Energy Changes in Chemical Reactions – Comprehensive Study Notes

2.1 ENERGY CHANGES

• Types of energy

  • Kinetic energy: due to motion

  • Potential energy: due to position

  • Electrical energy, light energy, thermal energy, chemical energy, mechanical energy, etc.

• Law of Conservation of Energy

  • Energy cannot be created or destroyed; it can only be transformed from one form to another.

  • In chemical processes, energy can transfer to or from the surroundings; total energy is conserved.

  • In a chemical reaction, energy changes often occur in the form of heat; energy can be transferred to the surroundings or absorbed from them.

• Energy transformations in daily devices

  • Electric motor: electrical energy → kinetic energy

  • Hair dryer: electrical energy → thermal energy

  • Battery: chemical energy → electrical energy

  • Automobile engine: chemical energy → mechanical energy

  • Solar cell: light energy → electrical energy

• Energy changes in chemical reactions

  • All chemical reactions involve energy changes due to breaking and forming bonds.

  • Even changing states of matter involves energy transfer (phase changes).

  • Examples of everyday state changes: Evaporation (boiling water), Melting (ice), Condensation (water droplets on cold glass), Sublimation (mothball/dry ice).

• Chemical energy changes and everyday applications

  • Photosynthesis: plants convert light energy, CO₂, and H₂O into stored chemical energy to make plant food.

  • Combustion of fossil fuels releases chemical energy as heat and light, powering electricity and transportation.

  • Except for nuclear energy, most energy resources involve chemical reactions.

• Learning outcomes (conceptual groundwork)

  • Distinguish reactions as spontaneous vs non-spontaneous; exothermic vs endothermic.

  • Use energy concepts to understand bond breaking and bond making.

  • Differentiate enthalpy (AH) and enthalpy changes of reactions; discuss types of enthalpy changes and how to measure them.

  • Apply Hess’s law to construct enthalpy cycles and calculate enthalpy changes.

  • Calculate enthalpy change of a reaction using bond energies (bond enthalpy).

• Spontaneous vs non-spontaneous processes

  • Spontaneous: occurs without added energy under given conditions.

  • Example: Dilute sulfuric acid + sodium hydroxide reacts spontaneously.

  • Non-spontaneous: need energy to proceed (e.g., Mg + O₂, Cu + O₂ require initiation energy).

  • Some reactions may start non-spontaneously but continue spontaneously once begun (e.g., rusting is exothermic and spontaneous; ice melting at room temperature is spontaneous and endothermic).

• Exothermic vs endothermic reactions

  • Exothermic: heat released to surroundings; products have lower energy than reactants.

  • Example: Burning candle, fuel combustion, neutralisation, respiration, rusting of iron.

  • Energy level behavior: energy of reactants > energy of products; temperature of surroundings rises.

  • Endothermic: heat absorbed from surroundings; products have higher energy than reactants.

  • Example: Dissolution of ammonium chloride in water, melting ice, cooking an egg, evaporation of liquid water, thermal decomposition of limestone.

• Energy transfer via bonds

  • Bond breaking is endothermic (absorbs energy).

  • Bond making is exothermic (releases energy).

  • Overall energy change depends on the balance between bonds broken and bonds formed.

  • Example explanation: When nitrogen reacts with oxygen to form NO, bonds in N≡N and O=O must break (absorb energy), then new N–O bonds form (release energy).

• Energy diagrams (conceptual)

  • Exothermic: products have lower energy; energy released to surroundings.

  • Endothermic: products have higher energy; energy absorbed from surroundings.

• Quick example illustrating bond energy balance

  • If energy absorbed in bond breaking < energy released in bond making → exothermic.

  • If energy absorbed in bond breaking > energy released in bond making → endothermic.

