Chapter - 3 Thermodynamics Notes | EduRev

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Class 12 : Chapter - 3 Thermodynamics Notes | EduRev

The document Chapter - 3 Thermodynamics Notes | EduRev is a part of the Class 12 Course Chemistry Class 12.
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Thermodynamic processes :

  • Isothermal process: T = constant
    dT = 0
    ΔT = 0
  • Isochoric process: V = constant
    dV = 0
    ΔV = 0
  • Isobaric process: P = constant
    dP = 0
    ΔP = 0
  • Adiabatic process: q = 0
    or heat exchange with the surrounding = O(zero)

IUPAC Sign convention about Heat and Work :
Work done on the system = Positive
Work done by the system = Negative 

1st Law of Thermodynamics
ΔU = (U- U1) = q + w

Law of equipartion of energy :
Chapter - 3 Thermodynamics Notes | EduRev (only for ideal gas)

Chapter - 3 Thermodynamics Notes | EduRev
where f= degrees of freedom for that gas. (Translational + Rotational)
f = 3 formonoatomic
= 5 for diatomic or linear polyatmic
= 6 for non - linear polyatmic

Calculation of heat (q): 
Total heat capacity:

Chapter - 3 Thermodynamics Notes | EduRev

Molar heat capacity :

Chapter - 3 Thermodynamics Notes | EduRev

Specific heat capacity (s):

Chapter - 3 Thermodynamics Notes | EduRev


WORK DONE (w) :
Isothermal Reversible expansion/compression of an ideal gas :

W = - nRT In (Vf/Vi)

Reversible and irreversible isochoric processes.
Since dV = 0
So dW = - Pext . dV = 0.

Reversible isobaric process:
W = P (Vf - Vi)
Adiabatic reversible expansion :
Chapter - 3 Thermodynamics Notes | EduRev

Reversible Work:
Chapter - 3 Thermodynamics Notes | EduRev

Irreversible Work :
Chapter - 3 Thermodynamics Notes | EduRev

Free expansion - Always going to be irrerversible and since Pext = 0
so dW = -Pext..dV = 0
If no. heat is supplied q = 0
then ΔE = 0 so ΔT = 0.

Application of 1st Law :
ΔU = ΔQ + ΔW ⇒ ΔW = -P ΔV
∴ ΔU = ΔQ -PΔV

Constant volume process
Heat given at constant volume = change in internal energy
∴du = (dq)v
du = nCvdT
Chapter - 3 Thermodynamics Notes | EduRev

Constant pressure process:
H ≡ Enthalpy (state function and extensive property)
H = U + PV

⇒ Cp - Cv = R (only for ideal gas)

Second Law Of Thermodynamics :
Chapter - 3 Thermodynamics Notes | EduRev fo r a spontaneous process.

Entropy (S):
Chapter - 3 Thermodynamics Notes | EduRev

Entropy calculation for an ideal gas undergoin a process :

Chapter - 3 Thermodynamics Notes | EduRev

Chapter - 3 Thermodynamics Notes | EduRev (only for an ideal gas)

Third Law Of Thermodynamics :
The entropy of perfect crystals of all pure elements & compounds is zero at the absolute zero of temperature.

Gibb’s free energy (G) : (State function and an extensive property)
Chapter - 3 Thermodynamics Notes | EduRev

Criteria of spontaneity:
(i) If ΔGsystem is (-ve) < 0 ⇒ process is spontaneous
(ii) If ΔGsystem is > 0 ⇒ process is non spontaneous
(iii) lf ΔGsystem  = 0 = 3 system is at equilibrium.

