HSC 2012-2015 Chemistry Marathon (archive) (7 Viewers)

[Drsoccerball]

re: HSC Chemistry Marathon Archive

You should know this. If you can't curium, nor can you helium, you have to barium. It appears in the HSC every year!

Ok, no.
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Outline the most optimal conditions for the production of ammonia in the Haber process. Explain why these conditions are suitable. In your response, make reference to Le Chatelier's principle on the system's equilibrium. (4)

This is a horribly worded question...

Covers up the bad joke with a question :L...
The optimal conditions for the production of Ammonia is that of around 300 degrees and around 200 atmospheric pressure.
Le Chatlier's principle states that a change in pressure, concentration or temperature would cause the equilibrium to shift in the direction that minimises the disturbance. An increase in pressure would cause the equilibrium to shift to the side with the least number of moles. Thus as pressure is high the equilibrium shifts to the right increasing the yield of ammonia. However, there is a specific amount of strain the pipes can hold thus the pressure is not increased to high. A decrease in temperature, due to the systems exothermic nature, would cause the equilibrium to shift to the right. However, due to the Kinetic theory this will slow down the rate of reaction and rate of production of Ammonia. Thus in order to get a compromise higher temperature is used. Thus showing that the suitable conditions are around 300 degrees and 200 times AP

 

[hawkrider]

re: HSC Chemistry Marathon Archive

Covers up the bad joke with a question :L...
The optimal conditions for the production of Ammonia is that of around 300 degrees and around 200 atmospheric pressure.
Le Chatlier's principle states that a change in pressure, concentration or temperature would cause the equilibrium to shift in the direction that minimises the disturbance. An increase in pressure would cause the equilibrium to shift to the side with the least number of moles. Thus as pressure is high the equilibrium shifts to the right increasing the yield of ammonia. However, there is a specific amount of strain the pipes can hold thus the pressure is not increased to high. A decrease in temperature, due to the systems exothermic nature, would cause the equilibrium to shift to the right. However, due to the Kinetic theory this will slow down the rate of reaction and rate of production of Ammonia. Thus in order to get a compromise higher temperature is used. Thus showing that the suitable conditions are around 300 degrees and 200 times AP

Solid response, except I'm not 100% sure about this bit. I thought that it has to be moderate temperature when monitored to ensure that (like you said in your answer) an appropriate compromise is maintained between maximum rate (achieved at higher temperature) and maximum yield (achieved at low temperature).

 

[Drsoccerball]

re: HSC Chemistry Marathon Archive

Solid response, except I'm not 100% sure about this bit. I thought that it has to be moderate temperature when monitored to ensure that (like you said in your answer) an appropriate compromise is maintained between maximum rate (achieved at higher temperature) and maximum yield (achieved at low temperature).

Yes sorry should be moderate temperature

 

[hawkrider]

re: HSC Chemistry Marathon Archive

Yes sorry should be moderate temperature

all g brah

tbh, it's those minute things in short answers that can ring alarm bells in markers and thus you'd get marked down for using incorrect terminology (as you would know - chem marking is anal af)

 

[Ekman]

re: HSC Chemistry Marathon Archive

Covers up the bad joke with a question :L...
The optimal conditions for the production of Ammonia is that of around 300 degrees and around 200 atmospheric pressure.
Le Chatlier's principle states that a change in pressure, concentration or temperature would cause the equilibrium to shift in the direction that minimises the disturbance. An increase in pressure would cause the equilibrium to shift to the side with the least number of moles. Thus as pressure is high the equilibrium shifts to the right increasing the yield of ammonia. However, there is a specific amount of strain the pipes can hold thus the pressure is not increased to high. A decrease in temperature, due to the systems exothermic nature, would cause the equilibrium to shift to the right. However, due to the Kinetic theory this will slow down the rate of reaction and rate of production of Ammonia. Thus in order to get a compromise higher temperature is used. Thus showing that the suitable conditions are around 300 degrees and 200 times AP

Just to add on a bit there, the presence of a catalyst (ferric oxide Fe2O3) and the constant addition of N2 and H2, and the removal of NH3 at the rate required according to the mole ratio, are also conditions required in the Haber process. The use of the catalyst is straight forward, it minimises the temperature compromise, allowing the reaction to continue at higher rates and producing higher yields. And the constant addition of N2 and H2, and the removal of the NH3 is due to LCP.

