Is this solubility question's answer really right?

[mrh125]

the reaction is Ca(IO3)2 <=> Ca2+ + +2IO3-

Why does adding acid increase the solubility of Ca(IO3)2 (s)? I've seen questions like this on practice mcats that say stuff like you have the reaction Na2SO4 <=> 2Na+ +SO42- and adding base increases solubility (so acid shouldn't uhm wouldn't adding acid make it more soluble?), so why would adding acid increase the solubility in this particular case? and how do you recognize how these problems work?

If it was Ca(OH)2 <=> Ca2+ +OH- I'd say that adding acid would increase its solubility but in this case i'm not sure.

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[Czarcasm]

the reaction is Ca(IO3)2 <=> Ca2+ + +2IO3-

Why does adding acid increase the solubility of Ca(IO3)2 (s)? I've seen questions like this on practice mcats that say stuff like you have the reaction Na2SO4 <=> 2Na+ +SO42- and adding base increases solubility (so acid shouldn't), so why would adding acid increase the solubility in this particular case? and how do you recognize how these problems work?

If it was Ca(OH)2 <=> Ca2+ +OH- I'd say that adding acid would increase its solubility but in this case i'm not sure.

Generally, when you're adding strong acid, you're adding both H+ and it's conjugate base. Consider HCl for example. It's very acidic and so we can assume it deprotonates 100%. If you add this acid with a solution of Ca(IO3)2, what you find happens is the conjugate base of the strong acid (in my example, chloride anions (Cl-)), will combine with the calcium cation and precipitate to the bottom of the flask. The solution recognizes that fewer calcium ions now exist in the solution and so to re-establish equilibrium, it will produce more product to compensate (essentially increasing the solubility of the substance).

An important point to consider is that for a compound to even precipitate out of a solution (making it more soluble in the process), it must exceed it's Ksp value -- anything lower will simply be soluble in the beaker and dissolve. So if we were to compare the solubility of CaCl2, we would likely see it's Ksp value is considerably lower than Ca(IO3)2.

 

[mrh125]

Generally, when you're adding strong acid, you're adding both H+ and it's conjugate base. Consider HCl for example. It's very acidic and so we can assume it deprotonates 100%. If you add this acid with a solution of Ca(IO3)2, what you find happens is the conjugate base of the strong acid (in my example, chloride anions (Cl-)), will combine with the calcium cation and precipitate to the bottom of the flask. The solution recognizes that fewer calcium ions now exist in the solution and so to re-establish equilibrium, it will produce more product to compensate (essentially increasing the solubility of the substance).

An important point to consider is that for a compound to even precipitate out of a solution (making it more soluble in the process), it must exceed it's Ksp value -- anything lower will simply be soluble in the beaker and dissolve. So if we were to compare the solubility of CaCl2, we would likely see it's Ksp value is considerably lower than Ca(IO3)2.

Thank you. Got it! So in the case of the the Na2So4 <=> SO42- + 2Na+ adding a base would make it precip more instead of making it more soluble correct, while adding an acid would make it more soluble? (replace Na with a non-soluble element if you want because i think that's soluble)

 

[Czarcasm]

Thank you. Got it! So in the case of the the Na2So4 <=> SO42- + 2Na+ adding a base would make it precip more instead of making it more soluble correct, while adding an acid would make it more soluble? (replace Na with a non-soluble element if you want because i think that's soluble)

Yep. NaOH, a strong base which would dissociate to Na+ and OH-. Would this increase the reactant or product? This will ultimately tell us if it will increase or decrease solubility (based on lechat's principle). The reaction will act to re-establish equilibrium. Because dissolving NaOH would produce more product (Na+), it would therefore increase the reverse reaction, decreasing (not increasing) the solubility of the compound as it stands.

Na2SO4 <=> 2Na+ +SO42- and adding base increases solubility (so acid shouldn't), so why would adding acid increase the solubility in this particular case?

It's been so long since I've reviewed this topic, but that might have to do with complex solubility problems -- that is the product of one reaction is the reactant of another.

