Q1. Which of the following best describes an Arrhenius base?

Background

Topic: Acid-Base Definitions

This question tests your understanding of the Arrhenius definition of acids and bases, which is foundational in General Chemistry.

Key Terms:

• Arrhenius base: A substance that increases the concentration of hydroxide ions (OH−) in aqueous solution.

• Arrhenius acid: A substance that increases the concentration of hydronium ions (H3O+) in aqueous solution.

Step-by-Step Guidance

1. Review the Arrhenius definitions for acids and bases.

2. Identify which option describes a substance that produces OH− ions in water.

3. Eliminate options that refer to proton donors/acceptors or electron donors, as these relate to other acid-base theories.

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Q2. Which of the following is amphoteric (able to react both as a base and as an acid)?

Background

Topic: Acid-Base Properties

This question tests your knowledge of amphoteric substances, which can act as either an acid or a base depending on the reaction.

Key Terms:

• Amphoteric: A substance that can act as both an acid and a base.

• Examples: Water (H2O) is a classic amphoteric molecule.

Step-by-Step Guidance

1. Recall the definition of amphoteric substances.

2. Consider each option and whether it can donate and accept protons.

3. Eliminate substances that only act as acids or only as bases.

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Q3. What is the conjugate base of HNO3?

Background

Topic: Conjugate Acids and Bases

This question tests your ability to identify conjugate bases formed when an acid loses a proton.

Key Terms:

• Conjugate base: The species formed when an acid loses a proton (H+).

• Formula:

Step-by-Step Guidance

1. Write the formula for HNO3.

2. Remove one H+ from HNO3 to find the conjugate base.

3. Identify the resulting ion.

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Q4. What is the conjugate base of H3PO4?

Background

Topic: Conjugate Acids and Bases

This question tests your understanding of polyprotic acids and their conjugate bases.

Key Terms:

• Polyprotic acid: An acid that can donate more than one proton.

• Conjugate base: Formed by removing one H+ from the acid.

Step-by-Step Guidance

1. Write the formula for H3PO4.

2. Remove one H+ to find the conjugate base.

3. Identify the resulting ion.

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Q5. Which of the following would you expect the concentration of acid to be nearly the same before and after dissociation?

Background

Topic: Acid Strength and Dissociation

This question tests your understanding of strong vs. weak acids and how their concentrations change upon dissociation.

Key Terms:

• Strong acid: Completely dissociates in water.

• Weak acid: Only partially dissociates.

• Dilute vs. concentrated: Refers to the amount of acid in solution.

Step-by-Step Guidance

1. Recall that strong acids dissociate almost completely, so their concentration changes significantly.

2. Weak acids dissociate very little, so their concentration remains nearly the same.

3. Consider how dilution affects dissociation for weak acids.

4. Match the options to these concepts.

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Q6. Which of the following is true regarding the Ka values of a diprotic acid?

Background

Topic: Acid Dissociation Constants (Ka)

This question tests your understanding of how the dissociation constants change for polyprotic acids.

Key Terms:

• Diprotic acid: An acid that can donate two protons.

• Ka1: The dissociation constant for the first proton.

• Ka2: The dissociation constant for the second proton.

Step-by-Step Guidance

1. Recall that the first proton is easier to remove than the second.

2. Compare the values of Ka1 and Ka2 for typical diprotic acids.

3. Eliminate options that suggest Ka1 is less than or equal to Ka2.

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Q7. Which of the following could act as a Lewis base?

Background

Topic: Lewis Acid-Base Theory

This question tests your understanding of Lewis bases, which donate electron pairs.

Key Terms:

• Lewis base: Electron pair donor.

• Lewis acid: Electron pair acceptor.

Step-by-Step Guidance

1. Identify which substances have lone pairs of electrons available for donation.

2. Eliminate options that lack lone pairs or are cations without available electrons.

3. Consider if more than one option fits the definition.

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Q8. What is the concentration of H3O+ in a solution which has a concentration of OH− equal to 8.9×10–9?

Background

Topic: Water Ionization and pH

This question tests your ability to use the ion product constant for water (Kw) to relate [H3O+] and [OH−].

Key Formula:

Step-by-Step Guidance

1. Write the expression for Kw.

2. Plug in the given [OH−] value.

3. Rearrange to solve for [H3O+].

4. Set up the calculation, but stop before computing the final value.

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Q9. The equilibrium constant Kc for the following reaction is > 1. H3PO4 + NO2− ⇆ H2PO4− + HNO2 Which of the following is correct?

