More questions on The Acidic Environment :
Question 1.
Spoiler
Is interpreting equilibrium graphs an important/necessary skill for the course?
Question 2.
Spoiler
For the outcome "Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier’s principle"
Why is Le Chaterlier's principle applied differently for an increase in concentration and volume of CO2?
"An increase in the concentration of CO2 (g) will shift the equilibrium to the right, converting carbon dioxide and water into carbonic acid in order to reduce the concentration of carbon dioxide."
I understand how this works for concentration, but for:
"An increase in the volume of CO2 (g) will shift the equilibrium to the right, converting carbon dioxide and water into carbonic acid in order to reduce the volume of carbon dioxide. Thus the system will attempt to counteract this change by favouring the backwards reaction."
It almost seems to me that LC's principle has been applied twice in the instance of increased volume (inferred in the first sentence then applied in the second). Why is this the case?
Question 2a. (Same outcome)
Spoiler
For the dissolution of CO2 in water, given by:
CO2 (g) +H2O (l) ⇌ H2CO3 (aq)
Will the rate of the forward reaction increase if the concentration of either reactant is increased?
1. I think equilibrium graphs are very important. Take for instance q12 from 2012 HSC, which is a multiple choice with equilibrium graphs. I also remember there was another long response on equilibrium graphs, which requires you to explain all 3 factors (temperature, pressure and concentration). So yes, I think it is very important and try your best to master it
2. Im not really sure if quote thing is right (increasing volume doesn't really 'shift' it to the right, I may be wrong though). concentration is given by the formula c=n/v. So by increasing the volume, you're basically decreasing the concentration of the gases. So by LCP, the system will try to increase the concentration, and this is the backwards reaction. Another way to think about this is that volume is inversely proportional to the pressure. And pressure is directly proportional the amount of gaseous moles. Thus, by increasing the volume, the pressure is decreased. So to increase the pressure, the equilibrium shifts to the side with the most GAS moles, and this is the backwards reaction. This explanation might be really confusing, so I hope someone else is able to clearly explain this
2a. Recall that LCP is not applied to pure solids and liquids, ie they do not affect the equilibrium at all. So increasing/decreasing the concentration of water has no effect on the equilibrium. So the rate of forward reaction will only increase if the concentration of CO2 is increased.
3. Basically, its not necessary for all the hydrogen in a molecule to be 'removable'. eg in ethnaoic, the H's bonded to the carbon won't ionise, only the one that is bonded with the oxygen. Similarly, If you look at the structure of citric acid, only the 3 hydrogens sticking out of the carboxylic acid groups are able to ionise, hence its tri-protic.