Equilibria Quiz

Dynamic Equilibrium: A situation in which the ❌ of a constant ❌ reaction mixture does not ❌ because both forward and backward reactions are proceeding at the same ❌.

A + B <---> C + D
At the ❌ of the reaction the forward rate is fast, because A and B are ❌. There is no ❌ because there is no C and D.
Then as the ❌ of C and D build up, the reverse reaction speeds up. At the same time the concentrations of A and B ❌ so the forward reaction ❌.
A point is reached where exactly the same ❌ are changing from A and B to C and D as are changing from C and D to A and B. ❌ has been reached.

NB. An equilibrium ❌ can have any proportions of reactants and products. The proportions may be changed depending on the ❌ of the reaction, such as temperature, ❌, and concentration. But at any given constant conditions the proportions of reactants and products ❌ change.

If you ❌ the concentration of one of the reactants, Le Chatelier's principle says that the ❌ will shift in the direction what tends to reduce the concentration of this reactant.
A(aq) + B(aq) <---> C(aq) + D(aq)
If you add some extra A (increase its ❌), the only way that the system can reduce the concentration of A is by ❌ it with B to form more C and D. So, adding more A uses up more B, ❌ more C and D, and moves the equilibrium to the ❌. You end up with a ❌ proportion of products in the reaction mixture than before.

Pressure changes only affect reactions involving ❌. Changing the overall pressure will only change the position of equilibrium of a gaseous reaction if there are a ❌ number of molecules on either side of the equation.

N204(g) <---> 2NO2
❌ the pressure of a gas means that there are more molecules of it in a given volume - it is equivalent to increasing the ❌ of a solution.
If you increase the pressure on this system, Le Chatelier's principle tells us that the position of equilibrium will move to ❌ the pressure. This meas it will more to the ❌ because fewer molecules exert less pressure.

N2(g) + ❌H2(g) <---> ❌NH3(g) [delta H = -92kJmol^-1]

Almost all ammonia is made by the ❌ process, in which the reaction above is the key step. A low temperature and high ❌ would give close to 100% conversion whereas low pressure and high temperature would give almost no ❌. However, a compromise is used to ensure the rate of reaction is ❌ viable and the conditions safe. Most plants run at a pressure of around 20000kPa (c. ❌ atmospheres) and a temperature of about ❌K (c. 450 celsius).
The raw materials for the Haber process are air (providing the nitrogen), water, and natural gas (methane). These provide the hydrogen by the following reaction:
❌(g) + H2O(g) ---> ❌(g) + 3H2(g)
The nitrogen and hydrogen are fed into a converter in the ratio ❌ and passed over an ❌ catalyst.
The conversion is not at maximum as the conditions are not optimum (a ❌ is used) and the gases flow continuously over the catalyst so they do not spend long enough in contact with the catalyst to reach ❌. Any unconverted reactants are ❌ through the converter to increase the yield.

Kc > 1 --> ❌
Kc = 1 --> ❌
Kc < 1 --> ❌