fresh N₂ (g).

Main uses: N₂(g) is used in the industrial manufacture of NH₃ (by Haber’s process.

NH₃ is an important starting material in the production of HNO₃ (by Ostwald’s process), nitrogenous fertilizers such as (NH₄)₂SO₄, NH₄NO₂, (NH₄)₃PO_4, etc, explosives, and pharmaceutical products.

Being an inert gas creates an unreactive atmosphere inside an electric bulb, and thermometer to prevent oxidation.

In many industrial pros, N₉ boils at —196 C. So, liquid N₂ is used cesses it is used to create an unreactive atmo- ² ………. ² to cool down the exothermic reaction, sphere, as N₂ is cheaper than He/Ar.

Packets of chips/pop-corn/ other food packets are made puffy using N₂ (g) so that oxidation and growth of bacteria are stopped.

The food products also remain fresh. That’s why N₂ is very important in food packaging.

Liquid N₂ has many real-life applications such as the preservation of biological specimens example: blood, cornea, eye, tissue samples, etc.

WBBSE Chapter 8 Inorganic Chemistry In The Laboratory And In Industry Properties Of N₂

Physical properties: N₂a colorless, odorless, and tasteless gas.

Density: Slightly less dense than air (vapor density of N₂ and air are respectively 14 and 14-4).

Solubility: N₂ (g) is sparingly soluble in water (1L of water dissolves 22 mL of N₂ at 0°C). 0 N₂ (g) freezes at – 210-1°C.

At high pressure, N₂ (g) can be liquified. N₂ boils at – 196°C.

Chemical properties: The molecular structure of nitrogen is However each N₂ molecule is very difficult to break and that’s why N₂ molecules have a very strong triple bond shared between two N atoms which behave like an inert/unreactive gas at normal temperature.

But at extremely high temperatures N₂ can react with other metals/non-metals. For example—

Reaction with hydrogen: Mixture of pure and dry N₂ (g) and H, (g) in the volume ratio 1 : 3 when passed over hot (450 – 500°C) finely divided Fe-catalyst and Mo-promoter under high pressure (> 200 atm) then NH₃ (g) is produced through an exothermic reaction.

This is the principle of industrial production of NH₃ by Haber’s process.

⇒ \(\mathrm{N}_2+3 \mathrm{H}_2 \underset{\substack{450-500^{\circ} \mathrm{C} \\ \text { above } 200 \mathrm{~atm}}}{\stackrel{\mathrm{Fe} \text { and } \mathrm{Mo}}{\rightleftharpoons}} 2 \mathrm{NH}_3+\text { Heat }\)

Reaction with magnesium: N₂ (g) puts off a burning candle as it is neither combustible nor a supporter of combustion (inert gas).

But a burning element can react with N₂ exactly what happens in the case of a burning magnesium ribbon placed in an Infilled gas jar.

Mg-ribbon burns and forms a yellowish powder called magnesium nitride (Mg₃N₂).

Chemical reaction:  \(3 \mathrm{Mg}(s)+\mathrm{N}_2(g) \stackrel{\text { Heat }}{\longrightarrow} \mathrm{Mg}_3 \mathrm{~N}_2(s)\)

Reaction with calcium carbide: Formation of nitrolim: When N₂(g) is passed over hot (~ 1100°C) calcium carbide dust (CaC₂), a brownish grey solid mixture is produced which contains calcium cyanamide (CaNCN) and carbon (C)…

noo°c nitro slim Commercially, this solid mixture (CaNCN + C) is called nitrolim- an inorganic compound used as a chemical fertilizer.

Nitrogen fixation: Nitrogen is present in abundance (~ 78%) in the atmosphere. Plants and animals cannot take nitrogen directly from the atmosphere.

So, anyhow N₂ needs to be converted into its usable forms (usually nitrogenous compounds). This process of conversion of atmospheric N₂ into useable forms is referred to as “nitrogen fixation.”

N₂ combines with O₂ in the presence of lightning at a temperature of 3000 – 5000°C to form nitric oxide (NO), and NO further oxides into NO₂ (nitrogen dioxide).

⇒ \(\mathrm{N}_2(g)+\mathrm{O}_2(g) \stackrel{\text { lightning }}{\longrightarrow} 2 \mathrm{NO}(g) ; 2 \mathrm{NO}(g)+\mathrm{O}_2(g) \stackrel{\text { lightning }}{\longrightarrow} 2 \mathrm{NO}_2(g)\)

During rainfall, NO₂ (g) reacts with water and produces both aq. nitric acid (HN0₃) and unstable aq. nitrous acid (HNO₂) and fall down to the soil.

⇒ \(2 \mathrm{NO}_2(g)+\mathrm{H}_2 \mathrm{O}(l) \rightarrow \mathrm{HNO}_3 \text { (aq.) }+\mathrm{HNO}_3 \text { (aq.) }\)

Alkaline substances present in soil react with aq. HNO₃ forms nitrate (NO₃) and nitrite (NO₂) salts (soluble in water).

This is how atmospheric nitrogen is fixed/ converted into a usable form. Plants use these nitrates to make proteins.

Animals also eat plants to get proteins. Animal excreta and dead plants/animals again get converted to nitrates.

Denitrifying bacteria convert these nitrates back to atmospheric nitrogen. As a whole, % of N₂ in the air remains constant.

WBBSE Chapter 8 Inorganic Chemistry In The Laboratory And In Industry Industrial Manufacture Of HCl, HNO_3, And H₂SO₄

Manufacture of HCI (by Synthetic method): In the Castner-Kellner (commercial) process, caustic soda (NaOH) is manufactured by the electrolysis of aqueous sodium chloride (NaCI) solution. The by-products obtained in this process are H₂ (g) and Cl₂ (g).

Industrially hydrogen chloride gas is manufactured by burning /direct combustion of equal volumes of H₂ (g) and Cl₂ (g) in a combustion chamber made of silica (SiO₂).

⇒ \(\mathrm{H}_2(g)+\mathrm{Cl}_2(g) \rightarrow 2 \mathrm{HCl}(g)\)

The HCI (g) is passed through a cooling chamber. Then the cooled gas is allowed to absorb water. Finally, a saturated solution of hydrochloric acid is prepared.

⇒ \(\mathrm{HCl}(g) \stackrel{\mathrm{H}_2 \mathrm{O}}{\longrightarrow} \mathrm{HCl} \text { (aq.) }\left[\mathrm{H}^{+} \text {(aq.) }+\mathrm{Cl}^{-} \text {(aq.) }\right]\)

2. Manufacture of HNO₃ (by Ostwald’s process):

The reaction\(\mathrm{NH}_3(g) \rightarrow \mathrm{HNO}_3 \text { (aq.) }\) takes place in 3 steps:

Step 1: Oxidation of NH₃(g): A mixture of NH₃ (g) and dry pure air in a 1:10 volume ratio is passed very fast over catalyst platinum-rhodium (90:10) gauze heated at about 800°C.

The time of contact is ~ 0-0014 seconds and NH₃ (g) gets oxidized by O₂ of air to produce nitric oxide vapor (NO).

The reaction is exothermic and reversible. Once the reaction starts, the heat evolves and continuous heating is not required.

⇒ \(4 \mathrm{NH}_3+5 \mathrm{O}_2 \stackrel{\mathrm{Pt}-\mathrm{Rh}}{\underset{800^{\circ} \mathrm{C}}{\rightleftharpoons}} 4 \mathrm{NO}(g)+6 \mathrm{H}_2 \mathrm{O}(g)+\text { Heat }\)

Step 2: Oxidation of NO vapor: Hot gaseous mixture (NO vapor + water vapor + excess O₂) is cooled to 50°C and NO vapor is further oxidized by O₂ sent in the oxidizing chamber to produce nitrogen dioxide gas (NO₂).

⇒ \(2 \mathrm{NO}(g)+\mathrm{O}_2(g) \longrightarrow 2 \mathrm{NO}_2(g)\)

Step 3: Absorption of NO₂(g): The absorption of NO₂ (g) by water sprayed from the top of the absorption tower produces nitric acid (HNO₃).

Initially, very dilute HNO₃ is produced which is recycled to absorb, more and more NO (g) till 68% HNO₃ is obtained.

Further concentration is done by distillation with a cone. H₂SO₄ at temperature ~120°C. This HN0₃ is almost 98% concentrated.

⇒ \(3 \mathrm{NO}_2(g)+\mathrm{H}_2 \mathrm{O}(I) \longrightarrow 2 \mathrm{HNO}_3 \text { (aq.) }+\mathrm{NO}(g)\)

⇒ \(\mathrm{HNO}_3(68 \%) \stackrel{+ \text { conc. } \mathrm{H}_2 \mathrm{SO}_4}{\underset{\text { Distillation }}{\longrightarrow}} \mathrm{HNO}_3(98 \%)\)

(3) Manufacture of H₂SO₄ (by Contact Process):

Step 1: Production of sulfur dioxide (SO₂):

⇒ \(\mathrm{S}(\mathrm{s})+\mathrm{O}_2(g) \rightarrow \mathrm{SO}_2(g)\)

Or, \(4 \mathrm{FeS}_2(s)+11 \mathrm{O}_2(g) \longrightarrow 2 \mathrm{Fe}_2 \mathrm{O}_3(s)+8 \mathrm{SO}_2(g)\)

By heating sulfur or iron pyrites in the presence of air (O₂) SO₂ (g) is prepared [heating in the presence of excess air is called Roasting].

