3.4 Unit 4 CHEM4 Kinetics, Equilibria and
Organic Chemistry
3.4.1 – Kinetics (1 page)
Simple rate equations
Understand and be able to use rate equations of the form Rate = k[A]m [B]n where m and
n are the orders of reaction with respect to reactants A and B (m, n restricted to values
1,2 or 0)
The rate equation tells us how the rate is affected by the concentration of
each reactant.
Equation: A + B C + D
Rate: k[A]m[B]n
k = rate constant
m and n = orders of reaction – how the concentration of each reactant
affects the rate
Square brackets = “concentration of..”
Zero order – concentration of reactant does not affect rate
First order – rate is proportional to concentration of reactant
Second order – rate is proportional to concentration of reactant squared
Overall order – m + n
Determination of rate equation
Be able to derive the rate equation for a reaction from data relating initial rate to the
concentrations of the different reactants
Find trial where concentration of only one reactant changes.
If concentration doubles and:
Rate stays the same – zero order
Rate doubles – first order
Rate quadruples (22) – second order
Be able to explain the qualitative effect of changes in temperature on the rate constant k
Higher temperature increases the rate constant.
When we increase the temperature the rate increases. The concentrations
of reactants stay the same therefore the rate constant, k, must be
increasing.
Understand that the orders of reactions with respect to reactants can be used to provide
information about the rate determining/limiting step of a reaction
The overall rate of a reaction is decided/controlled by the step with the
slowest rate – the rate determining step. When you measure the rate of a
reaction, you are actually measuring the rate of the rate determining step.
The orders of reaction can be used to provide information about the rate
determining step.
Rate determining/slowest step must include all the molecules that are in
the rate equation. BUT sometimes an intermediate may have been made
up of some of the molecules so look out for this.
,3.4.2 – Equilibria (1 page)
Equilibrium constant Kc for homogeneous systems
Know that Kc is the equilibrium constant calculated from equilibrium concentrations for a
system at constant temperature
Kc is the equilibrium constant, which is calculated from the equilibrium
concentrations.
Be able to construct an expression for Kc for a homogeneous system in equilibrium;
Be able to perform calculations involving such an expression
To calculate Kc for the reaction aA + bB cC + dD: (small letter = number
of moles)
Kc = [C]c [D]d
[A]a [B]b
This is using equilibrium concentrations NOT initial concentrations.
Be able to work out units of Kc by writing moldm-3 instead of the
substances on top and bottom and cancelling, and remember if you bring
a power to the top it becomes negative.
Qualitative effects of changes of temperature and concentration
Be able to predict the effects of changes of temperature on the value of the equilibrium
constant
Work out what the change in temperature will do to the position of the
equilibrium.
If the equilibrium shifts to the RIGHT – Kc will INCREASE.
If the equilibrium shifts to the LEFT – Kc will DECREASE.
Understand that the value of the equilibrium constant is not affected by changes either
in concentration or the addition of a catalyst
Changes in concentration or the addition of a catalyst does NOT affect Kc.
This is because:
Changes in concentration shift the equilibrium but then it will shift back.
You return to the original ratios but higher amounts so Kc stays the same.
Catalysts don’t affect the position of the equilibrium, they just affect the
rate of attainment; it increases the rate of the forward and backward
reactions equally so no effect on Kc.
, 3.4.3 – Acids and Bases (5 pages)
Brønsted-Lowry acid-base equilibria in aqueous solution
Know that an acid is a proton donor
A Brønsted-Lowry acid is a proton (H+) donor.
Know that a base is a proton acceptor
A Brønsted-Lowry base is a proton (H+) acceptor.
Know that acid-base equilibria involve the transfer of protons
Acid-base equilibria involve the transfer of protons (H+ ions).
Definition and determination of pH
Know that pH = -log10 [H+], where [ ] represents the concentration in moldm -3
pH = -log10 [H+]
Be able to convert concentration into pH and vice versa
[H+] = 10-pH
Be able to calculate the pH of a solution of a strong acid from its concentration
Calculate pH of a solution of a strong monoprotic acid:
The concentration of the acid is the same as the concentration of H+, so
put the concentration straight into the pH formula.
(Calculate concentration from pH – put pH into [H+] formula – then [H+] is
the same as [acid].
Calculate pH of a solution of a strong diprotic acid:
Each molecule of the acid will release 2 protons when it dissociates.
The concentration of H+ is double the concentration of the acid.
So multiply the concentration of the acid by 2 and then put into the pH
formula.
(Calculate concentration from pH – put pH into [H+] formula – then divide
by 2 to get [acid].
