Thermodynamics and Thermochemistry

Introduction & Basic Concepts

Thermodynamics studies energy forms and transformations. Based on empirical laws: Zeroth, First, Second, Third.

System: Observed part of universe. Surroundings: Rest. Boundary: Separator.

System Types

Isolated

No matter/energy exchange (e.g., thermos flask)

Closed

Energy only (e.g., pressure cooker)

Open

Matter & energy (e.g., open beaker)

Homogeneous/Heterogeneous

Homogeneous

One phase (e.g., solution)

Heterogeneous

Multiple phases (e.g., ice-water)

Properties & State Functions

Intensive/Extensive

Intensive

Size-independent (T, P, density)

Extensive

Size-dependent (V, E, H, S)

Extensive/mole → Intensive

Processes & Equilibrium

Process	Condition	Key Relation
Isothermal	ΔT=0	ΔE=0
Adiabatic	q=0	ΔE=w
Isobaric	ΔP=0	ΔH=q
Isochoric	ΔV=0	ΔE=q
Cyclic	Initial→Final	ΔE=ΔH=0
Reversible	Infinitesimal	Max work
Irreversible	Finite step	ΔS_univ >0

Equilibrium: Chemical (fixed comp.), Mechanical (const. P), Thermal (const. T).

Internal Energy (E) & First Law

E = E_trans + E_rot + E_vib + E_elec + E_nuc

Properties

Extensive, state function, ΔE path-independent

Ideal gas: f(T) only

First Law

ΔE = q + w (w on system)

Or ΔE = q - w (w by system)

Units: 1 J = 10^7 erg = 0.239 cal

Heat (q) & Work (w)

q

Absorbed: +q (endothermic)

Evolved: -q (exothermic)

w

On system: +w

By system: -w (expansion)

Zeroth & Heat Capacities

Zeroth Law: Basis of temperature; thermal equilibrium transitive.

Cp & Cv

Cp - Cv = R (gases); γ = Cp/Cv

Gas Type	Cv	Cp	γ
Monoatomic	3/2 R	5/2 R	1.67
Diatomic	5/2 R	7/2 R	1.40
Polyatomic	3R	4R	1.33

Enthalpy & Expansion

H = E + PV; ΔH = ΔE + Δn_g RT

Solids/liquids: ΔH ≈ ΔE

Isothermal Expansion (ΔT=0)

Reversible: w = -nRT ln(V2/V1) = 2.303 nRT log(V1/V2)

Irreversible: w = -P_ext (V2 - V1)

Adiabatic Expansion (q=0)

PV^γ = const; TV^{γ-1} = const; TP^{(1-γ)/γ} = const

w = ΔE = n Cv ΔT

Spontaneity & Second Law

Spontaneous: Occurs without work (e.g., diffusion, heat flow hot→cold).

Second Law Statements

Kelvin

No perpetual motion; can't convert heat fully to work

Clausius

Heat cold→hot requires work

Efficiency = 1 - T_low/T_high <1

Entropy (S)

Entropy: Measure of disorder/randomness. ΔS = q_rev / T

ΔS_univ = ΔS_sys + ΔS_surr >0 (spontaneous); =0 (eq.); <0 (non-spont.)

Changes

Exothermic

ΔS_surr >0

Endothermic

ΔS_surr <0

Phase Transition

ΔS_fus = ΔH_fus / T_fus; ΔS_vap = ΔH_vap / T_boil

Ideal Gas

ΔS = n Cv ln(T2/T1) + n R ln(V2/V1)

Adiabatic reversible: ΔS=0 (isoentropic)

Units: J/K or cal/K; S = k ln W (Boltzmann)

Free Energy & Third Law

G = H - TS; ΔG = ΔH - T ΔS

Properties

State function; ΔG=0 (eq.); <0 (spont.); >0 (non-spont.)

ΔG° = -2.303 RT log K

Spontaneity Criteria

ΔH<0, ΔS>0: Always spont.

ΔH>0, ΔS<0: Never spont.

ΔH<0, ΔS<0: Low T spont.

ΔH>0, ΔS>0: High T spont.

ΔH	ΔS	ΔG	Spontaneity	Example
-	+	Always -	All T	2O3(g) → 3O2(g)
+	-	Always +	No T	3O2(g) → 2O3(g)
-	-	Low T -	Low T	CaCO3(s) → CaO(s) + CO2(g)
+	+	High T -	High T	CaO(s) + CO2(g) → CaCO3(s)

Third Law

Perfect crystal entropy S=0 at 0 K. Absolute S calculable: S = 2.303 Cp log T

Limitations: Glasses, isotopes, disordered crystals (e.g., CO, N2O) have S>0 at 0 K.

