Sodium Thiosulfate and Hydrochloric Acid. Why is it zero order?
The reaction between Na2S2O3 and HCl, with respect to HCl, is zero order. Why is this? I investigated a reaction between 0.25M of Na2S2O3 and 0.5M, 0.7M, 0.9M, 1.1M, and 1.3M of HCl and it appears that the reaction rate is zero order, until we look at 0.5M, where the reaction rate sharply decreases.
Would I be right in saying that it is zero order and first-order, depending on the concentration? I made sure that the volumes and concentrations are stoichiometrically balanced and the range of concentrations for HCl are so that the max concentration is balanced, and lower concentrations are not enough. So when you are around the high-ish concentrations like 0.7M, 0.9M, 1.1M, and 1.3M, then the H+ ions are reacting as fast as they can, even though there are not enough, or just enough of them. However, if the concentration is dropped below a certain point, there is just not enough H+ ions to sustain a reaction where the HCl is not the limiting factor, and then the reaction rate drops. Am I right in thinking this?
Is there any more information that I am missing?
Would I be right in saying that it is zero order and first-order, depending on the concentration? I made sure that the volumes and concentrations are stoichiometrically balanced and the range of concentrations for HCl are so that the max concentration is balanced, and lower concentrations are not enough. So when you are around the high-ish concentrations like 0.7M, 0.9M, 1.1M, and 1.3M, then the H+ ions are reacting as fast as they can, even though there are not enough, or just enough of them. However, if the concentration is dropped below a certain point, there is just not enough H+ ions to sustain a reaction where the HCl is not the limiting factor, and then the reaction rate drops. Am I right in thinking this?
Is there any more information that I am missing?
Reply 1
Original post
by FightingFalcon16The reaction between Na2S2O3 and HCl, with respect to HCl, is zero order. Why is this? I investigated a reaction between 0.25M of Na2S2O3 and 0.5M, 0.7M, 0.9M, 1.1M, and 1.3M of HCl and it appears that the reaction rate is zero order, until we look at 0.5M, where the reaction rate sharply decreases.
Would I be right in saying that it is zero order and first-order, depending on the concentration? I made sure that the volumes and concentrations are stoichiometrically balanced and the range of concentrations for HCl are so that the max concentration is balanced, and lower concentrations are not enough. So when you are around the high-ish concentrations like 0.7M, 0.9M, 1.1M, and 1.3M, then the H+ ions are reacting as fast as they can, even though there are not enough, or just enough of them. However, if the concentration is dropped below a certain point, there is just not enough H+ ions to sustain a reaction where the HCl is not the limiting factor, and then the reaction rate drops. Am I right in thinking this?
Is there any more information that I am missing?
Would I be right in saying that it is zero order and first-order, depending on the concentration? I made sure that the volumes and concentrations are stoichiometrically balanced and the range of concentrations for HCl are so that the max concentration is balanced, and lower concentrations are not enough. So when you are around the high-ish concentrations like 0.7M, 0.9M, 1.1M, and 1.3M, then the H+ ions are reacting as fast as they can, even though there are not enough, or just enough of them. However, if the concentration is dropped below a certain point, there is just not enough H+ ions to sustain a reaction where the HCl is not the limiting factor, and then the reaction rate drops. Am I right in thinking this?
Is there any more information that I am missing?
You are absolutely right! Your analysis is spot on. Let me break it down for you:
1.
At high concentrations of HCl (0.7M, 0.9M, 1.1M and 1.3M) the reaction is indeed zero order. This means that the rate of the reaction is independent of the concentration of HCl. In this regime, the reaction rate depends on how fast the H^+ ions react with Na2S2O3, rather than on the availability of H^+ ions.
2.
However, at lower concentrations of HCl (0.5 M) the reaction rate decreases rapidly.
3.
It can be seen that the availability of H^+ ions is limiting the reaction. Your explanation of this behaviour is correct: at high concentrations of HCl, there are sufficient H^+ ions to sustain the reaction and other factors limit the reaction rate.