Step 2:  Oxidation of sulfur dioxide(SO₂) to sulphur trioxidfe (SO₃):  (most important step:)

In the presence of a solid vanadium pentoxide catalyst (V₂O₅) at the optimum temperature of 450^°C and high pressure -2 atm, pure and dust-free SO₂ (g) is oxidized by dry pure O₂ (g) to form SO₃ (g).

Step 3: Absorption of SO₃ (g): [SO₃ is not diluted directly by adding water into \(\mathrm{SO}_3 \mathrm{O}\mathrm{SO}_2+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{H}_2 \mathrm{SO}_4\)  (not done).

Because too much heat is evolved for SO₃ and H₂SO₄ vapors form an acid mist.]

After cooling, SO₃ (g) is dissolved in 98% concentrated H₂SO₄ sprayed from the top of the absorption chamber.

Fuming sulphuric acid or pyro-sulphuric acid (H₂S₂O₇) is formed. Its commercial name is Oleum,

⇒ \(\left.\mathrm{SO}_3+\mathrm{H}_2 \mathrm{SO}_4 \text { ( } 98 \% \text { conc. }\right) \rightarrow \mathrm{H}_2 \mathrm{~S}_2 \mathrm{O}_7(l)\)

Then, after dilution with water, sulphuric acid of any desired concentration can be produced. \(\mathrm{H}_2 \mathrm{~S}_2 \mathrm{O}_7(I)+\mathrm{H}_2 \mathrm{O}(I) \rightarrow 2 \mathrm{H}_2 \mathrm{SO}_4 \text { (aq.) }\)

Note:

Oleum is a more powerful oxidizing agent than a con. H₂SO₄.

Solid crystals used in industrial manufacturing processes (Ostwald’s process/Contact Process) are usually taken in dust form not in bigger granules in order to get a greater surface area.

For it, the activity of catalysts/ rate of reaction increases because of greater adsorption.

WBBSE Chapter 8 Ionic And Covalent Bonding What Is Chemical Bond? Why Chemical Bonds Are Formed?

By natural process, atoms of the same element or different elements have a tendency to combine with each other to form different molecules.

So, there exists a force of attraction between the atoms that holds them together in the molecule in order to get stability.

A chemical bond means a chemical force of attraction by which two or more ions/atoms/molecules/elements are held together to form a stable state.

We know that valency is the combining capacity of an atom of a particular element with other elements.

WBBSE Notes For Class 10 Physical Science And Environment

On the basis of the electronic concept of valency, we will learn about two types of chemical bonds, which are—

1. Ionic or electrovalent bonds and
2. covalent or molecular bonds.

Ionic bonds are formed through the complete transfer of electrons from the outermost shells between a metal and a non-metal.

For example solid NaCI (Na⁺cr), MgCI₂ (Mg²⁺2Cr), CaO (Ca²⁺ O^2—), etc. Here (Na⁺cr) is an ionic compound.

On the other hand, covalent bonds are formed through the mutual sharing of one or more electron pairs between two or more non-metals.

For example: H₂(H-H), O₂(O=O), N₂(N=N), etc. H₂ and O₂ are covalent compounds.

1. Metals have a tendency to lose electrons.
2. Non-metals have a tendency to accept electrons.
3. Metals get oxidized, while non-metals get reduced.
4. As ionization potential is inversely proportional to the tendency to lose electrons, metals should have low ionization potential. As electron affinity is directly proportional to the tendency of accepting electrons, so non-metals should have high electron affinity.

WBBSE Chapter 8 Ionic And Covalent Bonding Properties Of Ionic Compounds

The crystalline structure of sodium chloride where Na⁺ (cations) and Cl¯ (anion) ions are held together by a very strong electrostatic force of attraction, mainly because of which different bulk properties in ionic solids are observed, such as:

Physical state: Ionic compounds are generally hard and brittle solids.

In ionic compounds, particles are only ions (both cations & anions) due to the transfer of electrons.

Melting point (m.p.) and boiling point (b.p.): Because of the very strong electrostatic force of attraction, ionic compounds have high m.p. & high b.p. Lot of temperature (energy) is required for melting & boiling of ionic solids.

For example: for solid NaCI m. p. is = 820 e and b. p. = 1600°C.

Electrical conductivity: They conduct electricity in a molten state or aqueous state but not in a solid state.

Reason: In solid-state (Na⁺CP) — the Na⁺ & Cl^– ions are tightly bound with each other and the ions are not free to move to conduct electricity.

But when Na⁺CP is molten (which means heated enough) or in the aqueous state (which means NaCI in water), the Na⁺ and Cl” ions break down and become mobile.

Know: Water molecules have a property to lessen the force between two charges and this is known as dielectric property.

Exception: CaF₂, Ba₃(PO₄)₂—for these compounds, the electrostatic force of attraction is so huge that water cannot separate out the ions.

Solubility: Usually highly soluble in water. You know from your daily experience that NaCI is such a compound.

But they are insoluble in organic solvents like alcohol/acetone/benzene/toluene/CCI₄/CS₂ etc. because water is itself a polar solvent.

WBBSE Chapter 8 Ionic And Covalent Bonding Discovery Of Noble Gases Ionic Bonding

Discovery of noble gases: We know that Earth’s atmosphere contains very little amount of noble gases mainly argon.

In the last decade of the 19th century, English scientist Rayleigh and Scottish scientist Ramsay were able to separate out the noble gases or inert gases (He, Ne, Ar, Kr, Xe, and Rn) which are chemically inactive.

In the first half of the 20th century, Lewis and Kossel separately predicted that except He atoms (which have 2 electrons), other noble gas atoms have 8 electrons in their outermost / valence shell.

Electronic configuration: ₂He (2), ₁₀Ne (2, 8), _lgAr (2, 8, 8), ₃₆Kr (2, 8, 18, 8), ₅₄X(2, 8, 18, 18, 8), ₈₆Rn(2, 8, 18, 32, 18, 8).

According to Gilbert Newton Lewis, an American physical chemist, noble gas atoms have complete outermost shells, which made them stable in comparison to rest atoms.

So we can say that for this special type of electronic configuration, the noble-gas atoms are especially stable.

Thus, the stability of an atom is decided by its electronic configuration.

There are two stability rules-

Octet Rule: If an atom has 8 electrons in an outermost shell then the atom becomes stable,

Duplet Rule: If an atom has only one shell and 2 electrons in the outermost shell then also the atom becomes stable.

₃Li, ₁H follows the duplet rule to become stable. All the atoms other than noble-gas atoms are trying to follow the octet rule to acquire stability.

An atom can achieve an octet or a duplet structure in two ways:

1. Either by transfer of electrons or
2. By mutual sharing of electron pairs.

The respective type of chemical bonding is

1. ionic bonding and
2. covalent bonding.

Let us now see the electronic configuration of elements of the third period and their corresponding ions

Note:

Na, Mg, and Al try to acquire electronic configurations like their nearest inert gas Ne

S (Group-14) and Cl (Group-17) try to acquire electronic configuration like their nearest inert gas Ar.

Si (Group-14), and P (Group-15) do not form ions.

Ionic bonding: Kossel established the existence of Na⁺ & CP ions in ionic compounds. NaCI from the idea of conductivity and other experiments.

It has also been established that ionic bonding takes place when a metal and a non-metal try to acquire an electronic configuration like their nearest inert gas.

Let us take some examples:

Bonding in NaCI (Na⁺CP): Electronic configuration: _nNa = 2 + 8 + 1 and ₁₇CI = 2 + 8 + 7

Neither Na nor Cl has stable electronic configurations. To avail stability, the Na atom needs either to give away le^_ or to accept 7 more e¯s in the outermost shell.

Similarly, the C1 atom needs either to accept le¯or give away all the 7 e¯s from the outermost shell. For the Na atom, it is easier to give away le¯ than to accept 7 e¯s (because less energy is utilized).

A similar reason is applicable to the Cl atom. So for attaining a stable electronic configuration, the Na atom will lose le¯ from its valence cell to become the Na⁺ ion & Cl atom will gain that single e^– to become the CP ion. Basically, this is a complete transfer of electrons.

These Na⁺ and CP ions attract each other by strong electrostatic Schematic diagram force and they form ionic compound NaCI.

the relative position of Na⁺ and CP ions in the schematic diagram of the NaCI crystal lattice)

Here, the no. of electrons lost or gained by the atoms at the time of formation of the ionic bond is called electrovalency.

In this example, the electrovalency of Na⁺= 1 and that of Cl¯ = -1.

Note :

When metals of Group 1 and 2 react with non-metals of Group 16 and 17, ideal ionic bonding is formed,

Due to the strong electrostatic force of attraction between cations and anions, the ionic lattice structure of NaCI is very stable.

