Explore topic-wise InterviewSolutions in .

Surroundings:- They represents the rest of the universe which surrounds the system.

The system and surroundings are separated by a real or imaginary boundary.

System:- It is the specific part of the universe in which energy changes are taking place. The size of the system may vary from very small to very large.

Heat capacity of a system is defined as quantity of heat required to raise temperature by 1°.

(a) +ve (because gases are highly random than liquid)

ΔS = S(g) - S(e) = +ve

(b) +ve (because ions are in more random state in aqueous solution as compared to solid state.)

ΔS = s(aq) - S(s) = +ve

(c) -ve (because gases are highly more random than solid)

ΔS = S(s) - S(g) = -ve

(d) Zero

Diamond and graphite both are solid and in crystalline state.

Heat of formation of a compound is the change in enthalpy produced when one mole of the compound is formed from its elements denoted as ΔH_f.

Enthalpy of formation is defined as enthalpy change when 1 mole of the substance is formed from the constituting elements in their standard states e.g.

H₂(g) + 1/2O₂(g) → H₂O(l); Δ_fH = -286 kJ/mol

Enthalpy of reactions is defined as enthalpy change when reactants react completely to form products according to balanced chemical reaction,

e.g. N₂(g) + 3H₂(g) → 2NH₃(g) 

ΔH = -92.0 kJ

Thermodynamics is a branch of chemistry that deals with the study of interconversion of with other forms energy during physical and chemical changes.

Heat of a reaction is the change in enthalpy produced when the number of moles of the reactants as represented in the balanced chemical equation have completely reacted with each other.

There are three states of matter. The three  states of matter are inter-convertible. A solid on heating gets converted into a liquid and a liquid into a gas.

Each physical form of any substance is a different phase, whenever one phase changes into another, certain amount of heat is either evolved or absorbed. Most commonly observed phase transitions are

Solid → Liquid Melting (Fusion)

Liquid → Vapour Vaporisation

Solid → Vapour Sublimation

No, the heats released in the two reactions are not equal. The heat released in any reaction depends upon the reactants, products and their physical states. Here in reaction (i), the water produced is in the gaseous state whereas in reaction (ii) liquid is formed. As we know, that when water vapors condensed to from water, heat equal to the latent heat of vaporization is released. Thus, more heat is released in reaction (ii).

The heat of solution is defined as “the change in enthalpy of the system when one mole of a substance is dissolved in a specified quantity of solvent at a given temperature”.

When one mole of the solid melts at its melting point reversibly the heat absorbed is called molar heat of fusion. The entropy change is given by 

∆S = ΔH_f / T_f

Where ∆H_f = molar heat of fusion, T_f is melting point.

The total heat content of a system is called enthalpy.

ΔH = ΔE + Δn_gRT

ΔH = -93 x 10³ J

T = 300 K

R = 8.314 Jk^-1 mol^-1

Δn_g = 2 - 4 = -2

Now ΔE = ΔH - Δn_gRT

= -93 x 10³ - (-2) x 8.314 x 300

= -93000 x +4988.4

= -88011.6 J = -88.0 kJ

When Δn = 0, ΔH = ΔE

The heat of transition is defined as the change in enthalpy when one mole of an element changes from one allotropic form to another.

The molar heat of fusion is defined as “the change in enthalpy when one mole of a solid substance is converted into the liquid state at its melting points’.

ΔH = ΔE + RTΔn, ΔH = Enthalpy change, ΔE = Internal energy change Δn = Number of gaseous product number of gaseous reactant or number of moles.

A chemical reaction involves the breaking of bonds in reactants and formation of new bonds in products. The heat of reaction (enthalpy change) depends on the values of the heat needed to break the bond formation .Thus (Heat of reaction = (Heat needed to break the bonds in reactants – Heat liberated to from bonds in products).

ΔH^O = Bond energy in (to break the bonds) X Bond energy out (to form the bonds) = Bond energy of reactants – Bond energy of products.

1. Sublimation is a process when a solid changes directly into gaseous state without changing into liquid state. 

2. Molar heat of sublimation is defined as the change in enthalpy when one mole of a solid is directly converted into the gaseous state at its sublimation temperature.

∆G = ∆H – T∆S 

When the reaction is carried out at 0°K 

or ∆S = 0 

∆G = ∆H 

H₂O_(s) ⇌ H₂O_(I) 

∆_fusH = 6 kJ mol^-1 = 6000 Jmol^-1 

∆_fusH = 6 kJ mol^-1 

6000 J mol^-1

A chemical reaction involves the breaking of bonds in reactants and formation of new bonds in products. The heat of reaction (enthalpy change) depends on the values of the heat needed to break the bonds and heat liberated in the bond formation. Thus (Heat of reaction = Heat needed to break the bonds in reactant - Heat liberated to form bonds in products)

ΔH° = Bond energy in (to break the bonds) x Bond energy out (to form the bonds)

= Bond energy of reactants - Bond energy of products.