 

[Librah]

re: HSC Chemistry Marathon Archive

Moderate temperature could mean 1x10^8 degrees celsius. Some things go up to temperatures in the billions.

 

[Drsoccerball]

re: HSC Chemistry Marathon Archive

Just to add on a bit there, the presence of a catalyst (ferric oxide Fe2O3) and the constant addition of N2 and H2, and the removal of NH3 at the rate required according to the mole ratio, are also conditions required in the Haber process. The use of the catalyst is straight forward, it minimises the temperature compromise, allowing the reaction to continue at higher rates and producing higher yields. And the constant addition of N2 and H2, and the removal of the NH3 is due to LCP.

These wern't mentioned due to the nature of the question. Unless you could relate the use of a catalyst to le chatiliers principle (which you can through lower activation energy) i don't see the use in this question. Maybe if the question was how is the process monitored these can be mentioned? but a low mark allocation also lead me to believe that what i wrote is all you need. Correct me if im wrong

 

[leehuan]

re: HSC Chemistry Marathon Archive

In the exam I'd mention it just in case. However, for the question I wrote, I decided to let the catalyst slip by because it would've been a fifth or sixth mark. And yes, should there be a need to mention the catalyst, do remember to relate it back to the activation energy.

Nice response nonetheless!
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This time, you are given unknown solutions of sodium phosphate and sodium sulfate. You are given nitric acid, barium nitrate and sodium hydroxide solution to distinguish between the anions.

(I don't see enough of a point in asking for the method again.)

(i) Explain why the nitric acid is used to acidify solutions before precipitating. (2)
(ii) Write suitable chemical equations to account for the expected precipitates. (2)
(iii) Use Le Chatelier's principle to predict whether there will be more phosphate ions, or hydrogen phosphate ions present when nitric acid is added. (Ignore dihydrogen phosphate and phosphoric acid for these intensive purposes.) (2)
(iv) Hence, explain why the acidified mixture of sodium phosphate needs added sodium hydroxide for the precipitate to form. Contrast this to that of the mixture involving sodium sulfate. (5)
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I promise, next question I get to write won't be anion testing.

 

[porcupinetree]

re: HSC Chemistry Marathon Archive

This time, you are given unknown solutions of sodium phosphate and sodium sulfate. You are given nitric acid, barium nitrate and sodium hydroxide solution to distinguish between the anions.

(I don't see enough of a point in asking for the method again.)

(i) Explain why the nitric acid is used to acidify solutions before precipitating. (2)
(ii) Write suitable chemical equations to account for the expected precipitates. (2)
(iii) Use Le Chatelier's principle to predict whether there will be more phosphate ions, or hydrogen phosphate ions present when nitric acid is added. (Ignore dihydrogen phosphate and phosphoric acid for these intensive purposes.) (2)
(iv) Hence, explain why the acidified mixture of sodium phosphate needs added sodium hydroxide for the precipitate to form. Contrast this to that of the mixture involving sodium sulfate. (5)
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I promise, next question I get to write won't be anion testing.

Here's my attempt; keep in mind that I'm fairly new to learning all of the ion tests / solubility rules.

i.If we add HNO3, and then add Ba(NO3)2, the precipitate formed by sulfate (BaSO4) is insoluble in HNO3, whereas the phosphate precipitate (Ba3(PO4)2) is soluble and will not form a visible solid.

ii.Na2SO4 (aq) + Ba(NO3)2 (aq) --> BaSO4 (s) + 2NaNO3 (aq)
2Na3PO4 (aq) + 3Ba(NO3)2 (aq) --> Ba3(PO4)2 (aq) + 6NaNO3 (aq)

iii.The equation we are concerned with is:
H+(aq) + PO43-(aq) <--> HPO42-(aq)
According to Le Chatelier’s principle, adding H+ in the form of nitric acid will shift the equilibrium to the right, forming more hydrogen phosphate.

iv.The acidified mixture of sodium phosphate has many hydrogen phosphate ions present, which will not form a precipitate with Ba2+. However, if we can reduce the number of hydrogen phosphate ions and increase the number of phosphate ions, we can get a significant amount of Ba3(PO4)2 to form. Adding sodium hydroxide effectively decreases the number of H+ in the above equilibrium reaction, which according to LCP will shift the equilibrium to the left, forming more phosphate ions (along with H+), allowing for significantly more Ba3(PO4)2 to form.
I’m not entirely sure about how this is different to the sulfate solution, however, I am willing to guess that the equilibrium reaction:
H+ + SO42- <--> HSO4- naturally lies more to the left.

