CHEM 1005 Exam 4 Study Guidance

Q1. Which of the following statements is true about strong acids?

Background

Topic: Acid Strength and Properties

This question tests your understanding of what defines a strong acid and how it behaves in aqueous solution compared to weak acids.

Key Terms:

• Strong acid: An acid that completely ionizes in water.

• pH: A measure of hydrogen ion concentration; lower pH means higher acidity.

• Conjugate base: The species formed when an acid donates a proton.

Step-by-Step Guidance

1. Recall that a strong acid is defined by its ability to completely ionize in water, meaning all acid molecules donate their protons to water.

2. Consider what happens to the pH of a solution when a strong acid is dissolved compared to a weak acid of the same concentration.

3. Think about the properties of the conjugate base of a strong acid—does it tend to be strong or weak?

4. Review the difference between acid strength (degree of ionization) and acid concentration (amount of acid dissolved).

Try solving on your own before revealing the answer!

Final Answer: B. Strong acids are 100% ionized in water.

Strong acids dissociate completely in aqueous solution, meaning every molecule donates its proton. Their conjugate bases are very weak, and strong acid solutions have lower pH than weak acids of the same concentration.

Q2. Calculate the pH of a strong acid solution formed by mixing 30.0 mL of 0.050 M HBr with 20.0 mL of 0.015 M HNO3.

Background

Topic: Acid-Base Calculations, pH of Mixed Solutions

This question tests your ability to calculate the pH of a solution formed by mixing two strong acids of different concentrations and volumes.

Key Terms and Formulas:

• Strong acid: Completely ionizes in water, so [H+] equals the acid's concentration.

• pH formula: $pH = -\log_{10}[H^+]$

• Dilution formula: $n = M \times V$ (moles = molarity × volume)

Step-by-Step Guidance

1. Calculate the moles of H+ contributed by each acid:

   • For HBr: $n_{HBr} = 0.050 \; M \times 0.0300 \; L$

   • For HNO3: $n_{HNO_3} = 0.015 \; M \times 0.0200 \; L$

2. Add the moles of H+ from both acids to get the total moles of H+ in the mixture.

3. Calculate the total volume of the mixed solution: $V_{total} = 30.0 \; mL + 20.0 \; mL = 50.0 \; mL = 0.0500 \; L$

4. Find the final concentration of H+: $[H^+] = \frac{\text{total moles of } H^+}{0.0500 \; L}$

Try solving on your own before revealing the answer!

Final Answer: E. The pH of this solution is 1.44

After calculating the total moles of H+ and dividing by the total volume, you use the pH formula to find the answer.

Q3. Arrange the acids HBrO, HBrO4, and HBrO2 in order of increasing acid strength.

Background

Topic: Oxyacid Strength Trends

This question tests your understanding of how the number of oxygen atoms affects the strength of oxyacids with the same central atom.

Key Terms:

• Oxyacid: An acid with hydrogen, oxygen, and another element (the central atom).

• Acid strength trend: For oxyacids with the same central atom, more oxygen atoms generally mean a stronger acid.

Step-by-Step Guidance

1. List the formulas and count the number of oxygen atoms in each acid:

   • HBrO: 1 oxygen

   • HBrO2: 2 oxygens

   • HBrO4: 4 oxygens

2. Recall the trend: For acids with the same central atom, acid strength increases with the number of oxygens.

3. Arrange the acids from least to most oxygens to determine the order of increasing acid strength.

Try solving on your own before revealing the answer!

Final Answer: A. HBrO < HBrO2 < HBrO4

More oxygens stabilize the conjugate base, making the acid stronger.

Q4. A 500-mL buffer solution is made up of 0.500 M in fulminic acid (HCNO) and 0.600 M in sodium fulminate (NaCNO). What is the pH of this buffer solution? Ka of HCNO is 3.51 × 10–4

Background

Topic: Buffer Solutions and the Henderson-Hasselbalch Equation

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

Key Formula:

Henderson-Hasselbalch Equation:

$pH = pK_a + \log_{10}\left(\frac{[\text{A}^-]}{[\text{HA}]}\right)$

• $[\text{A}^-]$ = concentration of conjugate base (NaCNO)

• $[\text{HA}]$ = concentration of weak acid (HCNO)

• $K_a$ = acid dissociation constant

Step-by-Step Guidance

1. Calculate $pK_a$ from $K_a$:

   $pK_a = -\log_{10}(3.51 \times 10^{-4})$

2. Identify $[\text{A}^-] = 0.600 \; M$ and $[\text{HA}] = 0.500 \; M$.

3. Plug these values into the Henderson-Hasselbalch equation:

   $pH = pK_a + \log_{10}\left(\frac{0.600}{0.500}\right)$

Try solving on your own before revealing the answer!

Final Answer: B. The pH of this solution is 3.53

After calculating $pK_a$ and the log ratio, you find the pH using the Henderson-Hasselbalch equation.

Q5. Which one of the following pairs when mixed will form a buffer solution?

Background

Topic: Buffer Solutions

This question tests your understanding of what combinations of acids and bases can create a buffer solution.

Key Terms:

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

• Conjugate acid-base pair: A weak acid and its salt (conjugate base), or a weak base and its salt (conjugate acid).

Step-by-Step Guidance

1. Identify which pairs consist of a weak acid and its conjugate base, or a weak base and its conjugate acid.

2. Eliminate pairs that contain strong acids or bases, or two weak acids/bases that are not conjugates.

3. Recall that a buffer is most effective when both components are present in significant amounts.

Try solving on your own before revealing the answer!

Final Answer: A. LiF and HF

HF is a weak acid and LiF provides its conjugate base (F–), forming a buffer.

Q6. A sample of solid Ba(OH)2 was stirred in water at room temperature until the solution contained as much dissolved Ba(OH)2 as it could hold. A 100-mL sample of this solution was withdrawn and titrated with 1.50 M HCl. It required 50.0 mL of the acid solution for neutralization. What is the concentration of the initial solution of Ba(OH)2 in water, at room temperature?

Background

Topic: Titration and Stoichiometry

This question tests your ability to use titration data to determine the concentration of a saturated solution of a base.

Key Terms and Formulas:

• Neutralization reaction: $Ba(OH)_2 + 2HCl \rightarrow BaCl_2 + 2H_2O$

• Moles: $n = M \times V$

Step-by-Step Guidance

1. Calculate the moles of HCl used in the titration:

   $n_{HCl} = 1.50 \; M \times 0.0500 \; L$

2. Use the stoichiometry of the reaction (2:1 ratio) to find the moles of Ba(OH)2 in the 100-mL sample.

3. Calculate the concentration of Ba(OH)2 in the original solution using the moles found and the sample volume (0.100 L).

Try solving on your own before revealing the answer!

Final Answer: D. The concentration is 0.501 M

After finding the moles of HCl, divide by 2 to get moles of Ba(OH)2, then divide by 0.100 L to get the concentration.