Step-by-Step Guidance for College-Level Acid-Base Chemistry Questions

Q1. What mass of NaCHO2 must be added to a 0.750 L solution of 0.350 M formic acid to make a buffer with a pH of 3.745?

Background

Topic: Buffer Solutions and Henderson-Hasselbalch Equation

This question tests your understanding of how to prepare a buffer solution with a specific pH by adding a conjugate base to a weak acid.

Key Terms and Formulas

• Buffer: A solution that resists changes in pH when small amounts of acid or base are added.

• Henderson-Hasselbalch Equation: $\text{pH} = \text{p}K_a + \log\left(\frac{[\text{A}^-]}{[\text{HA}]}\right)$

• Formic acid (HA), sodium formate (NaCHO2, A-)

Step-by-Step Guidance

1. Write the Henderson-Hasselbalch equation for the buffer: $\text{pH} = \text{p}K_a + \log\left(\frac{[\text{NaCHO}_2]}{[\text{HCHO}_2]}\right)$

2. Find the $\text{p}K_a$ for formic acid (look up or use provided value).

3. Plug in the given pH (3.745) and the concentration of formic acid (0.350 M) into the equation.

4. Rearrange the equation to solve for the concentration of sodium formate needed.

5. Calculate the number of moles of sodium formate required using the volume of the solution.

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Q2. Which solution has the highest pH?

Background

Topic: Strong vs. Weak Acids

This question tests your ability to compare the acidity of different solutions and identify which is least acidic (highest pH).

Key Terms

• Strong acid: Completely dissociates in water (e.g., HBr, HI, HCl, HClO4).

• Weak acid: Partially dissociates (e.g., HF).

• pH: Measure of acidity; higher pH means less acidic.

Step-by-Step Guidance

1. Identify which acids in the list are strong and which are weak.

2. Recall that strong acids dissociate completely, resulting in lower pH.

3. Compare the dissociation of HF (weak acid) to the others (strong acids).

4. Determine which solution will have the least concentration of hydronium ions.

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Q3. Rank the following in order of decreasing acid strength in aqueous solution: HBr, HOCl, HOBr, HOI.

Background

Topic: Acid Strength and Molecular Structure

This question tests your understanding of how molecular structure and electronegativity affect acid strength.

Key Terms

• Acid strength: Related to the tendency to donate a proton.

• Electronegativity: More electronegative atoms stabilize the conjugate base, increasing acid strength.

• Binary acids (HBr) vs. oxyacids (HOCl, HOBr, HOI).

Step-by-Step Guidance

1. Recall that for binary acids, acid strength increases with larger, less electronegative halogens.

2. For oxyacids, acid strength increases with more electronegative central atoms.

3. Compare the electronegativity of Cl, Br, and I.

4. Rank the acids based on these principles.

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Q4. What is Ka for a weak monoprotic acid if a 0.020 M solution of the acid has a pH of 3.28 at 25°C?

Background

Topic: Acid Dissociation Constant (Ka)

This question tests your ability to calculate the Ka of a weak acid from its concentration and pH.

Key Terms and Formulas

• Ka: Acid dissociation constant.

• pH: $\text{pH} = -\log[\text{H}^+]$

• ICE Table: Used to track concentrations during dissociation.

Step-by-Step Guidance

1. Calculate $[\text{H}^+]$ from the given pH: $[\text{H}^+] = 10^{-\text{pH}}$

2. Set up an ICE table for the dissociation of the acid: $\text{HA} \rightleftharpoons \text{H}^+ + \text{A}^-$

3. Express the equilibrium concentrations in terms of $[\text{H}^+]$.

4. Write the Ka expression: $K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}$

5. Plug in the values and set up the calculation for Ka.

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Q5. What is the hydronium-ion concentration of a 0.0041 M LiOH solution?

Background

Topic: Strong Bases and pH Calculations

This question tests your ability to relate the concentration of a strong base to the hydronium ion concentration.

Key Terms and Formulas

• LiOH: Strong base, dissociates completely.

• $[\text{OH}^-] = 0.0041$ M

• $K_w = [\text{H}^+][\text{OH}^-] = 1.0 \times 10^{-14}$ at 25°C

Step-by-Step Guidance

1. Write the relationship: $K_w = [\text{H}_3\text{O}^+][\text{OH}^-]$

2. Plug in the known values for $K_w$ and $[\text{OH}^-]$.

3. Rearrange to solve for $[\text{H}_3\text{O}^+]$.

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Q6. Which of the following is the most effective buffer system for a pH value of 4.45?

