Equilibrium... (1 Viewer)

[Aerlinn]

Quick question.
The equilibirium constant for the reaction given by the equation:
2HI(g) <---> H2(g) + I2(g) is 48.8 at 455C. An equilibrium mixture in a 2.0L vessel at this temperature contains 0.200mol of H2 and 0.100mol of I2
a. Calculate the concentration of HI in this mixture
Hmm... is it crucial to use the equilibrium constant in this question... wHy?!

b. Another mixture was prepared by placing 4.0mol of HI in a 2.0L vessel at 330C. At equilibrium 0.44mol of H2 and 0.44mol of I2 were present. Calculate the value of the equilibrium constant at this temperature.

(Ans. for a is 0.011M, Ans. for b is 0.020)...


~~~
Wanted to make some things clear--
About the influence of various things on equilibrium...
Adding extra reactant or product/ varying concentration-- Affects position of equilibrium, not Kc.
Changing pressure--Affects position of equilibrium, not Kc.
Dilution-- Affects position of equilibrium, not Kc.
Changing temp-- Affects both.
Adding catalysts-- Affects neither
Correct me if I'm wrong.

Are Kc and K are essentially the same thing?
Also, why does an inert gas, even though it reduces pressure, not have an affect on equilibrium when reducing volume reduces pressure too, and DOES have an affect?

 

[Aerlinn]

During the Contact process for sulfuric acid manufacture, sulfur dioxide is converted to sulfur trioxide at temperatures of 400-500C:
2SO2(g) + O2 <---> 2SO3(g); dH= -197kJ mol
-In practice, increasing pressure (on this reaction) is usually performed at atmospheric pressures. Suggest why.
(Uh... To compare so you can ascertain how much pressure you've added?)
-During the process, sulfur trioxide is removed from the reaction mixture by converting it to sulfuric acid. The remaining gases are then recycled to the reaction vessel. Explain the reason for recycling the gases.
-What factors would have influenced the choice of the reaction temperature... and why?

A question that left me stumped :S
Help= great

 

[xiao1985]

re post#1:

a)K = [prod]/[reactant] = [H2][I2]/[HI]^2 = 48.8

[H2]=0.2/2=0.1 M ; [I2] = 0.1/2=0.05 M ; [HI]^2 = 0.1*0.05/48.8 = 1.02e-4 M, [HI] = 0.010M (answer is wrong)
yes, it is crucial to use mole constant for this q... how else can you find [HI]??


b) 0.44 mol of H2 and I2 present = 0.88moles of HI reacted = 3.12 moles of HI remaining...
0.44mol of H2 and I2 = 0.22 M of [H2] and [I2] ; 3.12 mol of HI = 1.56 of [HI]
K=[H2][I2]/[HI]^2 = 0.020

statements: what if they are not wrong?

Kc vs K: I would think so... I've never heard of Kc before though... what's Kc?

Inert gas: I would believe that introduction of inert gas would INCREASE pressure... i don't quite get the question...



POST#2

- it's much more expensive to increase pressure (use a highly firm rubber, and high powered piston)
- again economic reasons.... if the residue unreacted stuff is recycled, you won't have to chuck it away (garbage cost) and you make more stuff which is valuable (H2SO4)
- balancing reaction rate with reaction yield

 

[Aerlinn]

Thanx

how else can you find [HI]??

I wasn't thinking too much first time round and for some reason decided, 'cause it gives you the no. moles and the volume, I found the 'the HI concentration' (which wasn't really the answer) without using K. I can see it gets you totally the wrong answer... but... why? ^^

Well, inert gases do increase pressure, and so does reducing volume. So my question is, why does reducing volume effect equilibrium, but inert gases don't? (because apparently pressure affects equilibrium)

what's Kc?

Kc= a equilibrium constant symbol, like K apparently... hmm...

balancing reaction rate with reaction yield

What sure I entirely get what you mean with this one...

 

[Aerlinn]

pH buffers

Came across...
'Control of pH in the blood is achieved using different buffers. One of the important buffers is made up of carbonic acid (H2CO3) and the hydrogen carbonate ion (HCO3-):
H2CO3(aq) + H20(l) <---> HCO#-(aq) + H3O+(aq)
The action of this buffer system becomes clear is we consider what happens when an acid or base is added to blood:
-If H2O+ ions were added, a net back reaction occurs, removing most of these ions.
-If OH- ions were added, they react with H3O+ ions. A net forward reaction results, producing more H3O+ ions. The OH- ions have effectively reacted with the carbonic acid.'

The second dot point ^, is somewhat confusing. Especially the last sentence...

 

[xiao1985]

Re: pH buffers

1st post:
inert gas imo should influence the equilibrium position... but in this case, you are reacting 2moles of gas <---> 2 moles of gas, which means increasing/decreasing pressure has no effect on system, holding other variable constant... which means, even if you add inert gases to the chamber to increase pressure, it won't influence the equilibrium position.

2nd post:
first of all, it's H3O+ (yes, it makes a big differnce between H3O+ and H2O+.)
2ndly, adding H3O+ (say you added HCl, 10M, 15mL) to the system will cause the equilibrium to go backwards to minimise the change of H3O+ due to le chatelier principle.

effect: minimise pH change

adding OH- react with H3O+: H3O+ + OH- --> 2 H2O. the infamous acid base neutralisation reaction.
which means adding OH-, effectively, you are removing H3O+ from the system, cause the equilibrium to shift to the right, to minimise the used-up of H3O+

effect: minimise pH change