If a buffer solution is 0.190 M in a weak acid (Ka = 5.8 x 10^-5) and 0.550 M in its conjugate base, what is the pH?

Question

If a buffer solution is 0.190 M in a weak acid (Ka = 5.8 x 10^-5) and 0.550 M in its conjugate base, what is the pH?

What if it said "if a buffer solution is ___ M in a weak base (Kb = ____) and ___ M in its conjugate acid, what is the pH? Would the steps from the first problem still be identical to this problem?

Answer 1

The steps would be the same. Use the Henderson-Hasselbalch equation for both.

Answer 2

Thank you! we havn't covered this in lecture yet so I wasn't sure what the formula was called

Answer 3

To find the pH of a buffer solution, we can use the Henderson-Hasselbalch equation, which is given by:

pH = pKa + log ([A-] / [HA])

In the given problem, we have a buffer solution consisting of a weak acid and its conjugate base. The acid dissociation constant (Ka) for the weak acid is given as 5.8 x 10^-5. The concentrations of the weak acid and its conjugate base are 0.190 M and 0.550 M, respectively.

Using the Henderson-Hasselbalch equation, we can calculate the pH as follows:

1. Calculate the pKa for the weak acid:
pKa = -log(Ka) = -log(5.8 x 10^-5)

2. Calculate the ratio of the concentration of the conjugate base ([A-]) to the concentration of the weak acid ([HA]):
[A-] / [HA] = 0.550 / 0.190

3. Substitute the values into the Henderson-Hasselbalch equation and solve for pH:
pH = pKa + log ([A-] / [HA])

Now, let's address your second question. If the problem provides a weak base and its conjugate acid, instead of a weak acid and its conjugate base, the steps to find the pH will be slightly different. The Henderson-Hasselbalch equation still applies, but the pKa and concentration ratios will be based on the Kb and concentrations of the weak base and its conjugate acid rather than the weak acid and its conjugate base.

Answer 4

To find the pH of a buffer solution, you need to use the Henderson-Hasselbalch equation, which is given by:

pH = pKa + log([A-]/[HA])

In the case of the given buffer solution, the weak acid is the acid (HA) and its conjugate base is A-. The pKa is the negative logarithm of the acid dissociation constant (Ka).

Let's calculate the pH using the given values:

Ka = 5.8 x 10^-5

pKa = -log(Ka) = -log(5.8 x 10^-5) = 4.24

[A-] = 0.550 M

[HA] = 0.190 M

pH = 4.24 + log(0.550/0.190)

= 4.24 + log(2.895)

= 4.24 + 0.461

= 4.70

So, the pH of the buffer solution is 4.70.

Now, let's consider the second problem where a buffer solution has a weak base and its conjugate acid. The steps to calculate the pH would be slightly different, but the concept remains the same.

In this case, we would use the Henderson-Hasselbalch equation with the Kb and pKb values:

pOH = pKb + log([HA-]/[A])

To find the pH, we can use the fact that pH + pOH = 14:

pH = 14 - pOH

The rest of the steps would be similar to the first problem. Therefore, the steps from the first problem would not be identical, but the overall process of using the Henderson-Hasselbalch equation to calculate the pH of a buffer solution would still be applicable.