THERMODYANAMICS class 11th chemistry

1. THERMODYANAMICS
The branch of science which deals with the quantitative
relationship between heat and other forms of energies is
called thermodynamics.

2. Internal Energy (E or U):
 It is the total energy stored within the substance. It is the sum of many
types of energies like vibrational energy, translational energy. etc.
 It is an extensive property and a state function.
 Its absolute value cannot be determined but experimentally change in
internal energy (ΔU) can be determined by
ΔU = U2 – U1
or ΔU = Up– UR
Sign Convention:
a) ΔU = +ve , if energy is absorbed by the system (U1>U2)
b) ΔU = -ve , if energy is evolved (given) by the system (U1<U2)
Units: Joule, ergs
(1 Joule = 107 ergs)

3. Work (W) :
Work is said to be done when an object is displaced in he direction of force
on application of force.( mechanical work)
Work = Force X Displacement
Pressure – Volume work (or Work of expansion/ contraction)
 If the system involves gaseous substances and there is a difference of
pressure between system and surroundings. Work is referred as pressure –
volume work (W=pV).
 It is the work done when the gas expands or contracts against the
external pressure P. Work of expansion = - Pext . ΔV
Work of compression = Pext . ΔV
Sign Convention: a) Work done by the system = -ve
b) Work done on the system = +ve
Units: Joule, ergs

4. i) Work of Irreversible expansion against constant pressure B under
isothermal conditions
W = – Pext. ΔV
(ii) Work of reversible isothermal expansion of an
ideal gas:

5. (iii) Free expansion of an ideal gas( Expansion against vacuum):
if an ideal gas expands against vacuum, the irreversible expansion is called free expansion.
as Pext = 0,
Thus for reversible as well as irreversible expansion
Work done = 0

6. Heat (q):
Another mode of energy exchanged between the system and surroundings, as a
result of temperature difference between them.
Sign Convention:
q is +ve, if heat is absorbed by the system.
q is –ve, if heat is given out ( evolved ) by the system.
Units: Joule, Calories
1 J = 0.239cal, 1 cal = 4.184 J
Both work and heat are path function.

7. First Law of Thermodynamics
(or Law of Conservation of energy):
 Energy can neither be created nor destroyed although
it can be converted from one form to the other.
Mathematically, ΔU = q + W
where, ΔU = internal energy change
q = heat added to system
W = work added to system
Sign convention:
(i) q is + ve = heat is supplied to the system
(ii) q is – ve = heat is lost by the system
(iii) W is + ve = work done on the system
(iv) W is – ve =work done by the system

8. Alternative statements:
 The total energy of the universe remains constant, although it may undergo
transformation from one form to the other.
 The energy of an isolated system is constant.
 Whenever a certain amount of some form of energy disappears, an exactly
equivalent amount of some other form of energy must be produced.
 It is impossible to prepare a perpetual motion machine i.e. a machine which
would produce work without consuming energy.

9. Important Conclusions:
Mathematical statement of first law of thermodynamics
ΔU = q + W
a) if work is the work of expansion, W = - PΔV
Then ΔU = q – PΔV
q = ΔU + PΔV
b) for an isothermal process, ΔU = 0, then q = -W.
heat absorbed by the system is equal to the work done by the system.
c) For adiabatic process, q = 0, then ΔU = Wad

10. Enthalpy (H)
 It is the heat content of the system.
 It is the sum of internal energy and pressure volume work done of the system.
Mathematically, H = U + PV
 It is a state function and extensive property.
 Like U, absolute value of H also cannot be known, ΔH is determined
experimentally. ΔH = H2 – H1
 If H1 = enthalpy in initial state, H2 = enthalpy in final state
then H1 = U1 + PV1 and H2 = U2 + PV2
H2 – H1 = (U2 – U1) + P(V2 – V1)
ΔH = ΔU + P ΔV
 For exothermic reaction (the reaction in which heat is evolved), ΔH = -ve
 For endothermic reaction (the reaction in which heat is absorbed), ΔH = +ve.

11. Enthalpy change:
Acc. To first law of thermodynamics
Δ U = q + W
q = ΔU – W
At constant pressure, W = - P ΔV and heat absorbed is represented by qp
qp = ΔU + P ΔV
or, qp = ΔH ( ΔH = ΔU + P ΔV)
Thus, enthalpy change of the system is equal to the heat absorbed or evolved by
the system at constant pressure.

12. Internal Energy change:
Acc. To first law of thermodynamics
Δ U = q + W
q = ΔU – W
q = ΔU + P ΔV (since w = - P ΔV )
At constant volume, ΔV = 0 and heat absorbed is represented by qv
qv = ΔU
Thus, internal energy change of the system is equal to the heat absorbed or
evolved by the system at constant volume.

