1 Fossum-Reyes Chapter 14 – Chemical Kinetics How fast do chemical processes occur? There is an enormous range of time scales. Kinetics also sheds light on the reaction mechanism (exactly how the reaction occurs). Why is it important? Increase the shelf life of products like food, beverages, pharmaceutical drugs, etc. (How?) – slowing down some reactions. Remove hazardous pollutants from the environment (How?) – speeding up some reactions (biodegradable materials). 14.1 Factors that affect Reaction Rates 1. Physical state of reactants (s, l, g, aq). Molecules must come in contact with each other in order to react, therefore: – The more homogeneous the mixture of reactants, the faster the molecules can react. – Gases tend to react faster than liquids, which react faster than solids. – For solids, however, surface area is important (powders react faster than chunks). 2. Reactant Concentrations. Generally, the larger the concentration of reactant molecules, the faster the reaction will proceed (Why?). The likelihood that reactant molecules will collide increases. 3. Reaction Temperature. Higher T leads to higher kinetic energy, molecules move faster and collide more often and with greater energy (i.e. more molecules have enough energy to react). – chemist’s rule of thumb: ΔT = 10 °C → 2(rate) 4. The presence/concentration of a catalyst. – Catalysts speed up reactions by changing the mechanism of the reaction. – Catalysts are not consumed during the course of the reaction. 14.2 Reaction Rates [] ( ) For the hypothetical reaction: A → B ‘A’ will decrease and ‘B’ will increase (every 1 A that reacts produces 1 B). [] [] The rate of reaction is positive (+) by convention; however, the change of reactants and products can be positive (+) or negative (–).
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1
Fossum-Reyes
Chapter 14 – Chemical Kinetics
How fast do chemical processes occur?
There is an enormous range of time scales.
Kinetics also sheds light on the reaction mechanism (exactly how the reaction occurs).
Why is it important?
Increase the shelf life of products like food, beverages, pharmaceutical drugs, etc. (How?)
– slowing down some reactions.
Remove hazardous pollutants from the environment (How?)
– speeding up some reactions (biodegradable materials).
14.1 Factors that affect Reaction Rates
1. Physical state of reactants (s, l, g, aq). Molecules must come in contact with each other in order to
react, therefore:
– The more homogeneous the mixture of reactants, the faster the molecules can react.
– Gases tend to react faster than liquids, which react faster than solids.
– For solids, however, surface area is important (powders react faster than chunks).
2. Reactant Concentrations. Generally, the larger the concentration of reactant molecules, the faster the
reaction will proceed (Why?). The likelihood that reactant molecules will collide increases.
3. Reaction Temperature. Higher T leads to higher kinetic energy, molecules move faster and collide
more often and with greater energy (i.e. more molecules have enough energy to react).
– chemist’s rule of thumb: ΔT = 10 °C → 2(rate)
4. The presence/concentration of a catalyst.
– Catalysts speed up reactions by changing the mechanism of the reaction.
– Catalysts are not consumed during the course of the reaction.
14.2 Reaction Rates
[ ]
( )
For the hypothetical reaction: A → B
‘A’ will decrease and ‘B’ will increase (every 1 A that reacts produces 1 B).
[ ]
[ ]
The rate of reaction is positive (+) by convention; however, the change of
reactants and products can be positive (+) or negative (–).
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Fossum-Reyes
The average rate over a certain time interval depends on the time interval chosen.
The instantaneous rate refers to the rate at a particular moment.
To determine the instantaneous rate from a graph of concentration vs time:
1. Draw a tangent to the curve at the desired time point.
2. Determine the slope of the tangent.
For a chemical reaction, the rate decreases as the time goes on (i.e. the slope gets less steep).
Reaction Rate Changes Over Time
Because the concentration of the reactants decreases.
At some time the reaction stops, either because the reactants run out or because the system has reached
equilibrium.
Problem 1 Using Figure 14.4, determine the instantaneous rate of disappearance of C4H9Cl at t = 300 s.
Relating rates
Rates of products and reactants can be related: N2O5(g) 2 NO2(g) + ½ O2(g)
As 1 mol N2O5(g) is consumed, 2 mol NO2(g) and ½ mol of O2(g) are formed.
So [NO2] changes twice as fast as [N2O5].
So rate [N2O5] = ½ rate [NO2].
[ ]
[ ]
[ ]
Reaction Rates and Stoichiometry
To generalize, then, for the reaction:
[ ]
[ ]
[ ]
[ ]
Instantaneous Rate Example
Using [H2], the instantaneous rate at 50 s is:
Using [HI], the instantaneous rate at 50 s is
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Problem 2 a. How is the rate at which ozone disappears related to the rate at which oxygen appears in the reaction?
2 O3(g) 3 O2(g)?
b. If the rate at which O2 appears, [O2]/t, is 6.0 10–5
M/s at a particular instant, what is the rate of change of
O3?
c. What is the overall rate at this instant?
Measuring the Reaction Rate
To measure the reaction rate we need to measure the concentration of at least one component (reactant or
product) in the mixture at many points in time.
- Continuous monitoring for reactions that take place in less than an hour.
- Sampling of the mixture at various times for reactions that happen in a long time.
Possible tests:
Measure the absorbance of visible or UV light, pH, electrical conductivity, and pressure of a gas.
14.3 Concentration and Rate Laws
We can gain information about the rate of a reaction by seeing how the rate changes with changes in
concentration.
The rate law of a reaction is the mathematical relationship between the rate of the reaction and the
concentrations of the reactants.
The rate law must be determined experimentally!! (There is no way to predict it!!)
For the reaction:
N2O5(g) 2 NO2(g) + ½ O2(g)
It was determined experimentally that when [N2O5] doubles, the reaction rate doubles. If [N2O5] triples, the rate
triples. If [N2O5] is halved, the rate is halved.
Therefore: rate α [N2O5].
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Fossum-Reyes
N2O5(g) 2 NO2(g) + ½ O2(g)
The rate law for this reaction: [ ].
Where is the rate constant, which depends on the reaction and on T. (It is determined by experiment.)
In general:
[ ] [ ] [ ]
are called the orders for each reactant (usually 0, 1, 2, but they can be fractions or negative).
Again: is called the rate constant.
The rate law must be determined experimentally, and it is not related to the balanced equation. For:
The rate law is:
[ ][ ][ ]
exponents ≠ coefficients
(except by coincidence)
2 NO(g) + O2(g) → 2 NO2(g)
The rate law for the reaction is: Rate = k[NO]2[O2]
The reaction is:
second order with respect to [NO],
first order with respect to [O2],
and third order overall (order of the reaction)
For:
[ ][ ][ ]
First order in [ ]
First order in [ ] Second order in [ ]
Fourth order overall (add all exponents).
For: ( ) ( ) ( ) [ ]
Zero order in NH3
Experimental determination of Rate Law: one method is the Method of Initial Rates.
Initial rate – rate at the very beginning of the reaction (≈1-2% complete. Convenient because: initial [ ] are
known, avoids complications with products and side reactions).
Approach: vary initial [reactants] and see how the rate is affected.