The ionic product of water, Kw
Know that water is weakly dissociated
Water weakly dissociates to hydroxonium ions and hydroxide ions:
H2O + H2O H3O+ + OH-
This simplifies to:
H2O H+ + OH-
Know that Kw = [H+] [OH-]
Kc expression for above equilibrium:
Kc = [H+] [OH-]
[H2O]
Organic Chemistry
3.4.1 – Kinetics (1 page)
Simple rate equations
Understand and be able to use rate equations of the form Rate = k[A]m [B]n where m and
n are the orders of reaction with respect to reactants A and B (m, n restricted to values
1,2 or 0)
The rate equation tells us how the rate is affected by the concentration of
each reactant.
Equation: A + B C + D
Rate: k[A]m[B]n
k = rate constant
m and n = orders of reaction – how the concentration of each reactant
affects the rate
Square brackets = “concentration of..”
Zero order – concentration of reactant does not affect rate
First order – rate is proportional to concentration of reactant
Second order – rate is proportional to concentration of reactant squared
Overall order – m + n
Determination of rate equation
Be able to derive the rate equation for a reaction from data relating initial rate to the
concentrations of the different reactants
Find trial where concentration of only one reactant changes.
If concentration doubles and:
Rate stays the same – zero order
Rate doubles – first order
Rate quadruples (22) – second order
Be able to explain the qualitative effect of changes in temperature on the rate constant k
Higher temperature increases the rate constant.
When we increase the temperature the rate increases. The concentrations
of reactants stay the same therefore the rate constant, k, must be
increasing.
Understand that the orders of reactions with respect to reactants can be used to provide
information about the rate determining/limiting step of a reaction
The overall rate of a reaction is decided/controlled by the step with the
slowest rate – the rate determining step. When you measure the rate of a
reaction, you are actually measuring the rate of the rate determining step.
The orders of reaction can be used to provide information about the rate
determining step.
Rate determining/slowest step must include all the molecules that are in
the rate equation. BUT sometimes an intermediate may have been made
up of some of the molecules so look out for this.
,3.4.2 – Equilibria (1 page)
Equilibrium constant Kc for homogeneous systems
Know that Kc is the equilibrium constant calculated from equilibrium concentrations for a
system at constant temperature
Kc is the equilibrium constant, which is calculated from the equilibrium
concentrations.
Be able to construct an expression for Kc for a homogeneous system in equilibrium;
Be able to perform calculations involving such an expression
To calculate Kc for the reaction aA + bB cC + dD: (small letter = number
of moles)
Kc = [C]c [D]d
[A]a [B]b
This is using equilibrium concentrations NOT initial concentrations.
Be able to work out units of Kc by writing moldm-3 instead of the
substances on top and bottom and cancelling, and remember if you bring
a power to the top it becomes negative.
Qualitative effects of changes of temperature and concentration
Be able to predict the effects of changes of temperature on the value of the equilibrium
constant
Work out what the change in temperature will do to the position of the
equilibrium.
If the equilibrium shifts to the RIGHT – Kc will INCREASE.
If the equilibrium shifts to the LEFT – Kc will DECREASE.
Understand that the value of the equilibrium constant is not affected by changes either
in concentration or the addition of a catalyst
Changes in concentration or the addition of a catalyst does NOT affect Kc.
This is because:
Changes in concentration shift the equilibrium but then it will shift back.
You return to the original ratios but higher amounts so Kc stays the same.
Catalysts don’t affect the position of the equilibrium, they just affect the
rate of attainment; it increases the rate of the forward and backward
reactions equally so no effect on Kc.
, 3.4.3 – Acids and Bases (5 pages)
Brønsted-Lowry acid-base equilibria in aqueous solution
Know that an acid is a proton donor
A Brønsted-Lowry acid is a proton (H+) donor.
Know that a base is a proton acceptor
A Brønsted-Lowry base is a proton (H+) acceptor.
Know that acid-base equilibria involve the transfer of protons
Acid-base equilibria involve the transfer of protons (H+ ions).
Definition and determination of pH
Know that pH = -log10 [H+], where [ ] represents the concentration in moldm -3
pH = -log10 [H+]
Be able to convert concentration into pH and vice versa
[H+] = 10-pH
Be able to calculate the pH of a solution of a strong acid from its concentration
Calculate pH of a solution of a strong monoprotic acid:
The concentration of the acid is the same as the concentration of H+, so
put the concentration straight into the pH formula.
(Calculate concentration from pH – put pH into [H+] formula – then [H+] is
the same as [acid].
Calculate pH of a solution of a strong diprotic acid:
Each molecule of the acid will release 2 protons when it dissociates.
The concentration of H+ is double the concentration of the acid.
So multiply the concentration of the acid by 2 and then put into the pH
formula.
(Calculate concentration from pH – put pH into [H+] formula – then divide
by 2 to get [acid].
The ionic product of water, Kw
Know that water is weakly dissociated
Water weakly dissociates to hydroxonium ions and hydroxide ions:
H2O + H2O H3O+ + OH-
This simplifies to:
H2O H+ + OH-
Know that Kw = [H+] [OH-]
Kc expression for above equilibrium:
Kc = [H+] [OH-]
[H2O]