• Quick numerical example (conceptually)

  • Reaction: H₂(g) + Cl₂(g) → 2 HCl(g)

  • Given: bond energies to break: HH = 436 kJ, ClCl = 242 kJ; to form: 2× H–Cl = 431 kJ

  • Total energy absorbed for bond breaking = 436 + 242 = 678 kJ

  • Total energy released for bond making = 2 × 431 = 862 kJ

  • Since 678 kJ < 862 kJ, the reaction is exothermic.

2.2 ENTHALPY CHANGES IN CHEMICAL REACTIONS

• Enthalpy concept

  • Enthalpy (H) is the total chemical energy content of a system.

  • Enthalpy change AH (or ΔH) is the heat change at constant pressure during a reaction.

  • Units: $AH \; (\text{or } \Delta H) = \text{kJ mol}^{-1}$

  • AH describes heat transferred between system and surroundings under constant pressure.

  • Sign convention:

  • Endothermic: AH is positive (products higher energy).

  • Exothermic: AH is negative (energy released).

• Standard conditions

  • Standard temperature: $T^{\circ} = 298\ \text{K}$ (25 °C)

  • Standard pressure: $P^{\circ} = 1\ \text{atm}$ (101 kPa; 760 mmHg)

• Thermochemical equations

  • Balanced chemical equation with physical states and the standard enthalpy change (AH°).

  • The enthalpy change can be written as AH° or AH°rxn.

  • The enthalpy change depends on the states of reactants and products; states are indicated as (s), (l), or (g).

  • Example thermochemical equations (illustrative):

  • HgO(s) → Hg(l) + 1/2 O₂(g) ; AH° = +90.7 kJ mol⁻¹

  • C(graphite) + O₂(g) → CO₂(g) ; AH° = -394 kJ mol⁻¹

• Standard enthalpy changes of reactions (AH°)

  • AH° is the enthalpy change when the molar quantities stated in the chemical equation react under standard conditions.

  • AH° can be exothermic (negative) or endothermic (positive).

  • Example relationships (illustrative values from the source):

  • H₂(g) + ½ O₂(g) → H₂O(l) ; AH° = -286 kJ mol⁻¹

• Key standard enthalpy concepts

  • Standard enthalpy change of formation, AH°f: enthalpy change when 1 mole of a compound is formed from its elements in their standard states.

  • Standard enthalpy changes of combustion, AH°c: enthalpy change when 1 mole of a substance is burned completely in O₂ under standard conditions (always negative).

  • Example: CH₄(g) + 2 O₂(g) → CO₂(g) + 2 H₂O(l) ; AH°c ≈ -891 kJ mol⁻¹

  • Standard enthalpy change of neutralisation, AH°neut: enthalpy change when an acid and base react to produce 1 mole of water under standard conditions (negative).

  • Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l) ; AH°neut ≈ -57.1 kJ mol⁻¹

• Formation, combustion, and neutralisation (typical values from the text)

  • Formation of NO₂ from N₂ + O₂: AH°f(NO₂) ≈ +33 kJ mol⁻¹

  • Formation of CO from C(graphite) + 1/2 O₂(g) → CO(g): AH°f(CO) ≈ -110 kJ mol⁻¹

  • Combustion of carbon (graphite): AH°c ≈ -394 kJ mol⁻¹

  • Combustion of methane: AH°c ≈ -891 kJ mol⁻¹

  • Combustion of ethanol: AH°c ≈ -1367 kJ mol⁻¹

  • Formation of ethanol: AH°f(ethanol) ≈ -235 kJ mol⁻¹

  • Formation of NaHCO₃(s): AH°f ≈ -951 kJ mol⁻¹

• Example 2: Write thermochemical equations for the following reactions

  • (i) Combustion of carbon (graphite): AH°c = -394 kJ mol⁻¹

  • (ii) Combustion of methane: AH°c = -891 kJ mol⁻¹

  • (iii) Combustion of ethanol: AH°c = -1367 kJ mol⁻¹

  • (iv) Formation of ethanol: AH°f(ethanol) ≈ -235 kJ mol⁻¹

  • (v) Formation of NaHCO₃(s): AH°f ≈ -951 kJ mol⁻¹

• Measuring enthalpy changes

  • Direct measurement via calorimetry is not always feasible; use calorimetry to measure q and relate to AH.

