Reaction Kinetics Flashcards
Reaction Kinetics
Study of the rates of chemical reactions, including the factors that affect them and the mechanisms by which the reactions occur
Rate of reaction
- Change in concentration of a particular reactant or product per unit time
- mol dm⁻³ s/min/h⁻¹
- Positive quantity
- Can be obtained from concentration-time graphs
Instantaneous rate
Rate at a particular time → gradient of concentration-time graph
Initial rate
- Instantaneous rate at time t=0
- Can be approximated by average rate provided 1) time interval small enough 2) time interval start from t=0
Average rate
Change in concentration of a reactant/product over that time interval
Rate equation/rate law
- A mathematical expression that relates the rate of reaction to the concentration of each reactant raised to the appropriate power. It shows the exact dependence of the reaction rate on the concentrations of all the reactants
- E.g. rate = k[A][B]
- Can only be obtained by experiment
Rate constant/velocity constant
- The constant of proportionality in the rate equation of the reaction
- E.g. rate ∝ [A] rate = k[A]
- Units depend on overall order of reaction
- Constant for a given reaction at a particular temperature, ↑ with temperature or in presence of catalyst
Order of reaction with respect to a reactant
- The power to which the concentration of that reactant is raised in the rate equation
- Exact dependence of rate on concentration of a given reactant
- Found experimentally
- Can be whole number, fraction, positive or negative
- Can help us to work out the reaction mechanism
Overall order of a reaction
Sum of the powers of the concentration terms in the rate equation
Half-life (t½)
Time taken for the concentration of a reactant to decrease to half its initial value
First-order reaction
- Reaction rate directly proportional to the concentration of a single reactant raised to the power of one
- Rate equation: rate = k[A]
- Units of k: s⁻¹
- t½ is constant at constant temperature, independent of initial concentration → ln2/k
- t½ can be found by plotting [reaction]/[product] against time
Second-order reaction
- Reaction rate proportional to the product of the concentrations of 2 reactants or to the concentration of a single reactant raised to the power of 2
- Rate equation: rate = k[A][B] or k[A]²
- Units of k: mol⁻¹ dm³ s⁻¹
- t½ not constant
Zero-order reaction
- Reaction rate independent of the concentration of the reactant
- Rate equation: rate = k
- Units of k: mol dm⁻³ s⁻¹
- t½ not constant
Pseudo-order reactions (3)
- Presence of large excess of reactant
- Reactant is also the solvent
- Presence of a catalyst
Presence of large excess of reactant
- E.g. rate = k[A][B]
- If initial [A]»_space; initial [B] → [A] will hardly change during reaction relative to the change in [B] → effectively constant
- rate = k’[B], where k’ = k[A]
- Reaction exhibits pseudo-first-order kinetics
- k’ → pseudo-first-order rate constant
Reactant is also the solvent
- E.g. rate = k[A][H₂O], where A (aq)
- As water is used as solvent → concentration remains essentially constant
- rate = k’[A], where k’ = k[H₂O]
Presence of a catalyst
- E.g. rate = k[A][Catalyst]
- Catalyst increases rate of reaction but not consumed by the reaction → regenerated → regarded as essentially constant
- rate = k’[A], where k’ = k[Catalyst] = constant
Finding order of reaction from initial rates data
- Compare initial rates of reactions at different known initial concentrations
- Compare experiments 1 and 2, when initial [X] is the same and initial [Y] x 2, initial rate x 2
- rate ∝ [Y]
- Order of reaction with respect to Y is 1
Finding order of reaction from concentration-time graphs
- From the graph,
- 1st t½ = time taken for [A] to drop from X to 0.5 X = 5 min
- 2nd t½ = time taken for [A] to drop from Y to 0.5 Y = 5 min
- Since t½ is (relatively) constant, the reaction is first order with respect to A
- t½ = ln2/k, k = ln2/t½
Experimental techniques to measure reaction rates
- Method of isolation → continuous measurement
2. Method of initial rates → one reading from each experiment
Method of initial rates
- Physical methods e.g. colorimetry
2. Clock reaction
Clock reaction
- Certain reactions accompanied by prominent visual changes e.g. forming precipitate/obvious change in colour
- Measure time taken for visual change to occur
Reaction between thiosulfate ions and hydrogen ions
- S₂O₃²⁻(aq) + 2H⁺(aq) → S(s) + SO₂(g) + H₂O(l)
- Yellow precipitate of sulphur → suspension
- Time taken for fixed amount of sulphur to form → obscure cross
Why is the volume of water varied in each experiment?
