Weak Acid

1.) The pH of a 0.060M weak monoprotic acid HA is 3.44. Calculate the Ka of the acid.
2.) The pH of 0.100M solution of weak monoprotic acid HA is 2.85. What is the Ka of the acid?

I have shown you how to do these below.

If you want to post your answer I can check it.

To calculate the Ka of a weak acid, we can use the formula:

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

where [HA] is the concentration of the weak acid, [A-] is the concentration of the conjugate base, and pH is the measured pH of the solution.

Let's solve the first question step-by-step:

1.) The pH of a 0.060M weak monoprotic acid HA is 3.44. Calculate the Ka of the acid.

Step 1: Identify the unknowns:
We need to find the Ka value of the acid.

Given:
pH = 3.44
[HA] = 0.060M

Step 2: Calculate the concentration of the conjugate base [A-]:
Since the weak acid (HA) is monoprotic, it dissociates into an equal concentration of its conjugate base (A-).
[A-] = [HA] = 0.060M

Step 3: Calculate the Ka value:
Ka = 10^(-pH) * [HA] / [A-]
= 10^(-3.44) * 0.060M / 0.060M
= 3.98 * 10^(-4)

Therefore, the Ka value of the acid is approximately 3.98 * 10^(-4).

Let's move on to the second question:

2.) The pH of a 0.100M solution of weak monoprotic acid HA is 2.85. What is the Ka of the acid?

Step 1: Identify the unknowns:
We need to find the Ka value of the acid.

Given:
pH = 2.85
[HA] = 0.100M

Step 2: Calculate the concentration of the conjugate base [A-]:
Since the weak acid (HA) is monoprotic, it dissociates into an equal concentration of its conjugate base (A-).
[A-] = [HA] = 0.100M

Step 3: Calculate the Ka value:
Ka = 10^(-pH) * [HA] / [A-]
= 10^(-2.85) * 0.100M / 0.100M
= 1.26 * 10^(-3)

Therefore, the Ka value of the acid is approximately 1.26 * 10^(-3).

To calculate the Ka (acid dissociation constant) of a weak acid given the pH of its solution, we need to use the concept of the Henderson-Hasselbalch equation.

The Henderson-Hasselbalch equation is expressed as follows:

pH = pKa + log [A-]/[HA],

where pH is the measure of acidity, pKa is the negative logarithm of the Ka, and [A-] and [HA] represent the concentrations of the acid's conjugate base and the acid, respectively.

Now let's solve the first question step by step:

1.) The pH of a 0.060M weak monoprotic acid HA is 3.44. Calculate the Ka of the acid.

Step 1: Identify the given values:
- pH = 3.44
- [HA] = 0.060 M

Step 2: Calculate the concentration of the conjugate base A- using the equation:
[A-] = 10^(-pH) × [HA]

[A-] = 10^(-3.44) × 0.060
= 0.000378 M

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

3.44 = pKa + log (0.000378/0.060)

Rewriting the equation:
pKa = 3.44 - log (0.000378/0.060)
pKa = 3.44 - log (0.0063)

Step 4: Calculate Ka from pKa using the equation:
Ka = 10^(-pKa)

Ka = 10^(-(3.44 - log (0.0063)))

With these steps, you can determine the Ka of the acid based on its pH.

Now let's move on to the second question:

2.) The pH of a 0.100M solution of weak monoprotic acid HA is 2.85. What is the Ka of the acid?

Using the same steps:
Step 1: Identify the given values:
- pH = 2.85
- [HA] = 0.100 M

Step 2: Calculate the concentration of the conjugate base A- using the equation:
[A-] = 10^(-pH) × [HA]

[A-] = 10^(-2.85) × 0.100
= 0.00168 M

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

2.85 = pKa + log (0.00168/0.100)

Rewriting the equation:
pKa = 2.85 - log (0.0168)

Step 4: Calculate Ka from pKa using the equation:
Ka = 10^(-pKa)

Ka = 10^(-(2.85 - log (0.0168)))

By following these steps, you can determine the Ka of the weak acid based on its pH.