Definition: To acquire the nearest noble-gas-like stable electronic configuration, cations, and anions are formed through the transfer of electrons from neutral metal and non-metal atoms.

The electrostatic force of attraction which holds oppositely charged ions together is called the ionic bond. The new compound so-formed is called an ionic compound.

This type of combining capacity is called ionic valency or electrovalency.

Representation of ionic bonding in terms of electron dot (•) / cross (x) structure:

Note: Li⁺ & H ions in the LiH compound fulfill the duplet rule in their outermost shells with no question of making an octet.

Why for ionic compounds, the concept of formula mass is more appropriate than molecular mass?

Ionic compounds are made up of huge no. of cations and anions arranged in a 3-dimensional stable ionic structure (strong electrostatic force of attraction).

They contain ions only no molecules are there.

So the concept of molecular mass is not appropriate in the case of ionic compounds-rather, formula mass is more appropriate.

For example, not saying molecular mass of NaCl = 23 + 35-5 = 58-5, we will say, the formula mass of NaCI = 58-5 or molar mass of NaCI = 58-5 g/mol.

Solid NaCI contains Na⁺and Cl” ions which are not free to move due to the existence of a strong electrostatic force of attraction between the ions even when an external potential difference is applied across a NaCI crystal.

But when NaCI is in a molten or aqueous state, then the ions become free to move and can conduct electricity.

That’s why we say that ionic solids have low intrinsic electrical conductivity in contrast to that in their molten or solution states.

WBBSE Chapter 8 Ionic And Covalent Bonding Properties Of Covalent Compounds

In order to achieve a stable electronic configuration like the nearest noble gases, atoms of two or more non-metals come together.

This coming together and mutual sharing of one or more electron pairs leads to the formation of a chemical bond known as covalent bonding.

The new compound so formed is called covalent compound and this type of combining capacity is called covalency.

Some common compounds in which covalent bonding is found to exist are naphthalene, sugar, water, ethanol, methane, chloroform, carbon dioxide, carbon monoxide, hydrogen chloride, ammonia, etc.

Let us see some of their physical properties:

From this table, we come to know about some common properties of covalent compounds, which are—

Physical state: Covalent compounds exist as gases, liquids, or soft solids because of the very weak intermolecular force of attraction (weak Van der Waals force) existing between the molecules.

Due to this fact, a covalent bond is weak in comparison to an ionic bond. In covalent compounds, particles are molecules with no existence of ions.

Melting point (m.p.) and boiling point (b.p): Covalent compounds have low m.p. and low b.p.

Because of the very weak intermolecular force of attraction, solid covalent compounds can easily melt and boil, a liquid covalent compound is volatile in nature and a gaseous covalent compound exists in vapor form.

Solubility: Generally, covalent compounds are insoluble in polar solvents like water, but soluble in non-polar solvents like alcohol, and ether.

Exception: Cane sugar, glucose, H₂S, NH₃ HCI (g), ethanol, etc are soluble in water.

Electrical conductivity: They are non-conductors in a solid, molten, or aqueous state. Because covalent compounds do not contain ions to conduct electricity.

Exception: Aqueous conduct electricity in very little proportion.

WBBSE Chapter 8 Ionic And Covalent Bonding Covalent Bonding

According to the proposal of G. N. Lewis (1916) about electron pairs, the concept of covalent/ chemical bonding has grown up.

Covalent bonding is said to exist between non-metal atoms and covalent compounds are formed through mutual sharing of electrons.

Here we will discuss Lewis dot diagrams (where only the valence electrons are represented by a dot (•) / cross (x) (2D structure). Let us understand covalent bonding with the help of examples:

Bonding in F₂molecule: Electronic configuration: _gF = 2 + 7.

Each F atom has 7 e^_s in its valence shells. They need 1 more e¯ to attain stable e-configuration like the nearest noble gas Ne (complete octet).

Two F atoms come together and they share le each in a valence shell to complete the octet. Ultimately a molecule of F₂ is formed.

Since a single pair of e” is shared between two F atoms, the bond is known as a single covalent bond and is represented by a single dash (—) between two F atoms.

Bonding in H₂0 (water molecule): Electronic configuration: _XH = 1 & ₈0 = 2 + 6.

Each H atom needs 7 e¯ in its valence shell to complete the duplet and the 0 atom needs 2 more e¯s to complete the octet. So, in a water molecule (H₂0) two single covalent bonds are formed.

Bonding in CO₂ molecule: Electronic configuration : ₆C = 2 + 4 and _gO = 2 + 6. Each  O atom needs 2 more e¯s to complete the octet and the C atom needs 4 more e¯s to complete the octet.

So, when two O atoms and one C atom come together, each O atom will share 2 e¯s from the valance shell with 2 e¯s of the C atom in the valance shell so that all atoms acquire stable electronic configuration.

Since each O atom shares 2e¯s with the C atom, so two double covalent bonds are produced keeping C at the center and O atoms at two ends.

Bonding in N₂ molecule: Electronic configuration: ₇N = 2 + 5.

N atom has 5 e¯s in its valence shell. It needs 3 more e¯s to complete its octet.

So, when two N atoms come together, each N atom will share its 3 e¯s with another N atom and becomes stable.

So, each N atom shares 3 e¯ s with each other. Since the two N atoms share 3 pairs of e¯s, there is a triple covalent bond between two N atoms.

Few more examples of covalent bonding:

Note: (1)  The no.of Electrons shared by an atom for the formation of shared electron pairs to covalent compounds is called the covalency of that particular atom.

For example: As H or Cl Or F Atom Shares One explain Pair so the covalency of H or Cl or F is 1. O atom shares two pairs of electrons,

So its Covalency =2 n atom Shares 3 Pairs Of electrons, So its Covalency = 2 N atom dot structure of O-does does not support the experimental result of the paramagnetic property of O₂ molecules paramagnetic property refers to the ability to get attracted to the magnetic field).

WBBSE Chapter 8 Electricity And Chemical Reactions Electrolytes

Electrolytes Definition: Electrolytes are the chemical compounds that ionize or dissociate into their ions [cations (+) and anions (-)] in a molten/aqueous state and the solutions so-produced have the ability to conduct electricity (d.c.).

Take some table salt (NaCI, an ionic compound) and heat over 801°C, it melts and the ions Na⁺ and CP already present in it, get released.

Similarly, in its aqueous solution, H₂O molecules occupy some positions of NaCI molecules and because of the dielectric property of water, the electrostatic force of attraction between Na⁺ and Cl” ions is decreased so the ions get free.

Ionic compounds (NaCI, KCI, CuSO₄, CuCI₂, NaOH, etc.) in a molten/aqueous state can conduct electricity through these ions.

They behave like conductors but are different from metals. But solid NaCI is not an electrolyte. Because in the solid state, a strong electrostatic force of attraction exists between the Na⁺ and CP ions the ions are not free to move.

WBBSE Notes For Class 10 Physical Science And Environment

In the case of covalent compounds like HCI (g), no ions exist but in an aqueous solution, it forms HCI (strong acid) which can be dissociated into H⁺ and CP ions. This is called ionization. Other such examples are HF, HBr, HI, NH_3, etc.

Also, there are some covalent compounds like (pure) distilled water, sugar solution, alcohol, ether, benzene, kerosene, molten sulfur, ethanol, urea, and glycerine which do not dissociate into ions and cannot conduct electricity even by little amount.

They are called nonelectrolytes. They behave like non-conductors of electricity because they do not belong to free electrons (like metals) and ions (like electrolytes).

WBBSE Chapter 8 Electricity And Chemical Reactions Strong And Weak Electrolytes

This classification is based on the degree of ionization or the efficiency of dissociation that an electrolyte exhibits in its aqueous solution.

A greater degree of ionization corresponds with a stronger electrolyte. We can classify strong electrolytes, weak electrolytes, and non-electrolyte by measuring the electrical conductivity of their aq. solutions.

See the experimental- 

Observation of the experiment: Aq. solution of NaCI or KCI or H₂SO₄ or NaOH can conduct electricity very well so they are strong electrolytes.

On the other hand, acetic acid CH₃COOH aq.) or ammonium hydroxide (aq.) can also conduct electricity but not as well as a strong electrolyte they are weak electrolytes.

The non-electrolytes like ethanol/ benzene/ glycerine show zero conductivity.

Strong electrolyte: If a substance is molten/aq. the state completely dissociates into ions such that the solution consists entirely of ions, and molecules left, and can conduct electricity very well is called a strong electrolyte.

Strong acids (H₂S0₄, HN0₃, HCI), strong soluble bases (NaOH, KOH), and soluble salts (NaCI, K₂S0₄, CuS0₄, CuCI₂) are strong electrolytes.

Ionization reaction:

⇒ \(\mathrm{NaCl} \text { (aq.) } \rightarrow \mathrm{Na}^{+}+\mathrm{Cl}^{-} \text {. } \mathrm{CuSO}_4 \text { (aq.) } \rightarrow \mathrm{Cu}^{2+}+\mathrm{SO}_4^{2-} . \mathrm{H}_2 \mathrm{SO}_4 \text { (aq.) } \rightarrow 2 \mathrm{H}^{+}+\mathrm{SO}_4{ }^{2-} \text {. }\)

Since every molecule of a strong electrolyte dissociates into ions, so a single arrow (→) is used in the dissociation reaction. It’s a chemical change.