 

[leehuan]

re: HSC Chemistry Marathon Archive

The first part was independent to the rest of the question. It was more of a method to increase accuracy of results.

(i) Carbon dioxide is a substance naturally present. In solutions, it can form carbonate ions, which introduces an impurity to the mixture. The nitric acid will react with these carbonate ions and cause it to form carbon dioxide again. As the impurity is removed, more reliable and accurate results are obtainable.

Problem with (ii) - Barium phosphate is insoluble.

Edit: Also, your guess to part (iv) involving the sulfate is correct. But I think you probably wouldn't have been able to get the entire question out with the error in (ii)

 

[porcupinetree]

re: HSC Chemistry Marathon Archive

Problem with (ii) - Barium phosphate is insoluble.

My textbook says that it is soluble in HNO3

 

[leehuan]

re: HSC Chemistry Marathon Archive

I'm not sure if I'd say it's "soluble in HNO3", rather I thought it was just limited traces of it to be noticeable.

However, the idea is that when the NaOH is added, because the NaOH will react with the HNO3, the equilibirum system between the phosphate ion and its conjungate acid (hydrogen phosphate) will shift to produce more PO4(3-). Because there are now more phosphate ions, there will be a more noticeable reaction between the Na3PO4 and the Ba(NO3)2 and the precipitate can now be seen.

 

[Drsoccerball]

re: HSC Chemistry Marathon Archive

Explain how the Farsh process can be used to obtain Sulfur. 4 Marks

 

[Crisium]

re: HSC Chemistry Marathon Archive

Explain how the Farsh process can be used to obtain Sulfur. 4 Marks

Never heard of the Farsh process before

But I have heard of the Frasch process if you want me to answer the question on that

 

[Dan036]

re: HSC Chemistry Marathon Archive

I'm quite confused on how it quickly changed from chemical monitoring into industrial chemistry.

 

[leehuan]

re: HSC Chemistry Marathon Archive

I'm nowhere near industrial chemistry

But personally, I'm not sure if posting questions on the option topics is a good idea, since it's not like every school does industrial chemistry.

 

[Drsoccerball]

re: HSC Chemistry Marathon Archive

Never heard of the Farsh process before

But I have heard of the Frasch process if you want me to answer the question on that

nah man whats that

 

[leehuan]

re: HSC Chemistry Marathon Archive

I'll let Crisium do that question.
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I'm still not too far into chemical monitoring and management yet, so I'll have to go backwards

Explain why previous definitions of acids and bases needed to be replaced, and account for how the Bronsted-Lowry theory became the most accepted theory today, giving examples of Bronsted-Lowry acids and bases. (5)

 

[porcupinetree]

re: HSC Chemistry Marathon Archive

I'll let Crisium do that question.
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I'm still not too far into chemical monitoring and management yet, so I'll have to go backwards

Explain why previous definitions of acids and bases needed to be replaced, and account for how the Bronsted-Lowry theory became the most accepted theory today, giving examples of Bronsted-Lowry acids and bases. (5)

Lavoisier was the first to properly define what an acid is; suggesting that they contained oxygen, however this was flawed as many acids such as HCl do not contain oxygen, and bases such as NaOH. Davy showed that HCl didn't contain oxygen and proposed that acids contained replaceable hydrogen after noticing the displacement of hydrogen from acids by metals, however this concept failed to explain bases such as NH3. Next, Arrhenius proposed that acids ionised in solution to form H+ and an anion, while bases formed OH- and a cation. This was an improvement upon earlier definitions but could not explain bases such as NH3. The Bronsted Lowry theory has become the most accepted theory today as it improved upon and replaced these earlier definitions by asserting that acids (such as HCl) are proton donors and bases (such as NaOH or NH3) are proton acceptors, solving many issues encountered by using earlier definitions such as the classification of NH3 as a base.

 

[porcupinetree]

re: HSC Chemistry Marathon Archive

By including relevant chemical equations in your answer, justify the procedure you used to prepare an ester in a school laboratory. 6 marks