Background

Topic: Buffer Systems and pKa

This question tests your ability to select the best buffer system based on the desired pH and the pKa values of the acid/base pairs.

Key Terms and Formulas

• Buffer: Most effective when pH ≈ pKa.

• pKa: $\text{p}K_a = -\log K_a$

Step-by-Step Guidance

1. Calculate the pKa for each acid/base pair using the given Ka values.

2. Compare each pKa to the desired pH (4.45).

3. Identify which buffer system has a pKa closest to 4.45.

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Q7. A solution has a pH of 10.20 at 25°C. What is the hydroxide-ion concentration at 25°C?

Background

Topic: Relationship between pH, pOH, and Ion Concentrations

This question tests your ability to convert pH to pOH and then to hydroxide ion concentration.

Key Terms and Formulas

• $\text{pH} + \text{pOH} = 14$ at 25°C

• $[\text{OH}^-] = 10^{-\text{pOH}}$

Step-by-Step Guidance

1. Calculate pOH: $\text{pOH} = 14 - \text{pH}$

2. Plug in the given pH (10.20) to find pOH.

3. Calculate $[\text{OH}^-]$ using $[\text{OH}^-] = 10^{-\text{pOH}}$.

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Q8. What is the percent ionization of a 1.5 M HC2H3O2 solution (Ka = 1.8 × 10–5) at 25°C?

Background

Topic: Weak Acid Ionization and Percent Ionization

This question tests your ability to calculate the percent ionization of a weak acid using its concentration and Ka.

Key Terms and Formulas

• Percent ionization: $\frac{[\text{H}^+]}{[\text{HA}]_\text{initial}} \times 100\%$

• Ka: Acid dissociation constant.

• ICE Table: Used to determine equilibrium concentrations.

Step-by-Step Guidance

1. Set up an ICE table for the dissociation of acetic acid.

2. Write the Ka expression: $K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}$

3. Assume $x$ is the amount ionized; set up the equation and solve for $x$.

4. Calculate percent ionization using $\frac{x}{1.5} \times 100\%$.

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Q9. What is the pH of a solution that is 0.048 M in HA and also 0.0032 M in NaA? (Ka = 4.1 × 10–6)

Background

Topic: Buffer Solutions and Henderson-Hasselbalch Equation

This question tests your ability to calculate the pH of a buffer solution using the concentrations of acid and conjugate base.

Key Terms and Formulas

• Henderson-Hasselbalch Equation: $\text{pH} = \text{p}K_a + \log\left(\frac{[\text{A}^-]}{[\text{HA}]}\right)$

• pKa: $\text{p}K_a = -\log K_a$

Step-by-Step Guidance

1. Calculate pKa from the given Ka.

2. Plug the concentrations of HA and NaA into the Henderson-Hasselbalch equation.

3. Set up the calculation for pH.

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Q10. List if the following are acidic or basic solutions: a. Sodium Acetate (NaC2H3O2), b. FeCl3, c. Ammonium Acetate (NH4C2H3O2)

Background

Topic: Salt Hydrolysis and Solution pH

This question tests your ability to determine whether a salt solution is acidic, basic, or neutral based on the properties of its ions.

Key Terms

• Salt hydrolysis: Some salts form acidic or basic solutions depending on their constituent ions.

• Ka and Kb: Used to determine the strength of acidic or basic ions.

Step-by-Step Guidance

1. For each salt, identify the ions and whether they are acidic or basic.

2. Use the provided Ka and Kb values to compare the relative strengths.

3. Determine if the solution will be acidic, basic, or neutral based on the ion strengths.

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Q11. A 225 mL buffer solution is 0.485 M in HCN and 0.385 M in LiCN. Calculate the initial pH after the addition of 0.0350 moles of NaOH and after the addition of 0.0225 moles of HCl. The Ka for HCN is 6.2 × 10–10

Background

Topic: Buffer Capacity and pH Changes

This question tests your ability to calculate pH changes in a buffer solution after adding strong acid or base.

Key Terms and Formulas

• Buffer: Solution of weak acid and its conjugate base.

• Henderson-Hasselbalch Equation: $\text{pH} = \text{p}K_a + \log\left(\frac{[\text{A}^-]}{[\text{HA}]}\right)$

• pKa: $\text{p}K_a = -\log K_a$

Step-by-Step Guidance

1. Calculate the initial moles of HCN and LiCN in the buffer.

2. Determine how the addition of NaOH or HCl will change the moles of acid and base in the buffer.

3. Update the concentrations based on the new moles and total volume.

4. Use the Henderson-Hasselbalch equation to set up the calculation for the new pH after each addition.

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