Multiple Choice
Consider the titration of a 26.0-mL sample of 0.175 M CH3NH2 (Kb=4.4×10⁻⁴) with 0.150 M HBr. Determine the pH at the equivalence point.
18. Aqueous Equilibrium
Titrations: Weak Base-Strong Acid
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What is the initial pH for the titration of 30.0 mL of 0.850 M NH3 with 0.100 M HCl. Kb for NH3 = 1.8 × 10−5.
A
2.40
B
4.14
C
4.66
D
9.34
E
11.60
Verified step by step guidance1
Start by identifying the species involved in the solution. Here, we have ammonia (NH3), which is a weak base, and its concentration is given as 0.850 M.
To find the initial pH, we need to calculate the pOH first, since NH3 is a base. Use the formula for the base dissociation constant (Kb) to find the concentration of OH⁻ ions: \( K_b = \frac{[NH_4^+][OH^-]}{[NH_3]} \).
Assume that the initial concentration of OH⁻ is x. Therefore, \([NH_4^+] = x\) and \([OH^-] = x\). Substitute these into the Kb expression: \( 1.8 \times 10^{-5} = \frac{x^2}{0.850 - x} \).
Since Kb is small, we can assume \( x \ll 0.850 \), simplifying the expression to \( 1.8 \times 10^{-5} = \frac{x^2}{0.850} \). Solve for x to find the concentration of OH⁻ ions.
Convert the concentration of OH⁻ ions to pOH using the formula \( pOH = -\log[OH^-] \). Finally, use the relationship \( pH + pOH = 14 \) to find the initial pH of the solution.
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