Background

Topic: Acid-Base Equilibria

This question tests your understanding of equilibrium constants and relative acid/base strengths.

Key Terms:

• Kc: Equilibrium constant for the reaction.

• Acid strength: Stronger acids have larger Ka values.

• Base strength: Stronger bases have larger Kb values.

Step-by-Step Guidance

1. Recall that Kc > 1 means products are favored.

2. Compare the acid strengths of H3PO4 and HNO2.

3. Consider the base strengths of NO2− and H2PO4−.

4. Eliminate options inconsistent with the equilibrium favoring products.

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Q10. Which of the following substances will give an aqueous solution with a pH closest to 7?

Background

Topic: Acid-Base Properties of Salts

This question tests your ability to predict the pH of solutions formed from various substances.

Key Terms:

• Neutral salt: Formed from strong acid and strong base.

• Acidic/basic salt: Formed from weak acid/base and strong base/acid.

Step-by-Step Guidance

1. Identify which substances are salts, acids, or bases.

2. Recall that salts from strong acid and strong base yield neutral solutions.

3. Eliminate options that would produce acidic or basic solutions.

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Q11. Which of the following statements is true about buffers?

Background

Topic: Buffer Solutions

This question tests your understanding of buffer properties and their resistance to pH changes.

Key Terms:

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

• Buffer capacity: Greatest when [acid] = [conjugate base].

Step-by-Step Guidance

1. Recall the conditions for maximum buffer capacity.

2. Consider the effect of concentration ratios on buffer effectiveness.

3. Eliminate statements that contradict buffer theory.

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Q12. What is the pH of a buffer that consists of 0.55 M HCOOH and 0.63 M HCOONa (pKa of HCOOH = 3.74)?

Background

Topic: Buffer Calculations (Henderson-Hasselbalch Equation)

This question tests your ability to calculate the pH of a buffer using the Henderson-Hasselbalch equation.

Key Formula:

• [A−] = concentration of conjugate base (HCOONa)

• [HA] = concentration of acid (HCOOH)

• pKa = 3.74

Step-by-Step Guidance

1. Identify the concentrations of acid and conjugate base.

2. Write the Henderson-Hasselbalch equation.

3. Plug in the values for pKa, [A−], and [HA].

4. Set up the calculation for pH, but stop before computing the final value.

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Q13. For a Na2HPO4/KH2PO4 solution with pH 7.40, what is the buffer component concentration ratio? (For phosphoric acid Ka1 = 7.2×10–3; Ka2 = 6.3×10–8; Ka3 = 4.2×10–13)

Background

Topic: Polyprotic Buffer Calculations

This question tests your ability to use the Henderson-Hasselbalch equation for a buffer involving a polyprotic acid.

Key Formula:

• Use Ka2 for the HPO42−/H2PO4− pair.

• pKa = –log(Ka2)

Step-by-Step Guidance

1. Calculate pKa for Ka2.

2. Write the Henderson-Hasselbalch equation for the buffer.

3. Plug in the pH and pKa values.

4. Set up the equation to solve for the concentration ratio, but stop before calculating the final value.

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Q14. The indicator thymolphthalein has Ka = 7.9×10–11. Over what approximate pH range does it change color?

Background

Topic: Acid-Base Indicators

This question tests your understanding of how the pKa of an indicator relates to its color change range.

Key Formula:

Indicator color change occurs around pKa ± 1.

Step-by-Step Guidance

1. Calculate pKa from the given Ka.

2. Determine the pH range by adding and subtracting 1 from the pKa.

3. Match the calculated range to the options provided.

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Q15. Does any solid Cu(OH)2 form when 0.075 g KOH is dissolved in 1.0 L of 1.0×10–3 M Cu(NO3)2?

Background

Topic: Solubility and Precipitation

This question tests your ability to compare the reaction quotient (Q) to the solubility product constant (Ksp).

Key Formula:

Compare Q to Ksp to predict precipitation.

Step-by-Step Guidance

1. Calculate the concentration of OH− from the mass of KOH added.

2. Write the expression for Q.

3. Plug in the concentrations to calculate Q.

4. Set up the comparison between Q and Ksp, but stop before stating the result.

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Q16. Hypochlorous acid (HOCl) is a weak acid with pKa = 7.54. Suppose a 45.00 mL sample of 0.240 M HOCl solution is titrated with a 0.250 M standard KOH solution. What is the pH at the equivalence point?