4.
However, at lower concentrations, the availability of H^+ ions becomes the limiting reactant and the reaction rate decreases. Edit: A limiting reactant is the first reactant in a chemical reaction that produces the least amount of product, either by being completely consumed and by stopping the chemical reaction.
5.
To further support your conclusion, you could try plotting the reaction rate against the concentration of HCl. If the reaction is zero order, the plot shall be a horizontal line. If the reaction is first order, the plot shall be a straight line with a positive slope.
Kind regards from Italy!
Attachment not found
Sandro
Reply 2
Original post
by NitrotolueneYou are absolutely right! Your analysis is spot on. Let me break it down for you:
Well done on your observation and analysis! Bravo!😀
Kind regards from Italy!
Sandro
1.
At high concentrations of HCl (0.7M, 0.9M, 1.1M and 1.3M) the reaction is indeed zero order. This means that the rate of the reaction is independent of the concentration of HCl. In this regime, the reaction rate depends on how fast the H^+ ions react with Na2S2O3, rather than on the availability of H^+ ions.
2.
However, at lower concentrations of HCl (0.5 M) the reaction rate decreases rapidly.
3.
It can be seen that the availability of H^+ ions is limiting the reaction. Your explanation of this behaviour is correct: at high concentrations of HCl, there are sufficient H^+ ions to sustain the reaction and other factors limit the reaction rate.
4.
However, at lower concentrations, the availability of H^+ ions becomes the limiting reactant and the reaction rate decreases.
5.
To further support your conclusion, you could try plotting the reaction rate against the concentration of HCl. If the reaction is zero order, the plot shall be a horizontal line. If the reaction is first order, the plot shall be a straight line with a positive slope.
Kind regards from Italy!
Attachment not found
Sandro
Thank you for your help!
Reply 3
For my data table that I am inserting into the report, my data samples all have 4 significant figures. Therefore my average has 4 significant figures. Therefore my uncertainty of the mean (max-min)/2 and percent uncertainty ((uncertainty of the mean/average)*100) both have 4 significant figures. However, should I show them as 4 significant figures or would it be better style to round uncertainty and percent uncertainty to 2 decimal places?
Reply 4
Original post
by FightingFalcon16The reaction between Na2S2O3 and HCl, with respect to HCl, is zero order. Why is this? I investigated a reaction between 0.25M of Na2S2O3 and 0.5M, 0.7M, 0.9M, 1.1M, and 1.3M of HCl and it appears that the reaction rate is zero order, until we look at 0.5M, where the reaction rate sharply decreases.
Would I be right in saying that it is zero order and first-order, depending on the concentration? I made sure that the volumes and concentrations are stoichiometrically balanced and the range of concentrations for HCl are so that the max concentration is balanced, and lower concentrations are not enough. So when you are around the high-ish concentrations like 0.7M, 0.9M, 1.1M, and 1.3M, then the H+ ions are reacting as fast as they can, even though there are not enough, or just enough of them. However, if the concentration is dropped below a certain point, there is just not enough H+ ions to sustain a reaction where the HCl is not the limiting factor, and then the reaction rate drops. Am I right in thinking this?
Is there any more information that I am missing?
Would I be right in saying that it is zero order and first-order, depending on the concentration? I made sure that the volumes and concentrations are stoichiometrically balanced and the range of concentrations for HCl are so that the max concentration is balanced, and lower concentrations are not enough. So when you are around the high-ish concentrations like 0.7M, 0.9M, 1.1M, and 1.3M, then the H+ ions are reacting as fast as they can, even though there are not enough, or just enough of them. However, if the concentration is dropped below a certain point, there is just not enough H+ ions to sustain a reaction where the HCl is not the limiting factor, and then the reaction rate drops. Am I right in thinking this?
Is there any more information that I am missing?
I reckon your analysis is correct. It is well known that once a reagent is in excess, its concentration changes relatively little and so the reaction order with respect to it is for all intents and purposes zero for the reasons you have outlined. It’s why large excesses of all but one reagent are used in continuous monitoring methods when trying to find the order of reaction with respect to a given species.
