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Chemistry Sem 2 > Reaction Kinetics > Flashcards

Reaction Kinetics Flashcards

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1
Q

Reaction Kinetics

A

Study of the rates of chemical reactions, including the factors that affect them and the mechanisms by which the reactions occur

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2
Q

Rate of reaction

A
  • 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
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3
Q

Instantaneous rate

A

Rate at a particular time → gradient of concentration-time graph

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4
Q

Initial rate

A
  • 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

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5
Q

Average rate

A

Change in concentration of a reactant/product over that time interval

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6
Q

Rate equation/rate law

A
  • 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
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7
Q

Rate constant/velocity constant

A
  • 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
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8
Q

Order of reaction with respect to a reactant

A
  • 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
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9
Q

Overall order of a reaction

A

Sum of the powers of the concentration terms in the rate equation

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10
Q

Half-life (t½)

A

Time taken for the concentration of a reactant to decrease to half its initial value

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11
Q

First-order reaction

A
  • 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
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12
Q

Second-order reaction

A
  • 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
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13
Q

Zero-order reaction

A
  • Reaction rate independent of the concentration of the reactant
  • Rate equation: rate = k
  • Units of k: mol dm⁻³ s⁻¹
  • t½ not constant
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14
Q

Pseudo-order reactions (3)

A
  1. Presence of large excess of reactant
  2. Reactant is also the solvent
  3. Presence of a catalyst
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15
Q

Presence of large excess of reactant

A
  • E.g. rate = k[A][B]
  • If initial [A]&raquo_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
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16
Q

Reactant is also the solvent

A
  • 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]
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17
Q

Presence of a catalyst

A
  • 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
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18
Q

Finding order of reaction from initial rates data

A
  • 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
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19
Q

Finding order of reaction from concentration-time graphs

A
  • 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
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20
Q

Experimental techniques to measure reaction rates

A
  1. Method of isolation → continuous measurement

2. Method of initial rates → one reading from each experiment

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21
Q

Method of initial rates

A
  1. Physical methods e.g. colorimetry

2. Clock reaction

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22
Q

Clock reaction

A
  • Certain reactions accompanied by prominent visual changes e.g. forming precipitate/obvious change in colour
  • Measure time taken for visual change to occur
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23
Q

Reaction between thiosulfate ions and hydrogen ions

A
  • 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
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24
Q

Why is the volume of water varied in each experiment?

A
  • Keep total volume of reaction mixture constant

- Initial concentration of each reactant is directly proportional to its volume used

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25
Q

What is the relationship between the rate of reaction and the time t taken for the cross to be obscured?

A
  • rate ∝ 1/t
  • Inversely proportional since amount of sulfur produced to obscure printed material is kept the same for all the experiments
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26
Q

Reaction between hydrogen peroxide and iodide ions in acidic medium

A
  • 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
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27
Q

Method of isolation

A
  1. Chemical method → sampling, quenching and titration

2. Physical methods → continuous measurement of physical property

28
Q

Sampling, quenching and titrimetric analysis

A
  1. Starting the reaction
  2. Sampling → use pipette → aliquot portion
  3. Quenching followed by titration → quenching agent/ice-cold water
  4. Plot graph → determine order of reaction/instantaneous rate (gradient)
29
Q

Measuring colour intensity at regular time intervals

A
  • 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
30
Q

Measuring electrical conductivity at regular time intervals

A
  • Number and types of ions present in a solution affect its electrical conductivity
  • Immerse 2 inert electrodes to measure
31
Q

Measuring volume of gas produced at regular time intervals

A
  • Use gas syringe
  • Graph of V against time
  • Graph of (V∞ - Vt) against time
  • (V∞ - Vt) ∝ reactant
  • Determine half-life, instantaneous rate
32
Q

Measuring mass of reaction mixture at regular time intervals

A

Reaction whereby a gas is produced and allowed to escape

33
Q

Measuring pressure at regular time intervals

A
  • If reaction involves a change in the number of moles of gas
  • Manometer
  • Constant-temperature water bath
34
Q

Reaction mechanism

A
  • Sequence of bond-making and bond-breaking steps that occur during the conversion of reactants to products
  • Can be used to deduce rate equation
35
Q

