comprised two large chambers, at the same temperature, in each of which the
four species A, B, C and D are always in equilibrium. Each of the four walls is
permeable to only one of the substances and by transferring reactants and
products reversibly and isothermally between the two boxes through the walls
van’t Hoff was able to deduce the same equilibrium expression (see Appendix
C). It was Haber who named the device the ‘equilibrium box’. Needless to
say, the law of mass action is a much simpler derivation but interested readers
can find the detail of van’t Hoff’s argument in older physical chemistry text
books e.g., Samuel Glasstone, Textbook of Physical Chemistry, 2nd Ed,
Macmillan, (1966).
45. Le Chatelier’s Principle, an empirical generalisation that describes what
happens to the equilibrium position (i.e., the composition of the equilibrium
mixture of a chemical reaction) when the reaction conditions are changed.
The Principle was originally stated as;
Every change in one of the factors of an equilibrium occasions a rearrangement of the system in such a direction that the factor in question
experiences a change in the sense opposite to the original change
A more approachable definition might be;
‘when a constraint is applied to a chemical system at equilibrium, the system
responds in such a way as to minimise the effect of the constraint’
Thus, in the general reaction, if at equilibrium we add more of A (and, or B),
then the system will respond by counteracting (minimising) the effect of the
increased concentration(s) by using A and B to form more of C and D. The
equilibrium is then said to move to the right of the chemical equation. This
behaviour follows directly from the definition of K eq in that if [A] (and, or
[B]) are suddenly increased the only way K eq can remain constant is to use
them to increase [C] and [D]. The reverse is true if we add more of C (and,
or D) moving the equilibrium to the left.
Similarly, if we were to increase the temperature of a system at equilibrium,
we would favour the reaction that absorbed heat (i.e., the endothermic reaction) and vice versa. The effect of pressure on a gaseous reaction too follows the
Principle. If we were to increase the pressure then the equilibrium will react to
favour the reaction (forward or reverse) that results in a smaller volume. In the
important ammonia reaction;
N 2ðgÞ þ 3H 2ðgÞ 2NH 3ðgÞ
increasing pressure will move the equilibrium to the right resulting in
more ammonia formation because the forward reaction results in four
volumes of gas reducing to two. For this reaction, the decomposition
76
D. Sheppard
four species A, B, C and D are always in equilibrium. Each of the four walls is
permeable to only one of the substances and by transferring reactants and
products reversibly and isothermally between the two boxes through the walls
van’t Hoff was able to deduce the same equilibrium expression (see Appendix
C). It was Haber who named the device the ‘equilibrium box’. Needless to
say, the law of mass action is a much simpler derivation but interested readers
can find the detail of van’t Hoff’s argument in older physical chemistry text
books e.g., Samuel Glasstone, Textbook of Physical Chemistry, 2nd Ed,
Macmillan, (1966).
45. Le Chatelier’s Principle, an empirical generalisation that describes what
happens to the equilibrium position (i.e., the composition of the equilibrium
mixture of a chemical reaction) when the reaction conditions are changed.
The Principle was originally stated as;
Every change in one of the factors of an equilibrium occasions a rearrangement of the system in such a direction that the factor in question
experiences a change in the sense opposite to the original change
A more approachable definition might be;
‘when a constraint is applied to a chemical system at equilibrium, the system
responds in such a way as to minimise the effect of the constraint’
Thus, in the general reaction, if at equilibrium we add more of A (and, or B),
then the system will respond by counteracting (minimising) the effect of the
increased concentration(s) by using A and B to form more of C and D. The
equilibrium is then said to move to the right of the chemical equation. This
behaviour follows directly from the definition of K eq in that if [A] (and, or
[B]) are suddenly increased the only way K eq can remain constant is to use
them to increase [C] and [D]. The reverse is true if we add more of C (and,
or D) moving the equilibrium to the left.
Similarly, if we were to increase the temperature of a system at equilibrium,
we would favour the reaction that absorbed heat (i.e., the endothermic reaction) and vice versa. The effect of pressure on a gaseous reaction too follows the
Principle. If we were to increase the pressure then the equilibrium will react to
favour the reaction (forward or reverse) that results in a smaller volume. In the
important ammonia reaction;
N 2ðgÞ þ 3H 2ðgÞ 2NH 3ðgÞ
increasing pressure will move the equilibrium to the right resulting in
more ammonia formation because the forward reaction results in four
volumes of gas reducing to two. For this reaction, the decomposition
76
D. Sheppard