Thermochemistry

Energy changes in reactions. Exothermic (ΔH<0, heat evolved); Endothermic (ΔH>0, heat absorbed).

ΔH_rxn = Σ ΔH_f(products) - Σ ΔH_f(reactants)

Factors Affecting ΔH

Physical State

Gas vs liquid changes ΔH (e.g., H2O(g) vs l)

Allotropes

C(diamond) vs C(graphite)

Temperature

Kirchhoff: ΔH_T2 = ΔH_T1 + ΔCp (T2 - T1)

P vs V

ΔH = ΔE + Δn_g RT

Types of Heats

Type	Definition	Example
Formation (ΔH_f°)	1 mol from elements (25°C, 1 atm)	NH3(g): -11 kcal/mol
Combustion	1 mol complete burn in O2	CH4(g): -192 kcal/mol
Neutralization	1 eq acid + base (dil. soln.)	Strong acid+base: -13.7 kcal/eq
Solution	1 mol solute in excess solvent	NH4Cl(s): +3.90 kcal/mol
Hydration	Anhyd. salt + H2O → hydrate	CuSO4·5H2O: -18.69 kcal/mol
Vaporization	Liquid → gas	H2O(l): +10.5 kcal/mol
Fusion	Solid → liquid	Ice: +1.44 kcal/mol
Sublimation	Solid → gas	I2(s): +14.8 kcal/mol
Precipitation	Electrolytes → sparingly sol. salt	BaSO4: -4.66 kcal/mol

Hess's Law & Calorimetry

Hess's Law: ΔH independent of path; total ΔH = sum of steps.

Applications: ΔH_f from combustion; slow reactions; allotropes; bond energies; resonance; lattice energy.

Calorimetry

Bomb (Const. V)

ΔE = [(m s + W)(T2 - T1)] / wt

Combustion

Organic compounds

Levoissier-Laplace: ΔH_decomp = - ΔH_form

Bond Energies

Bond Energy: Avg. energy to break 1 mol bonds in gas phase (kJ/mol).

ΔH_rxn = Σ BE(reactants) - Σ BE(products)

Examples: H-H: 433; Cl-Cl: 242; H-Cl: 431 kJ/mol

Resonance Energy

RE = Exp. ΔH_f - Calc. ΔH_f (from BE)

Important Points & Tips

Key JEE Points

ΔE = q + w; ΔH = ΔE + ΔnRT
Reversible work max; ΔS_univ >0 spontaneous
Cp = Cv + R; γ mono=5/3
Hess: Path independent; Bond E for ΔH calc.
ΔG = -RT ln K; Third Law: S=0 at 0K
Neutralization strong: -57 kJ/mol (H+ + OH- → H2O)

Do's

Use Kirchhoff for ΔT effects

Apply Hess for indirect ΔH

Sign conv.: q in +, w on +

Δn_g for gases only

Don'ts

Forget state in ΔH (gas/liquid)

Use BE for solids

Ignore ΔS in spontaneity

Confuse ΔE & ΔH

Objective Questions

1. Internal energy of ideal gas depends on:

(a) Volume (b) Temperature (c) Pressure (d) None

2. Series returning to initial state:

(a) Boyle's (b) Reversible (c) Adiabatic (d) Cyclic

3. 1 cal = ? J [CPMT 1988]

(a) 0.4184 (b) 4.184 (c) 41.84 (d) 418.4

4. ΔE for reversible isothermal cycle:

(a) 100 cal/° (b) Negative (c) 0 (d) Positive

5. Thermos with ice: [AIIMS 1992]

(a) Closed (b) Open (c) Isolated (d) Non-thermo

6. Intensive quantity: [IIT JEE 1993]

(a) Enthalpy, T (b) V, T (c) Enthalpy, V (d) T, Refractive index

7. Largest energy unit: [CPMT 1989]

(a) eV (b) Erg (c) Joule (d) Calorie

8. Order: 1 erg, 1 J, 1 cal [NCERT 1980]

(a) cal > J > erg (b) J > cal > erg (c) cal > J > erg (d) erg > cal > J

JEE Main Weightage

Typically 3-4 questions. Focus: Laws, ΔH/ΔE calcs, Hess, spontaneity, bond E.

Weightage Very High