# Destroyer GC 85 and 86

#### Shinpe

So in #85, it's asking, "which is the weakest base?"
One option is HClO4 and another ClO4-. and then it says the answer is ClO4- because HClO4 is an acid. Isn't base strength how well it will accomadate an extra H+?? I know HClO4 is not considered a base and is an acid, but doesn't that itself make it the WEAKEST base? because it will never want to take an extra H and become H2ClO4+???

In 86, I think this was on here a while ago, what is the normality of the a 2.5 solution of chromateCr2O7(2-) -> Cr3+.

I think in that thread some one said that the normality should be considered depending on the number of electrons in the balanced equation, but destroyer says the balancing was not necessary and u just need to look at oxidation states, Anyone have any comments/resources to defend either of these??

#### joonkimdds

##### Senior Member
10+ Year Member
7+ Year Member
for the 2nd question, whether u balance it or not, you should still get the same answer.

For the 1st one, I will answer after i came back .

#### Shinpe

for the 2nd question, whether u balance it or not, you should still get the same answer.

For the 1st one, I will answer after i came back .
If you balance the equation (with 7 H2Os on right and thus 14 H+ on the left), you see that you need 6e- on the left, so if you use that (which was suggested on this forum before), you'd get 6*2.5=15N, while if you just use the fact that you're going from +6 oxidation state to +3, you would just use 3*2.5=7.5N (which is what destroyer says is right).

#### sixkiller

If you balance the equation (with 7 H2Os on right and thus 14 H+ on the left), you see that you need 6e- on the left, so if you use that (which was suggested on this forum before), you'd get 6*2.5=15N, while if you just use the fact that you're going from +6 oxidation state to +3, you would just use 3*2.5=7.5N (which is what destroyer says is right).
Man, destroyer is wrong wrong wrong about that normality question.

Ask a chemistry professor at your school. The normality of a dichromate solution in acid (assuming the reaction goes to completion) is a textbook problem. If you're lucky you might be able to find a sample normality problem in an older gchem book that uses dichromate.

Dichromate is a single compound that reacts with 6 equivalents of electrons per mole....

Try using the destroyer logic when solving redox titration problems involving dichromate...it will fail everytime.

About the acid question. You are right, in theory perchloric acid can accept a proton, but such questions seem to assume that it can't. I've seen similar problems in other sources and that seems to be the predicated logic. Nothing you can do but just roll with it...you have to choose the "best" answer.

#### joonkimdds

##### Senior Member
10+ Year Member
7+ Year Member
I am a destroyer fanboy.

#### Shinpe

Man, destroyer is wrong wrong wrong about that normality question.

Ask a chemistry professor at your school. The normality of a dichromate solution in acid (assuming the reaction goes to completion) is a textbook problem. If you're lucky you might be able to find a sample normality problem in an older gchem book that uses dichromate.

Dichromate is a single compound that reacts with 6 equivalents of electrons per mole....

Try using the destroyer logic when solving redox titration problems involving dichromate...it will fail everytime.

About the acid question. You are right, in theory perchloric acid can accept a proton, but such questions seem to assume that it can't. I've seen similar problems in other sources and that seems to be the predicated logic. Nothing you can do but just roll with it...you have to choose the "best" answer.
Hmm would the fact that this problem doesn't say anything about dichromate "in acid" or titration, make any difference? 'cause normality depends on what reaction you're talking about right? (It doesnt really specificy a reaction here, pretty annoying). I'm sure you know what u'r talkin about though. I dont get this normality thing too well myself, have to look into a book.

And for second question, so we basically have to decide which ones CAN actually accept H+, and then among those, which one is worst at doing so?

Thanks

#### Orgolover

Can I just say " I freakin' hate that orgoman normality question" --- that was for the record. In general i hate the term "normality" My chemistry textbook used in university doesn't even use the term. It's utter crap. Another damn definition to make us remember. Since when has chemistry been subjected to such political bureaucracy.
Why couldn't they just let us figure out molarity and leave it at that . . !!!

#### sixkiller

Hmm would the fact that this problem doesn't say anything about dichromate "in acid" or titration, make any difference? 'cause normality depends on what reaction you're talking about right? (It doesnt really specificy a reaction here, pretty annoying). I'm sure you know what u'r talkin about though. I dont get this normality thing too well myself, have to look into a book.

And for second question, so we basically have to decide which ones CAN actually accept H+, and then among those, which one is worst at doing so?

Thanks
Yes, the reaction does matter, but for starters...dichromate (orange)is the dominant form in strong acid. Also, if I recall correctly, the solution balances out the half reaction in acidic solution.

If the solution were basic, chromate would be the dominant form (yellow). And in the case of chromate....we would multiply molarity by a factor of three because ONE mole of chromate reacts with 3 equivlants of electrons.

The best way to think of redox normality is the number of electrons that ONE mole of a compound either accepts or donates....

