A buffer contains 5.00 M acetic acid and 5.00 M acetate anion. Gaseous HCl (0.010 mole) is added to 1.00 L of this buffer solution (the total volume does not change). For this buffer solution, the initial pH is_______and the final pH after the addition of the HCl gas is_______ . For acetic acid, Ka = 1.8 x 10-5.
Help me fill in the blanks please? please explain.
Use the Henderson-Hasselbalch equation for part 1. For th second part, HCl will cause an increase in acetic acid concn and a decrease in acetate concn. Post your work if you get stuck.
To find the initial pH of the buffer solution, we can use the Henderson-Hasselbalch equation, which is given by:
pH = pKa + log([A-]/[HA])
Where:
pH = the pH of the solution
pKa = the negative logarithm of the acid dissociation constant (Ka)
[A-] = concentration of the conjugate base (acetate anion)
[HA] = concentration of the weak acid (acetic acid)
Given information:
[A-] = 5.00 M (concentration of acetate anion)
[HA] = 5.00 M (concentration of acetic acid)
Ka = 1.8 x 10^-5
Now, let's calculate the initial pH using the Henderson-Hasselbalch equation:
pH = pKa + log([A-]/[HA])
= -log(1.8 x 10^-5) + log(5.00/5.00)
= -(-4.74) + log(1.00)
= 4.74
Therefore, the initial pH of the buffer solution is 4.74.
To find the final pH after the addition of the HCl gas, we need to consider the reaction between acetic acid (HA) and HCl.
The balanced chemical equation is as follows:
HA + HCl → H2O + ACl
Since solid NaCl is formed as a result of neutralization, it does not participate in the pH calculations.
In this case, HCl reacts with acetic acid, causing the concentration of acetic acid ([HA]) to decrease, while the concentration of acetate anion ([A-]) increases.
Our goal is now to find the new concentrations of acetic acid and acetate anion after the reaction has reached equilibrium.
Given:
Initial concentration of acetic acid [HA] = 5.00 M
Initial concentration of acetate anion [A-] = 5.00 M
Initial volume of the buffer solution = 1.00 L
The moles of HCl added to the buffer solution is given as 0.010 mole, and the total volume of the solution remains constant.
Therefore, the new concentration of acetic acid ([HA]) can be calculated using the equation:
[HA] = (moles of acid / total volume) = (5.00 mole - 0.010 mole) / 1.00 L
The new concentration of acetate anion ([A-]) can be calculated by subtracting the change in the concentration of acetic acid from the initial concentration of acetate anion:
[A-] = (initial concentration of acetate anion - change in acetic acid concentration)
= 5.00 M - (0.010 mole / 1.00 L)
Using the new concentrations of [HA] and [A-], we can now calculate the final pH using the Henderson-Hasselbalch equation:
pH = pKa + log([A-]/[HA])
First, let's calculate the new concentrations of [HA] and [A-]:
[HA] = (5.00 mole - 0.010 mole) / 1.00 L = 4.99 M
[A-] = 5.00 M - (0.010 mole / 1.00 L) = 4.99 M
Using the Henderson-Hasselbalch equation:
pH = pKa + log([A-]/[HA])
= -log(1.8 x 10^-5) + log(4.99/4.99)
= -(-4.74) + log(1.00)
= 4.74
Therefore, the final pH of the buffer solution after the addition of the HCl gas is also 4.74.