NEET Chemistry - Chapter 4

Thermodynamics

Fresh NEET thermodynamics notes on system and surroundings, first law, internal energy, enthalpy, calorimetry, Hess law, entropy, Gibbs free energy, and spontaneity.

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Concept Block

1. System, Surroundings, and the First Law of Thermodynamics

Thermodynamics studies energy transformations in matter. The system is the part under study; the rest is the surroundings. Systems can be open (exchange both matter and energy), closed (exchange energy only), or isolated (exchange neither).

Process	Condition	Result
Isothermal	$T$ = constant	$\Delta U=0$ for ideal gas
Adiabatic	$q=0$	$\Delta U=w$
Isobaric	$P$ = constant	$q_p=\Delta H$
Isochoric	$V$ = constant	$w=0,\; q_v=\Delta U$

\Delta U=q+w\qquad(\text{First Law: energy is conserved})

w_{pV}=-P_{ext}\Delta V\quad\text{(work done by gas is negative in IUPAC convention)}

NEET trap: NCERT (IUPAC) sign convention:

w = -P_{ext}\Delta V

. Work done on the system is positive. Work done by the system (expansion) is negative. Don't confuse with physics convention.

Concept Block

2. Enthalpy, Heat Capacities, Hess's Law, and Thermochemistry

Enthalpy $H = U + PV$ is the state function that equals heat exchanged at constant pressure — the condition of most lab reactions. That's why $\Delta H$ appears in thermochemical equations.

\Delta H=\Delta U+\Delta n_g RT\qquad(\Delta n_g=\text{moles of gaseous products}-\text{moles of gaseous reactants})

C_p-C_v=R\quad(\text{for ideal gas})

Hess's Law: $\Delta H$ for a reaction is the same regardless of the path — only initial and final states matter. Add or subtract thermochemical equations algebraically.

Worked example (Hess): If C + O

_2

→ CO

_2

\Delta H_1 = -393

kJ, and CO + ½O

_2

→ CO

_2

\Delta H_2 = -283

kJ, then C + ½O

_2

→ CO has

\Delta H = \Delta H_1 - \Delta H_2 = -110

kJ mol

^{-1}

Standard state: 1 bar pressure, 298 K. The standard enthalpy of formation of any element in its standard state is zero by definition.

Concept Block

3. Standard Enthalpies: Formation, Combustion, Neutralisation, Bond Energy

NEET tests a menu of standard enthalpies. Know their definitions precisely.

Enthalpy Type	Definition	Sign
Formation ( $\Delta H_f°$ )	1 mol compound formed from elements in standard state	Usually −
Combustion ( $\Delta H_c°$ )	1 mol substance completely burned in O $_2$	Always −
Neutralisation ( $\Delta H_n°$ )	Strong acid + strong base → water: −57.1 kJ mol $^{-1}$	Always −
Atomisation ( $\Delta H_{at}°$ )	1 mol gaseous atoms from substance in standard state	Always +
Lattice enthalpy	1 mol ionic solid → gaseous ions	Always +

\Delta H_{rxn}°=\sum\Delta H_f°(\text{products})-\sum\Delta H_f°(\text{reactants})

\Delta H\approx\sum BE(\text{bonds broken})-\sum BE(\text{bonds formed})

NEET trap: Neutralisation enthalpy for weak acid or weak base is less than 57.1 kJ mol

^{-1}

because energy is used to ionise the weak acid/base. For strong acid + strong base, all the enthalpy is simply

H^+ + OH^- → H_2O

Concept Block

4. Entropy, Second Law, and Direction of Spontaneity

Entropy ( $S$ ) measures the degree of dispersal or disorder at the molecular level. The second law states: the total entropy of the universe increases in any spontaneous process.

\Delta S_{universe}=\Delta S_{system}+\Delta S_{surroundings}>0\quad(\text{spontaneous})

\Delta S=\frac{q_{rev}}{T}\quad(\text{at constant T})

Entropy order: $S_{solid}<S_{liquid}\ll S_{gas}$ . Processes that increase entropy:

Solid → liquid → gas (phase transitions)
Dissolution of ionic solids in water (usually)
Reactions that increase moles of gas ( $\Delta n_g > 0$ )
Temperature increase

NEET tip: Entropy of mixing is always positive for ideal solutions. When comparing

\Delta S

of reactions, focus on

\Delta n_g

— any increase in gaseous moles almost certainly gives positive

\Delta S

Concept Block

5. Gibbs Free Energy, Spontaneity Analysis, and Temperature Dependence

Gibbs free energy combines enthalpy and entropy into one criterion for spontaneity at constant $T$ and $P$ .

\Delta G=\Delta H-T\Delta S

\Delta G°=-RT\ln K\qquad(\text{links thermodynamics to equilibrium})

The four-case spontaneity table is a NEET favourite:

$\Delta H$	$\Delta S$	$\Delta G$	Spontaneity
−	+	Always −	Spontaneous at all $T$
+	−	Always +	Non-spontaneous at all $T$
−	−	− at low $T$	Spontaneous at low $T$ only
+	+	− at high $T$	Spontaneous at high $T$ only

NEET trap:

\Delta G &lt; 0

means the reaction is thermodynamically favourable (spontaneous). It says nothing about the rate — a reaction can be spontaneous but extremely slow (e.g., diamond → graphite).

Practice Tests

5 Chapter Tests of 25 Questions Each

Each test is original, NEET-aligned, and answer-backed. Use them as sectional revision instead of a single long mock so your weak subtopics become easier to identify quickly.

Test 1: First Law and Sign Convention

System types, heat, work, internal energy, and process basics.

Test 2: Thermochemistry

Enthalpy, Hess law, bond enthalpy, standard enthalpies, and calorimetry.

Test 3: Entropy and Gibbs Energy

Entropy, free energy, and spontaneity conditions.

Test 4: Numericals

Process numericals, heat capacity, calorimetry, and temperature-based spontaneity.

Test 5: Mixed NEET Drill

Integrated conceptual and numerical thermodynamics questions.

Open Practice Tests

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Move straight from chapter-wise questions into a subject test, then loop back into weaker areas instead of ending the session here.

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