Calculate the pH of a solution made by mixing 60.0 ml of 0.80 M N2H4 (a weak base) with 50.0 ml of 0.50 M HCl. (Kb for N2H4=1.7 x 10^-6)
Can I use the Henderson Hasselbach equation here?
No, you cannot use the Henderson-Hasselbalch equation directly in this case because it is applicable to buffer solutions, which are composed of a weak acid and its conjugate base or a weak base and its conjugate acid. In this problem, you are dealing with a mixture of a weak base (N2H4) and a strong acid (HCl).
To calculate the pH of the resulting solution, you need to consider the reaction between N2H4 and HCl. Since HCl is a strong acid, it will completely dissociate in water, whereas N2H4 is a weak base that will undergo a partial dissociation reaction.
First, let's write the balanced equation for the reaction:
N2H4 + HCl -> N2H5+ + Cl-
Now, let's calculate the number of moles of N2H4 and HCl in the given volumes:
moles of N2H4 = volume (in liters) × concentration
= 0.060 L × 0.80 M
= 0.048 moles
moles of HCl = volume (in liters) × concentration
= 0.050 L × 0.50 M
= 0.025 moles
Since the weak base (N2H4) reacts with the strong acid (HCl), the excess moles of HCl will determine the pH of the solution.
Next, let's calculate the concentration of the excess HCl:
excess moles of HCl = moles of HCl - moles of N2H4
= 0.025 moles - 0.048 moles
= -0.023 moles
Since the moles of HCl are now negative, it indicates that all the HCl has reacted with N2H4, and 0.023 moles of HCl is in excess.
To find the concentration (or molarity) of the excess HCl, divide the moles by the total volume of the solution (given by the sum of the initial volumes):
excess concentration of HCl = excess moles / total volume
= -0.023 moles / (0.060 L + 0.050 L)
= -0.023 moles / 0.110 L
= -0.21 M
Finally, to calculate the pH, take the negative logarithm (base 10) of the excess HCl concentration:
pH = -log10(excess HCl concentration)
= -log10(-0.21)
= 0.68
Note: The negative sign in the concentration arises because excess HCl is present in greater quantity than N2H4.