Reaction Energy and Reaction Kinetics

General ChemistryUnit 12

Driving Forces

• Enthalpy and Entropy• Enthalpy (heat of reaction) is the

amount of energy released or absorbed during a chemical reaction • Symbol is ΔH• Think of it as energy needed

Thermochemical Equations

• A thermochemical equation shows the energy (enthalpy) change in the reaction• Put in as reactant or product

2 H2 + O2 → 2 H2O + 483.6 kJ

• List behind as ΔH2 H2 + O2 → 2 H2O ΔH = -483.6 kJ

• If energy is released (product) the reaction is exothermic and ΔH is negative

• If energy is absorbed (reactant) the reaction is endothermic and ΔH is positive

Entropy

• Entropy is a measure of randomness, tendency toward disorder• Symbol is ΔS• More disorder = more entropy• If reaction leads to more disorder, the

entropy change (ΔS) is positive, if it becomes more ordered, ΔS is negative

• Example: melting ice, condensing water, cleaning your room (+,-,-)

Free Energy (ΔG)• Free energy combines enthalpy and

entropy to measure the spontanaeity of a reaction

• Gibbs Free Energy Equation:ΔG = ΔH - T ΔS (T is in Kelvin: +273 to ºC)

• If ΔG is negative, reaction is spontaneous

• If ΔG is positive, reaction is NOT spontaneous, but would be spontaneous in the reverse direction

ExampleFind ΔG for the reaction:

NH4Cl(s) → NH3(g) + HCl(g)

Using the following data:ΔH = 176 kJ, ΔS = 285 J/K, T = 25ºC

Solution: (Change to kJ and K) ΔG = 176 kJ – (298 K)(.285 kJ/K) ΔG = 176 kJ – 84.9 kJ

= 91 kJ NOT spontaneous

Comparison of Signs

ΔH ΔS ΔGSpontaneous?

- + - ALWAYS spont. + - + NEVER spont.

- - - / + Spont. at low T + + - / + Spont. at high T

Reaction Mechanisms

• Step-by-step sequence that occurs to create the products

• Intermediates may form that do not appear in overall reaction – they are used up in another step

• Homogeneous reaction: all reactants in same phase

• Heterogeneous reaction: reactants in different phases

• Rate-determining step: slowest step of reaction mechanism

Activation Energy

• Minimum energy to make the reaction go (form activated complex which allows reaction to proceed)

• Reaction needs:• Enough energy• Proper orientation of molecules –

must hit each other at correct spot

Energy Diagrams

Exothermic/Endothermic

Energy Example

1. Calculate the ΔH.20 kJ – 40 kJ = -20 kJ

2. Calculate the Ea.

100 kJ – 40 kJ = 60 kJ

3. Calculate the Ea‘.

100 kJ – 20 kJ = 80 kJ

Reaction Rate

• Rate can be defined in terms of molar concentration (M) for the disappearance of a reactant or the appearance of a product

• Concentration shown as:[HCl] = 0.1 means the molar concentration of HCl is 0.1 M

Factors Affecting Reaction Rate

• Nature of reactants• Concentration• Temperature• Catalysts

Nature of Reactants

• Ionic – almost instantaneous• Molecular – slower (bonds must

break and reform)• Surface area – rate increases with

greater surface area

Concentration

• Measured in molarity [A]• Increasing the concentration of

reactants increases the rate• Rate law:

Rate = k[A]m[B]n

The exponents m and n must be determined experimentally

Temperature

• Increasing the temperature gives more collisions between molecules

• This leads to the formation of more activated complexes and this causes the rate to increase

↑ T → ↑ collisions → ↑ complexes → ↑ rate

Catalysts• Catalyst – increase reaction rate without

being used up• Lower the activation energy Animation

• Heterogeneous – not in same phase as reactants, provides surface to give more effective collisions Catalytic Converter

• Homogeneous – In same phase as reactants, makes different activated complex, returns to original form at end of reaction

• Demo: Catalysts