• Spontaneous reaction is a reaction which tends to go without being driven by any external agency. Spontaneous reactions are the chemical equivalent of water flowing downhill.

Enthalpy changes and spontaneous reactions.

  • Most reaction are spontaneous are also exothermic. They have a negative standard enthalpy change of reaction.
  • Use sign of ΔH to decide whether or not a reaction is likely to go.
  • Exothermic reactions with high activation energy are also not feasible.

However, some endothermic processes are spontaneous too. This shows the limitation of using the enthalpy change to decide the likely direction of change or spontaneity of reaction.

A new factor is there to determine the spontaneity of the reaction. That is known as entropy.

Diffusion – spontaneous change.

After removing the barrier, the bromine diffuses from the right – hand bottle until the bromine molecules are evenly spread between the two bottles. Once mixed the air never unmix.

Disorder, or entropy(S)

Degree of disorder is known as entropy.

  • Boltz mann was shown the relation ship between entropy and number of ways of arranging the particles.

S = k lnW.

k is the Boltz mann constant and W is the number of ways of arranging the particles.

  • When considering chemical reaction it is essential to calculate the total entropy change in two parts: the entropy change of the system and the entropy change of surroundings.

ΔS = ΔSsystem + ΔSsurroundings.

  • Standard molar entropy, S, is the entropy per mole for a substance under standard conditions at 298K and 100kPa.
  • The units for standard molar entropy are joules per mole per Kelving (Jmol-1K-1)
  • If a perfectly ordered crystal at 0K the entropy is zero. The entropy of a chemical rises as the temperature rises.

The entropy change of the system.

  • ΔS = Sum of the standard molar entropies of the products – Sum of the standard molar entropies of the reactants.

The entropy change of the surroundings.

  • It is not enough to consider only the entropy of the system. What matters is the total entropy change, which is sum of the entropy changes of the system and the entropy change in the surroundings.

ΔSsurroundings = -

  • For an exothermic reaction, which transfers energy to the surroundings, ΔH is negative, so

–ΔH is positive.

The total entropy change.

~A reaction is only feasible if the total entropy change, ΔStotal is positive:

ΔStotal = ΔSsystem + ΔSsurroundings.

Thermodynamic stability and kinetic stability.

  • If the total entropy change, ΔStotal, is negative this indicates that the reaction doesnot occur. A chemical mixture of chemical is thermodynamically stable if there is no tendency for a reaction.
  • Some chemical reaction does not be spontaneous even it has positive entropy, because activation energy will be high. This condition is known kinetic stability.

Equilibrium constant and entropy.

The relation between entropy change and equilibrium constant is

ΔStotal = RlnK,

K is the equilibrium constant and R is the gas constant.

ΔStotal = ΔSsystem -

What happens to equilibrium constant when temperature increases in an exothermic reaction and endothermic reaction, also explain the position of equilibrium in exothermic and endothermic reaction on heating.

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Lattice energy and its importance.

Lattice enthalpy of an ionic crystal is defined as the enthalpy change for the formation of one mole of the ionic compound from gaseous ions under standard temperature(298K) and standard pressure100kPa. Thus, the lattice enthalpy of sodium chloride corresponds to the process;

Na+(g) + Cl-(g) NaCl(S)

Enthalpy of hydration ΔHhyd.

ΔHhyd is the heat change per mole for the hydration of the gaseous ion with enough water for there to be no further heat change on dilution.

M+(g) + aq M+(aq)

The ionic compound dissolves if ΔHhydration is equal or over comes lattice energy.

ΔHsolution = ΔHhydration of cation + ΔHhydration of anion –ΔHlattice of ionic compound.

Draw the Hess`s law cycle of dissolution of sodium chloride.

The decomposition of group 2 metal carbonates is used on a large scale to make oxides such as magnesium and calcium oxides.

MgCO3(s) MgO(s) + CO2(g)

ΔH=+117kJ/mol ΔSsystem = +175Jmol-1K-1

BaCO3(s) BaO(s) + CO2(g)

ΔH = + 268kJ/mol ΔSsystem = +172Jmol-1K-1

The carbonate of group 2 metals do not decompose at room temperature. They do decompose on heating.

  1. Why does the entropy of the system increase when a group 2 carbonate decomposes?
  2. a) Calculate the total entropy change for the decomposition of;

i)magnesium carbonate.

ii)Barium carbonate.

b) Are these two compounds stable or relative to decomposition into their oxide and carbon dioxide at room temperature(298K)?

3. Assuming that ΔH and ΔSsystem for the reaction do not vary with temperature, estimate the temperatures at which the two decomposition reactions become feasible.

Ihavandhoo school/A level note/How far?- Entropy Page 1