CHAPTER 6: CHEMICAL KINETICS – RATES OF REACTION
Key Ideas: p 358
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Introduction, p 358 : examples of chemical reactions, and where the rate of reaction is important.
6.1: Rate of Reaction:
Chemical kinetics : ______
-includes rates of reaction – times for reactions and concentration changes
-also includes models / theories to explain observations
-starting materials expensive – faster rate of reaction and more efficient reactions = more profits and cheaper medications
Describing Reaction Rates:
reaction rate = change of concentration per unit of time = c / t = mol / L / s = mol / L.s
-p 361 practice # 1,2.
-Graph of rate of consumption of a reactant – see p 362
-Graph of rate of production of a product – see p 363
-instantaneous rate of reaction = slope of tangent to line
-as concentration of reactant decreases, rate slows down
-rate may be based of decreasing concentration of reactant or increasing concentration of product – numbers of moles based on balanced chemical equations!
p 364 practice # 1, 2, 4.
Measuring Reaction Rates:
- Produce a gas: measure the volume of gas
- Involve ions: measure decreased ( reactant) or increased (product) conductivity of the solution
- Colour change – spectrophotometer used to measure changes in colour intensity
ex: ClO-1 (aq) + I-1 (aq) IO-1 (aq) + Cl-1 (aq)
colourless colourless yellow colourless
ex: H2 (g) + I2 (g) 2 HI (g)
colourless purple colourless
p 365 practice # 5, 6, 7, 8, 9.
p 366 practice # 1, 2, 3ab
6.2: Factors Affecting the Rate of Reaction
- Chemical nature of the reactants:
-group 1 metals react faster with oxygen in the air than gold
-the farther down the group, the faster the reaction with air
-corrosion: many metals react with water / air to oxidize – or with acids
-ions react faster than molecules ex: glucose molecules or iron III ions with permanganate ions
- Concentration:
-the greater the concentration, the faster the reaction rate
- Temperature:
-usually the higher the temperature, the faster the reaction ex: cooking, reptile metabolism
- Presence of a Catalyst:
-catalyst: - speeds up reaction
-same amount of substance at the end as in the beginning, whether or not it is a reactant or not
-works by providing an alternative pathway with a lower activation energy
-so at a lower temp, more molecules have enough energy to overcome energy barrier
-so reaction happens faster
-ex: transition metals and their oxides – see table p 370
Homogeneous catalysts: reagents and catalyst in same phase: all gases, or all liquids
-NO2 is used as a catalyst in the production of SO3, to make sulfuric acid
Heterogeneous catalysts: catalysis happens at the point between two phases
-ex: MnO2 to catalyze the decomposition of hydrogen peroxide
Inhibitors: slow reactions down, by occupying active sites on catalyst or on the reactant
-can be used with catalyst to carefully control the rate of reaction
Catalytic Converters and Engine Exhaust
-transition metal oxides act as heterogeneous catalysts to reduce / remove unwanted products of incomplete combustion:
-products of combustion of gasoline include CO2, H2O and unburned hydrocarbons, CO and NO from the reaction of N2 and O2 in the hot engine.
-NO promptly reacts with O2 once out of the exhaust pipe to form NO2 – contributes to acid rain and the formation of smog and causes respiratory problems for all air-breathing organisms
-2 NO N2 + O2 so less acid pollution is emitted
-hydrocarbons + CO + x O2 CO2 + H2O + Heat ( converter gets hot!!)
Catalysis in Industry:
-vital for savings in energy costs – reaction can happen at a lower temperature
-ex: series of reactions to get Zn from ZnS:
- ZnS (s) + H2SO4 + ½ O2 ZnSO4 + S (s) + H2O (l) too slow!!!!
Fortunately, FeS is also in Zn ore naturally, and the FeS reacts the same way, but more quickly.
FeS (s) + H2 SO4 + ½ O2 FeSO4 + S (s) + H2O (l)
Then the FeSO4 acts like a catalyst by taking part in the next reaction, and being regenerated in the
last reaction.
2 FeSO4 (aq) + H2SO4 + ½ O2 Fe2(SO4)3 (aq) + H2O (l)
Fe2(SO4)3 (aq) + ZnS (s) ZnSO4 (aq) + 2 FeSO4 (aq) + S (s)
-N.B. – catalysts may take part in the reaction, but are regenerated in the process so at the end the catalyst still exists- like iron II sulfate above
-specific catalyst needed for a specific reaction
-choice of catalysts used to be a trial and error process, involving transition metals and their compounds usually
-with knowledge to the structure of the reaction and the catalysis, a more efficient process of choosing a catalyst should soon be a reality.
Enzymes: natural catalysts in living organisms
-life wouldn’t be able to exist without them!
-ex: lactase to digest lactose ( milk sugar )
- Surface Area:
-the more finely divided the surface, the greater the contact between reactant molecules and the faster the reaction
-ex: burning a block of wood vs burning wood shavings vs burning saw dust
-especially for heterogeneous systems – G with S, L with S, G with L
p 371 practice: 1, 2, 3, 4, 5.
section 6.2 practice: # 1.
6.3: Rate Laws and Order of Reaction
Empirical Determination of Rate Laws
-rate law: ______
-need to be determined for any specific reaction by experimental evidence
-if a A + b B products, then rate [A]m [B]n or rate = k [A]m [B]n
-both m and n and K have to be determined from experimental evidence
-the exponents and k may be any real number and differ for each reaction
Order: if exponent is 1, we say the reaction is a first order reaction with respect to that reactant
-if exponent is 2, then second order reaction with respect to that reactant, etc.
Overall reaction order: determined by the sum of the exponents: m + n
Determining Exponents of the Rate Law:
-examine experimental data – change in concentration per second measured under identical conditions – the only factor changed for each trial is the concentration of one reactant
-while conc of one reactant remains the same how does rate change with the change in conc of the other reactant?
-if conc doubles and rate doubles, then exponent is 1 : 21 = 2
-if conc doubles and rate is 4 times faster, then exponent is 2: 22 = 4
-see examples p 375-376