[alexpos]

1.) What is the pH of a buffer solution that is 0.10 NH3 and 0.10 NH4+? What is the pH if 12mL of 0.20 M hydrochloric acid is added to 125 mL of buffer?

For the first part, i figured that NH3 is the base and NH4+ is the conjugate acid. Therefore, the equil. reaction is NH3 + water --> NH4+ OH-. Using the given concentrations, i found Kb, pOH, and then Ph.

For the second part, i'm unsure. I figured out the moles of NH3 and NH4+ by multiplying each molarity by 0.125 L. Then i did the same for H3O+ ion (formed by HCl) using 12mL as the volume. However, the problem appears as I'm unsure whether the addition of HCl will depleat the amount of NH4+ or NH3. Should i subtract the moles of the hydronium ion from the moles of NH3 and add it to NH4+? Or the other way around?


2.) A chloride salt (rantidinium chloride) is presnet in Zantac. Should a solution of rantidinium chloride be acidic, basic, or neutral?

This is my logic: Since chloride ion forms an acidic solution with Hyrdogen (that is, its a conjugate base of a strong acid), then the solution should be acidic. But i need to consider rantidinium as well, and see whether that is an anion of a strong base, in which case the solution would be neutral. But how do i do that?

3.) I have to determine whether hydrolysis would occur with the following:

a.) NO3-
b.) OCl-
c.) NH2NH3+
d.) Br-

    a.) Since NO3- is a conjugate base of a strong acid, no hydrolysis would occur.
   b.) Since OCl- is a conjugate base of a weak acid (HClO), hydrolysis does happen.
   c.) I have no idea!
   d.) Since HBr is a strong acid, no hydrolysis happens.

My questions are: i.) Why does hydrolysis not happen when an ion is a conjugate acid/base of a strong acid/base?  and ii.) Part C.


Thank you.

[sdekivit]

the first thing you need to do with buffer solutions is to write down your equilibrium. In the first question, the equilibrium is:

NH4(+) + H2O <--> NH3 + H3O(+)

Then you know the ratio of NH3 and NH4(+) in the solution. In this case, this is 1 --> Ka = [H3O(+)] or pKa = pH

when we add 12 mL 0,20 M HCl, we add 12 * 0,2 = 2,4 mmol HCl, thus 2,4 mmol H3O(+). This will react with the NH3. You know the volume of your buffer and the concentrations of NH3 and NH4(+) (0,1 * 125 = 12,5 mmol), thus finally we have 12,5 - 2,4 = 10,1 mmol NH3 and 12,5 + 2,4 = 14,9 mmol NH4(+)

thus then we get: pKa = pH + log 10,1/12,4

[Borek]

Should i subtract the moles of the hydronium ion from the moles of NH3 and add it to NH4+?

Yes.

2.) A chloride salt (rantidinium chloride) is presnet in Zantac. Should a solution of rantidinium chloride be acidic, basic, or neutral?

Whatever rantidinium is, it is in protonated form - thus it will behave like an acid. Cl^- will behave as a base, but so weak, that you may neglect its hydrolysis.

I have to determine whether hydrolysis would occur with the following:

a, b and d - these are probably the answers your teacher expects.

NH₂NH₃⁺

Pretty tough question. For N₂H₄ pK_b1 = 6, pK_b2 = 14.9. Thus protonated form is a very weak base - it will hydrolize, but very, very slightly. At the same time protonated form will have a tendency to dissociate and to rise pH. Overall effect is that 0.05M solution will have a pH pf 8.00 (calculated with BATE). See comment below.

Why does hydrolysis not happen when an ion is a conjugate acid/base of a strong acid/base?

If not for the question c I will not touch the subject, as it is generally teached that strong acids and bases salts do not hydrolize just because they are strong acids and bases.

But that's not true.

Every acid/base hydrolize. The question is not whether but to what extent. The stronger acid/base the weaker the hydrolyzis, but it is always there. Just it is often neglectable.

Nitric acid is the weakest between strong, with pK_a=-1. 1M solution should have pH=0.00 (ionic strength ignored), but if you do exact calculations, you will find that pH=0.04. Small difference, neglectable, but still this 0.04 unit difference is an effect of hydrolysis, nothing else.