Kinetics The study of reaction rates. Spontaneous reactions are reactions

that will happen - but we can’t tell how fast.

Graphite will spontaneously turn to diamond– eventually.

Reaction mechanism- the steps by which a reaction takes place.

Review- Collision Theory Particles have to collide to react. Have to hit hard enough Things that increase this increase rate High temp – faster reaction High concentration – faster reaction Small particles = greater surface area

means faster reaction Catalyst

Reaction Rate

Rate = Conc. of A at t2 -Conc. of A at t1t2- t1

Rate =[A]t

Change in concentration per unit time For this reaction

N2 + 3H2 2NH3

As the reaction progresses the concentration H2 goes down

Concentration

Time

[H[H22]]

N2 + 3H2 → 2NH3

As the reaction progresses the concentration N2 goes down 1/3 as fast

Concentration

Time

[H[H22]]

[N[N22]]

N2 + 3H2 → 2NH3

As the reaction progresses the concentration NH3 goes up 2/3 times

Concentration

Time

[H[H22]]

[N[N22]]

[NH[NH33]]

N2 + 3H2 → 2NH3

Calculating Rates Average rates are taken over long

intervals Instantaneous rates are determined by

finding the slope of a line tangent to the curve at any given point because the rate can change over time

Derivative.

Average slope method

Concentration

Time

[H[H22]]

tt

Instantaneous slope method.

Concentration

Time

[H[H22]]

ttd[Hd[H22]]

dtdt

Defining RateWe can define rate in terms of the

disappearance of the reactant or in terms of the rate of appearance of the product.

In our example N2 + 3H2 2NH3

[H2] = 3[N2] t t

[NH3] = -2[N2] t t

Rate Laws Reactions are reversible. As products accumulate they can begin

to turn back into reactants. Early on the rate will depend on only the

amount of reactants present. We want to measure the reactants as

soon as they are mixed. This is called the Initial rate method.

Two key points The concentration of the products do

not appear in the rate law because this is an initial rate.

The order (exponent) must be determined experimentally,

can’t be obtained from the equation

Rate LawsRate Laws

You will find that the rate will only depend on the concentration of the reactants. (Initially)

Rate = k[NO2]n

This is called a rate law expression. k is called the rate constant. n is the order of the reactant -usually a

positive integer.

2 NO2 2 NO + O2

The rate of appearance of O2 can be said to be.

Rate' = [O2] = k'[NO2] t

Because there are 2 NO2 for each O2

Rate = 2 x Rate' So k[NO2]

n = 2 x k'[NO2]n

So k = 2 x k'

2 NO2 2 NO + O2

What are the units for k?It depends upon the order of the reaction.

Rate = mol/L sec

Zero Order: the value of x = 0

Rate = k[A]0 = k

therefore k=mol/L sec

What are the units for k?First Order: the value of x = 1

Rate = k[A]1 =

mol/L sec = k (mol/L)

therefore k=1/sec

What are the units for k?Second Order: the value of x = 2

Rate = k[A]2 =

mol/L sec = k (mol/L)2

therefore k=L/mol sec

What are the units for k?Third Order: the value of x = 3

Rate = k[A]3 =

mol/L sec = k (mol/L)3

therefore k=L2/mol2 sec

Types of Rate Laws Differential Rate law - describes how

rate depends on concentration. Integrated Rate Law - Describes how

concentration depends on time. For each type of differential rate law

there is an integrated rate law and vice versa.

Rate laws can help us better understand reaction mechanisms.

Determining Rate Laws The first step is to determine the form of

the rate law (especially its order). Must be determined from experimental

data. For this reaction

2 N2O5 (aq) 4NO2 (aq) + O2(g) The reverse reaction won’t play a role

because the gas leaves

[N[N22OO55] (mol/L) ] (mol/L) Time (s) Time (s)

1.001.00 00

0.880.88 200200

0.780.78 400400

0.690.69 600600

0.610.61 800800

0.540.54 10001000

0.480.48 12001200

0.430.43 14001400

0.380.38 16001600

0.340.34 18001800

0.300.30 20002000

Now graph the data

00.10.20.30.40.50.60.70.80.9

1

0 200

400

600

800

1000

1200

1400

1600

1800

2000

To find rate we have to find the slope at two points

We will use the tangent method.

00.10.20.30.40.50.60.70.80.9

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0 200

400

600

800

1000

1200

1400

1600

1800

2000

At .80 M the rate is (.88 - .68) = 0.20 =- 5.0x 10 -4 (200-600) -400

00.10.20.30.40.50.60.70.80.9

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0 200

400

600

800

1000

1200

1400

1600

1800

2000

At .40 M the rate is (.52 - .32) = 0.20 =- 2.5 x 10 -4 (1000-1800) -800

Since the rate at twice as fast when the concentration is twice as big the rate law must be..

First power

Rate = -[N2O5] = k[N2O5]1 = k[N2O5] t

We say this reaction is first order in N2O5

The only way to determine order is to run the experiment.