Class 11 Thermodynamics Short notes

  March 12, 2023    

 

Thermodynamics

 

  # It gives physicality or spontaneity of chemical reaction.

 

#  Reversible process  :-

    1) negligible difference between two opposing force

    2)  Quassi States 

    3)  gradual change

    4) always remain almost in equilibrium state.

 

#  Irreversible process :-

      1) large difference between two opposing force

      2)  No Quassi states

      3) sudden change

 

            The universe = the system +  the surroundings  

 

 

Work done :-

  Wirr = - Pext  x  ΔV 

  Wrev = - 2.303 nRt log Vf / vi 

 

Wrev > Wirr        ( expansion ) 

Wrev < Wirr        ( compression )

 

Law of thermodynamics :- 

1st law : Energy Conservation 

2nd law : Sponteinity 

3rd Law  : For a perfectly crystalline substance , entropy is zero at K 

4th or Zereth Law :  Heat transfer occures from higher to lower temperature.

 

Internal Energy - ( E or u )

All possible energies operating in the system 

 

# Gravitational E is not considered in Int. E 

# Absolute value of Int.E cannot be calculated , Thus changes in E is Calculated .

 

ΔE = q + w 

 

ΔE = q - Pext . ΔV                  (irev) 

ΔE = q - 2.303 nRt log Vf/Vi     ( rev) 

 

Isothermal :  ( ΔT=0)

ΔE=0 (Internal energy is thhe function of temprature)

q= -wà w done by the system

where q is amount of q given

 

Isobaric: (ΔP=0)

ΔE= q-pext. ΔV

 

Isochoric:( ΔV=0)

ΔE=qv àheat transferred at constant volume is called ΔE.

ΔEf-ΔEi=qvà q amt given to the system

Tf>Ti à T increase

 

Adiabatic : ( q = 0 )

Δ E = W

E2 – E1  = W à W amt of word is done

E2 > E1  

T2 > T1

# Under Vacuum / Free Expansion à W.D is zero

ΔH = qp à Heat Transfer at Const Pressure

 

SPONTINETY : 1) Enthalpy (H)

                              2) Entropy (S)

                              3) Gibbs Free (G)

 

     Enthalpy (H) –

 

ΔH = ΔE + Δ(p.v) à ΔH = ΔE + (p2v2 – p1v1)

 

ΔH = ΔE + PΔV +ΔP.V

 

At constant p , (Δp=0)

 

 

ΔH=ΔE+PΔV

Gaseous State , constant T&P , V is changing , n will be changing

 ΔH = ΔE+Δn(g) Rt

 

 

# Heat Capacity : ( C )

 

Amt of Heat required à Temp  by 1* C

Q = CT

 

# Special Heat Capacity : ( Cs )

Amt of heat required à Temp  of 1g substance by 1*C

Cs = C/mass

Q = mass × Cs  × ΔT

 

 

# Molar Heat Capacity : ( CM )

Amt of heat required à Temp  of 1 Mole substance by 1*C

C m C/mole             q = CΔT

q = mole * CH  * ΔT

 

At constant P ,

  ΔH = qp = nCpΔT

 

At constant v ,

 ΔH = nCvΔT = qv

 

# Gibb’s Free Energy (G) –

 

ΔG = ΔH – Δ(Ts)

ΔG = ΔH – TΔS -ΔTs

 

At constant Temp, (ΔT=0)

                ΔG = ΔH – TΔS

 

If ΔG = -ve à spontaneous

If ΔG = +ve à Non-spontaneous

If ΔG = 0 à Equilibrium

 

ΔG = ΔG⸰ +RT ln 

 

At Equilibrium à ΔG = 0 & Q = K à Equi. Constant

     ΔG⸰ = -RT lnK

 

·       ΔG⸰ ΔH⸰ - TΔS⸰

             Δ G⸰ =  -RT lnK

            ΔH⸰ - TΔS⸰ = -RT lnK

           ΔH⸰/RT - ΔS⸰/R  =   lnK

 

 

Ln K = ΔH⸰/R * 1/T +ΔS⸰/R

 

  Y =  m x + c



ΔG = ΔP.N – ΔT.S

 

At  constant T à ΔG = Δ.P.V

At Constant p à Δ G = -ΔT.S

 

   S = -ΔG/T

 

ΔG = ΔH – TΔS

ΔG = ΔH +TΔ( SG/ST)P

 

 

 

# Thermo-Chemistry :

 

1)     STANDARD ENTHAPLY OF FORMATION : ΔH⸰f   or ΔfH⸰

 

Amt of Energy released à 1 mole of substance formed from element in their standard state ( Most stable state )

·       Standard state : C Graphite , P4 (White) , S8 (orthorhombic) , H+ (aq) , O2 (g)

I2 (8) , Br2 (l) , Cl2 (g).