  • Key formula for calorimetry (constant pressure): $q = m c \Delta T$

  • If the system is at constant pressure, AH ≈ q (for the reaction).

• Enthalpy changes in solution

  • Enthalpy of solution AHsol: heat change when 1 mole of substance dissolves in water.

  • Experimental setup: use a calorimeter; record initial and final temperatures; compute q.

  • Example calculation (NH₄NO₃ dissolution):

  • Given: mass of water m = 100 g, c = 4.18 J g⁻¹ K⁻¹, mass NH₄NO₃ = 7.10 g, initial T = 18.2 °C, final T = 12.8 °C.

  • ΔT = Tfinal − Tinitial = 12.8 − 18.2 = −5.4 °C

  • q = m c ΔT = 100 × 4.18 × (−5.4) ≈ −2263 J ≈ −2.263 kJ

  • Molar mass NH₄NO₃ = 80.0 g/mol ⇒ n = 7.10 g / 80.0 g/mol ≈ 0.08875 mol

  • AHsol = q/n ≈ (−2.263 kJ) / 0.08875 mol ≈ −25.43 kJ mol⁻¹

  • Since ΔT is negative for the solution’s temperature rise, the dissolution is endothermic for the system (the water cooled; the solution absorbed heat). The sign convention here indicates AHsol > 0 for the dissolution process in this setup.

• Example 3: Neutralisation enthalpy (HCl + NaOH)

  • Setup: 50 cm³ of 2.0 M HCl mixed with 50 cm³ of 2.0 M NaOH; final temperature rise measured.

  • Given: initial temps HCl = 17.5 °C; NaOH = 17.9 °C; final temp = 31.0 °C; c(H₂O) = 4.18 J g⁻¹ °C⁻¹; volumes assumed additive and dilute so water-like.

  • Average temperature ≈ 17.7 °C; ΔT = 31.0 − 17.7 = 13.3 °C

  • Mass of solution ≈ 100 g (50 g HCl solution + 50 g NaOH solution; treat as 1 g/mL).

  • q = m c ΔT = 100 × 4.18 × 13.3 ≈ 5.56 kJ

  • Moles reacting: 0.050 L × 2.0 M = 0.10 mol of HCl (and equal moles of NaOH)

  • AH for neutralisation per mole of water formed: ΔH ≈ −5.56 kJ / 0.10 mol = −55.6 kJ mol⁻¹

  • Conclusion: Exothermic neutralisation; heat released during the reaction.

• Bomb calorimeter

  • Used to determine the heat of combustion by burning a sample in an oxygen-rich, sealed vessel and measuring the temperature rise.

  • Highly accurate standard enthalpy changes of combustion (AH°c) are determined using a bomb calorimeter.

• Review questions (conceptual focus)

  • Identify energy transformations in: (i) electric bulbs, (ii) fuel cells, (iii) electric generators, (iv) steam boilers, (v) gas turbines.

  • Predict exothermic vs endothermic for various processes: burning a candle, digestion, cooking an egg, condensation, cracking alkanes.

  • Classify a given reaction as exothermic or endothermic using bond energies and AH signs.

2.3 HESS'S LAW

• Core idea

  • The enthalpy change for a chemical reaction is the same irrespective of the route taken from reactants to products.

  • Path independence allows calculation of AH for difficult reactions by combining known AH values for simpler steps.

• Enthalpy cycle (illustrative)

  • Route I: A → B directly; Route II: A → C → B, etc.