REACTION 1: Na2SO4 --> Na+ and SO4-
REACTION 2: Na+ and HCl --> NaCl + [H+]

This is a poor example I made up. I'm sure someone can think of something better, but in this case, adding strong base like NaOH would not only increase the product of reaction 1 (decreasing solubility), but it would also increase the solubility of reaction 2 (more reactant, therefore will shift forward to produce product). In fact, adding strong base would neutralize some of the acid, shifting reaction 2 to a greater degree and making it more soluble.

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For these types of problems, never tell yourself "acid makes this more soluble" or "base makes that less soluble." Instead focus on how adding strong acid or base will effect the reactants or products of the compound dissolving -- and then consider what will happen according to lechat's principle. This will ultimately help you answer all of these questions very easily.

 

[Dirtybird]

^^Don't mean to be rude but I'm pretty sure Cl has NOTHING to do with this question.

I would have seen this and saw that H+ would probably react withwith IO3 and form HIO3. I had to look this up and yeah, HIO3 does exist and is called iodic acid. But you could probably have figured that a dissociated salt = conjugate base of an acid.
IO3- + H+ --> HIO3
So with H+ present and creating more HIO3 (there is likely already some present due to IO3 reacting with water), there is less IO3 in the right side of the original solubility equation (Ca(IO3)2 <=> Ca2+ + 2IO3-), and due to Le Chatelier's principle, the equation will shift right and more Ca(IO3)2 will dissolve.

For the case with Na2SO4, you also have a case where (SO4)2- is probably plucking off H+ from water to form HSO4-. This equation would be (SO4)2- + H2O --> HSO4- + OH-
If you add bass, LC principle applies again and the shift makes more (SO4)2-
When there is more (SO4)2-, then this original solubiltiy equation shift to the left: Na2SO4 <=> 2Na+ +SO42-
Thus, NA2SO4 becomes less soluble due to the increased SO42- created by OH- presence (same concept as common ion principle).

Hope this makes sense, I spent a good while wrapping my head around this problem too (seen it twice already).

 

[Czarcasm]

^^Don't mean to be rude but I'm pretty sure Cl has NOTHING to do with this question.

It's cool. I only used HCl as an example. All strong acids come with a conjugate base. In my example, that conjugate base was chloride anion, which if present in the solution, would dissolve with calcium ions. That's how I typically approach these types of problems. But like you said, you can also focus on the increase of [H+] concentration and how the system will consume that (by reacting with the counter base: IO3-), to re-establish equilibrium. Either scenario works and will help you answer the question.

The reason I instead focused on the salt however is because (from a solubility perspective) this ultimately is the key to solving these problems correctly. It is the salt that will dissolve out of the solution and therefore increase solubility of the original compound.

 

[MDforMee]

Generally, when you're adding strong acid, you're adding both H+ and it's conjugate base.

I would say that unless the question says otherwise, you can't assume that an "acid" implies that a coupled proton and base are administered at the same time.

Anyways, for this problem it is helpful to think of charge instead of what may or may not form. Simply, if a + acid is added to an equilibrium of
neutral <--> negative + positive
then the reaction will favor the --> direction, and since Ksp (solubility product) is an intrinsic property that's equal to
k1/k-1 = (negative)(positive)/neutral
then a decrease in the negative form must therefore accompany a decrease in the neutral form to maintain the constant ratio, aka the reaction proceeds to the right.

Test writers won't want you to memorize the solubility tables, either.

You should also know that temperature is what effects equilibria and rates/directions, and that ions don't. G = -RTlnK and dG = dH - TdS blah blah blah exergonic endergonic blah blah then go google and wiki yourself up some physical chemistry and learn it for yourself

 

[Jepstein30]

I would say that unless the question says otherwise, you can't assume that an "acid" implies that a coupled proton and base are administered at the same time.

If it's a strong acid, that's essentially what you are doing since it will 100% deprotonate and you will be left with H+ and its conjugate base.

 

[MDforMee]

If it's a strong acid, that's essentially what you are doing since it will 100% deprotonate and you will be left with H+ and its conjugate base.

Yes and no. Yes strong acids dissociate (almost) completely, but no because the point of what I wrote is to avoid having to deal with bases interacting with other things in the solution.