 

è ΔH⸰f for species in std. State is ‘Zero’

 

2)     STANDARD ENTHALPY OF COMBUSTION: ΔH⸰C or ΔCH⸰

è  Amt of E released à 1 mole subs.

 

 

3)     STANDARD ENTHALPY OF ATOMISATION :  ΔH⸰ atomization

 

E Changes occurs when à All covalent bond broken from 1 mole of covalent comp.

 

 

4)     STANDARD BOND ENTHALPY : ΔH⸰ Bond

 

E change à 1 mole of covalent bond is broken

      ΔHatom = n Δ H⸰bond

 

5)     STANDARD ENTHALPY OF NEUTRALISATION : Δ H⸰N

 

E change à 1 equivalent of H+ is completely neutralized by 1 1 equi of OH-

 

·       Δ Hn = 57.1 / mole  or 13.6 K cal/mole

·       NaoH + HF à Naf + H2O   ; Δ H⸰ > 57.1k ; High Hydration E of F- ion

 

 

6)     STANDRAD LATTIC ENTHALPY : Δ H⸰ Lattic

 

·       E Change à 1 mole of ionic bond broken into constituent ions in Gaseous State

 

 

7)     STANDARD ENTHALPY OF SOLUTION : Δ H⸰ Solution

 

·       Δ H⸰ solution = Δ H⸰ Hydration + Δ H⸰ Lattic

 

8)     STANDARD ENTHALPY OF FUSION : Δ H⸰ Fusion

 

·       E Change à 1 mole solid converted into 1 mole liquid at const. T&P

 

 

9)     STANDARD ENTHALPY OF SUBLIMATION : Δ H⸰ sublimation

 

·       E. change à 1 mole solid converted into 1 mole gas at constant  T&P

 

 

10)   STANDARD ENTHALPY OF VAPOURIZATION : Δ H⸰ vaporization

 

·       E. change à 1 mole liquid converted into 1 mole gas at const. T&P

 

11)  STANDARD ENTHALPY OF REACTION : Δ H⸰r

 

·       E. change à in any reaction

Δ H⸰r = ∑ Δ HR

 

 

 

LAWS

 

1)     Laplace lavoiser law :

 

              à   Exothermic        , vice versa     

 Endothermic

 

2)     Hess’ law of Summation :

 

i)              Δ H1            ΔH3 = ΔH1 + ΔH2

ii)             Δ H2                

iii)            Δ H3

 

 

3 ) Kirrchoff’s Equation :

·       Δ H2 - Δ H1 = Δcp ( ΔT2 – ΔT1)

 

 

4)     Clausius – Cleperyon Equation :

·       Log p2/p1 = ΔH/2.303R ( 1/T1  -  1/T2 )

 

5)     Troton’s Rule :

 

Δ H⸰ Vapourisation/TB  = 88J/K mole  

 

è Not application to liquids which shows strong intermolecular forces such as H – bonds

 

 

 

6)     Calorific Value :

 

 Δ H⸰ Combustion / Mol.wt

 

 

# Bomb Calorimeter –

     q a ΔT

    q = c ΔT

 

ΔE = qv  = C ΔT ( const. v)

 

# W.D in Rev. Adiabatic Process –

Wrev =  n.r/r-1

 

 

# States Function : 

è Depends only on State of the system & not the path required to achieve particular state

·       Eg : P ,Vol , Temp , ΔH , ΔE , etc

 

# Path Function :

è Depends on path

·       W.D , q  , etc

 

# Intensive Properties :

Do not depend on mass

 

# Extensive Properties :

Depend on mass