I'm a tad confused by the question, The protonation constant for the acid HA has the value K=2*10^6.

(a) what is the principal species at pH 6.00?

(b) what is the principal species at pH 9.00?

(c) what is the quotient [A-]/[HA] at pH 6.3 and 7.5?

In solution:

HA <----> H^+ +A^-

Ka=[H^+][A^-]/[HA]

I've never heard the question asked this way either, but I think for A and B, the author of the question wants to know if A^-/H^+ or HA are the dominant species. The lower the pH, the more A^- you have; the higher the pH, the less A^- you have. This should be enough so that you can answer the first two by yourself.

For C:

Use the Henderson-Hasselbach equation

pH=pka-+log[A^-/HA]

pka=-log[Ka]

Solve for [A^-/HA] ratio:

10^(pH-pka)=[A^-/HA]

I must confess that I've never heard the term "protonation constant" used. Logic tells me this might be the reciprocal of Ka, the dissociation constant but I don't know that. Can you shed some light? If we can define what the protonation constant is I can help.

2) At 298K, what are the molarities of pure water?

Look up the density of water at 25C and use that to calculate the mass of 1000 mL. Ifound density at 25 C = 0.99705 g/mL; therefore 1000 mL has a mass of 997.05g.Then mols H2O = grams/molar mass = 997.05/18.015 = about 55.35 M

To answer these questions, we need to understand some concepts related to the protonation constant and the pH of a solution.

First, let's define the protonation constant of an acid. The protonation constant, denoted as K, is a measure of how easily an acid donates a proton (H+) in a chemical reaction. It is also known as the acid dissociation constant (Ka). A higher value of K indicates a stronger acid, meaning it easily donates protons.

The pH of a solution measures its acidity or basicity. It is a logarithmic scale that ranges from 0 to 14, where pH below 7 indicates acidity, pH above 7 indicates basicity, and pH equal to 7 indicates neutrality.

Now, let's address each question step by step:

(a) What is the principal species at pH 6.00?
To determine the principal species at pH 6.00, we need to compare the pH with the pKa value, which is the negative logarithm of the protonation constant (pKa = -log(K)). In this case, pKa = -log(2*10^6).

If the pH is less than the pKa, the acid is predominantly in its protonated form (HA), while if the pH is greater than the pKa, the acid is predominantly in its deprotonated form (A-).

Now, let's calculate pKa using the given protonation constant K:
pKa = -log(2*10^6)
Using logarithmic identity: log(a*b) = log(a) + log(b)
pKa = -(log(2) + log(10^6))

Since log(10^6) = 6, we can simplify further:
pKa = -(log(2) + 6)
pKa ≈ -9.30

Given that the pH is 6.00, which is greater than the pKa (-9.30), the principal species at pH 6.00 will be A-, indicating that the acid has deprotonated.

(b) What is the principal species at pH 9.00?
Following the same logic as before, we compare the pH with the pKa value. In this case, the pH is 9.00, which is greater than the pKa (-9.30).

Given that the pH is greater than the pKa, the principal species at pH 9.00 will also be A-, indicating that the acid has deprotonated.

(c) What is the quotient [A-]/[HA] at pH 6.3 and 7.5?
To calculate the quotient [A-]/[HA], we need to use the Henderson-Hasselbalch equation, which relates the pH and the ratio of the deprotonated form (A-) to the protonated form (HA) of an acid.

The Henderson-Hasselbalch equation is given by:
pH = pKa + log([A-]/[HA])

We have the pKa value from earlier, which is -9.30.

For pH 6.3:
6.3 = -9.30 + log([A-]/[HA])
log([A-]/[HA]) = 6.3 + 9.30
log([A-]/[HA]) ≈ 15.6
[A-]/[HA] ≈ 10^15.6

Similarly, for pH 7.5:
7.5 = -9.30 + log([A-]/[HA])
log([A-]/[HA]) = 7.5 + 9.30
log([A-]/[HA]) ≈ 16.8
[A-]/[HA] ≈ 10^16.8

So, the quotient [A-]/[HA] at pH 6.3 is approximately 10^15.6, and at pH 7.5 is approximately 10^16.8.