More chemisty questions!!?

Question

1. The harbour process is used industrially to synthesize ammonia from nitrogen and hydrogen. The thermochemical equation for this is : N2(g) +3H2(g) <-> 2 NH3(g) + 92kJ

a) Use the le chateliers principle to explain why the following changes result in an increase in the yeild of ammonia
i)a decrease in temperature
ii)an increase in pressure

B) even though a reduction in temperature favours the formation of ammonia , describe why, industrially the temperatre of the reactuib chamber is increased.

2. Consider the folliowing reaction: N2(g) + 2O(g)<=> 2NO2 (g)
[n2]= 8.0 mol/l [o2]=2.0 mol/l [NO2]=4.0 mol/l
**Determine the equilibrium constant.
3. Write the equilibrium expressions (K) for each of the following reactions:
i) H2(g)+F2(g) <=> 2NF(g)
ii)4NO(g) +3O2(g)<=> 2N2O5(g)
iii)BaCO3(s)<=>BaO(s)+CO2(g)

The equilibrium constants for three differnet reactionsare
i) Ke= 1.5x10^12
ii)Ke= 0.15
iii)ke=4.3x10^.15
In which reactionis a)The ration the bigest b)smallest

Answer 1

You should pay more attention in class. You don't even have the name right.

The name of the process is "Haber". A person's name is important and not to be tampered. This person's name was Fritz Haber. It is important because the Nazis destroyed him emotionally and now you have destroyed his name.

Spens some time on Google and you will get the answers you need.

The Haber process is run at high pressures because that stress caused more ammonia to be produced. A low temperature would help too, but it slows down the reation rate too much, so it is run at a high temp. despite Le Chatelier's Principle