Reversible Reactions

1. Lots of chemical reactions are reversible- they go both ways. To show a reaction's reversible, you stick in a reversible reaction arrow (see equation below). Here's an example:

This reaction can go in either direction, forwards or backwards.

2. As reactants get used up, the forward reaction slows down- and as more product is formed, the reverse reaction speeds up.

3.After a while, the forward reaction will be going at exactly the same rate as the backward reaction. The amounts of reactants and products won't be changing any more, so it'll seem like nothing's happening. It's a bit like you're digging a hole, while someone else is filling it in exactly the same speed. This is called a dynamic equilibrium.

4. A dynamic equilibrium can only happen in a closed system. This just means nothing can get in or out.

If you change the concentration, pressure or temperature of a reversible reaction, you're going to alter the position of the equilibrium. This just means you'll end up with different amounts of reactants and products at equilibrium.

If the position of equilibrium moves to the left, you'll get more reactants.

If the position of equilibirum moves to the right, you'll get more products.

Le Chatlier's orinicple tells you how the position of equilibirum will change if a condition changes:

If there's a change in concentration, pressure or temperature, the equilibrium will move to help counteract the change.

So basically, if you raise the temperature, the position of equilibrium will shift to try to cool things down. And if you raise the pressure or concentration, the position will shift to try to reduce it again.

Catalysts have no effect on the position of equilibrium.

They can't increase yield- but they do mean equilibrium is reached faster.

2SO2 + O2 <-> 2SO3

1. If you increase the concentration of a reactant, (e.g. SO2 or O2), the equilibrium tries to get rid of the extra reactant. It does this by making more products (SO3). So the equilibrium's shifted to the right.

2. If you increase the concentration of the product (SO3), the equilibrium tries to remove the extra products. This makes the reverse reaction go faster. So the equilibrium shifts to the left.

3. Decreasing the concentrations has the opposite effect.

(changing this only affects equilbria involving gases)

1. Increasing the pressure shifts the equilibrium to the side with fewer molecules. This reduces the pressure.

2. Decreasing the pressure shifts the equilibrium to the side with more molecules. This raises the pressure.

E.g. 2SO2 + O2 <--> 2SO3

There's 3 moles on the left, but only 2 on the right. So an increase in pressure shifts the equilibrium to the right.

1. Increasing the temperature means adding heat. The equilibrium shifts in the endothermic direction (positive enthalpy change) to absorb this heat.

2. Decreasing the temperature removes heat. The equilibrium shifts in the exothermic direction (negative enthalpy change) to try and replace the heat.

3. If the forward reaction's endothermic, the reverse reaction will be exothermic, and vice versa.

E.g. 2SO2 + O2 <--> 2SO3

This reaction's exothermic in the forward reaction. If you increase the temperature, the equilibrium shifts to the left to absorb the extra heat.

Right, you've got to be able to apply this Le Chatelier's Principle stuff to industrial processes, like the production of ethanol and methanol. 

1. Ethanol is produced via a reversible exothermic reaction between ethene and steam.

2. The reaction is carried out at a pressure of 60-70 atmospheres and a temperature of 300 degress Celsius, with a catalyst of phosphoric acid.

1. Because it's an exothermic reaction, lower temperatures favour the forward reaction. This means that at lower temperature more ethane and steam is converted into ethanol- you get a better yield.

2. But lower temperatures mean a slower rate of reaction. You'd be daft to try to get a really high yield of ethanol it it's going to take you 10 years. So the 300 degress Celsius is a compromise between maximum yield and a faster reaction.

3. Higher pressure favour the forward reaction, so a pressure of 60-70 atmospheres is used- high pressure moves the reaction to the side with fewer molecules of gas. Increasing the pressure also increases the rate of reaction.

4. Cranking up the pressure as high as you can sounds like a great idea so far. But high pressure are expensive to produce. You need stronger pipes and containers to withstand the pressure. And, in this process, increasing the pressure can also cause side reactions to occur.

5. So the 60-70 atmospheres is a compromise between maximum yield and expense. In the end, it all comes down to minimising the costs.

1. Only a small proportion of ethene reacts each time the gases pass through the catalyst.

2. To save money and raw materials, the unreacted ethene is seperated from the liquid ethanol and recycled back into the reactor. Thanks to this around 95% of the ethene is eventually converted to ethanol.

1. Methanol is also made industrially in a reversible reaction. It's made from hydrogen and carbon monoxide.

2H2 + CO <--> CH3OH

Industrial conditions- pressure: 50-100 atmospheres, temperature: 250 degress Celsius,

catalyst: mixture of copper, zinc oxide and aluminium oxide

2. Just like with the production of ethanol, the conditions used are a compromise between keeping costs low and yield high.

1. Methanol is mainly used to make other chemicals, but both methanol and ethanol can also be used as fuels for cars- either on their own, or added to petrol.

2. Ethanol and methanol are thought of as greener than petrol - they can be made from renewable resources and they produce fewer pollutants (like NO2 and CO).

3. Methanol and ethanol can both be carbon neutrak fuels (pretty much).

Something is carbon neutral if it has no net annual carbon (greenhouse gas) emissions to the atmosphere.