The enthalpy (H) and energy (E) are related by H = E + pV. For a change in the state of the system, we have

Initial state H₁ = E₁ + p₁V₁

Final state H₂ = E₂ + p₂V₂

Enthalpy change = ΔH = Final enthalpy - Initial enthalpy

= H₂ - H₁

= (E₂ + p₂V₂) - (E₁ + p₁V₁)

= (E₂ - E₁) + (p₂V₂ - p₁V₁)

or, ΔH = ΔE + (p₂V₂ - p₁V₁)

When there is no change in the pressure p₁ = p₂ and when there is no change in the volume V₁ = V₂. Therefore, p₁V₁ = p₂V₂ and ΔH = ΔE. That is enthalpy change (ΔH) is equal to the energy change when a chemical reaction takes place at constant pressure and at constant volume.

Characteristics of work:

1. Work is defined as the force (F) multiplied by the displacemcnt (x). 

– w = F.x   … (1)

The – ve sign is introduced to indicate that the work has been done by the system by spending a part of its internal energy.

2. Work is a path function. 

3. Work appears only at the boundary of the system. 

4. Work appears during the change ¡n the state of the system. 

5. Work brings a permanent effect in the surroundings.

6. Units of work: The SI unit of work is the joule (J) or Kilojoule (KJ).

7. If work done by the system, the energy of the system decreases, hence by convention work is taken to be negative (- w).

8. If work done by the system, the energy of the system increases, hence by convention work is taken to be positive (+ w).

The process in which the system and surroundings can be restored to the initial state from the final state without producing any changes in the thermodynamics properties of the Universe is called a reversible process. 

Example, H₂ + I₂ ⇌ 2HI

(i) The standard enthalpy of formation of carbon dioxide is -393.5 kilojoule per mole of CO₂. That is, ΔH°_f (CO₂, g) = -393.5 kJ mol^-1

(ii) The standard enthalpy of combustion of carbon is -393.5 kilo joule per mole of carbon. That is

ΔH°_comb (C) = -393.5 kJ mol^-1

First law of thermodynamics:

• The total energy of an isolated system remains constant though it may change from one form to another.
• Whenever energy of a particular type disappears equivalent amount of another type must be produced.
• Total energy of a system and surroundings remains constant.
• Energy can neither be created nor destroyed, but may be converted from one form to another.
• The change in the internal energy of a closed system is equal to the energy that passes through its boundary as heat or work.
• Heat and work are equivalent ways of changing a systems internal energy.

According to the first law of thermodynamics,

dq = dU + PdV

Dividing both sides by dT, we have

dq / dT = (du + PdV) / dT

At constant volume dV = 0, then

dq / dT = (dU / dT) V

C_v = (dU / dT) v

Thus the heat capacity ai constant volume (C_v) is defined as the rate of change of internal energy with respect to temperature at constant volume.

Mathematical statement of the First law of Thermodynamics is 

∆U = q + w

Case 1: 

For a cyclic process involving isothermal expansion of an ideal gas

∆U = 0 

∴ q = -w

In other words, during a cyclic process, the amount of heat absorbed by the system is equal to work done by the system.

Case 2:

For an isochoric process (no change in volume) there is no work of expansion.

∆V = 0 

w = 0 

∆U = 0

In other words, during isochoric process, the amount of heat supplied to the system is converted to its internal energy.

Case 3:

For an adiabatic process there is no change in heat .

i.e., q O. 

Hence, q = 0 

∆U = w 

In other words, in an adiabatic process, the decrease in internal energy is exactly equal to the work done by the system on its surroundings.

Case 4:

For an isobaric process. There is no change in the pressure. P remains constant. 

Hence, ∆U = q + w

∆U = q – P∆V 

In other words, in an isobaric process a part of heat absorbed by the system is used for PV expansion work and the remaining is added to the internal energy of the system.

Amount of heat required to vaporise one mole of a liquid at constant pressure and temperature is called its molar enthalpy of vaporization.

Energy required q = C_m x n x Δ T

= 24(J/mol/K) X 2.22(mol) X 20 (K )

[Δ T = 55°C-35°C =20 °C=20 K]

= 1066.8 J 

or 1.07 KJ

ΔS is positive.

(a) 300 K 

∆G = ∆H – T∆S 

At 300 K.

∆G = 30000 J mol^-1 – 300 J mol^-1 x 100 J K mol^-1

∆G = 0 

above 300 K; 

∆G will be negative and reaction becomes spontaneous.

(d) CaCO_3(s) → CaO_(s) + CO_2(g)

In CaCO_3(s) → CaO_(s) + CO_2(g) entropy change is positive in (a) and (b) entropy change is negative; in (c) entropy change is zero.