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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
Slide 2/18
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
Slide 3/18
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
Slide 4/18
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
Slide 5/18
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
Slide 6/18
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
Slide 7/18
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] — +=
Slide 8/18
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
Slide 9/18
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
Slide 10/18
e
• 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
Slide 11/18
e
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 ….
Slide 12/18
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
Slide 13/18
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
Slide 14/18
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
Slide 15/18
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
Slide 16/18
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
Slide 17/18
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
Slide 18/18
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.
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