For instance, if HCl is added, but you have Ag+ present in the solution somewhere, then the insoluble AgCl will form and shift your equilibrium. This is why I wrote that test takers probably won't want you to memorize the solubility tables, but to be safe instead of sorry, don't assume anything.

http://www.chem.sc.edu/faculty/morgan/resources/solubility/

 

[Jepstein30]

But you *still* have to deal with bases interacting with other things in the solution whether or not you act as if you're adding a strong acid or a proton and its conjugate base. It's the same thing considering they almost completely (for MCAT purposes, completely) dissociate.

If you add HCl and have Ag+ present.. the same shift will occur as if you added equal amounts of H+ and Cl- with Ag+ present... will it not?

 

[MDforMee]

But you *still* have to deal with bases interacting with other things in the solution whether or not you act as if you're adding a strong acid or a proton and its conjugate base. It's the same thing considering they almost completely (for MCAT purposes, completely) dissociate.

If you add HCl and have Ag+ present.. the same shift will occur as if you added equal amounts of H+ and Cl- with Ag+ present... will it not?

You shouldn't practice bad chemistry, it will lead to mistakes. Stick with what the question states, and don't assume anything.

 

[Jepstein30]

You shouldn't practice bad chemistry, it will lead to mistakes. Stick with what the question states, and don't assume anything.

I fail to see where its bad chemistry. It's skipping a step, sure.. but it's the same thing.. is it not?

Again, does the same thing not occur if:
1) You add HCl and have Ag+ present
2) You added equal amounts of H+ and Cl- with Ag+ present

You said we shouldn't assume HCl = H+ and Cl- for this very reason when that makes absolutely zero sense..

It's not an assumption that HCl will dissociate almost entirely into H+ and Cl-.. so thinking of a strong acid (or base) in terms of what it dissociates into is actually a better way to approach such problems. You'll be more likely to pick up that the free chlorine anion reacts with Ag+.

 

[Czarcasm]

You shouldn't practice bad chemistry, it will lead to mistakes. Stick with what the question states, and don't assume anything.

I dunno ... very few people have the capability of mustering up a 38+ on their MCAT. S/he's one of them. 😉

 

[MDforMee]

I fail to see where its bad chemistry. It's skipping a step, sure.. but it's the same thing.. is it not?

Again, does the same thing not occur if:
1) You add HCl and have Ag+ present
2) You added equal amounts of H+ and Cl- with Ag+ present

You said we shouldn't assume HCl = H+ and Cl- for this very reason when that makes absolutely zero sense..

It's not an assumption that HCl will dissociate almost entirely into H+ and Cl-.. so thinking of a strong acid (or base) in terms of what it dissociates into is actually a better way to approach such problems. You'll be more likely to pick up that the free chlorine anion reacts with Ag+.

You're missing the point. It's to avoid having side reactions in the case that test writers write different types of questions...

maybe they'll ask if the Ag+ precipitates without telling you the identity of the conjugate base?

I can think of lots of ways that test writers can make solubility rules tricky. Someone without much chemistry experience may assume that strong acid means that a chloride conjugate must always be present for strong acids... there are other strong acids, you know. Like sulfuric acid.

Point being, don't assume anything.

 

[Jepstein30]

This is turning into a someone has to be right argument, and you're one of those people that has to be right, am I right?

Uh not at all. In fact, it's pretty much the opposite.

I'm saying there's no difference. You're just saying there is without actually explaining why, jumping to ridiculous claims like "bad chemistry" and incorrect reasoning involving complex equilibrium. I'm not even sure when the last post you made that was actually on topic vs. just trying to be right.. you usually have to be try to show why you're right, which you haven't.

This forum is about helping others on the MCAT so I'd rather not have blatantly wrong advice on these forums. There is absolutely zero difference between thinking of a strong acid as a strong acid or its conjugate and a proton. If there is, I'd personally like to know and like to the actual rasoning to be out there rather than devolving to stuff like you've been posting since I pointed out your reasoning is wrong.

Again, there is zero difference. And if your argument is now, "well it doesn't matter either way so I guess someone has to be right".. well.. isn't that zero difference then?