- Keep total volume of reaction mixture constant
- Initial concentration of each reactant is directly proportional to its volume used
What is the relationship between the rate of reaction and the time t taken for the cross to be obscured?
- rate ∝ 1/t
- Inversely proportional since amount of sulfur produced to obscure printed material is kept the same for all the experiments
Reaction between hydrogen peroxide and iodide ions in acidic medium
- Reaction 1: H₂O₂(aq) + 2I⁻(aq) + 2H⁺(aq) → I₂(aq) + 2H₂O(l)
- Reaction 2: 2S₂O₃²⁻(aq) + I₂(aq) → S₄O₆²⁻(aq) + 2I⁻(aq)
- Small but constant constant amount of sodium thiosulfate added to reaction mixture
- Since 2 relatively faster than 1, [I₂] effectively 0 as long as there is still some S₂O₃²⁻
- When S₂O₃²⁻ completely used up, 2 stops but 1 continues → I₂ produced in 1 will be present in solution → starch indicator
Method of isolation
- Chemical method → sampling, quenching and titration
2. Physical methods → continuous measurement of physical property
Sampling, quenching and titrimetric analysis
- Starting the reaction
- Sampling → use pipette → aliquot portion
- Quenching followed by titration → quenching agent/ice-cold water
- Plot graph → determine order of reaction/instantaneous rate (gradient)
Measuring colour intensity at regular time intervals
- When one of the reactants/products is coloured
- Concentration ∝ colour intensity
- Calibration curve of solutions of known concentrations
- Colorimetric analysis
- Light source → filter to select wavelengths absorbed by solution → reaction mixture → detector
Measuring electrical conductivity at regular time intervals
- Number and types of ions present in a solution affect its electrical conductivity
- Immerse 2 inert electrodes to measure
Measuring volume of gas produced at regular time intervals
- Use gas syringe
- Graph of V against time
- Graph of (V∞ - Vt) against time
- (V∞ - Vt) ∝ reactant
- Determine half-life, instantaneous rate
Measuring mass of reaction mixture at regular time intervals
Reaction whereby a gas is produced and allowed to escape
Measuring pressure at regular time intervals
- If reaction involves a change in the number of moles of gas
- Manometer
- Constant-temperature water bath
Reaction mechanism
- Sequence of bond-making and bond-breaking steps that occur during the conversion of reactants to products
- Can be used to deduce rate equation
Single-step reactions
- A reaction that takes place in a single step and is termed an elementary reaction → cannot be broken down into simpler steps
- Reaction mechanism identical to stoichiometric equation
- Rate equation can be deduced directly from stoichiometric equation
Multi-step reactions
- Takes place by 2 or more steps that are likely to proceed at different rates
- Rate of overall reaction depends on rate-determining steps
Rate-determining step
Slowest step in the reaction mechanism of a multi-step reaction and determines the overall reaction rate → steps with highest activation energy
Proposed reaction mechanism must meet the following criteria: (3)
- Must agree with stoichiometric equation
- Must be consistent with the observed kinetics
- Must include intermediates detected during reaction
Deducing rate equation from a given reaction mechanism
- Consider the reactants that participate in the slow step and all the fast steps before it
- Intermediates should not appear
Theories of reaction rate (2)
- Collision theory
2. Transition state theory
Collision theory
Reactant particles must collide:
- in order to react → effective collisions
- in favourable orientation → correct collision geometry
- with a certain min amt of energy (Ea)
Activation energy
Minimum amount of energy that the reactant particles must possess before their collisions can result in a reaction
Kinetic stability vs thermodynamic stability
- Kinetic feasibility → depends on Ea → whether it proceeds at observable rate
- Thermodynamic feasibility → depends on ΔG⦵ → whether it can occur
Maxwell-Boltzmann distribution curves (2)
- Of molecular speeds
2. Of kinetic energies