The best way to confirm it would be to repeat the experiment, but make it so the concentrations of HCl you use are less than 0.5 M.
Quickly doing a google search, I have found a YouTube video where they determined the orders of reaction with respect to each reagent and they found that the orders of reaction with respect to both Na2S2O3 and HCl were 1.
https://m.youtube.com/watch?v=SZFcAa19YFQ
Reply 5
Another reference well worth looking at: https://www.docbrown.info/page03/ASA2rates27.htm
I had completely forgotten about the possibility of there being a complex rate law. Unless you are an undergrad, I seriously doubt you’ll need to worry about these.
I had completely forgotten about the possibility of there being a complex rate law. Unless you are an undergrad, I seriously doubt you’ll need to worry about these.
Reply 6
Original post
by TypicalNerdAnother reference well worth looking at: https://www.docbrown.info/page03/ASA2rates27.htm
I had completely forgotten about the possibility of there being a complex rate law. Unless you are an undergrad, I seriously doubt you’ll need to worry about these.
I had completely forgotten about the possibility of there being a complex rate law. Unless you are an undergrad, I seriously doubt you’ll need to worry about these.
Ha. Ha. Ha. Highschool. Ha.
And in Australia. Ha. Funny.
And I am out of lessons to work on the assessment so no chance of repeating the experiment.
The order of reaction should be both zero and one. What concentrations were the people in the youtube video doing?
Also I think there is some miscommunication. I am investigating the reaction with respect to HCl. I did the stoichiometry and ensured that HCl was not in excess. Yet the order was still zero. Only the lowest concentration showed signs of a first order. So all concentrations except 1.3M HCl are limiting. 1.3M was control.
Reply 7
Original post
by FightingFalcon16Ha. Ha. Ha. Highschool. Ha.
And in Australia. Ha. Funny.
And I am out of lessons to work on the assessment so no chance of repeating the experiment.
The order of reaction should be both zero and one. What concentrations were the people in the youtube video doing?
Also I think there is some miscommunication. I am investigating the reaction with respect to HCl. I did the stoichiometry and ensured that HCl was not in excess. Yet the order was still zero. Only the lowest concentration showed signs of a first order. So all concentrations except 1.3M HCl are limiting. 1.3M was control.
And in Australia. Ha. Funny.
And I am out of lessons to work on the assessment so no chance of repeating the experiment.
The order of reaction should be both zero and one. What concentrations were the people in the youtube video doing?
Also I think there is some miscommunication. I am investigating the reaction with respect to HCl. I did the stoichiometry and ensured that HCl was not in excess. Yet the order was still zero. Only the lowest concentration showed signs of a first order. So all concentrations except 1.3M HCl are limiting. 1.3M was control.
Do they actually make you do steady state and pre-equilibrium approximations etc in high school in Australia? In most other places those would be first year undergrad level topics.
I was aware you were interested in the order with respect to HCl as opposed to the order with respect to Na2S2O3 - I made sure to include sources that had some input on what both should be. Different sources say different things and it seems no one can agree whether the order of reaction with respect to HCl is 0, 0.5 or 1.
So presumably the concentrations you listed in your first post were the stock solutions and not the concentrations of each species in the final solution?
In the video, each solution has a volume of 4 cm^3. Given the volumes and the initial concentrations of the reagents used specified in the video, the concentrations of each substance in each solution are as follows:
Solution 1: [Na2S2O3] = 0.1 M, [HCl] = 3 M
Solution 2: [Na2S2O3] = 0.1 M, [HCl] = 1.5 M
Solution 3: [Na2S2O3] = 0.05 M, [HCl] = 3 M
Solution 4: [Na2S2O3] = 0.025 M, [HCl] = 1.5 M
(The undiluted solutions of Na2S2O3 and HCl have concentrations of 0.2 M and 6 M respectively)
So in all four solutions prepared in the video, one would expect HCl to be in quite a large excess. However, using excel they have reached the conclusion that HCl has a reaction order of 1. This is done by calculating ratios of initial rates etc - possibly a completely different approach to how you’ve done your experiment.