  • According to Hess’s law: AH(Route I) = AH(Route II) = sum of AH along Route II

  • If Route II has steps with AH1, AH2, AH3, then AH(Route II) = AH1 + AH2 + AH3

• Constructing an enthalpy cycle 1) Write the equation for the enthalpy change of the target reaction. 2) Draw the enthalpy cycle using data provided; mark directions of energy changes. 3) Label enthalpy changes along the arrows (AH values). 4) Compute AH for the target reaction via two routes (clockwise vs anticlockwise) and apply Hess’s law.

  • Note: The standard enthalpy of formation of an element in its standard state is zero.

• Bond energy and enthalpy changes

  • Bond enthalpy (bond energy) E is the energy required to break one mole of a specific bond in a molecule.

  • Breaking a bond requires energy (AH > 0); forming a bond releases energy (AH < 0).

  • The energy to form or break bonds is used to estimate AH for reactions via bond energies.

  • Example: Br–Br bond energies: E(Br–Br) breaking = +193 kJ mol⁻¹; Br–Br formation = −193 kJ mol⁻¹.

• Using bond enthalpies to approximate AH

  • Hess’s law can be applied using bond enthalpies to estimate AH for reactions that are difficult to measure directly.

  • Limitations: bond enthalpy values are average values and gas-phase, etc.

• Practical takeaway

  • For reactions like C–H, O=O, etc., sum of bonds broken minus sum of bonds formed gives AH (approximate).

Notation and key formulas (recap)

• Energy balance in a reaction at constant pressure:

  • $\Delta H = q_p$

  • where q is heat at constant pressure and AH is the enthalpy change.

• Enthalpy change of reaction:

  • $\Delta H<em>{rxn} = H</em>{products} - H_{reactants}$

• Standard conditions:

  • $T^{\circ} = 298\ \text{K}, \quad P^{\circ} = 1\ \text{atm}$

• Enthalpy of formation, combustion, neutralisation (examples):

  • Formation: $\Delta H_f^{\circ}$

  • Combustion: $\Delta H_c^{\circ}$

  • Neutralisation: $\Delta H_{neut}^{\circ}$

• Bond enthalpy concept:

  • Bond breaking: AH > 0

  • Bond forming: AH < 0

  • Example bond energy notation: $E(\text{Br–Br}) = +193\ \text{kJ mol}^{-1}$ for breaking; $E(\text{Br–Br}) = -193\ \text{kJ mol}^{-1}$ for forming

Connections and implications

• Energy changes underpin predictions of reaction spontaneity, heat release/absorption, and feasibility of processes in chemistry and daily life.

• Hess’s Law enables calculation of AH for complex synthesis or decomposition routes using known data (formation, combustion, neutralisation).

• Bond energies provide quick, approximate estimates of AH when direct measurements are impractical.

• Calorimetry (q = mcΔT) links experimental temperature changes to AH, enabling practical determination of enthalpy changes in reactions and solutions.

• Standard enthalpy concepts connect to real-world data: energy content of fuels, heat of neutralisation, and energy conversions in biological and industrial processes.

Quick reference values (typical from the chapter)

• Standard enthalpy of combustion (examples):

  • Carbon (graphite): AH°c ≈ -394 kJ mol⁻¹

  • Methane (CH₄): AH°c ≈ -891 kJ mol⁻¹

  • Ethanol (C₂H₅OH): AH°c ≈ -1367 kJ mol⁻¹

• Standard enthalpy of formation examples:

  • NO₂(g): AH°f ≈ +33 kJ mol⁻¹

  • CO(g): AH°f ≈ -110 kJ mol⁻¹

  • H₂O(l) from H₂ and O₂: AH°f ≈ -286 kJ mol⁻¹

• Neutralisation: AH°neut ≈ -57.1 kJ mol⁻¹ for HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

• Formation of NaHCO₃(s): AH°f ≈ -951 kJ mol⁻¹

• Practical example: dissolution of NH₄NO₃ in water is endothermic with AHsol ≈ +25.4 kJ mol⁻¹ (per mole of NH₄NO₃).