Define the acid dissociation constant, K a .

K a is the equilibrium constant for the partial ionization of a weak acid HA in water: K a = [A - ][H + ] / [HA] A larger K a indicates a stronger weak acid. pK a = –log K a ; a larger pK a means a weaker acid.

True or False? Weak acids completely ionize in aqueous solution.

False. Weak acids only partially ionize; the equilibrium lies to the left, so most acid molecules remain un-ionized. This is what distinguishes them from strong acids, which fully ionize.

How is an ICE table used to find the pH of a weak acid?

An ICE table tracks I nitial concentrations, C hange (using +x / –x) and E quilibrium concentrations. The equilibrium row is substituted into the K a expression to solve for x = [H 3 O + ], then pH = –log(x). This approach avoids assuming x is negligible unless Ka is very small relative to the initial concentration.

Define the base dissociation constant, K b .

K b is the equilibrium constant for the partial ionization of a weak base B in water: K b = [BH + ][OH - ] / [B] A larger K b indicates a stronger weak base. pK b = –log K b ; a larger pK b means a weaker base.

Why does a weak acid solution have a higher pH than a strong acid of the same concentration?

A weak acid only partially ionizes, so [ H 3 O + ] is much lower than the initial acid concentration. A strong acid of the same concentration fully ionizes, producing far more H 3 O + ions. Because pH = –log[H 3 O + ], the lower [H 3 O + ] in the weak acid solution gives a higher pH.

True or False? A larger pK a value indicates a stronger acid.

False. A larger pK a indicates a weaker acid. pK a = –log K a , so a higher pK a corresponds to a smaller K a , meaning less ionization and a weaker acid.

The percentage ionization of a weak acid is calculated as: % ionization = ([H 3 O + ] / .......... ) × 100, where the denominator is the .......... concentration of the acid.

The percentage ionization of a weak acid is calculated as: % ionization = ([H 3 O + ] / [HA] ) × 100, where the denominator is the initial concentration of the acid.

What is the net ionic equation for a strong acid reacting with a strong base?

H 3 O + (aq) + OH - (aq) → 2H 2 O (l) or equivalently H + (aq) + OH - (aq) → H 2 O (l). The spectator ions cancel. The pH of the resulting solution is determined by which reagent (if any) is in excess.

True or False? When a weak base is in excess after reaction with a strong acid, a buffer solution is formed.

True. The excess weak base (B) and the conjugate acid salt (BH + ) produced by the reaction form a buffer. Its pOH can be calculated using the Henderson–Hasselbalch equation: pOH = pK b + log([BH + ]/[B]).

Define a buffer solution in the context of acid-base reactions.

A buffer is a solution that resists significant changes in pH when small amounts of acid or base are added. It consists of a weak acid and its conjugate base (or a weak base and its conjugate acid) in comparable concentrations, typically formed when one reactant is in partial excess.

In a weak acid–weak base reaction, how do you predict which side of the equilibrium is favoured?

The equilibrium favours the side containing the weaker acid and weaker base. If K > 1: products are the weaker acid/base — reaction favours products. If K < 1: reactants are the weaker acid/base — reaction favours reactants. Comparing K a and K b values for the species on each side determines the direction.

True or False? When a strong base is in excess after reaction with a weak acid, the pH is determined using the Henderson–Hasselbalch equation.

False. When the strong base is in excess (not the weak acid), the pH is determined from the moles of excess OH - and the total solution volume: [H + ] = K w / [OH - ]. The Henderson–Hasselbalch equation applies only when both the weak acid and its conjugate base are present in significant amounts (buffer condition).

Why does a weak acid–strong base reaction produce a slightly basic solution at the equivalence point?

At the equivalence point , all the weak acid has been converted to its conjugate base A - , which hydrolyzes water (A - + H 2 O ⇌ HA + OH - ), producing OH - ions and raising the pH above 7. The solution is basic because the conjugate base has a greater tendency to accept protons than to donate them.

When excess OH - ions remain after a strong acid–strong base reaction, the pH is calculated from the .......... of excess H + or OH - ions and the .......... of the combined solution.

When excess OH - ions remain after a strong acid–strong base reaction, the pH is calculated from the moles of excess H + or OH - ions and the total volume of the combined solution.

Define the conjugate base of an acid.

The conjugate base is the species formed when an acid donates a proton (H + ). For acid HA: HA ⇌ A - + H + A - is the conjugate base. The stronger the acid, the weaker (more stable) its conjugate base.

Why does acid strength increase down Group 17 for binary acids (HF → HI)?

Going down Group 17, the H–X bond becomes longer and weaker ( bond enthalpy decreases: HF 567 kJ mol -1 , HI 299 kJ mol -1 ), making the bond easier to break and proton release more favorable. HI therefore ionizes far more readily than HF and is a much stronger acid.

True or False? A stronger acid has a more stable conjugate base.

True. A more stable conjugate base is less likely to recapture a proton, so the equilibrium lies further to the right (more ionization). Greater stability of A - therefore corresponds to greater acid strength of HA.

How does the inductive effect explain why trichloroacetic acid is a stronger acid than acetic acid?

The three electronegative Cl atoms in trichloroacetic acid (CCl 3 COOH) withdraw electron density from the O–H bond via the inductive effect , weakening it and stabilizing the conjugate base (CCl 3 COO - ) by dispersing its negative charge. This makes H + far easier to donate than in acetic acid (pK a 0.6 vs. 4.8).

