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Unformatted text preview: CHAPTER 5 ENERGETICS
(IB TOPICS 5 AND 15) SUMMARY
Energetics deals with heat changes in chemical reactions. It is measured in joules, J. Units of
energy: joule; 1 cal = 4.184 J. In an exothermic reaction, heat is given out to the environment
and the temperature of the system increases. Endothermic reaction is one in which heat is taken
in from environment and the temperature of the system decreases. The reaction mixture is called
the system and anything around the system is called the surrounding or environment. ΔH°, Standard enthalpy change is heat change (given out or evolved) under standard
conditions of 101.3 kPa pressure and a temperature of 298 K. Only ΔH can be measured, not
H for the initial or final state of a system. ΔHºreaction is the sum of the change in enthalpies of
products minus sum of the change in enthalpies of reactants.
ΔH°reaction = ΣΔH°products –ΣΔH°Reactants
Thermochemical equations include a balanced equation and an indication of the heat flow:
H2 (g) + ½O2 (g) → H2O (l); ΔHº = –285.8 kJ mol–1 1. Coefficients refer to number of moles.
2. States of substances must be specified.
3. ΔH ∝ amount of substance that reacts or is produced in a reaction.
Standard Enthalpy Change of Formation, ΔHºf is the enthalpy change when one mole of a
compound is formed from its elements in their stable states under standard conditions of 298
K and 101.3 kPa pressure. By definition the enthalpy of formation of any element in its stable
form at 298 K and 101.3 kPa pressure equals zero. This is because no energy change is
required to form an element from itself. Enthalpy change of a reaction is the sum of the ΔHºf of
products minus sum of ΔHºf of the reactants. ΔHºrxn = Σ ΔHºf, Products – Σ ΔHºf, Reactants
ΔHneut: Enthalpy of neutralization is the heat evolved when one mol of H3O+ reacts with 1 mol OH– ions. Heat is evolved because bond formation takes place: H3O+ + OH– → 2H2O. ΔHlatt: Lattice enthalpy is the energy required for the conversion of one mol of an ionic solid
into gaseous ions or energy given out when one mole of ionic solid is formed from ions in
For NaCl it is: NaCl(s) → Na+(g) + Cl–(g); ΔHlatt = +
OR Na+(g) + Cl–(g) → NaCl (s); ΔHlatt = –
Calorimetry is the technique of measuring heat changes in physical processes (such as melting
or heat gained by a solid) and chemical reactions. In order to calculate heat change, we need:
Specific heat, c: the energy required to change the temperature of 1 g of the substance by 1°C.
Note: in case of aqueous solutions, we assume solutions are dilute enough, and use the specific
heat of water. Units of specific heat: J g−1°C−1 (=J g−1 K−1 since change in 1°C = change in 1 K).
Mass of the substance (in aqueous solutions, the mass of solution, not of solute), and ΔT, the
change in temperature. Thus: the heat change, Q = mcΔT.
The equation Q = mcΔT can be used to calculate:
1. Heat change of a substance, given its mass, specific heat and temperature change.
2. ΔH°reaction in aqueous solution using experimental data on temperature changes,
quantities of reactants and mass of solution (where it is assumed that that the solution is
dilute enough so that specific heat of solution = specific heat of water = 4.18 J g−1 C−1).
© IBID Press 2007 1 CHAPTER 5 ENERGETICS
(IB TOPICS 5 AND 15) SUMMARY
Note that if a calorimeter is used that also absorbs energy, then its mass, m and specific heat, c
(or heat capacity, m x c) must also be known:
ΔHsol = heat absorbed by solution + heat absorbed by calorimeter
= (msolution × csol × ΔTsol) + (mflask × cflask × ΔTflask) According to Hess's Law, ΔH of a reaction is independent of any intermediate steps. Hess’s law
is a special case of the law of conservation of energy.
Bond Enthalpy is the energy required when one mole of bonds is broken in the gaseous state
(bond enthalpies are valid only in the gaseous state). Bond breaking needs energy and is an
endothermic process. Bond making releases energy and it is an exothermic process. Average
bond energy is the average energy required to break a mole of the same type of bonds in the
gaseous state in a variety of similar compounds.
If bond enthalpies are given then: ΔHºrxn = Σ BEbonds broken - ΣBEbonds made (not products –
Born-Haber cycle, a special case of Hess’s law for the formation of ionic compounds involves
three key steps:
1. ΔH°f , standard change in enthalpy of formation of the ionic crystal, i.e., starting from the
elements in their stable states under standard conditions, e.g., Na (s) + ½ Cl2 (g) to form the
ionic solid, NaCl (s).
2. Starting from the same elements, produce gaseous ions, e.g., Na+ (g) and Cl− (g). This
involves a series of processes: convert Na (s) to Na (g) (sublimation) and then to Na+ (g)
(ionization energy). Similarly, convert ½ Cl2 (g) to Cl (g) (½ the BE), then form Cl− (g)
3. Lattice Enthalpy: Na+ (g) + Cl− (g) → NaCl (s)
Lattice enthalpy value from Born-Haber cycle is called experimental lattice energy value, as
experimental values of the other processes involved are used to arrive at it. On the other hand, a
theoretical value can be calculated using a (purely) ionic model and distance between the ions
using laws of electrostatics. Comparison between the theoretical and experimental lattice
enthalpies gives an indication of whether the substance is ionic (if there is good agreement) or
has covalent character.
Entropy, S is a property that relates to the degree of disorder or randomness in a system. In
general, systems in nature prefer greater disorder. The change in entropy, ΔS = Sp – SR. For a
reaction, ΔSº = Σ Sºproduct – Σ Sºreactants. By definition, ΔS, the entropy change is positive for
Free energy, G is the criterion for predicting the spontaneity of a reaction. Change in free
energy, ΔG accounts for both ΔH and ΔS and is determined by the Gibbs-Helmholtz equation:
ΔG° = ΔH° – TΔS°; ΔG° is the standard free energy change at 298 K and 101.3 kPa pressure,
ΔH° is the standard enthalpy change, ΔS° the standard entropy change and T is temperature in
Kelvin. ΔG is a measure of the driving force of a reaction, i.e., its tendency to proceed spontaneously. A
reaction or physical process can occur spontaneously if ΔG is negative. A spontaneous reaction
or process may be used to do work (ΔG it is the amount of energy available to do useful work)
on another system. For minimum energy ΔH = –, for maximum entropy ΔS = + Thus, ΔG is
negative for any spontaneous reaction. For a non-spontaneous reaction (or spontaneous in the
reverse direction) ΔG is positive and for an equilibrium process ΔG = 0.
(N.B. Shading indicates Topic 15 (AHL) material.)
© IBID Press 2007 2 ...
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This note was uploaded on 09/23/2010 for the course CS 001 taught by Professor Jix during the Spring '10 term at Riverside Community College.
- Spring '10