 

[MDforMee]

Uh not at all. In fact, it's pretty much the opposite.

I'm saying there's no difference. You're just saying there is without actually explaining why, jumping to ridiculous claims like "bad chemistry" and incorrect reasoning involving complex equilibrium. I'm not even sure when the last post you made that was actually on topic vs. just trying to be right.. you usually have to be try to show why you're right, which you haven't.

This forum is about helping others on the MCAT so I'd rather not have blatantly wrong advice on these forums. There is absolutely zero difference between thinking of a strong acid as a strong acid or its conjugate and a proton. If there is, I'd personally like to know and like to the actual rasoning to be out there rather than devolving to stuff like you've been posting since I pointed out your reasoning is wrong.

Again, there is zero difference. And if your argument is now, "well it doesn't matter either way so I guess someone has to be right".. well.. isn't that zero difference then?

Like I said, you can't let it go and move on. There are reasonable explanations (and reasons) for both of our sides, so if you can't drop it and see the bigger picture, I'm going to chalk this up to pointlessness.

 

[Jepstein30]

Says the guy who keeps posting as well haha

I would really like to hear why it's different. You very well could be correct but nothing you've written so far explains why. So if you are correct, I'd actually like to know why so I can learn something useful today.

Or you can keep posting about things completely unrelated... up to you.

 

[MDforMee]

Says the guy who keeps posting as well haha

I would really like to hear why it's different. You very well could be correct but nothing you've written so far explains why. So if you are correct, I'd actually like to know why so I can learn something useful today.

Or you can keep posting about things completely unrelated... up to you.

This is the jeppy baby I like to hear talking to me. Give me a kiss, snookie pookie.

Like I said, test writers may try and confuse people with solubility rules, so its best not to assume that specific ions are present unless otherwise stated. I don't see what's unclear about this statement.

 

[Jepstein30]

That's completely irrelevant and different.

Obviously, if a question says H+ is added, only consider H+. If a question says H2SO4 is added though, you can consider it as H+ and HSO4-.

No-one said you should simply add on a random conjugate into the question this thread is about, that would be foolish. What was said was "Generally, when you're adding strong acid, you're adding both H+ and its conjugate base". Is that not 100% true?

 

[Jepstein30]

Nice to see you're going back and editing your posts though haha.. that's good stuff. Helps when you want to completely change what you're talking about.

If you are saying we shouldn't just throw in a Cl-, then yep.. I agree, that's obvious. We shouldn't also just decide to throw in a +2 in a random math problem either.

But when you add a strong acid, you are adding a proton and ITS conjugate base. The conjugate base of that specific strong acid. Not necessarily Cl-.

 

[MDforMee]

Nice to see you're going back and editing your posts though haha.. that's good stuff. Helps when you want to completely change what you're talking about.

If you are saying we shouldn't just throw in a Cl-, then yep.. I agree, that's obvious. We shouldn't also just decide to throw in a +2 in a random math problem either.

But when you add a strong acid, you are adding a proton and ITS conjugate base. The conjugate base of that specific strong acid. Not necessarily Cl-.

Stop fishing for new responses from me in this thread then, pookums.

If you don't, I'm going to make you reply to something that's bogus and then counter it with how bad your response is.

 

[Jepstein30]

Well, I guess you've moved on from actually talking about chemistry given you're just using ad hominem now. Not to mention I should've kept quoting your posts as they keep mysteriously changing haha.. no more mention of H2SO4 in that last one? Weird how that keeps happening..

Enjoy convoluting one of the simplest things in general chemistry though.. cause remember HCl isn't H+ and Cl-!

 

[MDforMee]

Well, I guess you've moved on from actually talking about chemistry given you're just using ad hominem now. Not to mention I should've kept quoting your posts as they keep mysteriously changing haha.. no more mention of H2SO4 in that last one? Weird how that keeps happening..

Enjoy convoluting one of the simplest things in general chemistry though.. cause remember HCl isn't H+ and Cl-!

email me anytime at [email protected]

 

[Jepstein30]

email me anytime at [email protected]

True immaturity always shines through.