Summary
NCERT Class 11 Chemistry Chapter 5 introduces thermodynamics—the study of energy transformations in chemical and physical processes using the first and second laws, internal energy, enthalpy, entropy, and Gibbs energy to predict spontaneity and equilibrium.
Chapter 5 covers thermodynamic principles governing chemical reactions, including system definitions, the first law (∆U = q + w), and enthalpy changes. Key concepts include internal energy as a state function, enthalpy versus internal energy relationships, calorimetry measurements, and applications like Hess's law for calculating reaction enthalpies. The chapter also explores spontaneity through entropy (disorder measure) and Gibbs energy (∆G = ∆H – T∆S), relating free energy changes to equilibrium constants and reaction direction at different temperatures.
Key points & formulas
- 01System classified as open (matter and energy exchange), closed (energy only), or isolated (no exchange) separated by physical or imaginary boundaries from surroundings.
- 02Internal energy U is a state function; changes via work (w) and heat (q) follow ∆U = q + w (first law of thermodynamics).
- 03Enthalpy H = U + pV; at constant pressure ∆H = qp (heat of reaction); relates ∆U and ∆H via ∆H = ∆U + ∆ngRT.
- 04Standard enthalpy of formation (∆fH) for elements in reference states = zero; reaction enthalpy calculated as ∆rH = Σ∆fH(products) − Σ∆fH(reactants).
- 05Entropy S measures disorder/randomness; spontaneous processes have ∆Stotal = ∆Ssys + ∆Ssurr > 0 (second law).
- 06Gibbs energy ∆G = ∆H − T∆S determines spontaneity: ∆G < 0 (spontaneous), ∆G > 0 (non-spontaneous), ∆G = 0 (equilibrium); related to equilibrium constant K via ∆rG° = −RT ln K.
Frequently asked questions
01What is the difference between internal energy (U) and enthalpy (H) in thermodynamics?
Internal energy U is the total energy of a system and is a state function. Enthalpy H = U + pV is another state function used for reactions at constant pressure. At constant pressure, the heat absorbed equals the enthalpy change (∆H = qp). For solids and liquids, the difference between ∆H and ∆U is small, but for gases, ∆H = ∆U + ∆ngRT, where ∆ng is the change in moles of gaseous reactants and products.
02How does Hess's Law help calculate enthalpy changes for reactions?
Hess's Law states that if a reaction occurs in several steps, its standard enthalpy change equals the sum of the standard enthalpies of intermediate reactions. Since enthalpy is a state function, ∆rH is independent of the reaction pathway. This allows chemists to calculate ∆rH for reactions that are difficult to measure directly by combining known reactions and their enthalpy values.
03What determines whether a chemical reaction is spontaneous—enthalpy alone or entropy as well?
Neither enthalpy nor entropy alone determines spontaneity. Gibbs energy ∆G = ∆H − T∆S does. A reaction is spontaneous when ∆G < 0, which requires both favorable enthalpy (∆H < 0 exothermic) and entropy changes. Endothermic reactions (∆H > 0) can still be spontaneous if T∆S is large enough to make ∆G negative, explaining why reactions become spontaneous at higher temperatures.
04Is the NCERT Class 11 Chemistry Chapter 5 PDF free to download?
Yes, the NCERT Class 11 Chemistry Chapter 5 PDF is free to download. NCERT textbooks are freely available on the official NCERT website and through authorized educational platforms.
More chapters in Chemistry Part I
This is the complete Chemistry Part I Chapter 5 as published by NCERT — every diagram, solved example, and exercise included, free. Browse all NCERT Class 11 textbooks.
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