Physical interpretation of ΔG :
 The maximum amount of non-expansional (compression) work which can be performed.
Chapter - 3 Thermodynamics Notes | EduRev

Standard Free Energy Change (ΔG°) :
1. ΔG° = -2.303 RTIog10K
2. At equilibrium ΔG = 0.
3. The decrease in free energy (-ΔG) is given as :
Chapter - 3 Thermodynamics Notes | EduRev
4. Chapter - 3 Thermodynamics Notes | EduRevfor elemental state = 0

5. Chapter - 3 Thermodynamics Notes | EduRev

Thermochemistry:
Change in standard enthalpy Chapter - 3 Thermodynamics Notes | EduRev

= heat added at constant pressure. =  CPΔT.
If Hproducts > Hreactants

  • Reaction should be end other micas we have to give extra heat to reactants to get these converted into products and if
    Hproducts > Hreactants
  • Reaction will be exothermic as extra heat content of reactants will be released during the reaction. Enthalpy change of a reaction :
    Chapter - 3 Thermodynamics Notes | EduRev
    Chapter - 3 Thermodynamics Notes | EduRev
    = positive - endothermic
    = negative - exothermic

Temperature Dependence Of ΔH : (Kirchoff's equation):
For a constant volume reaction
Chapter - 3 Thermodynamics Notes | EduRev

where ΔCP = Cp (products) - Cp (reactants).
For a constant volume reaction
Chapter - 3 Thermodynamics Notes | EduRev

Enthalpy of Reaction from Enthalpies of Formation :
The enthalpy of reaction can be calculated by
Chapter - 3 Thermodynamics Notes | EduRev  is the stoichiometric coefficient.

Estimation of Enthalpy of a reaction from bond Enthalpies :
Chapter - 3 Thermodynamics Notes | EduRev

Resonance Energy:
Chapter - 3 Thermodynamics Notes | EduRev


CHEMICAL EQUILIBRIUM
At equilibrium :
(i) Rate of forward reaction = rate of backward reaction 

(ii) Concentration (mole/litre) of reactant and product becomes constant.

(iii) ΔG = 0.
(iv) Q = Keq

Equilibrium constant (K) :
Chapter - 3 Thermodynamics Notes | EduRev

Equilibrium constant in terms of concentration (Kc) :
Chapter - 3 Thermodynamics Notes | EduRev

Equilibrium constant in terms of partial pressure (Kp):
Chapter - 3 Thermodynamics Notes | EduRev

Equilibrium constant in terms of mole fraction (Kx) :
Chapter - 3 Thermodynamics Notes | EduRev

Relation between Kp & Kc :
Kp = Kc.(RT)Δn.

Relation between KP & KX :
KP = KX(P)Δn

Chapter - 3 Thermodynamics Notes | EduRevΔH = Enthalpy of reaction

Relation between equilibrium constant & standard free energy change :
ΔG° = - 2.303 RT log K 

Reaction Quotient (Q):
The values of expression Q = Chapter - 3 Thermodynamics Notes | EduRev

Degree of Dissociation (α):
α = no. of moles dissociated / initial no. of moles taken
= fraction of moles dissociated out of 1 mole.
Note: % dissociation = α x 100
Observed molecular weight and Observed Vapour Density of the mixture :
Observed molecular weight of An(g) =
 Chapter - 3 Thermodynamics Notes | EduRev
Chapter - 3 Thermodynamics Notes | EduRev

External factor affecting equilibrium :
Le Chatelier's Principle:
If a system at equilibrium is subjected to a disturbance or stress that changes any of the factors that determine the state of equilibrium, the system will react in such a way as to minimize the effect of the disturbance.

Effect of concentration:

  • If the concentration of reactant is increased at equilibrium then reaction shift in the forward direction .
  • If the concentration of product is increased then equilibrium shifts in the backward direction

Effect of volume:

  • If volume is increased pressure decreases hence reaction will shift in the direction in which pressure increases that is in the direction in which number of moles of gases increases and vice versa.
  • If volume is increased then, for
    Δn > 0 reaction will shift in the forward direction
    Δn < 0 reaction will shift in the backward direction
    Δn = 0 reaction will not shift.

Effect of pressure:
If pressure is increased at equilibrium then reaction will try to decrease the pressure, hence it will shift in the direction in which less no. of moles of gases are formed.

