Reaction Kinetics

• Rate of reaction is the change in concentration of products/reactants per unit time

R = Unit: mol dm^-3 s^-1

 

 

Order of Reaction

• For zero order, rate of reaction not dependant on conc. of that reactant but it is need for completion of reaction. Reaction continues until all of that reactant is used up.
• Recognize shapes above in graphs given to work out order of reactant and construct rate equation. Calculate kk using data from graph(s).

Constructing rate equation using conc. data:

Using following data from the experiments construct a rate equation and calculate rate constant.

Solution

Relationship of Temperature and k

• Increasing the temperature, increases the value of k
• When temp. is increased, the k.e. of reacting molecules increases resulting in more successful collisions
• Reactants change faster to products therefore conc. of reactants decreases.
• Using fraction above, numerator increases and denominator gets smaller therefore k increases.

 

Reaction Mechanism

Rate-determining step

•  the slowest step in a reaction mechanism
• many reactions consist of a series of separate stages
• each of these stages has its own rate and hence its own rate constant
• the slowest step is known as the rate determining step

 
 

In a multistep reaction:

• The rate of reaction is dependent on the slowest step that needs the highest activation energy.
• Rate equation includes only reactants that are present in the rate-determining step.
• The orders with respect to the reactants are the moles of the reactants in the rate determining step

Example 1

H₂O₂ + 2H3O+ + 2I¯ ——> I₂ + 4H₂O

Step 1 H₂O₂ + I¯ ——> IO¯ + H₂O slow

Step 2 IO¯ + H₃0⁺ ——> HIO + H₂O fast

Step 3 HIO + H₃0⁺ + I¯ ——> I₂ + 2H₂O fast

The rate determining step is STEP 1 as it is the slowest

 

Example 2

The reaction 2N₂O₅ —> 4NO₂ + O₂

Step 1     N₂O₅ —> NO₂ + NO₃
slow

Step 2     NO₂ + NO₃ —> NO + NO₂ + O₂
fast

Step 3     NO + NO₃ —> 2NO₂
fast

 

The rate determining step is STEP 1 as it is the slowest. The rate equation for the reaction is rate = k [N₂O₅]

 

Measuring Reaction Rates

Sampling method that involves taking small sample of a reaction mixture at various times and then carrying out chemical analysis on sample.

Chemical Analysis

C₄H₉Br + OH^– C₄H₉OH + Br^–

• Sampled removed at various times and quenched – stopping/slowing down reaction (e.g. cooling in ice)
• OH^– conc. can be found using titration
• Plot graph and calculate rate of reaction

 
 

Continuous

Method that involves monitoring a physical property over a period of time

Change in Volume of Gas Produced

Mg + 2HCl MgCl₂ + H₂

• Measure change in volume of gas using a gas syringe
• Take down readings at regular intervals
• Plot graph and calculate rate of reaction

 

Changes in Colour

CH₃COCH₃ + I₂ CH₃COCH₂I + HI

• I₂ starts brown, fades through orange to yellow to colourless as iodine used up.
• Colorimeter measures amount of light absorbed as it passes through solution; recorded as absorbance.
• ‎Before experiment, create calibration curve by finding absorbance of different conc. I₂ and plot a graph of concentration against absorbance
• During experiment, measure absorbance from meter at regular intervals, and use calibration curve to convert values into concentrations

 

Catalysts

They work by providing an alternative reaction pathway with a lower Activation Energy

Lowering Ea results in there being a greater area under the curve showing that more molecules have energies in excess of the Activation Energy. Catalysts remain chemically unchanged at the end of the reaction.

Types

1. Homogeneous Catalysts
2. Heterogeneous Catalysts

 

They do not affect the position of the equilibrium but they do affect the rate at which equilibrium is attained.

 

Homogeneous Catalysis

• Catalyst and reactants in same physical state
• Catalyst takes part in reaction:
  • Forms intermediate with reactant
  • Intermediate breaks down giving product
• Rate of reaction dependant on conc. of catalyst

Examples of Homogenous Catalysis

 

Iodine-Peroxydisulfate Reaction

S₂O₈^2- + 2I^– > 2SO₄^2- + I₂

• Mechanism of the catalysed reaction:
  • Reaction 1: reduction of Fe³⁺ to Fe²⁺ ions by I^– ions

    2I^– + 2Fe³⁺ > I₂ + 2Fe²⁺

  • Reaction 2: oxidation of Fe²⁺ to Fe³⁺ ions by S₂O₈^2- ions
  • S₂O₈^2- + 2Fe²⁺ > 2SO₄^2- + 2Fe³⁺

Oxides of Nitrogen and Acid Rain

‎

• SO₃ can then react with water to form H₂SO₄

 

Heterogeneous Catalysis

• Catalyst and reactants in different physical state
• Catalyst provides surface on which reaction occurs
  • Reactant particles adsorbed on surface on collision with catalyst
  • Molecular rearrangement occurs – bonds in reactants break and new bonds in product formed
  • Product molecules desorbed from surface
• Rate of reaction dependant on surface area of catalyst

 

Catalysis is thought to work in three stages

1. Adsorption – formation of bonds with surface weakens bonds in gas molecules makes a subsequent reaction easier
2. Reaction – adsorbed gases may be held on the surface increases chances of favourable collisions
3. Desorption – the products are then released from the active sites (catalyst (diffusion away from surface)

 

HARD
Hetero = Adsorption + Reaction + Desorption

 

Examples of catalysts

• Metals
  • Ni, Pt hydrogenation
  • reactions Fe Haber Process
• Oxides
  • AlO₃ dehydration
  • reactions V₂O₅ Contact Process

   

  

• Honeycomb structure containing small beads coated with Pt, Pd and Rh

 

Possible catalytic process

• Adsorption of NO_x and CO
• Weakening of covalent bonds within NO_x and CO
• Formation of new bonds between
  • Adjacent N atoms form N₂
  • CO and O atoms form CO₂
• Desorption of N₂ and CO₂