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Inorganic chemistry - solution chemistry

Hi guys, got a couple of questions here,

1) Why does H3PO4 and H3PO3 have similar first acid dissociation constant values? (in other words, Pauling's rules doesn't seem to work out well)

2) PH3 is a weak base in water but a strong base in HCl. In water, ethanoic acid is a weak acid, but in liquid ammonia, it is a strong acid. Explain why.

3) Is the failure of ClO3- to disproportionate in acidic aq solution to ClO4- and Cl- a thermodynamic or kinetic phenomenon? Give reasons.
- I said it is thermodynamic, from calculation, delta G for this disproportionation process is greater than 0, hence not spontaneous.

4) pKa of the haloacids increases down group 17. Explain.
I know how to explain about the overlap of AOs gets worse down the group due to bigger AOs, hence weaker bond, so deprotonation ability increases, hence stronger acid.

- Are there any other reasons that would explain the trends?

5) N(SiH3)3 and NF3 are much weaker lewis bases than NMe3.
- Weaker lewis bases means the lone pair on nitrogen is less available as electron pair donor. I am thinking of fluorine being electron withdrawing inductively, whereas d-p pi interaction means the lone pairs get delocalized into the d-p pi system, so less available for protonation. Are these reasons correct?

6) NH3 is a stronger Bronstead base thatn N(CH3)3 in water but is a weaker Bronstead base in the gas phase. Explain.

Thanks. I really need help in making sure I understand the right contents.

Reply 1

shengoc

Anyone willing to share their ideas, please?

Reply 2

MrCynical

OK, but I'm a bit rusty on this stuff, so I'm just thinking as I type here.

2)For ethanoic acid, this is just looking at the formula for pKa. In liquid ammonia Ka=[CH3COO-][(NH4)+]/[CH3COOH][NH3]. In water it'll be =[CH3COO-][(H3O)+]/[CH3COOH][H2O]. Which is more stable: NH4+ or H3O+? NH4+, so [NH4+] is larger than [H3O+], Ka is larger in liquid ammonia than in water, pKa is lower, and ethanoic acid is more acidic in ammonia. Flip the argument round for PH3. Kb = [PH3][H+]/[PH4+]. In HCl [H+] is going to be higher than in water. Kb therefore larger, pKb lower, so stronger base. (Could look at conjugate acids if you don't like Kb).

3)I don't have the numbers to hand. If delta G is positive, then yes that would be thermodynamic.

4)I think that's the standard explanation

5)Electronegativity for this one. F is highly electronegative, pulling electron density away from the lone pair and making it less available. Methyls have a slight net inductive effect as C has a higher electronegativity than H, increasing electron density at the lone pair. Si has a lower electronegativity than H, so effect reversed.

6)I'm not sure on this one - my first thought is lone pairs on water molecules can stabilise the NH4+ in solution, but aren't going to be available in the gas phase. Steric effect of methyls keeping water away as a reason why this doesn't happen for N(CH3)3?