NJC Energetics H1 Lecture Notes Student

National Junior College SH1 H1 Chemistry 1 Content ● Enthalpy changes (ΔH): ΔH of formation; combustion; neutralisation; bond energy; lattice energy ● Hess’ Law Learning Outcomes: Candidates should be able to (a) explain that most chemical reactions are accompanied by energy changes, principally in the form of heat usually associated with the breaking and forming of chemical bonds; the reaction can be exothermic (ΔH negative) or endothermic (ΔH positive) (b) construct and interpret an energy profile diagram, in terms of the enthalpy change of the reaction and of the activation energy (c) explain and use the terms: (i) enthalpy change of reaction and standard conditions, with particular reference to: formation; combustion; neutralisation (ii) bond energy (ΔH positive, i.e. bond breaking) (iii) lattice energy (ΔH negative, i.e. gaseous ions to solid lattice) (d) calculate enthalpy changes from appropriate experimental results, including the use of the relationship: heat change = mcΔT (e) explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy (f) apply Hess’ Law to carry out calculations involving given simple energy cycles and relevant energy terms (restricted to enthalpy changes of formation, combustion and neutralisation), with particular reference to: (i) determining enthalpy changes that cannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion (ii) average bond energies [construction of energy cycles is not required] Chemical Energetics All Rights Reserved. No part of this publication may be reproduced or transmitted in any form or by any means, electronic or mechanical, including photocopy, recording or any other information storage and retrieval system, without prior permission in writing from the copyright owner. Copyright © 2022 National Junior College

National Junior College SH1 H1 Chemistry 2 Note: It is not possible to measure the absolute enthalpy of the reactant or product in a system directly. We can only measure changes in enthalpy of the system when it undergoes chemical or physical changes. 1 Introduction The study of energy changes which t ake place during chemical reactions is called chemical energetics or thermochemistry. 1.1 Enthalpy Change, H Learning Objective: Able to explain that most chemical reactions and physical processes are accompanied by changes in heat energy which is a result of the breaking and forming of chemical bonds; ΔH = Hfinal − Hinitial ● When there is a net release in energy, the reaction is exothermic (Δ H, negative) ● When there is a net absorption in energy, the reaction is endothermic (Δ H, positive). ● The sign in ΔH enable you to deduce whether the reactant or product is more stable (lower in enthalpy (H)) The following are concepts and terms related to understanding of enthalpy change: Enthalpy, H ● A measure of energy content in any species, given the symbol H. ● Forming stronger bonds (stronger forces of attraction) reduces energy content (smaller value of H) => greater stability. Enthalpy Change, ΔH Enthalpy change refers to the overall energy absorbed or released in a chemical reaction for 1 mole of a particular substance or molar quantities indicated in a written equation. Its units is usually expressed in kJ mol−1. In all chemical reactions, bonds are broken and formed. (a) bond breaking absorbs energy (i.e. endothermic) (b) bond forming releases energy (i.e. exothermic) ∆H = sum of energy absorbed during bond breaking + sum of energy released during bond forming ● ΔH is either determined experimentally by measuring the temperature change or via application of Hess’ law. Heat Change, q Heat change refers to the overall energy absorbed or released in a chemical reaction. Its units is usually expressed in J or kJ Heat change, q = ∆Hr × no. of moles of limiting reagent

National Junior College SH1 H1 Chemistry 3 1.2 Exothermic / Endothermic reactions Exothermic reaction Endothermic reaction Resultant magnitude of energy sum of energy absorbed to break chemical bonds < sum of energy released from forming chemical bonds. sum of energy absorbed to break chemical bonds > sum of energy released from forming chemical bonds. ΔH ΔH < 0 (−ve) Hsystem decrease ΔH > 0 (+ve) Hsystem increase Relative energy levels of reactants and products (Energy Level Diagram) ● Eproduct < Ereactant ● Products energetically more stable than reactants. Enthalpy / kJ mol–1 ● Eproduct > Ereactant ● Products energetically less stable than reactants. Enthalpy / kJ mol–1 Examples NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l) ΔH = –57.3 kJ mol–1 (neutralisation reaction) NH4NO3(s) + aq → NH4NO3(aq) ΔH = +26.0 kJ mol–1 (dissolving of aqueous ammonium nitrate) More examples Combustion, freezing, condensation, thermite reaction Photosynthesis, melting, vapourisation Note: ● Reactions which are energetically feasible may still not take place because the reaction is too slow i.e. such reactions are kinetically not feasible due to high activation energy (energetically feasible VS kinetically feasible). Reactants Products Heat is released (H = negative) Reactants Products Heat is absorbed (H = positive)

National Junior College SH1 H1 Chemistry 4 Note: ΔH is defined differently for different type of reaction. 1.3 Thermochemical Equation A thermochemical equation is a balanced chemical equation that shows the state symbols of substances and the associated enthalpy change. Example: CH4(g) + 2O2(g) ⎯→ CO2(g) + 2H2O(l) ΔH = –890.4 kJ mol–1 Note: Units of heat energy, q is in J or kJ, as compared to unit of enthalpy change which is in J mol−1 or kJ mol─1 Worked Example 1 Given the following reaction: Fe2O3(s) + 3CO(g) ⎯→ 2Fe(s) + 3CO2(g) ΔHr = –27 kJ mol ̶ 1 Calculate the amount of heat released, q, when (a) 1 mol each of CO(g) and Fe2O3(s) are added together. CO is the limiting agent. According to the above written equation, when 3 moles of CO react, 37 kJ of heat energy is released. ∴ for 1 mole of CO, heat energy released = 27 ÷ 3 = 9 kJ (b) 4 mol of Fe(s) is formed. According to the above written equation, when 2 moles of Fe is formed, 37 kJ of heat energy is released. ∴ for 4 mole of Fe, heat energy released = 27 × 2 = 54 kJ Note: Standard conditions for energetics ≠ standard temperature and pressure for gases (273 K, 1 bar) Note: Enthalpy level of elements in their standard states = zero Note: C(s) in the form of diamond is not the most stable allotropic form of carbon at standard conditions 1.4 Standard Conditions for ΔH ● Changes in energy content of a system are most easily compared when there is a set of reference conditions. ● Hence most enthalpy changes are quoted under standard conditions of: o 298 K or 25 °C o 105 Pa or 1 bar ● Standard conditions are denoted by the symbol ꝋ in superscript. (e.g. Δ𝑯𝒄 ꝋ) ● The element in its natural physical state and its most s