Single-step reactions

A
  • 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
36
Q

Multi-step reactions

A
  • Takes place by 2 or more steps that are likely to proceed at different rates
  • Rate of overall reaction depends on rate-determining steps
37
Q

Rate-determining step

A

Slowest step in the reaction mechanism of a multi-step reaction and determines the overall reaction rate → steps with highest activation energy

38
Q

Proposed reaction mechanism must meet the following criteria: (3)

A
  1. Must agree with stoichiometric equation
  2. Must be consistent with the observed kinetics
  3. Must include intermediates detected during reaction
39
Q

Deducing rate equation from a given reaction mechanism

A
  • Consider the reactants that participate in the slow step and all the fast steps before it
  • Intermediates should not appear
40
Q

Theories of reaction rate (2)

A
  1. Collision theory

2. Transition state theory

41
Q

Collision theory

A

Reactant particles must collide:

  1. in order to react → effective collisions
  2. in favourable orientation → correct collision geometry
  3. with a certain min amt of energy (Ea)
42
Q

Activation energy

A

Minimum amount of energy that the reactant particles must possess before their collisions can result in a reaction

43
Q

Kinetic stability vs thermodynamic stability

A
  • Kinetic feasibility → depends on Ea → whether it proceeds at observable rate
  • Thermodynamic feasibility → depends on ΔG⦵ → whether it can occur
44
Q

Maxwell-Boltzmann distribution curves (2)

A
  1. Of molecular speeds

2. Of kinetic energies

45
Q

Maxwell-Boltzmann distribution of molecular speeds

A
  • 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
46
Q

Maxwell-Boltzmann distribution of kinetic energies

A
  • 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)
47
Q

Transition state theory

A
  • 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
48
Q

Factors affecting rate of reaction (4)

A
  1. Physical states of reactants
  2. Concentration of reactants
  3. Temperature
  4. Catalyst
49
Q

Physical states of the reactants

A
  • 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)
50
Q

Concentration of reactants

A
  • As [reactants] increase → reactant particles come closer together → frequency of effective collisions ↑
  • Partial pressures of gaseous reactants (p ∝ c)
51
Q

Temperature

A
  • 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
52
Q

Catalyst

A
  • 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
53
Q

Arrhenius equation

A
  • Temperature and Ea dependence of the rate constant k is shown
  • When T↑ or Ea↓, k↑
54
Q

Inhibitor

A

Substance which decreases the rate of a chemical reaction

55
Q

Promoter

A

Substance which enhances the efficiency of a catalyst

56
Q

Catalyst poison

A

Substance which inhibits the effectiveness of a catalyst

57
Q

Types of catalysis

A
  1. Homogeneous → catalyst and reactants in same phase

2. Heterogeneous → catalyst and reactants in different phase

58
Q

Why are transition metal ions effective homogeneous catalysts?

A
  • Can exist in diff oxidation states

- Can undergo conversion from one oxidation state to another relatively easily

59
Q

Heterogeneous catalysis (5)

A
  1. Diffusion
  2. Adsorption → onto active sites on catalyst surface
  3. Reaction
  4. Desorption
  5. Diffusion
60
Q

Adsorption (2)

A
  1. Weakens covalent bonds in reactant molecules → lower Ea
  2. ↑ concentration of reactant molecules at catalyst surface, allows them to come into close contact with proper orientation for reaction
61
Q

Catalytic converters

A
  • 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
62
Q

Autocatalysis

A
  • Product of reaction acts as catalyst for the reaction

- Test by adding autocatalyst at the start of the reaction

63
Q

Enzymes

A

Proteins which catalyse the chemical reactions in living systems → biological catalysts

64
Q

Properties of enzymes

A
  1. Nature and size → globular proteins with active sites
  2. Efficiency → required in very small amounts, very effective
  3. Specificity → very specific to particular reaction/type of reaction
  4. Temperature → most effective at body T, 37°C
  5. Sensitivity to pH → narrow pH range
65
Q

Enzymatic action

A
  • Enzyme-substrate complex
  • Very specific
  • Lock and key model
  • Induced-fit model
66
Q

Factors affecting rate of enzyme-catalysed reaction (4)

A
  1. Temperature
  2. pH
  3. Concentration of enzyme
  4. Concentration of substrate: at low [substrate] → 1st order, at high [substrate] → 0 order (similar for heterogenous catalyst)

Chemistry Sem 2 (5 decks)

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