Background

Topic: Weak Acid-Strong Base Titration

This question tests your ability to calculate the pH at the equivalence point for a weak acid titrated with a strong base.

Key Formula:

At equivalence, the solution contains the conjugate base (OCl−), which hydrolyzes:

Use and

Step-by-Step Guidance

1. Calculate the moles of HOCl and KOH to find the concentration of OCl− at equivalence.

2. Write the hydrolysis equation for OCl−.

3. Calculate Kb for OCl− using Kw and Ka.

4. Set up the equation for [OH−] and pOH, but stop before calculating pH.

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Q17. Find the solubility of AgI in 2.5 M NH3. (Ksp of AgI = 8.3×10–17; Kf of Ag(NH3)2+ = 1.7×107)

Background

Topic: Solubility and Complex Ion Formation

This question tests your ability to calculate solubility in the presence of a complexing agent.

Key Formula:

Combine the solubility and formation reactions:

(Ksp)

(Kf)

Overall:

Step-by-Step Guidance

1. Write the overall reaction and calculate Kx.

2. Set up the solubility expression using Kx and the NH3 concentration.

3. Plug in the values for Kx and [NH3].

4. Set up the equation for solubility, but stop before calculating the final value.

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Q18. A 50.0 mL volume of 0.50 M Fe(NO3)3 is mixed with 125 mL of 0.25 M Cd(NO3)2. If aqueous NaOH is added to the mixture, which ion precipitates first?

Background

Topic: Precipitation and Ksp Values

This question tests your ability to predict which ion will precipitate first based on solubility product constants.

Key Terms:

• Ksp: Solubility product constant.

• Lower Ksp means less soluble and precipitates first.

Step-by-Step Guidance

1. Compare the Ksp values for Fe(OH)3, Fe(OH)2, and Cd(OH)2.

2. Identify which has the lowest Ksp.

3. Determine which ion will reach its precipitation threshold first as NaOH is added.

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Q19. A 50.0-mL sample of 0.50 M HCl is titrated with 0.50 M NaOH. After 28.0 mL of NaOH have been added to the acid, which of the following statements is true?

Background

Topic: Strong Acid-Strong Base Titration

This question tests your understanding of titration progress and pH changes.

Key Terms:

• Equivalence point: Where moles of acid = moles of base.

• Before equivalence: Solution is acidic.

Step-by-Step Guidance

1. Calculate the moles of HCl and NaOH added.

2. Determine if the equivalence point has been reached.

3. Assess whether the solution is acidic, neutral, or basic at this stage.

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Q20. Which of the following titration mixtures will produce an equivalence point pH which is less than 7?

Background

Topic: Acid-Base Titration Equivalence Points

This question tests your ability to predict the pH at the equivalence point based on the acid and base strengths.

Key Terms:

• Strong acid + weak base: Equivalence point pH < 7.

• Strong acid + strong base: Equivalence point pH = 7.

• Weak acid + strong base: Equivalence point pH > 7.

Step-by-Step Guidance

1. Identify which mixtures involve a strong acid and a weak base.

2. Recall the rules for equivalence point pH.

3. Match the options to these rules.

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Q21. What happens if we add NaC2H3O2 to the following solution? HC2H3O2 ⇆ H+ + C2H3O2−

Background

Topic: Common Ion Effect

This question tests your understanding of how adding a common ion affects equilibrium and pH.

Key Terms:

• Common ion effect: Adding an ion already present shifts equilibrium.

• Le Chatelier's Principle: System shifts to counteract the change.

Step-by-Step Guidance

1. Identify the common ion being added (C2H3O2−).

2. Predict how the equilibrium will shift.

3. Determine the effect on [H+] and pH.

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Q22. HF is in solution with NaF. If HCl is added, which ion will react with the additional hydrogen ions from the HCl to prevent a change in pH?

Background

Topic: Buffer Action and Common Ion Effect

This question tests your understanding of how buffers work to resist pH changes.

Key Terms:

• Buffer: Contains a weak acid and its conjugate base.

• F−: Conjugate base of HF, reacts with added H+.

Step-by-Step Guidance

1. Identify the buffer components (HF and F−).

2. Determine which ion can react with added H+ to maintain pH.

3. Recall the buffer reaction: F− + H+ → HF.

Try solving on your own before revealing the answer!