Weak electrolyte: If a substance in aq. solution partially ionizes (which means only a few molecules break into ions and most stay as neutral molecules) which still can conduct electricity but not as well as a strong electrolyte then it is called a weak electrolyte.

Weak acids (CH₃COOH, H₂C0₃), weak bases (NH₄OH), and some salt (AgCI) are examples.

Ionization reaction:

⇒ \(\mathrm{CH}_3 \mathrm{COOH} \text { (aq.) } \rightleftharpoons \mathrm{H}^{+}+\mathrm{CH}_3 \mathrm{COO}^{-} . \mathrm{NH}_4 \mathrm{OH} \text { (aq.) } \rightleftharpoons \mathrm{NH}_4^{+}+\mathrm{OH}^{-} .\)

In such solutions, an equilibrium is established between the ions and the non-ionized neutral molecules.

This is represented by a double arrow () between the ions and the non-ionized molecules of weak electrolytes.

WBBSE Chapter 8 Electricity And Chemical Reactions Mechanism Of Electrical Conduction In Molten Solution State

Take molten/aq. solution of table salt (NaCI) in a glass bowl. In the solution state, NaCI contains plenty of Na⁺ (cation) and Cl^– (anion) ions. \(\left[\mathrm{NaCl} \text { (aq.) } \rightarrow \mathrm{Na}^{+}+\mathrm{Cl}^{-}\right]\)

Two conducting rods (electrodes) are inserted into the solution.

The rod connected with the + ve terminal of the battery is the anode (A) and the rod connected with the – ve terminal of the battery is the cathode (C).

[The electrodes are usually made of inert material (platinum/graphite) to ensure that the electrodes itself do not involve in the electrolytic reactions.] Now let us see what happens when the electrodes (say 9V or 12V).

⇒ \(2 \mathrm{NaCl} \text { (aq.) } \stackrel{\text { electricity }}{\longrightarrow} 2 \mathrm{Na}^{+}+2 \mathrm{Cl}^{-} \text {. }\)

The Na⁺ ions get attracted towards the – ve electrode (cathode) and Cu ions are attracted by the +ve electrode (anode).

Thus, the cations and anions move in opposite directions for the purpose of electrical discharge.

At the cathode, each Na⁺ ion captures le^_ to produce metal Na (it is a reduction reaction) and at the anode, each Cl¯ ion gives up le¯ and becomes neutral Cl.

Atom (it is an oxidation reaction). The single Cl atoms then pair up and form Cl₂ gas. Electron flows from the anode to the cathode in the external circuit.

At the Cathode:

⇒ \(\begin{array}{r}
2 \mathrm{Na}^{+}+2 e^{-} \rightarrow 2 \mathrm{Na}(s) \ldots . . \\
\quad \text { (reduction reaction) }
\end{array}\)

At the anode:

⇒ \(2 \mathrm{Cl}^{-} \rightarrow 2 \mathrm{Cl}+2 e^{-}, 2 \mathrm{Cl} \rightarrow \mathrm{Cl}_2(g)….. (oxidation reaction)\)

So, solid sodium is formed at the cathode, and chlorine gas being given released at the anode (chemical change).

Here, the coefficients used in the equation (for overall electrical neutrality of the entire system) indicate mol numbers.

In this reaction, at the anode 2 mol Cl^– ion removes 2 mol e¯, and finally, 1 mol Cl₂ gas is produced. Similarly, at the cathode 2 mol Na⁺ ion gain 2 mol e¯, and finally, 2 mol Na (s) is produced.

That’s why, the mechanism of electrical conduction metals and electrolytes are not exactly the same.

In electrical conduction through molten/aq. electrolytes only the anions and cations (not electrons like conductors) take part, and during electrolysis, they move in opposite directions.

The difference in electrical conductivity between metals and electrolytes:

WBBSE Chapter 8 Electricity And Chemical Reactions Electrolysis

Electrolysis Definition: Electrolysis is a chemical dissociation/ionization of an electrolyte (chemical change) in a molten/aqueous state by passing direct current (d.c.)-

Electrolysis Example: 

⇒ \(\mathrm{CuSO}_4 \text { (aq.) } \stackrel{\text { electricity }}{\longrightarrow} \mathrm{Cu}^{2+}+\mathrm{SO}_4^{2-} ; \quad \mathrm{HCl} \text { (aq.) } \stackrel{\text { eléctricity }}{\longrightarrow} \mathrm{H}^{+}+\mathrm{Cl}^{-}\)

The ionic compounds in their solution state, ionize into cations and anions which conduct electricity.

In the case of covalent compounds like HCI (g)- no free ions are there, which in an aqueous state dissociates into H⁺ and Cl¯ ions.

In the case of electrolysis, electrical energy → is chemical energy. Just the opposite for an electric cell where chemical energy → electrical energy.

In the case of electrolysis, the source of electricity must be d.c. (obtained from a cell/battery of low voltage). No heavy current is to be used. No. a.c.

Electrolysis is a type of Redox reaction. Always oxidation (loss in e¯ anion) takes part at the anode and reduction (gain in e¯ by cation) at the cathode.

With respect to electron transfer, the cathode can be defined as the electrode towards which cations move and the anode is the electrode towards which anions move.

Electrolysis of acidified water using Pt-electrode: Already learned that pure water has no free ions to conduct electricity. Few drops of oil. H₂SO₄ can ionize pure water.

Then pure water becomes to be an electrolyte.

Electrolytic cell: Electrolysis of acidified water forms a cell, called Hoffman Voltameter.

The arrangement in which two Pt-electrodes and a 12 V d.c. supply batteries are used. (Pt is inert and cannot react with the solution).

Ionization:

⇒ \(\mathrm{H}_2 \mathrm{O}(\mathrm{H}-\mathrm{OH}) \rightarrow \mathrm{H}^{+}+\mathrm{OH}^{-}, \mathrm{H}_2 \mathrm{SO}_4 \rightarrow 2 \mathrm{H}^{+}+\mathrm{SO}_4{ }^{2-}\)

When more than one anions/cations come into the electrolytic solution, then only one of them gets preferentially discharged at the electrode.

This factor depends on the position of anion/cation in the electrochemical series, from which we can determine the tendency to lose/ gain e¯s.

This tendency is as follows:

At Cathode:

⇒ \(\begin{aligned}
& \mathrm{Ag}^{+}>\mathrm{Cu}^{2+}>\mathrm{H}^{+}>\mathrm{Fe}^{2+}>\mathrm{Zn}^{2+}>\mathrm{Al}^{3+}>\mathrm{Mg}^{2+} \\
& >\mathrm{Na}^{+}>\mathrm{Ca}^{2+}
\end{aligned}\)

At anode:

⇒\(\mathrm{I}^{-}>\mathrm{Br}^{-}>\mathrm{Cl}^{-}>\mathrm{OH}^{-}>\mathrm{NO}_3^{-}>\mathrm{SO}_4^{2-}>\mathrm{F} .\)

As soon as the switch is on, a lot of bubbles are seen to be formed in two tubes.

This happens because – both the an-ions OH and SO₄^2¯ get attracted towards the anode while cations H⁺ are attracted by the cathode.

According to the electrochemical series, OH¯ has a greater tendency to lose e” than SO₄^2¯. Thus, at the anode, each OH¯ ion loses le¯ (oxidation), and at the cathode, each H⁺ ion gains le¯ (reduction).

So electron flow continues from anode to cathode in the external circuit and constitutes a flow of electric current from anode to cathode in the electrolyte solution.

At anode:

⇒ \(\begin{aligned}
& \mathrm{OH}^{-} \rightarrow 1 e^{-}+\mathrm{OH} \text { (unstable); } \mathrm{OH}+\mathrm{OH} \rightarrow \mathrm{H}_2 \mathrm{O}+\mathrm{O} \text { (nascent); } \mathrm{O}+\mathrm{O} \rightarrow \mathrm{O}_2 \\
& 4 \mathrm{OH} \rightarrow 4 e^{-}+4 \mathrm{OH} ; 4 \mathrm{OH} \rightarrow 2 \mathrm{H}_2 \mathrm{O}+\mathrm{O}_2 \uparrow \text { (oxidation) }
\end{aligned}\)

At Cathode: 

⇒ \(\begin{aligned}
& \mathrm{H}^{+}+1 e^{-} \rightarrow \mathrm{H} ; \mathrm{H}+\mathrm{H} \rightarrow \mathrm{H}_2 \\
& 4 \mathrm{H}^{+}+4 e^{-} \rightarrow 2 \mathrm{H}_2 \uparrow \text { (reduction) }
\end{aligned}\)

Observation: 

At the – ve and + ve electrodes, the gases collected are detected as H₂ and O₂ respectively. The ratio of the volume of H₂ and O₂ is 2 :1.