Maxwell-Boltzmann distribution of molecular speeds
- Gas contains large no. of particles in rapid motion
- Each particle has diff speed, collisions change speeds
- Low T → distribution narrow with peak at low speeds
- As T↑ → peak moves to higher speeds and distribution broadens out
Maxwell-Boltzmann distribution of kinetic energies
- KE of particles always changing due to collisions
- Distribution always constant under same conditions
- Area always the same at diff T as total no. of particles remains the same
- As T↑ → maximum of curve displaced to right and takes on smaller value & greater spread of KE (distribution curve broadens)
Transition state theory
- Concerned with what actually happens during a collision
- Given sufficiently energetic collision & correct molecular orientation, reactant species → unstable transition state → potential energy maximum
- Transition state ≠ intermediate → cannot be isolated
- Energy profile diagram → peak is transition state
Factors affecting rate of reaction (4)
- Physical states of reactants
- Concentration of reactants
- Temperature
- Catalyst
Physical states of the reactants
- Reactant particles must mix in order to collide and react
- Solution vs solid
- Molecules (cleavage of covalent bond → higher Ea) vs ions
- Solid form vs finely divided sate (available surface area per unit volume)
Concentration of reactants
- As [reactants] increase → reactant particles come closer together → frequency of effective collisions ↑
- Partial pressures of gaseous reactants (p ∝ c)
Temperature
- For most reactions, rate x 2 for every 10K rise in T
- When T↑, average KE of reactant particles ↑
- Significantly more reactant particles having energy ≥ to Ea
- Increase in effective collisions
- Larger rate constant
Catalyst
- A substance which increases the rate of reaction without itself undergoing any permanent chemical change
- Provides an alternative reaction pathway with lower Ea
- Significantly more reactant molecules have energy ≥ Ea
- ↑ effective collision frequency
- Larger rate constant k
Arrhenius equation
- Temperature and Ea dependence of the rate constant k is shown
- When T↑ or Ea↓, k↑
Inhibitor
Substance which decreases the rate of a chemical reaction
Promoter
Substance which enhances the efficiency of a catalyst
Catalyst poison
Substance which inhibits the effectiveness of a catalyst
Types of catalysis
- Homogeneous → catalyst and reactants in same phase
2. Heterogeneous → catalyst and reactants in different phase
Why are transition metal ions effective homogeneous catalysts?
- Can exist in diff oxidation states
- Can undergo conversion from one oxidation state to another relatively easily
Heterogeneous catalysis (5)
- Diffusion
- Adsorption → onto active sites on catalyst surface
- Reaction
- Desorption
- Diffusion
Adsorption (2)
- Weakens covalent bonds in reactant molecules → lower Ea
- ↑ concentration of reactant molecules at catalyst surface, allows them to come into close contact with proper orientation for reaction
Catalytic converters
- Remove 3 main pollutants (CO, NOₓ and unburnt hydrocarbons)
- Ceramic honeycomb structure coated with platinum (Pt), palladium (Pd) and rhodium (Rh)
- Maximise surface area
- Catalysts deactivated by lead
Autocatalysis
- Product of reaction acts as catalyst for the reaction
- Test by adding autocatalyst at the start of the reaction
Enzymes
Proteins which catalyse the chemical reactions in living systems → biological catalysts
Properties of enzymes
- Nature and size → globular proteins with active sites
- Efficiency → required in very small amounts, very effective
- Specificity → very specific to particular reaction/type of reaction
- Temperature → most effective at body T, 37°C
- Sensitivity to pH → narrow pH range
Enzymatic action
- Enzyme-substrate complex
- Very specific
- Lock and key model
- Induced-fit model
Factors affecting rate of enzyme-catalysed reaction (4)
- Temperature
- pH
- Concentration of enzyme
- Concentration of substrate: at low [substrate] → 1st order, at high [substrate] → 0 order (similar for heterogenous catalyst)
