Multiple Choice
What is the pH of a 0.145 M solution of (CH3)3N, given that the Kb of (CH3)3N is 6.4 × 10⁻⁵?
17. Acid and Base Equilibrium
pH of Weak Bases
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What is the pH of a one liter solution that is 0.100 M in NH3 and 0.100 M in NH4Cl after 1.4 g of NaOH has been added? Kb for NH3 is 1.8 × 10^-5.
A
8.50
B
7.00
C
9.25
D
11.00
Verified step by step guidance1
Identify the components of the buffer solution: NH3 (ammonia) is a weak base and NH4Cl (ammonium chloride) is its conjugate acid. This is a buffer solution.
Calculate the moles of NaOH added: Use the molar mass of NaOH (approximately 40.00 g/mol) to convert 1.4 g of NaOH to moles.
Determine the effect of NaOH on the buffer: NaOH will react with NH4+ to form NH3 and water. Calculate the change in moles of NH4+ and NH3 after the reaction.
Use the Henderson-Hasselbalch equation to find the pH: The equation is \( \text{pH} = \text{pK}_a + \log \left( \frac{[\text{base}]}{[\text{acid}]} \right) \). First, calculate \( \text{pK}_a \) from \( \text{K}_b \) using \( \text{pK}_a + \text{pK}_b = 14 \).
Substitute the concentrations of NH3 and NH4+ into the Henderson-Hasselbalch equation to calculate the pH of the solution after the addition of NaOH.
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