Effect of inert gas addition :
(i) Constant pressure: If inert gas is added then to maintain the pressure constant, volume is increased. Hence equilibrium will shift in the direction in which larger no. of moles of gas is formed
Δn > 0 reaction will shift in the forward direction
Δn < 0 reaction will shift in the backward direction
Δn = 0 reaction will not shift.
(ii) Constant volume: Inert gas addition has no effect at constant volume.

Effect of Temperature:
Equilibrium constant is only dependent upon the temperature.
If plot of lnk vs 1/T is plotted then it is a straight line with slope = Chapter - 3 Thermodynamics Notes | EduRev and intercept = 

Chapter - 3 Thermodynamics Notes | EduRev

  • For endothermic (ΔH > 0) reaction value of the equilibrium constant increases with the rise in temperature 
  • For exothermic (ΔH < 0) reaction, value of the equilibrium constant decreases with increase in temperature 
  • For ΔH > 0, reaction shifts in the forward direction with increase in temperatutre 
  • For ΔH < 0, reaction shifts in the backward direction with increases in temperature. 
  • If the concentration of reactant is increased at equilibrium then reaction shift in the forward direction . 
  • If the concentration of product is increased then equilibrium shifts in the backward direction

Vapour Pressure of Liquid
Relative Humidity = Chapter - 3 Thermodynamics Notes | EduRev

Thermodynamics of Equilibrium :
ΔG = ΔG°+ 2.303 RT log10Q
Vant Hoff equation- Chapter - 3 Thermodynamics Notes | EduRev


IONIC EQUILIBRIUM
OSTWALD DILUTION LAW:

  • Dissociation constant of weak acid (Ka), Chapter - 3 Thermodynamics Notes | EduRev If α « 1 , then 1 - α ≌ 1 or Ka = cα2 or α = Chapter - 3 Thermodynamics Notes | EduRev
  • Similarly for a weak base,Chapter - 3 Thermodynamics Notes | EduRev Higher the value of Ka/ Kb, strong is the acid/base.

Acidity and pH scale :
∴ pH = - log aH+  (where aH+ is the activity of H+ ions = molar concentration for dilute solution).
[Note : pH can also be negative or > 14]
Chapter - 3 Thermodynamics Notes | EduRev


PROPERTIES OF WATER :

  1. In pure water [H+] = [OH-] so it is Neutral.
  2. Moler concentration / Molarity of water = 55.56 M.
  3. Ionic product of water (Kw) :
    Kw= [H+][OH-] = 10-14 at 25° (experimentally)
    pH = 7 = pOH ⇒ neutral
    pH < 7 or pOH >7 ⇒ acidic
    pH>7or pOH<7 ⇒ Basic
  4. Degree of dissociation of water:
    Chapter - 3 Thermodynamics Notes | EduRev
  5. Absolute dissociation constant of water :
    Chapter - 3 Thermodynamics Notes | EduRev
    pK= pKb = - log (1.8 x 10-16) = 16 - log 1.8 = 15.74
    Ka x Kb = [H+] [OH-] = Kw
    ⇒ Note: for a conjugate acid- base pairs
    pKa + pKb = pKw = 14 at 25°C.
    pKa of H3O+ ions = -1.74
    pKb of OH- ions = -1.74.

pH Calculations of Different Types of Solutions :

