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Elementary Reactions & Collisions — AP CHEMISTRY Revision Notes & Topical Questions

5 min read 10 parts Lesson 1 of 14 Unit quiz

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Elementary Reactions & Collisions — AP CHEMISTRY notes: Describe elementary reactions, the collision model, energy profiles and the Arrhenius equation.A balanced equation gives the overall change but not how reactants become products.Elementary reaction = a single molecular event…

  • Unimolecular
  • Bimolecular
  • Termolecular
  • Rate constant
  • Reverse Ea
  • Taking logarithms, ln k

Section 1 of 10Elementary reactions

A balanced equation gives the overall change but not how reactants become products.
Elementary reaction
Elementary reaction a single molecular event involving the collision of small particles.
A set of elementary reactions that sums to the overall equation is a reaction mechanism.
Example
  • CO + NO2 → CO2 + NO occurs in two steps at low temperature (NO2 + NO2 → NO3 + NO, then NO3 + CO → CO2 + NO2).
NO3 is an intermediate because it appears in the steps but not in the overall equation.

Section 2 of 10Molecularity

Unimolecular
  • Unimolecular one reactant molecule.
Bimolecular
  • Bimolecular two reactant molecules.
Termolecular
  • Termolecular three reactant molecules, which is rare because simultaneous collision with correct orientation and energy is unlikely.
Example
  • CH3Br + OH− → CH3OH + Br− has a bimolecular first step and a unimolecular second step.

Section 3 of 10Rate laws of elementary steps

1
The order of each reactant in an elementary reaction equals its stoichiometric coefficient.
  • A → products gives Rate = k[A].
  • A + B → products gives Rate = k[A][B].
  • 2A → products gives Rate = k[A]2.
2
Rate laws are written from coefficients only for elementary steps, never from overall balanced equations.

Section 4 of 10Collision model

Molecules must collide to react, so more collisions per second give a faster reaction.
Higher concentration, higher temperature and greater surface area of solids all increase collision frequency.
Effective collisions have two requirements.
  • Energy at or above the activation energy (Ea).
  • Correct orientation, such as the carbon of CO striking an oxygen of NO2.
Rate constant
  • Rate constant k = p × Z × f.
  • Z is the collision frequency, p is the steric (orientation) factor and f is the fraction of collisions with energy of at least Ea.

Section 5 of 10Maxwell-Boltzmann distribution

Maxwell-Boltzmann distribution
The distribution shows the fraction of molecules with enough energy to break reactant bonds.
Raising the temperature shifts the peak to the right and flattens the curve, so a larger fraction has energy at or above Ea.
The area under the curve, the total number of particles, stays the same.
ChangeCurveFraction at or above EaRate
Higher temperatureFlatter, shifts rightIncreases greatlyIncreases
Lower temperatureTaller, shifts leftDecreasesDecreases
CatalystUnchanged; the Ea line moves leftIncreasesIncreases

Section 6 of 10Transition state theory and energy profiles

Collision theory does not explain the role of activation energy; transition state theory does.
The activated complex is an unstable, high-energy species in which bonds are partly broken and new bonds partly formed.
The activation energy is far smaller than the energy needed to break bonds completely, as for CO + NO2 where Ea is 134 kJ/mol.
An energy profile plots potential energy against reaction progress.
  • The difference between reactants and products is ΔH; products lower than reactants means exothermic.
  • The difference between the maximum and the reactants is Ea of the forward reaction.
  • The top of the barrier is the transition state.
Reverse Ea
Reverse Ea forward Ea − ΔH, so for CO + NO2 it is 134 − (−226) = 360 kJ/mol.
Rate depends on Ea, not on ΔH, so a highly exothermic reaction can be slow and an endothermic one fast.

Section 7 of 10The Arrhenius equation

Z is only slightly affected by temperature, whereas the fraction f = e−Ea/RT is strongly affected.
k
Ae−Ea/RT, where A is the frequency factor combining Z and p.
A higher Ea gives a smaller k because fewer molecules have the required energy.
Taking logarithms, ln k
Taking logarithms, ln k ln A − Ea/RT.
What the symbols mean: k = rate constant · A = pre-exponential (frequency) factor · Ea = activation energy (J mol−1) · R = 8.31 J mol−1 K−1 · T = temperature (K)
A plot of ln k against 1/T is a straight line with a negative slope of −Ea/R.
Temperature affects k but not the orders of reaction.

Section 8 of 10More from the specification

More from the specification
Kinetic energy
  • The energy of motion of a particle.
Maxwell-Boltzmann distribution
  • A graph showing how particle energies are spread at a given temperature.
Carbon monoxide
  • A toxic gas that binds to haemoglobin.

Section 9 of 10Key terms

TermMeaning
Elementary reactionSingle molecular event in a mechanism.
MolecularityNumber of reactant molecules in an elementary step.
Activation energyMinimum energy colliding molecules need to react.
Activated complexUnstable high-energy species at the transition state.
Arrhenius equationk = A e^(−Ea/RT).
Steric factorFraction of collisions with the correct orientation.

Section 10 of 10Exam technique

1
Writing a rate law for an overall reaction from its coefficients — only elementary steps follow coefficients
2
Saying a catalyst changes ΔH — it lowers Ea only
3
Saying a higher temperature increases the area under the Maxwell-Boltzmann curve — the area stays the same
4
Thinking a very exothermic reaction must be fast

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