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Chapter 5: Energetics
5.1 Exothermic and endothermic reactions
- Energetics deals with heat changes occurring during a chemical reactions.
- Enthalpy is the amount of heat energy contained in a substance. It is stored in the chemical bonds as potential energy. When substances react, the difference in the enthalpy between the reactants and products (at constant pressure) results in a heat change which can be measured.
- The reaction mixture is called the system and anything around the system is called the surroundings.
- Thermochemical equations give the balanced equation with the enthalpy change.
e.g. H2 (g) + ½O2 (g) ® H2O (l); DH q = –286 kJ mol–1
H2 (g) + ½O2 (g) ® H2O (g); DH q = –242 kJ mol–1
State symbols must be shown as DH q depends on the state of the reactants or products.
- In exothermic reactions heat is released to the surroundings.
- In endothermic reactions heat is absorbed from the surroundings.
- The standard enthalpy change (DH q) is the heat energy transferred under standard conditions (pressure 101.3 kPa, temperature 298 K). Only DH q can be measured, not H for the initial or final state of a system.
- The standard enthalpy change of combustion is the enthalpy change for the complete combustion of one mole of a substance in its standard state in excess oxygen under standard conditions. All combustion reactions are exothermic.
- The enthalpy of neutralization is the enthalpy change when one mol of H + (aq) reacts with one mol of OH– (aq) ions. The reaction is exothermic as bond formation takes place: H + (aq) + OH– (aq) ¯ H2O (l).
- Exothermic reactions have negative ΔH The temperature of the reaction mixture rises as the chemicals give out heat.
- Endothermic reactions have positive ΔH The temperature of the reaction mixture falls as the chemicals absorb heat.
An exothermic reaction: The products are more stable than the reactants as they have a lower enthalpy.
An endothermic reaction: The products are less stable than the reactants as they have a higher enthalpy.
5.2 Calculation of enthalpy changes
- Calorimetry is the technique of measuring heat changes in physical processes and chemical reactions.
- Heat changes can be calculated from the temperature changes:
heat change (q) = mass (m) ´ specific heat capacity (c) ´ temperature change (ΔT).
- The specific heat capacity is the amount of heat energy required to raise the temperature of unit mass (e.g. 1 kg or 1 g) of a substance, by 1°C or 1 K.
- and for reactions in aqueous solutions can be calculated if it is assumed that all the heat goes into the water.
|= –mH2O ´ cH2O ´ ΔTH2O/nfuelThe experiment is performed with a calorimeter which is a good conductor. This allows heat from the flame to pass to the water.||= –mH2O ´ cH2O ´ ΔTH2O/nlimiting reagentThe experiment is performed with a calorimeter which is an insulator of heat, which reduces heat losses from the system.|
If a calorimeter absorbs heat: Q = (mH2O ´ cH2O ´ ΔTH2O) + (mcalor ´ ccalor ´ ΔTcalor).
Heat loss and incomplete combustion can lead to systematic errors in experimental results.
5.3 Hess’s law
- Hess’s law states that the total enthalpy change for a reaction is independent of the route taken. It is a special case of the law of conservation of energy.
ΔH3 = ΔH1 + ΔH2
5.4 Bond enthalpies
- Average bond energy is the energy required to break one mole of the same type of bonds in the gaseous state averaged over a variety of similar compounds.
- Bond breaking absorbs energy and is endothermic. Bond making releases energy and is exothermic.
= Σ Ebonds broken – Σ Ebonds formed
When Σ Ebonds broken > Σ Ebonds formed : the reaction is endothermic.
When Σ Ebonds formed > Σ Ebonds broken : the reaction is exothermic.
15.1 Standard enthalpy changes of reaction
- The standard state of an element or compound is its most stable state under the standard conditions (pressure 101.3 kPa, temperature 298 K).
- The standard enthalpy change of combustion is the enthalpy change for the complete combustion of one mole of a substance in its standard state in excess oxygen under standard conditions.
- The standard enthalpy change of formation is the enthalpy change when one mole of a substance is formed from its elements in their standard states under standard conditions.
- The enthalpy of formation of any element in its stable state is zero, as there is no enthalpy change when an element is formed from itself.
|Using to find
|Using to find
calculated from or are more accurate than values based on bond enthalpies, which refer only to the gaseous state and are average values.
15.2 Born–Haber cycles
- The first electron affinity is the enthalpy change when one mole of gaseous atoms attracts one mole of electrons: X (g) + e– (g) ® X– (g) .
- The lattice enthalpy is the enthalpy change that occurs when one mole of a solid ionic compound is separated into gaseous ions under standard conditions. For example, for alkali metal halides: MX (s) ® M + (g) + X– (g) .
- depends on the attraction between the ions:
- an increase in the ionic radius of the ions decreases .
- an increase in ionic charge increases .
- The Born–Haber cycle is a special case of Hess’s law for the formation of ionic compounds. It allows the experimental lattice enthalpy to be calculated from other enthalpy changes.
- Theoretical lattice enthalpies can be calculated using a (purely) ionic model from the ionic charges and radii.
- Differences between the theoretical and experimental lattice enthalpies give an indication of the covalent character of the compound; the greater the difference the more covalent the compound.
Born–Haber cycle for NaCl
= 411 + 107 + ½( + 243) + 496 – 349 = + 786.5 kJ mol–1
- Entropy (S) is a property which quantifies the degree of disorder or randomness in a system.
- Ordered states have low S, disordered states have high S: S (s)<. S (l)< S (g).
- Generally matter and energy become more disordered, and Suniverse
- = ΣS q (products) – ΣS q (reactants).
- Gibbs’ free energy (G) is the criterion for predicting the spontaneity of a reaction or process: it is related to . It gives the energy available to do useful work and is related to the enthalpy and entropy changes of the system: .
- ∆Gsys <0 for a spontaneous process. ∆Gsys = 0 at equilibrium.
|Calculating (when T = 298 K)||Calculating (for all T)|
T is in K. As the units of S are J mol–1 K–1 and H are kJ mol–1 they need to be changed to be consistent.
- ∆Gsys and thus the direction of change varies with temperature.
At low temp: : exothermic reactions are spontaneous.
At high temp: : this allows some endothermic reactions to occur if .