  • Strong acid solution :
    (i) If concentration is greater than 10-6M
    In this case H+ ions coming from water can be neglected,
    (ii) If concentration is less than 10-6 M
    In this case H+ ions coming from water cannot be neglected
  • Strong base solution :
    Using similar method as in part (a) calculate first [OH-] and then use [H+] x [OH-] = 10-14
  • pH of mixture of two strong acids :
    Number of H+ ions from I-solution = N1 V1
    Number of H+ ions from II-solution = N2 V2
    Chapter - 3 Thermodynamics Notes | EduRev
  • pH of mixture of two strong bases :
    Chapter - 3 Thermodynamics Notes | EduRev
  • pH of mixture of a strong acid and a strong base :
    If N1V1 > N2V2, then solution will be acidic in nature and [H+] = N = Chapter - 3 Thermodynamics Notes | EduRev
    If N2V2 > N1V1, then solution will be basic in nature and [OH-] = N = Chapter - 3 Thermodynamics Notes | EduRev
  • pH of a weak acid(monoprotic) solution :
    Chapter - 3 Thermodynamics Notes | EduRev
    if a << 1 ⇒ (1 - α) = 1 ⇒ Ka = Cα2Chapter - 3 Thermodynamics Notes | EduRev (is valid if α < 0.1 or 10%)
    On increasing the dilution ⇒ Chapter - 3 Thermodynamics Notes | EduRev Chapter - 3 Thermodynamics Notes | EduRev
    (g) pH of a solution of a polyprotic weak acid :
    Chapter - 3 Thermodynamics Notes | EduRev
    RELATIVE STRENGTH OF TWO ACIDS :
    Chapter - 3 Thermodynamics Notes | EduRev
  • pH of a mixture of two weak acid(both monoprotic) solutions :
    (α1« 1) and (α2 « 1) ⇒ Chapter - 3 Thermodynamics Notes | EduRev
    Chapter - 3 Thermodynamics Notes | EduRev
     If water is again considered third weak acid in solution of two weak acid then
    Chapter - 3 Thermodynamics Notes | EduRev
  • pH of a mixture of weak acid(monoprotic) and a strong acid solution :
    If [SA] = C1 and [WA] =C2, then [H+] from SA = C1 and [H+] from WA = C2 Let HA is a weak acid.
    Chapter - 3 Thermodynamics Notes | EduRev
    If a strong acid of low cone is added in water then [H+] of solution can be calculated as
    Chapter - 3 Thermodynamics Notes | EduRev

SALT HYDROLYSIS :

Chapter - 3 Thermodynamics Notes | EduRev

Hydrolysis of ployvalent anions or cations
For [Na3PO4] = C.
Ka1 x Kh3 = Kw
Ka1 x Kh2 = Kw
Ka3 x Kh1 = Kw
Generally pH is calculated only using the first step Hydrolysis
Chapter - 3 Thermodynamics Notes | EduRev
Chapter - 3 Thermodynamics Notes | EduRev
So Chapter - 3 Thermodynamics Notes | EduRev
Hydrolysis of Amphiprotic Anion. (Cation is not Hydrolysed e.g. NaHCO3, NaHS, etc.)
Chapter - 3 Thermodynamics Notes | EduRev
Similarly for Chapter - 3 Thermodynamics Notes | EduRev  amphiprotic anions.
Chapter - 3 Thermodynamics Notes | EduRev and Chapter - 3 Thermodynamics Notes | EduRev

  • Chapter - 3 Thermodynamics Notes | EduRev

The

Chapter - 3 Thermodynamics Notes | EduRev

BUFFER SOLUTION :
(a) Acidic Buffer: e.g. CH3COOH and CH3COONa. (weak acid and salt of its conjugate base).
Chapter - 3 Thermodynamics Notes | EduRev [Henderson's equation]
(b) Basic Buffer: e.g. NH4OH + NH4CI. (weak base and salt of its conjugate acid).
Chapter - 3 Thermodynamics Notes | EduRev

Buffer capacity (index) :
Buffer capacity Chapter - 3 Thermodynamics Notes | EduRev
Buffer capacity = Chapter - 3 Thermodynamics Notes | EduRev


INDICATOR :
Chapter - 3 Thermodynamics Notes | EduRev
Chapter - 3 Thermodynamics Notes | EduRev ⇒ Chapter - 3 Thermodynamics Notes | EduRev


SIGNIFICANCE OF INDICATORS :
Extent of reaction of different bases with acid (HCI) using two indicators :

Chapter - 3 Thermodynamics Notes | EduRev

ISOELECTRIC POINT:
Chapter - 3 Thermodynamics Notes | EduRev

SOLUBILITY PRODUCT :
Ksp = (xs)x (ys)y = xx.yy.(s)x+y

CONDITION FOR PRECIPITATION :
If ionic product KI.P > Ksp precipitation occurs,
if KI.P = Ksp saturated solution (precipitation just begins or is just prevented).

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