22. Reaction Kinetics Flashcards
Define ‘rate of reaction’.
Change in concentration / time taken.
- units: mol/dm3 (min or h in slow reactions).
- this method gives an average rate, as the rate changes as [R] decreases.
How can you obtain a more accurate value for the rate of reaction?
Use shorter time intervals (Δt = 0) and draw tangents at different [R] values.
How can you monitor the rate of a very fast reaction?
Stopped-flow spectrophotometry.
- Small volume of reactant into a mixing chamber at very high speed.
- Reacting mixture moved into observation cell where the rate is monitored by the transmission of UV radiation through the sample.
- Graph automatically generated (rate against time).
Define ‘rate equation’.
Rate = k[A]m [B]n
- equation relating the initial rate of reaction to the concentrations of reactants involved.
- it can only be found experimentally.
Define ‘k’.
Rate constant - a proportionality constant relating the rate of a chemical reaction at a given temperature, to the concentration of the reactants involved.
Define ‘order of reaction’.
The power to which a reactant’s concentration is raised in the rate equation. It can be 0, 1, 2 or 3, but is rarely higher. A fractional order indicates the presence of a free radical.
How would you find a rate equation experimentally?
Vary the concentrations of one reactant whilst keeping the other(s) constant. Do this for each respective reactant and deduce the effect of each on the rate of reaction.
How would you calculate the units of ‘k’?
- Write the rate equation.
- Rearrange in terms of k.
- Substitute the units. Don’t forget s-1 for the rate.
- Solve to find the units of k.
Name three methods to deduce the order of reaction.
- Graphing rate against [R]
- Graphing [R] against time
- Successive half-lives (from [R] against time).
Describe the method of ‘graphing rate against [R]’.
- Zero-order = horizontal line
- First-order = linear increase (directly proportional)
- Second-order = upward curve (proportional to the square, cube etc).
Describe the method of ‘graphing [R] against time’.
- Zero-order = linear decrease (rate = gradient, constant rate of decline).
- First-order = shallow curve
- Second-order = deeper curve with longer tail.
Define ‘half-life’.
The time taken for [R] to decrease to half its original value. Independent of the original [R].
Describe the ‘successive half-lives (from [R] against time)’ method.
- Zero-order = successive half-lives decrease with time
- First-order = relatively constant
- Second-order = increase with time.
How would you find ‘k’ from the initial [R] and rate?
- Rearrange rate equation in terms of k.
- Substitute the values and units, solve.
How would you find ‘k’ from half-life data?
For a FIRST-ORDER reaction:
k = 0.693 / half-life (s)
How would you find the order of a reaction from monitoring the course of the reaction?
- Plot a graph of [R] or [P] against time.
- Draw tangents at different concentrations to find the rate.
- Plot a graph of rate against [R] or [P].
- Determine the overall order from the shape of the curve.
How would you find the order of a reaction from initial rate data?
- Conduct several experiments, varying initial [R].
- Find the initial rates by drawing a tangent at the start of each curve OR by measuring [R] after the reaction starts.
- For each [R], plot a graph of initial rate against [R].
Define ‘rate-determining step’.
The slowest step in a reaction, on which the overall rate depends. Reactants in this step will appear in the rate equation. They can be either pure reactants, or intermediates formed from reactants.
What is the reaction called based on the number of species in the rate equation?
One = unimolecular, two = bimolecular, three = tri.
- trimolecular is rare because it is highly unlikely for three particles to collide simultaneously, and in the correct orientation.
Define ‘homogeneous catalyst’.
A catalyst in the same phase as the reaction mixture.
Define ‘heterogeneous catalyst’.
A catalyst in a different phase to the reaction mixture.
How does homogeneous catalysis work?
Involves a change in oxidation states of the ions involved in catalysis. Ions of transition elements are useful due to their variable oxidation states.
Describe the iodine-peroxodiosulfate reaction (homogeneous).
S2O8 2- + 2I - -> 2SO4 2- + I2
- Both ions are negatively charged so there is repulsion.
- Fe3+ catalyst is attracted to the negative charge.
- 2Fe 3+ + 2I - -> 2Fe 2+ + I2 (reduction)
- 2Fe 2+ + S2O8 2- -> 2Fe 3+ + 2S2O4 2- (oxidation)
- The standard electrode potentials for each of these reactions must be between those involving the reactants.
Describe the formation of acid rain (homogeneous).
SO2 + 1/2 O2 -> SO3 (oxidation)
- SO2 + NO2 -> SO3 + NO
- NO + 1/2 O2 -> NO2