No need to go into it further for the DAT. If you want to however, try a quantitative analysis book or an older gchem book (preferably pre 90's).

#### Shinpe

Yes, the reaction does matter, but for starters...dichromate (orange)is the dominant form in strong acid. Also, if I recall correctly, the solution balances out the half reaction in acidic solution.

If the solution were basic, chromate would be the dominant form (yellow). And in the case of chromate....we would multiply molarity by a factor of three because ONE mole of chromate reacts with 3 equivlants of electrons.

The best way to think of redox normality is the number of electrons that ONE mole of a compound either accepts or donates....

No need to go into it further for the DAT. If you want to however, try a quantitative analysis book or an older gchem book (preferably pre 90's).
Yeah after looking around a bit too I agree with you.
The number of electrons makes a lot more sense to me than the change in oxidation number. I just hope the normality question I get is on simple acid/base stuff, not electrochemical cells and ****.

So here comes a question then, you have MnO4- + I- -> MnO2 + I2(don't feel like balancing it right now), but I remember the MnO4 half reaction would use 5 e- and I- produce 2e-, so if you had 1 M MnO4-, would it be 5N or 10N??

#### Pat Kelly

##### New Member
10+ Year Member
Destroyer is wrong - think of dichromate as a peroxide and you will see that you can't just compare the oxidative states of a single chromium ion.

#### sixkiller

Yeah after looking around a bit too I agree with you.
The number of electrons makes a lot more sense to me than the change in oxidation number. I just hope the normality question I get is on simple acid/base stuff, not electrochemical cells and ****.

So here comes a question then, you have MnO4- + I- -> MnO2 + I2(don't feel like balancing it right now), but I remember the MnO4 half reaction would use 5 e- and I- produce 2e-, so if you had 1 M MnO4-, would it be 5N or 10N??
Well, that's when we need to be careful. Multiply by five IF the final oxidation state of manganese is +2.

In the above problem, the final oxidation state of manganese is +4...therefore, we must multiply by a factor of 3.

*in strong acid, permanganate is reduced to Mn++, but in mildly acidic solution, it is reduced to Mn++++*

In most cases, you do not really have to balance the equation because the compound only has ONE atom that is being oxidized; then, by default, just examining the initial and final oxidation state will allow you to determine how many equivalents of electrons the compound accepts or releases.

The case of dichromate is a bit trickier because TWO chromium atoms are present in per dichromate ion. Just a little extra book keeping is involved.

So, it isn't absolutely necessary to write a balanced equation, but it does help to prevent sloppy errors.

Practice and eventually you don't have to always write balanced half reactions for redox titration problems.

#### Shinpe

Destroyer is wrong - think of dichromate as a peroxide and you will see that you can't just compare the oxidative states of a single chromium ion.

#### Shinpe

Well, that's when we need to be careful. Multiply by five IF the final oxidation state of manganese is +2.

In the above problem, the final oxidation state of manganese is +4...therefore, we must multiply by a factor of 3.

*in strong acid, permanganate is reduced to Mn++, but in mildly acidic solution, it is reduced to Mn++++*

In most cases, you do not really have to balance the equation because the compound only has ONE atom that is being oxidized; then, by default, just examining the initial and final oxidation state will allow you to determine how many equivalents of electrons the compound accepts or releases.

The case of dichromate is a bit trickier because TWO chromium atoms are present in per dichromate ion. Just a little extra book keeping is involved.

So, it isn't absolutely necessary to write a balanced equation, but it does help to prevent sloppy errors.

Practice and eventually you don't have to always write balanced half reactions for redox titration problems.
So it doesnt have anything to do with how many electrons are in the balanced equation, it's how many moles one mole of the compound accepts or gives off. In that case Mn is going from +7 to +4, and there's only 1 mole of Mn, so Normality=3*molarity??

#### sixkiller

So it doesnt have anything to do with how many electrons are in the balanced equation, it's how many moles one mole of the compound accepts or gives off. In that case Mn is going from +7 to +4, and there's only 1 mole of Mn, so Normality=3*molarity??
I'm not sure if I completely understand, but it's more or less the same thing.

In the final balanced equation for the titration it could be different, but it's the same if we're talking about the balanced HALF reaction.

This is because the half reaction is usually balanced with respect to ONE mole of the oxidizing or reducing agent.

#### Shinpe

I'm not sure if I completely understand, but it's more or less the same thing.

In the final balanced equation for the titration it could be different, but it's the same if we're talking about the balanced HALF reaction.

This is because the half reaction is usually balanced with respect to ONE mole of the oxidizing or reducing agent.
Oh oops, I had MnO4- -> Mn2+ in my head but was talking saying MnO2. So ok, the number of electrons in the half reaction then (5 if going to Mn2+, 3 if you're going to MnO2). and it does not matter that you have to multiply the half reaction by two to be able to cancel the electrons with the other half reaction?