Table of contents

1. Intro to General Chemistry3h 48m
2. Atoms & Elements4h 16m
3. Chemical Reactions4h 10m
4. BONUS: Lab Techniques and Procedures1h 38m
5. BONUS: Mathematical Operations and Functions48m
6. Chemical Quantities & Aqueous Reactions3h 53m
7. Gases3h 49m
8. Thermochemistry2h 37m
9. Quantum Mechanics2h 58m
10. Periodic Properties of the Elements3h 9m
11. Bonding & Molecular Structure3h 29m
12. Molecular Shapes & Valence Bond Theory1h 57m
13. Liquids, Solids & Intermolecular Forces2h 23m
14. Solutions3h 1m
15. Chemical Kinetics2h 53m
16. Chemical Equilibrium2h 29m
17. Acid and Base Equilibrium5h 1m
18. Aqueous Equilibrium4h 47m
19. Chemical Thermodynamics1h 50m
20. Electrochemistry2h 42m
21. Nuclear Chemistry2h 36m
22. Organic Chemistry5h 7m
23. Chemistry of the Nonmetals2h 39m
24. Transition Metals and Coordination Compounds3h 16m

18. Aqueous Equilibrium

Titrations: Weak Base-Strong Acid

General Chemistry

18. Aqueous Equilibrium

Titrations: Weak Base-Strong Acid

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Multiple Choice

What is the approximate pH at the equivalence point of a weak base-strong acid titration if 50.00 mL of aqueous ammonia (NH3) requires 37.10 mL of 0.1087 mol L^-1 HCl? pKb = 4.75 for ammonia. Report your answer to 2 decimal places.

4.75

5.25

7.00

9.25

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Verified step by step guidance

First, understand that at the equivalence point of a weak base-strong acid titration, the moles of acid added are equal to the moles of base present initially. Calculate the moles of HCl added using the formula: \( \text{moles of HCl} = \text{volume of HCl} \times \text{concentration of HCl} \).

Next, calculate the initial moles of ammonia (NH3) present using the formula: \( \text{moles of NH3} = \text{volume of NH3} \times \text{concentration of NH3} \). Since the moles of HCl added are equal to the moles of NH3 initially present, you can use this to find the concentration of the ammonium ion (NH4^+) formed at the equivalence point.

At the equivalence point, the solution contains ammonium ions (NH4^+), which is the conjugate acid of ammonia. Use the formula for the equilibrium constant of the conjugate acid, \( K_a = \frac{K_w}{K_b} \), where \( K_w \) is the ion-product constant of water (\( 1.0 \times 10^{-14} \) at 25°C) and \( K_b \) is given as 4.75.

Set up the expression for the dissociation of NH4^+ in water: \( NH4^+ \rightarrow NH3 + H^+ \). Use the expression for \( K_a \) to find \( [H^+] \) at equilibrium: \( K_a = \frac{[NH3][H^+]}{[NH4^+]} \). Assume \( [NH3] \approx [H^+] \) and \( [NH4^+] \) is the initial concentration of NH4^+.

Finally, calculate the pH using the formula: \( \text{pH} = -\log[H^+] \). This will give you the approximate pH at the equivalence point.