Electrolysis of CuSO₄ (aq.) using inert (Pt/graphite) electrode: Two Pt-electrodes are used as anode (+ ve) and cathode (- ve).

In place of Pt, graphite can also be used. Pt. is nonreactive but there is a little chance of oxidation of graphite.

Ionization: \(\mathrm{CuSO}_4 \text { (blue solution) } \rightarrow \mathrm{Cu}^{2+}+\mathrm{SO}_4^{2-}, \quad \mathrm{H}_2 \mathrm{O}(\mathrm{H}-\mathrm{OH}) \rightarrow \mathrm{H}^{+}+\mathrm{OH}^{-}\)

As soon as the switch is on, the -vely charged OH¯ ions move towards the anode (+ ve), and + very charged Cu²⁺ ions move towards the cathode (- ve).

Because according to the electrochemical series, OH¯ has a greater tendency to lose e¯ then SO₄^2+, and Cu²⁺ has a greater tendency to gain e¯ than H.

Then OH¯ and Cu²⁺ ions are discharged at the anode and cathode respectively.

At Anode:  \(\mathrm{OH}^{-} \rightarrow 1 e^{-}+\mathrm{OH} ; 4 \mathrm{OH}^{-} \rightarrow 4 e^{-}+4 \mathrm{OH} ; 4 \mathrm{OH} \rightarrow 2 \mathrm{H}_2 \mathrm{O}+\mathrm{O}, \uparrow \text { (oxidation) }\)

At Cathode: \(\mathrm{Cu}^{2+}+2 e^{-} \rightarrow \mathrm{Cu} \downarrow ; 2 \mathrm{Cu}^{2+}+4 e^{-} \rightarrow 2 \mathrm{Cu} \downarrow \text { (reduction) }\)

Observations :

Products are O₂ gas at the anode and Cu deposited at the cathode.

Pink/reddish brown Cu metal gets deposited at the cathode. So, cathode mass increases.

The presence of Cu²⁺ ions makes CuSO₄ While electrolysis, these Cu²⁺ ions from the solution get deposited slowly at the cathode, so the blue color of CuSO₄ solution fades away.

Electrolysis of CuSO₄ (aq.) using Cu electrodes: 

In the Previous case, we used Pt-electric rods in the case, and both the electrodes are of Cu.

In this case, both electrodes are of Cu. Now let’s see what difference is observed in the electrolysis of CuSO₄ solution.

Now let’s see what difference is observed in the electrolysis of CuSO₄ solution.

The ions present in the solution are

⇒\(\mathrm{Cu}^{2+}+\mathrm{SO}_4^{2-}\mathrm{H}^{+}, \mathrm{OH}^{-}\)

At cathode: \(\mathrm{Cu}^{2+} \text { (aq.) }+2 e^{-} \rightarrow \mathrm{Cu}(s) \text { [Reduction reaction] }\)

So Cu²⁺ ions on being discharged are deposited at the cathode as Cu-atoms.

At anode: \(\mathrm{Cu}(\mathrm{s}) \rightarrow \mathrm{Cu}^{2+}+2 e^{-} \text {[Oxidation reaction] }\)

Within the solution, Cu²⁺ ions exist and the anode itself is made up of Cu-atoms. In such a case, neither OH+ nor SO₄²⁺ will be discharged – rather, at the anode each Cu-atom losing 2e¯s goes into the solution as a Cu²⁺ ion.

As a whole, when a Cu² ion is deposited at the cathode, simultaneously from the anode another Cu²⁺ ion goes into the solution.

Ultimately, the amount by which anode mass decreases, and cathode mass increases by the same amount.

Also, the concentration of CuSO₄ solution effectively remains constant. So, the blue color of the CuSO₄ solution doesn’t fade away. (This process is used as electroplating).

WBBSE Chapter 8 Electricity And Chemical Reactions Applications Of Electrolysis

Electrolysis is mostly used in industry in many ways such as

In metal extraction,

In electrorefining,

In electroplating.

Electrolysis in the extraction of metals from their ores: Highly electropositive (active) metals like Na, Ca, and Al are extracted from their ores using electrolysis.

Al-Extraction: The principal ore of Al is bauxite (Al₂O₃-2H₂O).

It is purified to yield aluminum oxide (Al₂O₃) [white powder – its m.p. is very high (over 2000°C)].

Al₂O₃ is dissolved in molten cryolite (Na₃AlF₆) [an Al-compound with much lower m.p. than Al₂O₃ and the temperature is maintained at ~ 925°C.

Molten Al₂O₃ contains Al-ions (Al³⁺) and Oxide-ions (O^2- ).

Both the electrodes are made up of graphite. More than one graphite rod is used as the anode.

At the Anode: \(\mathrm{O}^{2-}-2 e^{-} \underset{\text { oxidation }}{\longrightarrow} 0 ; 2 \mathrm{O}^{2-} \rightarrow \mathrm{O}_2(g)+4 e^{-}\)

At the Cathode: \(\mathrm{Al}^{3+}+3 e^{-} \underset{\text { Reduction }}{\longrightarrow} \mathrm{Al}(\mathrm{s})\)

The O₂ molecules formed at the anode react with the graphite (carbon atoms) rods (Anode) and form  \(\mathrm{CO}_2 \text { gas. }\left[\mathrm{C}+\mathrm{O}_2 \rightarrow \mathrm{CO}_2(g)\right] \text {. }\)Due to this reason, the anode rods are regularly replaced.

Electro-refining of copper: Electro-refining is a process of purifying an impure metal to a high purity level to get pure metal using electrolysis.

Electrolytic solution: Aq. CuSO₄ + dil. H₂SO₄.

Cathode: A thin sheet of pure Cu.

Anode: Impure sheet of Cu. The ions present in the solution are Cu²⁺, SO₄^2-.

At the anode: From impure sheet \(\mathrm{Cu}(\mathrm{s}) \underset{\text { oxidation }}{\longrightarrow} \mathrm{Cu}^{2+}(\mathrm{aq} .)+2 e^{-}\)

At the cathode:

⇒ \(\mathrm{Cu}^{2+}(\text { aq. })+2 e^{-} \underset{\text { Reduction }}{\longrightarrow} \mathrm{Cu}(\mathrm{s})\)

The Cu²⁺ ions move between the anode and the cathode. When a Cu²⁺ ion is deposited cathode, at the same time, another Cu ion from an impure Cu sheet goes into the solution.

Thus, the concentration of CuSO₄ solution effectively remains the same.

As this happens, the size of the cathode sheet gets larger while the size of the anode sheet gradually decays, a sludge is left at the bottom of the anode which is called anode mud.

The anode als mud (Ag/Au) may contain many valuable met- Cu obtained from electro-refining and is about 99-99% pure.

Electroplating: Electroplating is an electrolytic process by which a thin layer of superior metal or noble metal like gold, silver, nickel, or chromium is deposited/coated on the surface of a base metal either for protection against corrosion or making attractive.

Basic principles of electroplating:

The article to be electroplated is cleaned first,

Metal on which electroplating is to be done is used as the cathode (- ve electrode). Because deposition always takes place at the cathode,

A pure block of the metal to be electroplated is used as the anode (+ ve electrode),

A solution of a salt of the metal is used as the electrolyte (the solution contains metal ions).

For example, in Cu-plating the electrolyte is CuS0₄(aq.); in silver plating, silver nitrate (AgNO₃)/potassium argento cyanide K [Ag (CN)₂J solution;

In gold-plating, potassium auro cyanide K [Au (CN)₂] solution; in nickel plating, a mixed solution of

⇒ \(\mathrm{NiSO}_4+\left(\mathrm{NH}_4\right)_2 \mathrm{SO}_4+\text { slight boric acio }\)

For getting uniform deposition, a low direct current (3A or 4A) for a longer time is passed through the solution. (No heavy current to be used or no a.c. to be used).

WBBSE Chapter 8 Physical And Chemical Properties Of Matter Periodic Table And Periodicity Of The Properties Of Elements

Brief History Of The Periodic Table

In the 18th century, only 31 elements were known. Up to 1865, more elements (63) were identified. At that time, the number of known elements was increasing, and separate studies of each element were becoming difficult.

Scientists were trying to classify the elements having identical properties in a single class/group.
The first attempt was made in 1829 by German chemist Johann Dobereiner.

He made groups of 3 elements in the order of increasing atomic mass such that the atomic mass of the middle element is equal to the average value of the atomic mass of the corner elements.

This group of elements was named Dobereiner’s traits He could identify only 3 triads for the elements known at that time.

For the first time, Dobereiner’s model gave an idea of the classification of elements. But all known elements could not be arranged in triads.

WBBSE Notes For Class 10 Physical Science And Environment

In some cases, even in the same group, this model was not obeyed.

For example F-19, CI-35.5, Br-80, the average atomic mass \(=\frac{19+80}{\cdot 2}=49 \cdot 5\) (it did not match). So, ultimately this model was rejected.

After a long time, in 1864, English scientist Newlands, making use of the octave concept of the music system was successful in arranging 8 elements in the order of increasing atomic mass, to make an octave.

