Unless otherwise stated, all images in this file have been reproduced from:

Blackman, Bottle, Schmid, Mocerino and Wille,     Chemistry, 2007 (John Wiley)

     ISBN: 9 78047081 0866

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e CHEM1002 [Part 2]

A/Prof Adam Bridgeman (Series 1)Dr Feike Dijkstra (Series 2)

Weeks 8 – 13

Office Hours: Monday 2-3, Friday 1-2Room: 543ae-mail: adam.bridgeman@sydney.edu.aue-mail: feike.dijkstra@sydney.edu.au

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e

Complexes IV

• The metals include many essential elements (such as Na+, K+ and Ca2+, some toxic (such as Hg2+ and Al3+) and some which are now being used in medicines (such as Pt2+)

• The essential metals have a variety of functions in the body (such as Na+/K+ in the nervous system, Fe2+/3+ in oxygen transport and Zn2+ in CO2 transport)

• The biological function is related to the oxidation number, the coordination type and the size of the atom

Summary of Last Lecture

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e

Lecture 16• Chemical Kinetics• Rate of Reaction• Rate Laws• Reaction Order• Blackman Chapter 14, Sections 14.1 - 14.3

Lecture 17• Half lives• The Temperature Dependence of Reaction Rates• Catalysis• Blackman Chapter 14, Sections 14.4 - 14.6

Chemical Kinetics I

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e Kinetics vs Thermodynamics

• Thermodynamics (ΔG, ΔunivS, Eo) tells us if a reaction favours the products or reactants

• It also gives the extent a reaction occurs (Keq)

• Thermodynamics says nothing about how fast or slow the reaction goes

• It gives us the equilibrium concentrations but not how long it takes to get to equilibrium

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e Reaction Rate

• The rate of a reaction is how fast the concentration of the molecules present change.

• Reaction rate: change in concentration of a product or a reactant per unit time.

• Rate is given by the gradient of concentration vs time graph

d[A]

= — — dt

reaction rate

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e Expressing Reaction Rates

• There are a number of ways to express the rate.

hydrolysis of cisplatin

[Pt(NH3)2Cl2] + H2O [Pt(NH3)2(H2O)Cl]+ + Cl-

d[reactant]

dt dtrate =

d[product] — +=

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x

• Express the rate of reaction for each reactant and product in the reaction:

4NH3(g) + 5O2(g) 4NO(g) + 6H2O (g)

d[NH3]

dtNH3:

d[O2]

dtO2:

d[NO]

dtNO:

d[H2O]

dtH2O:

Expressing Reaction Rates

- 14

- 15

+ 14

+ 16

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e Rates are Determined Experimentally

d[O3]C2H4 + O3 C2H4O + O2

• Rate is dependent of concentration of O3

dtrate = -

00.5

11.5

22.5

33.5

0 20 40 60time (s)

[O3]

(10-5

M)

fast

slow

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• k

rate = k [O3] [C2H4]

The Rate Law

• The rate of the reaction is proportional to the concentration [O3] and to that of [C2H4]

• This is expressed as a rate law:

• k increases with T

is the rate constant and is independent of concentration

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For the general reaction:

The Rate Equation

aA + bB + cC … mM + nN ….

dt= k [A]x [B]y [C]z …

-d[A]

• the rate law can only be determined by experiment, not from the stoichiometric equation

• x is the order of the reaction with respect to A,y is the order of the reaction with respect to B…

• the overall order of the reaction is given by x + y + z …

• there is no relationship between a and x, b and y ….

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x Example

• For the reaction:

2 NO (g) + O2 (g) 2 NO2 (g)

rate = k[NO]2[O2]

What is the order of reaction with respect to the reactants and the overall order of reaction?

• second order with respect to NO

• first order with respect to O2

• third order overall

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e Interpreting Rate Laws

• If x = 1, reaction is 1st order in A: rate α [A]1

• If [A] doubles, then rate goes up by factor of

= k [A]x [B]y [C]z …rate

• If x = 2, reaction is 2nd order in A: rate α [A]2

• If [A] doubles, then rate goes up by factor of

• If x = 3, reaction is 3rd order in A: rate α [A]3

• If [A] doubles, then rate goes up by factor of

two

four

eight

16

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x Using Data to Determine Order

ClO3-(aq) + 9 I- (aq) + 6 H+(aq) Cl- (aq) + 3 I3

- (aq) + 3 H2O(l)

rate = k [ClO3-]x [I-]y [H+]z

4 0.20 0.20 0.20 0.80

3 0.20 0.20 0.10 0.20

2 0.20 0.10 0.10 0.10

1 0.10 0.10 0.10 0.05

[ClO3-] / M [I-] / M [H+] / M initial rate / M s-1

x = 1 , y = 1 and z = 2 so

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x Using Data to Determine Order

Calculate the rate constant, k:

in experiment 1:

rate = 0.05 M s-1, [ClO3-] = 0.1 M, [I-] = 0.1 M, [H+] = 0.1 M

ClO3-(aq) + 9 I- (aq) + 6 H+(aq) Cl- (aq) + 3 I3

- (aq) + 3 H2O(l)

rate = k [ClO3

-] [I-] [H+]2

so,

(units of k depend on overall order)

0.05 M s-1 = k (0.10 M)(0.10 M)(0.10 M)2

k = 5.0 x 102 M-3 s-1

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x Question

NO2 (g) + CO (g) NO (g) + CO2 (g)

[NO2] / M [CO] / M Initial rate / M s-1

1 0.10 0.10 0.00502 0.40 0.10 0.0803 0.10 0.20 0.0050

• determine the rate equation and value of the rateconstant for this reaction

(zero order in [CO])

k = 0.5 M-1 s-1

rate = k[NO2]2

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e Summary: Chemical Kinetics I

Learning Outcomes - you should now be able to:

• Work out the order of the reaction with respect to each reactant from experimental data

• Work out the rate constant (including its units) from experimental data

• Hence, write down the rate law• Answer review problems 14.66 - 14.97 in Blackman• Complete the worksheet

Next lecture:

• Half lives, temperature dependence and catalysis

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x Practice ExampleThe data given in the table below were obtained for the reaction between nitric oxide and chlorine at 1400 K.

2NO(g) + Cl2(g) 2NOCl(g)

Experiment Number

Initial [Cl2] (mol L–1)

Initial [NO](mol L–1)

Initial Reaction Rate(mol L–1 s–1)

1 0.10 0.10 0.18

2 0.20 0.10 0.36

3 0.10 0.20 0.72

Deduce the rate law for this reaction.

Calculate the value of the rate constant.