(edited 1 year ago)
Reply 8
Original post
by TypicalNerdDo they actually make you do steady state and pre-equilibrium approximations etc in high school in Australia? In most other places those would be first year undergrad level topics.
I was aware you were interested in the order with respect to HCl as opposed to the order with respect to Na2S2O3 - I made sure to include sources that had some input on what both should be. Different sources say different things and it seems no one can agree whether the order of reaction with respect to HCl is 0, 0.5 or 1.
So presumably the concentrations you listed in your first post were the stock solutions and not the concentrations of each species in the final solution?
In the video, each solution has a volume of 4 cm^3. Given the volumes and the initial concentrations of the reagents used specified in the video, the concentrations of each substance in each solution are as follows:
Solution 1: [Na2S2O3] = 0.1 M, [HCl] = 3 M
Solution 2: [Na2S2O3] = 0.1 M, [HCl] = 1.5 M
Solution 3: [Na2S2O3] = 0.05 M, [HCl] = 3 M
Solution 4: [Na2S2O3] = 0.025 M, [HCl] = 1.5 M
(The undiluted solutions of Na2S2O3 and HCl have concentrations of 0.2 M and 6 M respectively)
So in all four solutions prepared in the video, one would expect HCl to be in quite a large excess. However, using excel they have reached the conclusion that HCl has a reaction order of 1. This is done by calculating ratios of initial rates etc - possibly a completely different approach to how you’ve done your experiment.
I was aware you were interested in the order with respect to HCl as opposed to the order with respect to Na2S2O3 - I made sure to include sources that had some input on what both should be. Different sources say different things and it seems no one can agree whether the order of reaction with respect to HCl is 0, 0.5 or 1.
So presumably the concentrations you listed in your first post were the stock solutions and not the concentrations of each species in the final solution?
In the video, each solution has a volume of 4 cm^3. Given the volumes and the initial concentrations of the reagents used specified in the video, the concentrations of each substance in each solution are as follows:
Solution 1: [Na2S2O3] = 0.1 M, [HCl] = 3 M
Solution 2: [Na2S2O3] = 0.1 M, [HCl] = 1.5 M
Solution 3: [Na2S2O3] = 0.05 M, [HCl] = 3 M
Solution 4: [Na2S2O3] = 0.025 M, [HCl] = 1.5 M
(The undiluted solutions of Na2S2O3 and HCl have concentrations of 0.2 M and 6 M respectively)
So in all four solutions prepared in the video, one would expect HCl to be in quite a large excess. However, using excel they have reached the conclusion that HCl has a reaction order of 1. This is done by calculating ratios of initial rates etc - possibly a completely different approach to how you’ve done your experiment.
Firstly, the state I live in has a really weird setup for the high school curriculum. This board has been recently implemented and they are still trying to find their feet. Some of the rules in place make you really wonder.
We have never even heard of steady state and pre-equilibrium approximations. (I'm going to go and study it myself though now.)
It seems like all of the sources online investigate the reaction with respect to Na2S2O3. Not HCl. And as you said, the HCl is in excess in your video, whereas in my experiment the Na2S2O3 is in excess. Hence the conclusion of a first-order reaction. The point at which both concentrations are balanced was 0.25M of Na2S2O3 and I thiinnk 1.2M HCl. So HCl was way in excess for that video.
Yes the concentrations were for the reactants.
Reply 9
Original post
by FightingFalcon16Firstly, the state I live in has a really weird setup for the high school curriculum. This board has been recently implemented and they are still trying to find their feet. Some of the rules in place make you really wonder.
We have never even heard of steady state and pre-equilibrium approximations. (I'm going to go and study it myself though now.)