Newlands found that every 8th element has chemical properties similar to the 1st element, the same for the 2nd and 9th elements, etc.

This model was successful up to Ca (only for lighter elements not for all elements known at that time). For the first time, Newlands gave the idea of periodicity/repetition of properties.

In 1869-70, German chemist Lothar Meyer and Russian chemist Dmitri Mendeleev were studying to find out the similarity in properties of the arrangement of elements.

Both scientists reached almost the same level of conclusion but in different ways.

Meyer’s investigation was based on a graphical analysis of atomic volume vs atomic mass

According to Meyer, elements at similar positions on the graph have similar properties.

Example: at the peaks he found alkali metals of Gr. 1 (Na, K, Rb), in descending slopes→ alkaline earth metals (Mg, Ca, Sr), in ascending slopes → halogens (Cl, Br, I).

Meyer classified only 28 elements. But from the graph, exact information about elements still remained unexplained.

In the same year, Mendeleev was studying to find out the relation between the atomic mass of

elements and their physical properties (density, LotherMeyersgraph melting point, boiling point) and chemical properties (how an element reacts with other elements).

Mendeleev’s Periodic Law: The physical and chemical properties of the elements are the periodic function of their atomic masses.

Based upon this law, Mendeleev arranged only 63 elements known at that time in order of increasing atomic mass in tabular form which he named Periodic While arranging so, he made 7 horizontal rows called Periods and represented these as periods 1, 2, 3 . . . 7.

He told in the same period moving from left →right, properties shift from metal-non-metal. Mendeleev’s original periodic table has 8 vertical columns which he named Groups

and numbered as Gr. I, II, III, VIII. After the discovery of noble gases (1891 – 1902), a separate Gr. (0) has been added.

He also left some gaps. The revised version of Mendeleev’s Periodic Table has 7 periods and 9 groups. He divided Grs. I → VII into sub-groups A, B.

But Gr. VIII has no sub-group it has 3 elements in a single period. He suggested that elements in the same sub-group have the same valency/identical properties.

Actually, Mendeleev’s analysis was based on the formulae of oxides and hydrides that the elements can form because 0₂ and H₂ are very much reactive and form compounds with most elements

Mendeleev’s Periodic Table completely changed old ideas. Although it was different from the modern one, it was a huge step in the classification of elements.

Mendeleev is considered the ‘Father of the Periodic Table.’

Few achievements of Mendeleev: In his periodic table, he left gaps below Al (26-98) and named it eka-aluminum (later it was discovered Ga) and below Si (28-09) and named it eka-silicon (later discovered element was Ge), etc.

In this way, he made a strong prediction of some undiscovered elements. Also, successfully predicted some of their physical and chemical properties.

In 4 misfit positions, he placed an element with a slightly greater atomic mass before an element with a slightly lower atomic mass.

For example Ar (39-95) and K (39-10); Co (58-93) and Ni (58-71); Te (127 60) and I (126-90); Th (232-04); and Pa (231).

Perhaps, this was done so that elements with similar properties could be grouped together.

After the discovery of noble gases, they could be placed in a new group without disturbing the existing order.

Improvements needed In Mendeleev’s Periodic Table: O In some cases, dissimilar elements were placed in the same group. For example : (Na, K, and Rb) were placed in the same Gr. I with (Cu, Ag).

In some cases, similar elements were placed in separate groups. Example : (Pt and Au).

Isotopes of the same element should have different positions in the periodic table.

No fixed position could be given to hydrogen because of its similarities with Gr IA elements (alkali metals) and Gr. VII B elements (halogens).

WBBSE Chapter 8 Physical And Chemical Properties Of Matter Modern Periodic Table

In 1911 Henry Moseley, while experimenting with X-rays taking 38 different elements as target material proved that the Frequency Characteristic To an Element Depends On the Atomic

a number of the element \((\sqrt{f} \propto Z)\), where Z = atomic number = a number of p’, present in the nucleus of an atom.

For a neutral atom, the number of ps = the number of e s. According to this concept, the chemical properties of the elements are more related to their atomic numbers than their atomic masses.

Modern Periodic Law: The physical and chemical properties of the elements are the periodic function of their atomic numbers.

Modern Periodic Table (Long Form of Periodic Table): Soon after the discovery/ of Moseley, all elements increased in the order of their atomic numbers instead of atomic masses. A group of scientists contributed to this work.

In 1920 Niels Bohr prepared the Modem Periodic Table, which is popularly known as the Long Form of the Periodic it is prepared on the basis of the electronic configuration of the elements.

The modem periodic table has 7 horizontal periods (1, 2, 3, 4, 5, 6, 7) and 18 vertical columns (1 – 18). As per the proposal of IUPAC (1984), the groups are numbered 1, 2, 3, 4, and 18 in place of Mendeleev’s sub-groups A, and B.

The elements mentioned in serial no. 1, 2, and 3, are the atomic number of different elements with every element the atomic number increasing by 1 means 1 extra p or 1 extra e¯ is added to the chemical structure.

Significance of Period: The number of shells (orbits) present in an atom is the period number. As in every next period, 1 more shell gets added.

Significance of Group: Elements having similar chemical properties or more specifically having similar electronic configurations are grouped together. Examples:

(1) Gr. I. elements (₁H – ₇₈Fr) have similar electronic configurations: \({ }_1 \mathrm{H}(\underline{1}),{ }_3 \mathrm{Li}(2, \underline{1}),{ }_{11} \mathrm{Na}(2,8, \underline{1}),{ }_{19} \mathrm{~K}(2,8,8, \underline{1})\)

.. All have 1 valence e¯ in the outermost shell. So they have the same valency = 1 and the same other chemical properties.

All Gr. 1 elements except are called alkali metals, because they form strong alkalis with water.

(2) Gr. 2 elements (₄Be – ₈₈Ra) have 2 valence e¯s: ₄Be = (2, 2), ₁₂Mg (2, 8, 2), ₂₀ Ca (2, 8, 8, 2) . . . So their valency = 2.

All Gr. 2 elements are called alkaline earth metals because they form weaker alkalis as compared to Gr. 1. elements. Their oxides are available on Earth.

(3) Gr. 17 elements (_gF – _g5 Al) have 7 valence e¯s and can accept le^_ in the outermost shell to complete the octet. ₉F (2, 7), ₁₇CI (2, 8, 7).

So they have the same valency = 1. These elements are called halogens (which means salt-producing).

(4) Gr. 18 elements (₂He –₈₆Rn) have fulfilled the outermost shell, for which they have zero (0) valency. ₂He(2), ₁₀Ne(2, 8), _lgAr (2, 8, 8).

These elements are called noble gases/inert gases because of their non-reacting power with other elements.

(5) Gr. 13 elements are elements of the B (Boron) family as B is the 1st member. Similarly, Gr. 14→ C (Carbon) family, Gr. 15→ N family,

Gr. 16→ O family or chalcogens (means ore forming).

Gr. 3-12 elements are called transition elements because the properties shift/transit from metal to non-metal.

Special feature:

Elements in Periodic 6 with atomic number 57 – 71 are called Lanthanides, because they start with ₅₇La (Lanthanum). They are called rare earth elements.

Elements in Period 7 with atomic number 89 – 103 are called Actinides, because they start with ₈₉Ac(Actinium). These are radioactive elements.

The success of the Modern Periodic Table: It rectified the anomalies in Mendeleev’s Periodic Table.

On the basis of the arrangement. of elements with increasing atomic number positions of 1 Co and Ni were corrected.

Isotopes have similar chemical properties. So, they belong to the same position as that of the element in the periodic table.

The position of hydrogen in the periodic, table is controversial. It belongs to some properties like Gr. 1 alkali metals and some properties like Gr. 17 halogens.

Example:

Because of these peculiarities, Mendeleev named hydrogen a ‘rogue element’. In general, hydrogen is considered as a non-metallic gas.

Some trans-uranic elements: These are elements having an atomic number greater than that of uranium i.e. Z > 92 in the periodic table.

Examples are Neptunium (Z = 93), Plutonium (Z = 94), Curium (Z = 96), Einsteinium (Z = 99), Mendeleevium (Z = 101), etc.

Except for Neptunium and Plutonium, others are man-made. The discovery of the element Copernicium (Z = 112) completes the transition element series.

WBBSE Chapter 8 Physical And Chemical Properties Of Matter Periodic Trends In The Periodic Table

When properties of elements in a period/group in the periodic table are compared, certain regularity is observed in their variations which are called periodic trends.

The properties repeat themselves (either increase or decrease) after regular intervals in the periodic table. This is a regular trend or pattern.

This trend is measured in two ways: Across a period (from left to right) and down a group (from top to bottom).

In our syllabus, we will study some periodic trends such as atomic radii, ionization energies, and oxidizing-reducing properties of elements of Gr.1-2 and 13-18.

1. Atomic radii/Atomic size: Atomic radius is the distance of the outermost from the nucleus of the atom (i.e. atom’s boundary).

It is so little that we cannot directly measure the radius of an isolated atom. But from experimental data collected from other experiments, the atomic radius can be derived indirectly.