13. Relationship between ΔH and ΔU
 ΔH = ΔU + P ΔV
 or ΔH = ΔU + Δn(g) RT
 Here, Δn(g) = change in the number of gas moles.
 Δn(g) = n2 – n1, n1 and n2 are no. of moles of gaseous reactants and
products

14. Limitations of first law of thermodynamics:
 It indicates only that for any process to occur, there is an exact equivalence
between the various forms of energy. But it provides no information
concerning the spontaneity or feasibility of the process i.e. whether the
process is possible or not.
 first law doesnot indicate whether heat can flow from a cold end to a hot
end or not. It ony said if that process occur, heat gained by one end would be
exactly equal to heat lost by other end.
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15. Thermochemical equation:
 A balanced chemical equation which also indicates the amount of heat evolved or
absorbed is called a thermochemical equation.
For example, formation of water is written as
H₂ (g) + ½ O₂(g) H₂O (l) + 285.8 kJ mol-1
or H₂ (g) + ½ O₂(g) H₂O (l) , ΔH = 285.8 kJ mol-1
Thus, 285.8 kJ of heat is produced when 1 mole of hydrogen reacts with 0.5 mole of
oxygen. If the quantities of reactants are doubled, the heat produced will also be doubled.

16. Heat of Reaction (or Enthalpy of reaction):
 The amount of heat evolved or absorbed in a chemical reaction when the number of
moles of the reactants as represented by the chemical equation have completely
reacted, is called the heat of reaction or enthalpy of reaction or enthalpy change.
 It is represented by ΔrH.
 Eg. CH4(g) + 2O2(g) CO2 (g) + 2H2O (l), ΔH = -890.4 kJ mol-1
C(s) + H2O(g) CO(g) + H2(g), ΔH = +131.4 kJ mol-1
ΔrH = Sum of enthalpies of products -- Sum of enthalpies of reactant
= ∑ a iH(products) - ∑ b iH(reactants)
where a I and b I are coefficient of products and reactants.

17. Standard Enthalpy of reaction):
 The standard enthalpy of a reaction is the enthalpy change accompanying the reaction
when all the reactants and products are taken in their standard states.
 A substance is said to be in the standard state when it is in the purest and most stable
form at 1 bar pressure and the specified temperature. The temperature is usually
taken as 25°C(298 K) ,though not always.
 For example, standard state of pure ethanol at 298 K is pure liquid ethanol at 298 K
and 1 bar pressure whereas standard state of iron at 500 K is pure iron at 500 K and I
bar pressure.
 Standard enthalpy of a reaction is represented by ΔrH0

18. Different Types of Heat/ Enthalpies of
reactions:
 Depending upon the nature of the reaction
o Enthalpy of combustion
o Enthalpy of neutralization
o Enthalpy of formation
o Enthalpy of solution
o Enthalpy of hydration etc.
 Depending upon the type of process involving a phase change
o Enthalpy of fusion
o Enthalpy of vaporisation
o Enthalpy of sublimation etc.

19. Enthalpy of combustion (ΔcH):
 It is defined as the heat change (usually the heat evolved) when I mole of substance is
completely burnt or oxidized in oxygen.
 e.g... CH4
(g) + 20₂ (g) CO₂ (g) + 2 H₂O (g), ΔcH = -890.4 kJ mol-1
 It is exothermic in nature.
Standard Enthalpy of combustion
It is the amount of heat evolved when one mole of the substance is completely
burnt under standard condition(298K, 1bar).
Standard enthalpy of combustion is represented by ΔcH0
eg. ΔcH0
value for combustion of butane C4H10 (a constituent of cooking gas)
C4
H10
(g) + 13/2 0₂ (g) 4CO₂ (g) + 5H₂O (g), ΔcH0
= – 2658kj mol-1

20. Enthalpy of formation(ΔfH :
 It is defined as the heat change, i.e., heat evolved or absorbed when I mole of the
substance is formed from its elements under given conditions of temperature and
pressure.
 It is usually represented by ΔfH.
 The enthalpy of formation under standard conditions(298K , 1bar) is called standard
enthalpy of formation. Represented by ΔfH0
 C(s) + 0₂ (g) CO₂ (g) , ΔcH = -393.5 kJ mol-1

21. Imp. Points to note:
 For elementary substances in the standard state, the standard enthalpy of formation taken as
zero.
 Eg. For O2(g), Br2(g), I2(g), C(s): their ΔfH0
= 0.
 Standard enthalpy of a reaction can be calculated by the formula
ΔrH0
=[Sum of standard enthalpy of products] - [Sum of standard enthalpy of reactants]