It seems like all of the sources online investigate the reaction with respect to Na2S2O3. Not HCl. And as you said, the HCl is in excess in your video, whereas in my experiment the Na2S2O3 is in excess. Hence the conclusion of a first-order reaction. The point at which both concentrations are balanced was 0.25M of Na2S2O3 and I thiinnk 1.2M HCl. So HCl was way in excess for that video.
Yes the concentrations were for the reactants.
We have never even heard of steady state and pre-equilibrium approximations. (I'm going to go and study it myself though now.)
It seems like all of the sources online investigate the reaction with respect to Na2S2O3. Not HCl. And as you said, the HCl is in excess in your video, whereas in my experiment the Na2S2O3 is in excess. Hence the conclusion of a first-order reaction. The point at which both concentrations are balanced was 0.25M of Na2S2O3 and I thiinnk 1.2M HCl. So HCl was way in excess for that video.
Yes the concentrations were for the reactants.
I see. That’s fair. The reason I asked about steady state and pre-equilibrium approximations was because with the complex rate laws they are used to derive, you can get some rather weird relationships between rate and concentrations where it is apparently first order with respect to one reagent and then not quite so straightforward with the order with respect to another reagent. This is likely why most experiments don’t tend to investigate the order with respect to HCl and why you haven’t found a lot.
A well-known example of the above is the reaction of H2 with Br2 where it is apparently first order with respect to H2, but the rate equation itself is a fraction with a [Br2] term on the numerator and the denominator. You can find the derivations online if you are interested.
You can work out the orders of reaction with respect to both using some of the results, but it does depend on what method of monitoring the rate is used. In the video, they managed it and on the Doc Brown website they have said the following:
“You would expect that the rate might be controlled by the interaction of the negative thiosulfate ion and a positive hydrogen ion. You would expect the interaction of oppositely charged ions to have a relatively low activation energy, so in the rate expression:
rate = k[S2O32–(aq)]t[H+(aq]h, you might expect the order t and h to be both 1.
•
t=1 is quoted on the web. and found to be so in most reliable experiments (but not all).
•
The reaction has been shown to be a multi-step complex mechanism, so what order h is I don't know?
•
I've come across references that indicates the order h could be 0–1 depending on the relative concentrations of thiosulfate and acid.
•
Whatever, the orders t and h can only be found by experiment and the mechanism is likely to be complex”
I believe the last part is highlighted in bold on the website since the mechanism hasn’t been confirmed and so finding the orders theoretically with the appropriate approximation cannot (yet) be done. The likelihood is the mechanism is complex and the rate law may not even take the form rate = k[HCl]^h[Na2S2O3]^t.
They also included a reference to a research paper from 1958 that concluded that both reagents have fractional orders (1.5 for thiosulphate, 0.5 for hydrochloric acid).
Reply 10
Original post
by TypicalNerdI see. That’s fair. The reason I asked about steady state and pre-equilibrium approximations was because with the complex rate laws they are used to derive, you can get some rather weird relationships between rate and concentrations where it is apparently first order with respect to one reagent and then not quite so straightforward with the order with respect to another reagent. This is likely why most experiments don’t tend to investigate the order with respect to HCl and why you haven’t found a lot.
A well-known example of the above is the reaction of H2 with Br2 where it is apparently first order with respect to H2, but the rate equation itself is a fraction with a [Br2] term on the numerator and the denominator. You can find the derivations online if you are interested.
You can work out the orders of reaction with respect to both using some of the results, but it does depend on what method of monitoring the rate is used. In the video, they managed it and on the Doc Brown website they have said the following:
“You would expect that the rate might be controlled by the interaction of the negative thiosulfate ion and a positive hydrogen ion. You would expect the interaction of oppositely charged ions to have a relatively low activation energy, so in the rate expression:
rate = k[S2O32–(aq)]t[H+(aq]h, you might expect the order t and h to be both 1.
I believe the last part is highlighted in bold on the website since the mechanism hasn’t been confirmed and so finding the orders theoretically with the appropriate approximation cannot (yet) be done. The likelihood is the mechanism is complex and the rate law may not even take the form rate = k[HCl]^h[Na2S2O3]^t.