It is expressed in a picometer (pm) unit which is even smaller than a nanometer (1 pm = 1CT¹² m).

Trends in atomic radii (along a period from left to right): Along a period, the number of shells remains the same.

A number of ps increasing i.e. nuclear charge increasing i.e. e¯s are attracted towards the nucleus by a greater force means atomic radii decrease.

Down the group (Top-Bottom): With, as we move down the group, the number of shells Number of valence e¯s remain the same.

This means outermost e¯s are going away from the nucleus, which means atomic radii increase.

The trend in atomic radii: Along a period from left to right decreases and down the group increases.

Ionization energy (I.E.):

Definition: Ionization (I.E.) energy is defined as the minimum amount of energy required to remove the outermost e¯ of an isolated neutral gaseous atom to form a cation.

i.E. is always + Ve because removal of e requires energy. It is an endothermic reaction.

I.E. is measured in eV / atom unit for an isolated atom. It means how much eV energy is required to remove e¯ from an atom.

Explanation: We know that e¯s are bound with the nucleus with a strong electrostatic force of attraction.

To remove the outermost e¯, energy is needed to be applied externally/ work to be done – which is called i.E.

Example: \(\underset{\text { (neutral) }}{\mathrm{M}^{\circ}} \rightarrow e^{-}+\underset{\text { (cation) }}{\mathrm{M}^{+}}(+\Delta \mathrm{H})\)

Another famous unit is KJ/mol for 1 mol gas is also used (leV/atom = 96 KJ / mol).

Trends in I.E. along a period (Left →Right): As we move along a period, atomic size increases meaning nuclear charge increases means attraction of e¯ increases means removal e¯ becomes tougher. So more I.E…….. is required for ₉F as compared to ₃Li.

Down the group (Top→Bottom): As we move down the group, the number of shells increases means atomic size increases means e¯ is going away from the nucleus means removal of e¯ becomes harder.

So less I.E. is required to produce K⁺ion than to produce Fl⁺ ion.

The general trend in I.E.: Along a period from left to right increases and down the group decreases. (There are some exceptional cases also. You will learn more in higher classes).

Electronegativity (E.N.):

Definition: The ability of an atom to attract shared pair of e^–s towards itself in a covalent bond within a molecule is called electronegativity.

Explanation: Let us take an example: Here, the shared pair of e¯s are attracted by both the atoms X and Y.

The atom which has more E.N. has a greater ability to attract the e pair. On the other hand, the atom which has less E.N. has a lesser ability to attract the e -pair.

So E.N. is a relative property as its idea is not related to an individual atom. So, it has no unit. It is measured in Pawling’s scale.

In the HCI molecule, H-nucleus weakly attracts the e¯pair, which is stronger for Cl-nucleus.
So, the E.N. of Cl is more than the E.N. of H.

As a result of this, both H and Cl atoms become partially charged – Cl gets a slight – ve charge (5-) and H gets a slight +i/e charge (8+). The chemical bond between H and Cl is called a polar bond.

In the F² molecule, the F atoms equally attract the e¯pair. So no partial +ve/-ve charge appears on either atom of F. So F² molecule has a non-polar bond.

Trends of E.N. in the Periodic Table along a period: As we move from left to right, atomic size decreases, and nuclear charge increases means the ability to attract shared e pair also increases. So, _gF is more E.N. than ₅B and ₅B is more E.N. than ₁H.

Note: F is the most electronegative element of all the elements.

Metallic and Non-metallic character: Elements that have more electronegativity have a tendency to gain e¯s.

They have non-metallic On the other hand, elements that have less electronegativity have a tendency to lose e¯s.

They have a metallic character. So in the periodic table, along a period metallic character decreases, and down the group metallic character increases.

Factors on which metallic and non-metallic character of elements depend are:

1. Atomic size
2. Nuclear charge.

Oxidizing / Reducing Properties: Oxidation is the loss of e¯ s and Reduction is the gain of e¯s.

Metals have a tendency to lose e¯s known as reducing property. Non-metals have a tendency to gain e¯s_r known as oxidizing property.

Within a compound, metal atoms have a relatively low attraction for the e¯s because of their low I.E., and low electronegativity.

Elements present on the right of the periodic table are strong oxidizing agents. Elements present on the left side of the periodic table are strong reducing agents.

General trend: Along a period, as we move from left to a right tendency to gain e¯ increases, and the tendency to lose e¯ decreases.

This means that oxidizing property gradually increases and the reducing property gradually decreases.

That is, properties shift from metallic to non-metallic. For the same reason, alkali metals of Gr. 1 are good reducing agents.

Halogens of Gr. 17 have a greater tendency to gain e¯s. Again down a group, reducing property increases, and oxidizing property decrease.

WBBSE Chapter 7 Atomic Nucleus Radioactivity

Radioactivity Definition: Radioactivity is a nuclear phenomenon of spontaneous emission of α or β and γ-radiations from the nuclei of heavier atoms with atomic number more than 82 (i.e., after 82Pb in the periodic table in order to attain a stable state from an unstable state.)

1. In 1896, French scientist Henry Becquerel, while conducting experiments with newly discovered.
2. X-rays to investigate how uranium salts are affected by this light and quite accidentally.
3. He discovered that the uranium salts spontaneously emit highly penetrating radiations (similar to X-rays) from within themselves.
4. The radiations can be registered on a photographic plate. Shortly afterward, Madam Curie and her husband Pierre Curie found that ‘pitchblende’ (an ore of uranium) emits a similar type of radiation.
5. They named this phenomenon natural radioactivity and commented that it is a property of the nucleus.
6. The phenomenon is ‘spontaneous’ and any physical or chemical condition cannot change the emission of radiation.
7. The process of such radiation is known as radioactive decay or disintegration. A nucleus can undergo radioactive -decay through the emission of α or β and γ rays.
8. Examples of some radioactive elements are uranium, radium, polonium, thorium, actinium, etc.

Read And Learn Also WBBSE Notes For Class 10 Physical Science And Environment

WBBSE Chapter 7 Atomic Nucleus Nature α,β And γ-Rays

Rutherford and his co-workers discovered that the radioactive rays are a combination of α, β, and γ-radiations.

Comparison of a, p, and y-radiations with reference to some highlighting properties:

How to know the charge of radioactive radiation? Radioactive radiations are affected by an electric field. It is observed that α-particles bent slightly towards the -ve plate.

So they are positively charged heavier particles, β-particles bent more towards -i-ve plate.

So they are negatively charged lighter particles, and γ-radiations remain undeflected so they are uncharged.

Effect on atomic number and mass number due to α, β, and γ-emi-ssion. In any nuclear reaction, total atomic no. and total mass no. always remain the same.

1. Due to the emission of an a-particle, the atomic number decreases by 2 units and the mass number decreases by 4 units. That is,

For example, 

2. Due to the emission of a (3-particle, the atomic number increases by 1 unit, but the mass number remains the same. That is,

For example,

3. Due to y-radiation, no change in mass and charge occurs.

Note:

The S.l. unit of radioactivity is Becquerel (Bq). lBq = one disintegration per second or one event per second.

Bigger units: 1 megabecqural or MBq = 10⁶ Bq and 1 gigabecqueral or GBq = 10⁹ Bq The common unit for measuring the activity of radioactive substances is curie.

1 curie = 3.70 x 10¹⁰ disintegration per second or 3.70 x10¹⁰ event per second.

Smaller units: 1 milli-curie = 10^-3 curie and 1 micro-curie = 10^-6curie. A unit Rutherford is equal to 10⁶ disintegrations per second.

Smaller units: 1 milli-rutherford = 10³ disintegrations per second and 1 micro-rutherford = 10⁶ disintegrations per second.

Note: Radioactivity is considered a nuclear phenomenon (not a chemical phenomenon).

Inside the nucleus of an atom, the repulsive electrostatic force between +vely charged protons (ps) tends to push them away from each other.

For lighter elements whose Z V 20(like \(\left.{ }_2^4 \mathrm{He},{ }_8^{16} \mathrm{O},{ }_7^{14} \mathrm{~N},{ }_{11}^{23} \mathrm{Na},{ }_{12}^{24} \mathrm{Mg},{ }_{20}^{40} \mathrm{Ca}\right)\)

(like for which these nucleons are held together by a strong force (other than electrostatic force) called nuclear force.

This nuclear force cancels out the electrostatic force of repulsion between Ps and keeps the nucleus stable.

The nuclear force is charge-independent, stronger than electrostatic force (~ 100 times), and has a very very short range.

But when the no. of ns exceeds the no. of ps for atoms with \(\mathrm{Z}>20 \text { (like }{ }_{26}^{56} \mathrm{Fe},{ }_{47}^{108} \mathrm{Ag} \text { etc.) }\) for which the ratio → 1.

Then the strong nuclear force cannot hold the ps and ns together and the nucleus becomes unstable.

To attain stability, the unstable nuclei fling out different particles or some energy. This is the reason why some elements are radioactive.

