Finding orders of reaction from graph.
Zero order reaction
 Rate = k [A]0
1.
 Changing the concentration of reactants will not affect the rate of
reaction.
2. First-order reaction
Rate = k [A]
 Doubling [reactant] will double the rate of reaction.
 Half-life, t1/2  is the time for the [reactant] to be halved.
 The [reactant] vs. time graph is a curve with a constant half-life.
 the half-life of first-order reaction is independent of concentration.
 For any other order, the half-life is not constant.
3. Second-order reaction.
 Rate = k [A]2
t3 > t2 > t1
 Doubling [reactant] will increase the rate of reaction four
fold.
Determination of a rate equation
Using Half-Life
11.3.2.7 be able to investigate how the rate of a
reaction changes with changing conditions
11.3.2.8 be able to carry out simple calculations
of rates and rate constants
Success criteria
• represents the rate of chemical reaction using
the equation of speed;
• knows the general equation of speed and can
determine the order of reaction graphically
and the method of initial velocities;
• characterizes the meaning of the rate
constant.
Terminology
rate of reaction
concentration
half-life
order of reaction
реакция
жылдамдығы
концентрациясы
жартылай ыдырау
мерзімі
реакция тәртібі
cкорость реакции
концентрация
период
полураспада
порядок реакции
Calculations using half-life of a first-order reaction.
 Example : First-order reaction of catalytic decomposition of H2O2.
 2H2O2(aq)  2H2O(l) + O2(g)
1. Calculate the half-life of this reaction.
2. Show that the reaction is a first-order reaction.
3. How long will it take for the concentration of H2O2 to
0.025 mol dm-3?
become
2. Plot a graph of concentration of R (y-axis) against time (x-axis). Use the
graph to determine the order of the reaction, and give an average value
for the half-life of the reaction.
Time / s
0
150
300
450
750
1150
1450
1750
[R]/mol dm-3
1.000
0.812
0.659
0.535
0.353
0.202
0.133
0.088
Calculation involving experimental & graphical
methods.
 E.g : RBr + NaOH  ROH + Na+Br The order of reaction may be deduced by plotting graphs of
[RBr] vs. time.
 #CStoh 88 graph
Individual work
• read the book (p.332)
Graphical determination of the
reaction order by half-life
Catalysis
Homogeneous
Catalyst
Heterogeneous
1. Homogeneous Catalysis
 Both catalyst and reactant are in the same physical state.
 Many homogeneous catalyst are transition metal ion, which act
as a catalyst by varying their oxidation states.
 The homogeneous catalyst takes part in the reaction  forms
intermediate
 Increase concentration of catalyst will increase the speed of ratedetermining step.
 E.g 1 : Oxidation of I-(aq) by S2O82-(aq)
 Uncatalysed reaction :
 S2O82-(aq) + 2I-(aq)  2SO42-(aq) + I2(aq)
 Reaction is kinetically slow due to high Ea.
 Reaction catalysed by Fe3+ ions :
 Step 1 : 2Fe3+ + 2I-  2Fe2+ + I2
 Step 2 : 2Fe2+ + S2O82-  2Fe3+ + 2SO42-
 The reaction can also be catalysed by Fe2+(aq).
 Step 1 : 2Fe2+ + S2O82-  2Fe3+ + 2SO42 Step 2 : 2Fe3+ + 2I-  2Fe2+ + I2
 Rate of reaction is faster in a catalysed reaction due to reaction
involving oppositely charged ions.
 E.g 2 : Catalytic oxidation of atmospheric SO2 by





atmospheric oxides of nitrogen.
When SO2 reacts with water in the atmosphere, it produces
H2SO4 (acid rain). This reaction is catalysed by NO2 from the
car exhaust fumes.
Step 1 : SO2 + NO2  SO3 + NO
Step 2 : SO3 + H2O  H2SO4
Step 3 : 2NO + O2  2NO2
Nitrogen act as a catalyst by varying its oxidation state form
+4 to +2 and back to +4.
2. Heterogeneous Catalysis.
 Catalyst and reactants are in different physical states.
 Heterogeneous catalyst are usually d-block transition elements.
 Works by providing a surface onto which reactants adsorbed and
products of the reaction are desorbed.
 Its catalytic activity depends on the availability of partially filled 3d
orbitals which allows reactants molecules to be adsorbed onto
catalyst surface.
 The adsorption weakens the bond in the reactant molecule 
lowers Ea and increase surface concentration of reactants.
 Reactant molecules are brought closer together  rate of reaction
increase.
 E.g 1 : Haber process.
 Catalyst : Finely divided solid iron.
 3H2(g) + H2(g)  2NH3(g)
 H2 and N2 molecules adsorbs on catalyst surface  forming
weak bonds with catalyst surface and weakens bond within the
H2 and N2 molecules (easily broken).
 More easily form new bond with each other and form NH3.
 Ea decreases, rate of reaction increases.
 The NH3 molecules desorbs form the catalyst surface.
Other examples of heterogeneous catalysis :
a) V2O5 as catalyst in Contact process.
b) Ni as catalyst in hydrogenation of alkene.
Autocatalysis
 Autocatalyst  a product of a chemical reaction that acts as a





catalyst.
Autocatalysed reaction is slow at first and becomes more rapid
with time (catalyst produced).
E.g : reaction of acidified KMnO4 with ethanedioate ions.
2MnO4- + 5C2O42- + 16H+  2Mn2+ + 10CO2 + 8H2O
Mn2+ ion as autocatalyst.
The concentration of catalyst increases with time.
Individual work
• Creating a mental map on the topic "The
speed of a chemical reaction"
Success Criteria:
- colorfully reflects the influence of various
factors on the speed of a chemical reaction;
- presents graphs for reactions of different
order;
- reflects the main calculation formulas.
Reflection
• What has been learned
• What remained unclear
• What is necessary to work on