Reaction kinetics

Factors affecting rate of reaction

Temperature
Concentration
Catalyst
Pressure
Total surface area

Definition of rate of reaction

Rate of reaction can be defined as the rate of change of concentration but it can also be calculated with the change in temperature, pressure, etc..

The formula for concentration is MOL/ VOLUME

Sampling

-Taking small samples of the reaction at different time intervals

-Analyse the data and calculate the average

-we can quench the reaction by cooling the sample with ice

chemical

Continuous

Physical property of the mixture is monitored over a period of time

Using colorimetry, conductivity meter or by measuring changes in volume or pressure of gas

physical

Colorimetry: colour of iodine fades in this reaction from purple to brown

Conductivity meter: electrical conductivity increases in this reaction because ions that can conduct electricity are present

Changes in pressure and volume can be measured with a pressure meter due to the increased amount of gases in the products of this reaction

Small changes of volume can be measured using a dilatometer

Graphical calculation of reaction

To calculate the gradient we draw a tangent forming a triangle and use rise/run to calculate the gradient. The gradient nearest to the origin is the initial gradient.

We can see that the gradient decreases as the time goes on, this is due to cyclopropane being the limiting reagent and the decreases in concentration resulting in less collisions of the reactants.

Graph of decreasing reactant

We can assume that:

Rate is directly proportional to the concentration
cyclopropane is doubled and rate is also doubled
cyclopropane falls to 1/3 if the rate falls by 1/3 since it is directly proportional

Rate equations

Rate of the reaction= k[cyclopropene]
k is a constant
rate of reaction can only be determined through experiments
Power is not from stoichiometric equation
concentration is not made out of reactants only and can sometimes not include a reactant

Order of reaction

The order of reaction in respect to a particular reactant is the power which the concentration of that reactant is raised in the rate equation

We can say in this reaction that:

First order with respect H2
Second order with respect with NO
Thus, overall order of reaction is the third order of reaction as the sum is 1+2

Unit of k

K can be written as rate/ concentrations and the unit depends on a and b which is the power of the concentration

Rate is mol dm^-3 s^-1

Concentration is mol dm^-3

reaction rate-concentration graph

If [ ]=2

zero order

A horizontal line

We can say that concentration is zero order with respect to [ ] so y=k

first order

Constant gradient

we can say that concentration is first order with respect with [ ] so y=2k

second order

Gradient increases

We can say that concentration is second order with respect with [ ] so y=4k

Concentration-time graph

zero order

Constant gradient

First order

Steep gradient

Second order

Steeper gradient

deeper and have a longer tail

Half life

Half-life, t1/2 is time taken for the concentration of a reactant to fall to half of its original value.

Half lives decreases over time

Constant half life

Half lives increases over time

Calculation k from half life

Formula:

Rate determining step

The overall rate of reaction depends on the slowest step
If a substance does not appear in the overall rate equation, it does not partake in the rate determining step

The rate determining step in this reaction is the slow step. The slow step involves one molecule of propanone and one hydroxide ion, so only these two species appear in the rate of reaction. The reaction is second overall, the first order with respect to propanone and first order with respect to hydroxide ions.

Catalysis

Homogeneous catalysis occurs when the catalyst is in the same phase in the reaction mixture such as hydrolysis of esters with hydrogen ions

Heterogeneous catalysis occurs when the catalyst is in a different phase in the reaction mixture such as decomposition of hydrogen peroxide catalysed by manganese oxide.

Homogeneous catalyst

Involves changes in the oxidation number
Catalyst is I-

Catalyst first forms an intermediate in the first step reaction and reacts again in the second step reaction forming the catalyst again so it is not used up in the overall reaction by providing an alternate pathway with lower activation energy

Peroxodisulfate ions, S2O8^2- oxidise iodide ions to iodine and the reaction is very slow

The peroxodisulfate and iodide ions both have a negative charge. Thus, in order to collide and react, these ions need a large amount of energy to overcome the repulsive forces when like charges collide with each other

Fe^3+ ions catalyse this reaction involving two reactions

Persulfate ions, arrow facing right since more positive

Iodine, arrow facing left since more negative

recall electrochemistry

Heterogeneous catalyst

Often involves gases molecules reacting at the surface of a solid catalyst
The mechanism of the catalyst can be explained using the theory of adsorbtion
•Adsorb=bond to the surface of the substance
Absorb=move right into the substance

Theory of adsorbtion in haber process

Nitrogen gas and hydrogen gas first diffuse to the surface of the iron
Adsorbtion occurs and the reactant molecules are chemically adsorbed to the surface of the iron. The bonds formed on the surface of the iron are:

-Strong enough to weaken the covalent bonds within the nitrogen and hydrogen molecules so the molecules can react together

-Weak enough to break and allow the products to leave the surface

The adsorbed nitrogen and hydrogen react together on the surface of the iron to form ammonia
Desorbtion then occurs because the bonds between the ammonia and the surface of the iron weaken and are eventually broken.
Diffusion occurs and ammonia diffuses away from the surface of the iron

Page updated

Google Sites

Report abuse