Schooling · Class 11 · Chemistry · Thermodynamics · Chapter on a page
CBSE Class 11 · Chemistry · NCERT Chemistry Part-I
Thermodynamics: chapter on a page
A cut of Shishya's notes for this chapter: the first paragraph of the introduction and the key ideas, formulas and facts, common mistakes and quick revision lists, word for word. Nothing is added or reworded; the worked examples are left out (they are on the chapter page).
About the notes: Written by Shishya's AI on 27 Sept 2026 from the chapter's title and class level, in Shishya's own words — not a copy or summary of the textbook. Read the official chapter for the book's own text, activities and exercises.
This note's current text has no applied fact-check on record.
Mind map
Drawn from the same notes by a fixed rule (no AI): each label is the item's own words — its bold term, its label before a colon, a short formula as written, or its first words.
The mind map as a list
- Key ideas
- System and surroundings
- State functions
- First law of thermodynamics
- Work in expansion
- Enthalpy (H)
- Standard enthalpy of formation (Δ_f…
- Hess's law
- Formulas and facts
- ΔU = q + w
- w = −P_ext × ΔV
- ΔH = q_p
- Standard enthalpy of reaction from…
- Hess's law
- Bond enthalpy approximation
- Gibbs equation
- Worked examples
- Applying the first law
- Enthalpy of reaction from formation…
- Predicting spontaneity with Gibbs…
- Watch out for
- Confusing sign of work
- Using ΔH where ΔU is required
- Forgetting to convert J to kJ or…
- Taking Δ_f H° of an element compound…
- Believing all exothermic reactions…
What this chapter is about
Thermodynamics is the branch of chemistry that deals with energy changes during physical and chemical processes. At Class 11, you study how heat and work are exchanged between a system and its surroundings, and how to predict whether a process will occur on its own. This knowledge helps you understand why reactions release or absorb heat, why ice melts at 0 °C, and why some reactions are spontaneous while others are not.
Key ideas
- System and surroundings: The part of the universe under study is the system; everything outside it is the surroundings. A system may be open (exchanges mass and energy), closed (exchanges only energy), or isolated (exchanges neither).
- State functions: Properties like internal energy (U), enthalpy (H), entropy (S) and temperature (T) depend only on the current state, not on the path taken. Work (w) and heat (q) are path functions.
- First law of thermodynamics: Energy can neither be created nor destroyed; it only changes form. Mathematically, ΔU = q + w, where ΔU is change in internal energy, q is heat absorbed by the system, and w is work done on the system.
- Work in expansion: For a gas expanding against external pressure, w = −P_ext × ΔV. Work done by the system is negative; work done on the system is positive.
- Enthalpy (H): Defined as H = U + PV. At constant pressure, the heat absorbed equals the enthalpy change: q_p = ΔH. This makes enthalpy the practical quantity for reactions in open beakers.
- Standard enthalpy of formation (Δ_f H°): Enthalpy change when one mole of a compound forms from its elements in their standard states (usually 1 bar, 25 °C). For any element in its most stable form, Δ_f H° = 0.
- Hess's law: The total enthalpy change for a reaction is the same whether it occurs in one step or several. This allows calculation of enthalpy changes that cannot be measured directly.
- Entropy (S) measures the randomness or disorder of a system. A spontaneous process in an isolated system increases total entropy.
- Gibbs energy (G): G = H − TS. At constant T and P, ΔG = ΔH − TΔS. If ΔG < 0, the process is spontaneous; if ΔG = 0, equilibrium exists; if ΔG > 0, the process is non-spontaneous.
Formulas and facts to remember
- First law: ΔU = q + w (energy conservation for any system).
- Work of expansion: w = −P_ext × ΔV, with P in Pa and V in m³ giving w in J.
- Enthalpy at constant pressure: ΔH = q_p.
- Standard enthalpy of reaction from formation enthalpies: Δ_r H° = Σ Δ_f H°(products) − Σ Δ_f H°(reactants).
- Hess's law: ΔH for overall reaction = sum of ΔH values for individual steps.
- Bond enthalpy approximation: Δ_r H ≈ Σ (bond enthalpies of bonds broken) − Σ (bond enthalpies of bonds formed).
- Gibbs equation: ΔG = ΔH − TΔS (T in kelvin).
- Spontaneity criterion: ΔG < 0 means the reaction can proceed on its own under those conditions.
Common mistakes
- Confusing sign of work: remember w is positive when work is done on the system, not when the gas expands → always check who does work on whom.
- Using ΔH where ΔU is required: ΔH = q only at constant pressure; at constant volume, q = ΔU → identify the conditions before applying a formula.
- Forgetting to convert J to kJ or vice-versa when mixing ΔH (often kJ) with TΔS (often J) → keep units consistent before adding.
- Taking Δ_f H° of an element compound as zero: only elements in their most stable standard form have Δ_f H° = 0, not all elements in any form → write the specific state.
- Believing all exothermic reactions are spontaneous: spontaneity depends on both ΔH and TΔS via Gibbs energy → always calculate ΔG.
Quick revision
- ΔU = q + w: energy in = energy stored + energy out.
- At constant pressure, heat exchanged equals ΔH.
- Hess's law lets you add enthalpy changes like algebra.
- ΔG = ΔH − TΔS decides spontaneity; negative means go.
- Standard enthalpy of formation for any element in its reference state is zero.
- Entropy increases when gases form or disorder rises.