Download Equilibrium Reactions

The collision Model
Chemists believe that for a chemical reaction to occur, the reactants
must collide with enough force to break bonds, allowing the
reactants to rearrange themselves to form products.
For example:
2BrNO(l)
2NO (g) +
Br2(g)
Molecules not correctly orientated, so no reaction
Molecules are correctly orientated, so reaction occurs
So for a chemical reaction to occur the following must happen:
•Reactants collide
•Collide at the right orientation
•Collide with enough violence
What factors affect the rate of a chemical reaction?
•Temperature
2NO2
N2O4
•Surface area (Grain explosion)
•Concentration (aq)
Mg + HCl
•Catalysts catalase in animal liver 2H2O2
•Volume (gases)
H2 + Cl2
2H2O + O2
This is the minimum energy required
Activation Energy (Ea) for a reaction to occur
An Endothermic Reaction
Catalyst
This is a substance that speeds up a reaction without
being consumed.
Enzymes are catalysts that are found in our bodies.
They allow our body to speed up the rate of reactions
that would be too slow at normal body temperature.
A good example is the breakdown of ozone, which is catalyzed by
chlorine.
Cl + O3

O + ClO
 Cl + O2
Sum:
Cl + O3 + O + ClO
O3 + O
ClO + O2
 ClO + O2 + Cl + O2

2O2
Equilibrium means balance or steadiness
In chemistry, an equilibrium reaction is when there are two
opposite reactions that are balanced.
We tend to think that all reactions go in only one direction. What
is called the forward reaction.
CaCO3 (aq) + HCl (aq)  CaCl2 (aq) + CO2 (g) + H2O (l)
In this reaction the reactants react to form products.
It goes to completion, or until one of the reactants runs out
But not reactions are like this!
Equilibrium Reactions:
This is reaction that has two directions, a forward and
reverse reaction.
An equilibrium reaction is reversible
Forward reaction
H2O(g) + CO (g)
Reverse reaction
H2 (g) + CO2 (g)
Dynamic Equilibrium:
If kept in a closed system, then both reactions will occur. If this
reaction is left then it will reach a point of equilibrium where the
rate of the forward reaction equals the rate of the reverse reaction.
So the concentrations of the reactants doesn’t change.
The equilibrium is dynamic, because even though the
concentrations stay the same and nothing appears to be
happening, both the forward and reverse reactions are continuing
to occur.
Changes in reaction rates of the forward and reverse reactions for:
H2 + CO
H2 +CO2
Rate of forward reaction decreases while reverse increases till the
concentrations reach a level at which the rate of the forward and
reverse reactions is the same. The system has reached equilibrium.
[ ] of A decreases while
[ ] of B increases till
equilibrium is reached.
Equilibrium is reached
when rate of forward
reaction is the same as
the reverse reaction.
The Equilibrium Constant:
When the system is at equilibrium, the concentrations remain
constant. So the ratio of the concentrations also remains constant.
E.g.
aA + bB  cC + dD
A, B, C, D = chemical species
a, b, c, d = coefficients
C D
Keq 
a
b
A  B 
c
d
[ ] = molarity in mol L-1
Keq = equilibrium constant
E.g. 4NH3 (g) + 7O2 (g)  4NO2 (g) + 6H2O (g)
NO2  H2 O

Keq 
4
7
NH3  O2 
4
6
You should always make sure that you write balanced equations
that include the physical states of the reactants and products.
Write equilibrium expressions for these reactions:
a) 2O3 (g)  3O2 (g)
b) H2 (g) + F2 (g)  2HF(g)
The equilibrium constant means that for a given reaction at a
given temperature the ratio of the concentration of reactants to
products is always be constant.
Calculating Equilibrium Constants:
The actual value of Keq is found experimentally. The individual
concentrations of all the reactants is calculated, and the
temperature recorded.
This is because the value of Keq will change with temperature.
What can Keq tell you?
Which direction is favored
If the value of Keq is one then ….
If the value of Keq is less than one ……
If the value of Keq is greater than one then ………
Results for three experiments for the reaction:
2NH3 (g) at 500 C
N2 (g) + 3H2 (g)
expt Initial Concentrations
Equilibrium
Concentrations
[NH3] [N2]
[NH3 ]2
K
3
[N2 ] [H 2 ]
[N2]
[H2]
[H2] [NH3]
I
1.000
1.000
0
0.921
0.763
0.157
0.0602
II
0
0
0
0.399
1.197
0.203
0.0602
III
2.00
1.00
3.00
2.59
2.77
1.82
0.0602
Calculating Keq for a reaction:
E.g. 4SO2 (g) + O2 (g)  2SO3(g)
Experiment 1
Initial
[SO2 ] = 2.00M
[O2 ] = 1.50M
[SO3 ] = 3.00M
Equilibrium
[SO2 ] = 1.50M
[O2 ] = 1.25M
[SO3] = 3.50M
Equilibrium constant for Experiment 1 =
Experiment 2
Initial
[SO2 ] = 0.500M
[O2 ] = 0.00M
[SO2 ] = 0.350M
Equilibrium
[SO2 ] = 0.590M
[O2 ] = 0.045M
[SO3 ] = 0.260M
Equilibrium constant for Experiment 2 =
Heterogeneous Equilibria:
A homogenous reaction is one in which all the substances are in the
same state.
A heterogeneous reaction is one in which all the substances are not in
the same state.
CaCO3 (s)
Calcium carbonate
CO2 (g)
+
carbon dioxide
CaO(s)
lime
When writing equilibrium constant expressions for Heterogeneous
equilibria, you don’t include pure solids or pure liquids.
Their concentrations don’t change
Keq = [CO2 ][ CaO ]
[ CaCO3 ]
Keq = [CO2 ]
Try writing an Keq expression for this reaction:
2H2O (l)
2H2 (g) +
O2(g)