It has been observed that the rate of decay of the radioactive elements remains unaffected:

By any physical change, such as a change in pressure, or temperature,

By any chemical change, such as excessive heating, cooling, oxidation, the action of strong electric and magnetic fields, etc.

3. How β-particles (fast-moving\({ }_{-1}^0 \text { es }\)) are emitted from the nucleus?

β particles are emitted from the nuclei which have too many ns and ps. In order to, maintain the conservation of electric charge, one of the ns Is transformed into ps and vice versa.

This is the reason for (β-decay from the nucleus. In the case of [β-decay, an n is changed into a p and an electron, accompanied by a charge less, massless particle antineutrino.

Although particles and cathode rays both are fast-moving electrons, they differ from one another.

Because; [β-particles are emitted from the nuclei of radioactive elements, whereas cathode rays originate from the orbital electrons.

4. After the emission of α or (γ-particles, the nuclei of radioactive elements acquire a state of excitation where they possess excess energy.

The excess energy is released in the form of electromagnetic radiation, known as the y-radiation so that the excited nuclei come to their ground state.

The y-radiation can be represented as follows:

Although X-rays and y-rays are similar types of electromagnetic radiation, they differ from one another in their origin.

X-rays are emitted due to the transition of electrons in their inner orbits, whereas y-radiations are emitted from the excited nuclei to come to their ground state.

Uses of radioactivity: Radioactivity is used in

(1) Medical field: In the treatment of thyroid, radioactive iodine\(\left({ }_{53}^{131} \mathrm{l}\right)\); in the treatment of blood cancer or leucemia or to detect brain tumor, radioactive phosphorus \(\left({ }_{15}^{32} \mathrm{P}\right)\) in the treatment of malignant cancer, radioactive cobalt \(\left({ }_{27}^{60} \mathrm{Co}\right)\); in the detection of blood circulation in heart, radioactive sodium \(\left({ }_{11}^{24} \mathrm{Co}\right)\) is used.

(2) To determine the age of dead microorganisms (plants, animals), radiocarbon \(\left({ }_{6}^{14} \mathrm{C}\right)\)

(3) The age of older rocks (dating) can be calculated using ²³⁸U, ⁴⁰K radioisotopes.

As rocks often contain traces of uranium, which eventually decays to lead).

(4) In agriculture, radioactive samples are used as tracers and also used in improving food production, improvement of soil fertility, and others.

WBBSE Chapter 7 Atomic Nucleus Mass Defect, Binding Energy, And Nuclear Fission

In the theory of relativity, Albert Einstein explained the equivalence of mass and energy by the equation: E = mc².

According to this law, when an amount of mass ‘m’-disappears, an equivalent amount of energy ‘E’ is released from the system to the surroundings.

Again, if the energy of a system is increased by the amount ‘E’ its mass will be increased by an amount \(m=\frac{E}{c^2}.\)

In the case of high-energy reactions such as nuclear fission and fusion, mass is converted into energy.

In general, the mass of a nucleus is slightly lower than the total mass of the nucleons (protons + neutrons) in the nucleus.

This difference in mass is termed the mass defect (Δm). Some mass disappears. during the formation of the nucleus (exothermic).

Binding energy is the amount of energy released when the nucleons bind together to form the nucleus of an atom.

Suppose that a nucleus has a rest massp\({ }_z^A X\).

If mp and mn are the mass of a proton and a neutron respectively, then the mass defect, \(\Delta m=\left[\mathrm{Z} \cdot m_p+(\mathrm{A}-\mathrm{Z}) m_n\right]-\mathrm{M}\)

Then according to Einstein’s mass-energy equivalence formula, the binding energy will be \(=\Delta m \cdot c^2=\left\{\left[Z \cdot m_p+\right.\right.\left.\left.(\mathrm{A}-\mathrm{Z}) m_n\right]-\mathrm{M}\right\} c^2\)

Nuclear fission: The process of splitting up the nucleus of a heavy atom after interacting with the projectile (mainly thermal neutron having energy ~ 0.025 eV) into two or lighter nuclei of comparable size with the release of energy is known as nuclear fission.

A typical example of a nuclear fission reaction is a single neutron \(\left({ }_0^1 n\right)\) in a uranium nucleus \(\left({ }_{92}^{235} \mathrm{U}\right)\) produces three neutrons accompanied by two comparatively lighter nuclei of barium and krypton.

The three neutrons produced in the reaction disintegrate three more \(\left({ }_{92}^{235} \mathrm{U}\right)\) nuclei. Then, nine neutrons are produced, which in turn can produce fission of nine more\(\left({ }_{92}^{235} \mathrm{U}\right)\) nuclei and so on.

Thus a chain reaction would start and within a few milliseconds, a tremendous amount of energy would be liberated. The daughter nuclei are also radioactive and they further emit [β-rays, and γ-rays.

Calculation of energy released: For nuclear fission reaction, all the masses are taken in a.m.u. or u (where 1 u= 1.66 x 10^-27Kg).

Total mass before reaction = 234.993 u + 1.00 87 u = 236.0022 u

Total mass after reaction = 140.8834 u + 91.9020 u + 3 x 1.0087 u = 235.8115 u

∴ Difference in mass (i.e., mass defect) = 236.0022 u – 235.8115 u = 0.1907 u

According to the equation E = me², 1 u mass is equivalent to 932-6 MeV energy.

Thus, one nucleus of ²³⁵U releases the energy = 0-1907 x 932-6 MeV = 177-85 MeV.

Can you imagine the fission of only 50 g of ²³⁵U? 3-55 x 10¹² J (very large amount) of energy is released  by which one 100 W bulb could run for about 1125 years

Atom bomb: The principle of nuclear fission (using U-235 or Pu-239) is used in the construction of an atom bomb.

In it, the fission reaction takes place in an uncontrolled manner so that a single fission causes more than fission and an enormous amount of energy (as I energy) is liberated within a few milliseconds. The how an atomic bomb is made.

During the final stage of World War II, two a-bombs were dropped on the Japanese cities Hiroshima (on Aug 6, 1945, at 8:11 am, named “Li Boy’) and Nagasaki (on Aug 9, 1945, at 11: 02 am, explosion and it’s named ‘Fat Man’).

The two bombings killed at least 129000 people, and destroyed a lot like burning effects, radiation sickness, genetic defects, loss of fertility of lands, and many others.

Nuclear reactors as peaceful use of nuclear energy: A nuclear reactor is a device in which a controlled nuclear chain reaction can be initiated to produce energy.

In a controlled nuclear reaction, a single fission causes one fission, and energy is liberated at a constant rate.

This energy is used at nuclear power plants for the generation of electricity. Today, about 450 or more nuclear reactors are used in about 30 nations around the world.

Physics Class 10 Wbbse

Nuclear accidents :

1. The Chornobyl accident was the worst nuclear power plant accident in history. It happened on 26 April 1986 in Ukraine (part of the former Soviet Union).

Operating the plant at very low power, the reactors became highly unstable, and hot (about 2000°C) radioactive fuel substances came out as a steam explosion.

It lifted the cover off the top of the reactor. A large area was contaminated within 36 hours of the accident.

The Fukushima accident (located in the northern part of Japan) occurred on 11 March 2011 following a massive earthquake (recorded magnitude 9-0) and tsunami (about 15 m high).

The plant’s emergency power generators in the basement were flooded, causing the leakage of radioactive substances into the environment.

The fuel rods were also melted with the leakage of deadly radiation. It is the worst nuclear disaster since the 1986 Chornobyl accident along with the release of a large amount of energy is called nuclear fusion.

The process of the formation of a heavy nucleus by the combination of two or lighter nuclei \(\left({ }_1^1 H,{ }_1^2 H,{ }_1^3 \mathrm{H},{ }_3^6\right. \text { Li, etc.) }\)

A typical example of a fusion reaction is:

⇒ \(4{ }_1^1 \mathrm{H} \longrightarrow{ }_2^4 \mathrm{He}+2{ }_{+1}^0 e+\text { Energy or, }{ }_1^2 \mathrm{H}+{ }_1^3 \mathrm{H} \longrightarrow{ }_2^4 \mathrm{He}+{ }_0^1 \mathrm{n}+\text { Energy }\)

The difference in mass between the parent nuclei and the daughter nuclei is released in the form of energy (approximately 17-6 MeV energy perfusion).

Physics Class 10 Wbbse

The nuclear fusion reactions take place at a very high temperature (about 107°C). Due to this fact, nuclear fusion is also called thermonuclear reaction.

Such a high temperature can result only from a nuclear fission reaction. But, once the nuclear fusion reaction starts, the energy release is so high enough to maintain a high temperature to fuse more nuclei and the reaction continues.

That’s why, we say nuclear fission initiates nuclear fusion. Mathematically, it can be proved that in fusion much more amount of energy is released than in fission reaction per unit mass of fuel.

Nuclear fusion is not a chain reaction.

Physics Class 10 Wbbse

Nuclear fusion takes place in the interior of the sun and other stars. Sun and other stars are composed of about 90% hydrogen and helium gas. Only 10% is other elements…

In a hydrogen bomb, the synthesis of a heavier atom from hydrogen results in the release of an enormous amount of energy.