They also included a reference to a research paper from 1958 that concluded that both reagents have fractional orders (1.5 for thiosulphate, 0.5 for hydrochloric acid).
A well-known example of the above is the reaction of H2 with Br2 where it is apparently first order with respect to H2, but the rate equation itself is a fraction with a [Br2] term on the numerator and the denominator. You can find the derivations online if you are interested.
You can work out the orders of reaction with respect to both using some of the results, but it does depend on what method of monitoring the rate is used. In the video, they managed it and on the Doc Brown website they have said the following:
“You would expect that the rate might be controlled by the interaction of the negative thiosulfate ion and a positive hydrogen ion. You would expect the interaction of oppositely charged ions to have a relatively low activation energy, so in the rate expression:
rate = k[S2O32–(aq)]t[H+(aq]h, you might expect the order t and h to be both 1.
•
t=1 is quoted on the web. and found to be so in most reliable experiments (but not all).
•
The reaction has been shown to be a multi-step complex mechanism, so what order h is I don't know?
•
I've come across references that indicates the order h could be 0–1 depending on the relative concentrations of thiosulfate and acid.
•
Whatever, the orders t and h can only be found by experiment and the mechanism is likely to be complex”
I believe the last part is highlighted in bold on the website since the mechanism hasn’t been confirmed and so finding the orders theoretically with the appropriate approximation cannot (yet) be done. The likelihood is the mechanism is complex and the rate law may not even take the form rate = k[HCl]^h[Na2S2O3]^t.
They also included a reference to a research paper from 1958 that concluded that both reagents have fractional orders (1.5 for thiosulphate, 0.5 for hydrochloric acid).
Oh wow that is a lot to take in one go! Yeah they teach you something in high school only to tell you later how the things taught in Highschoolers are either models and completely wrong in places or how oversimplified it is.
Thank you for all of that! I think I'll stick to the order being both 1 and 0 as I am limited to 2000 words. haha. If I started to explain the concepts you suggested I would blow over the limit again!
Thank youuu!
Reply 11
Original post
by FightingFalcon16Oh wow that is a lot to take in one go! Yeah they teach you something in high school only to tell you later how the things taught in Highschoolers are either models and completely wrong in places or how oversimplified it is.
Thank you for all of that! I think I'll stick to the order being both 1 and 0 as I am limited to 2000 words. haha. If I started to explain the concepts you suggested I would blow over the limit again!
Thank youuu!
Thank you for all of that! I think I'll stick to the order being both 1 and 0 as I am limited to 2000 words. haha. If I started to explain the concepts you suggested I would blow over the limit again!
Thank youuu!
I think that is completely fair. As long as it is a conclusion you can fully justify (which it would appear to be from your first post), then that would be the best approach.
Reply 12
Original post
by FightingFalcon16Ha. Ha. Ha. Highschool. Ha.
And in Australia. Ha. Funny.
And I am out of lessons to work on the assessment so no chance of repeating the experiment.
The order of reaction should be both zero and one. What concentrations were the people in the youtube video doing?
Also I think there is some miscommunication. I am investigating the reaction with respect to HCl. I did the stoichiometry and ensured that HCl was not in excess. Yet the order was still zero. Only the lowest concentration showed signs of a first order. So all concentrations except 1.3M HCl are limiting. 1.3M was control.
And in Australia. Ha. Funny.
And I am out of lessons to work on the assessment so no chance of repeating the experiment.
The order of reaction should be both zero and one. What concentrations were the people in the youtube video doing?
Also I think there is some miscommunication. I am investigating the reaction with respect to HCl. I did the stoichiometry and ensured that HCl was not in excess. Yet the order was still zero. Only the lowest concentration showed signs of a first order. So all concentrations except 1.3M HCl are limiting. 1.3M was control.
Hey I am doing a similar experiment changing the concentration of HCl. Would you mind if I have a look at your report, as I am a little confused at the moment
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