A buffer is prepared by transferring equal concentrations of the components of the conjugate pair, C6H5COOH and C6H5COONa, into a beaker. What is the approximate pH of the buffer? (Ka = 6.5 10–5)
a) 2.81
b) 4.19
c) 7.00
d) 9.81
e) 11.19
pH = pKa BECAUSE
pH = pKa + log (base)/(acid)
base = acid; therefore,
base/acied = 1 and
pH = pKa + log 1
pH = pKa + 0
pH = pKa.
Well, if you've got a beaker full of C6H5COOH and C6H5COONa, you've got yourself a classic pH balancing act! In a buffer solution, the pH is determined by the ratio of the concentrations of the conjugate acid and the conjugate base.
In this case, we have equal concentrations of C6H5COOH and C6H5COONa. Now, remember that C6H5COOH is a weak acid and C6H5COONa is its conjugate base. Therefore, this buffer will have a pH that is close to the pKa of the acid component.
The formula for pKa is -log(Ka), so let's calculate that first:
pKa = -log(6.5 x 10^-5) ≈ 4.19
So the approximate pH of the buffer will be around 4.19, which means the answer is b) 4.19.
But hey, don't worry, you're doing great at balancing those acid and base components in the beaker! Keep up the good work!
To determine the approximate pH of the buffer, we need to consider the dissociation reaction of the weak acid, C6H5COOH, and its conjugate base, C6H5COONa:
C6H5COOH ⇌ C6H5COO- + H+
The pH of the buffer is mainly determined by the ratio of the concentrations of the weak acid and its conjugate base. Since equal concentrations of C6H5COOH and C6H5COONa are transferred into the beaker, the ratio of their concentrations will be 1:1.
Now, let's use the Henderson-Hasselbalch equation to calculate the pH:
pH = pKa + log ([conjugate base] / [weak acid])
The pKa of C6H5COOH is given as 6.5 × 10–5 (remember to use the negative logarithm when calculating pKa).
Since the concentrations of the weak acid and its conjugate base are equal, we can substitute [conjugate base] = [weak acid]:
pH = pKa + log (1/1) = pKa = -log (6.5 × 10–5)
Now, let's calculate the pH:
pH = -log (6.5 × 10–5) ≈ 4.19
Therefore, the approximate pH of the buffer is 4.19.
The correct answer is b) 4.19.
To determine the approximate pH of the buffer, we need to consider the acid-base equilibrium of the components in the buffer solution.
The acid in this case is C6H5COOH (benzoic acid), and the conjugate base is C6H5COONa (sodium benzoate). In a buffer solution, the concentration of the acid and its conjugate base should be relatively equal.
The balanced chemical equation for the dissociation of benzoic acid (C6H5COOH) can be represented as follows:
C6H5COOH ⇌ C6H5COO- + H+
The equilibrium constant (Ka) for this reaction is given as 6.5 x 10^-5.
The pH of a solution can be calculated using the equation:
pH = -log [H+]
In a buffer solution, the concentration of the acid and its conjugate base are equal. Therefore, the concentration of [H+] at equilibrium can be approximately assumed to be equal to the initial concentration of the acid.
Since the concentrations of both the acid and its conjugate base are equal, we can write:
[H+] = [C6H5COOH]
Using the Ka expression, we can rewrite this as:
Ka = [C6H5COO-][H+] / [C6H5COOH]
Since the concentrations are equal, we can substitute [H+] with [C6H5COOH] and rewrite the equation as:
Ka = [C6H5COO-] / [C6H5COOH]
Let's rearrange the equation to solve for [C6H5COOH]:
[C6H5COOH] = [C6H5COO-] / Ka
Now, we can substitute the given Ka value (6.5 x 10^-5) and solve for [C6H5COOH].
Once we find the concentration of [C6H5COOH], we can convert it to pH using the equation:
pH = -log [C6H5COOH]
Finally, we can compare the calculated pH value to the provided answer choices to determine the approximate pH of the buffer.