True or False? Strong bases such as NaOH have weak conjugate acids.

True. Strong bases ionize completely, which means their conjugate acids have very little tendency to donate a proton. For example, the conjugate acid of OH - is H 2 O (pK a ≈ 15.7), which is an extremely weak acid.

Define the inductive effect in the context of acid strength.

The inductive effect is the withdrawal of electron density through sigma bonds by electronegative atoms or groups. In oxy acids, electronegative substituents on the carbon attached to –COOH pull electron density away from the O–H bond , weakening it and increasing the ease of proton donation, thus increasing acid strength.

In comparing binary acids across a period, the .......... of the H–X bond is the dominant factor; in comparing them down a group, the .......... of the H–X bond is the dominant factor.

In comparing binary acids across a period, the polarity of the H–X bond is the dominant factor; in comparing them down a group, the strength (bond enthalpy) of the H–X bond is the dominant factor.

pK a = –log(K a ), where K a is the acid dissociation constant. A low pK a indicates a strong (more ionized) acid. A high pK a indicates a weak (less ionized) acid. To recover K a from pK a : K a = 10 -pKa .

State the Henderson–Hasselbalch equation and explain each term.

pH = pK a + log([A - ] / [HA]) pK a : negative log of the acid dissociation constant [A - ] : molar concentration of the conjugate base [HA] : molar concentration of the weak acid This equation links the pH of a buffer to the pK a of the weak acid and the ratio of conjugate base to acid.

True or False? At the half-equivalence point of a weak acid–strong base titration, pH = pK a .

True. At the half-equivalence point, exactly half the weak acid has been neutralized, so [HA] = [A - ]. Substituting into Henderson–Hasselbalch: pH = pK a + log(1) = pK a + 0 = pK a .

How does pH differ from pK a as measures of acidity?

pH measures the acidity of a specific solution and depends on both the acid identity and its concentration, whereas pK a is a fixed property of the acid itself (independent of concentration) used to compare inherent acid strengths. A low pK a means a strong acid; a low pH means an acidic solution at that moment.

True or False? A buffer region appears on a strong acid–strong base titration curve.

False. A buffer region only appears when a weak acid or weak base is involved (e.g. strong base–weak acid or strong acid–weak base titrations). In a strong acid–strong base titration, there is no weak acid/conjugate base pair to resist pH change, so no buffer region forms.

The Henderson–Hasselbalch equation states that pH = pK a + log( .......... / .......... ), where the numerator is the concentration of the conjugate base and the denominator is the concentration of the weak acid.

The Henderson–Hasselbalch equation states that pH = pK a + log( [A - ] / [HA] ), where the numerator is the concentration of the conjugate base and the denominator is the concentration of the weak acid.

AP®ChemistryCollege BoardFlashcardsUnit 8: Acids & BasesAcids & Bases

Acids & Bases (College Board AP® Chemistry): Flashcards

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FrontIonization of Weak Acids & Weak Bases
Define the acid dissociation constant, K_a.

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Define the acid dissociation constant, K_a.
K_a is the equilibrium constant for the partial ionization of a weak acid HA in water:
K_a = [A^-][H⁺] / [HA]
A larger K_a indicates a stronger weak acid. pK_a = –log K_a; a larger pK_a means a weaker acid.
True or False?
Weak acids completely ionize in aqueous solution.
False.
Weak acids only partially ionize; the equilibrium lies to the left, so most acid molecules remain un-ionized. This is what distinguishes them from strong acids, which fully ionize.
How is an ICE table used to find the pH of a weak acid?
An ICE table tracks Initial concentrations, Change (using +x / –x) and Equilibrium concentrations. The equilibrium row is substituted into the K_a expression to solve for x = [H₃O⁺], then pH = –log(x). This approach avoids assuming x is negligible unless Ka is very small relative to the initial concentration.
Define the base dissociation constant, K_b.
K_b is the equilibrium constant for the partial ionization of a weak base B in water:
K_b = [BH⁺][OH^-] / [B]
A larger K_b indicates a stronger weak base. pK_b = –log K_b; a larger pK_b means a weaker base.
Why does a weak acid solution have a higher pH than a strong acid of the same concentration?
A weak acid only partially ionizes, so [H₃O⁺] is much lower than the initial acid concentration. A strong acid of the same concentration fully ionizes, producing far more H₃O⁺ ions. Because pH = –log[H₃O⁺], the lower [H₃O⁺] in the weak acid solution gives a higher pH.
True or False?
A larger pK_a value indicates a stronger acid.
False.
A larger pK_a indicates a weaker acid. pK_a = –log K_a, so a higher pK_a corresponds to a smaller K_a, meaning less ionization and a weaker acid.
The percentage ionization of a weak acid is calculated as: % ionization = ([H₃O⁺] / .......... ) × 100, where the denominator is the .......... concentration of the acid.
The percentage ionization of a weak acid is calculated as: % ionization = ([H₃O⁺] / [HA] ) × 100, where the denominator is the initial concentration of the acid.
What is the net ionic equation for a strong acid reacting with a strong base?
H₃O⁺ (aq) + OH^- (aq) → 2H₂O (l)
or equivalently H⁺ (aq) + OH^- (aq) → H₂O (l).
The spectator ions cancel. The pH of the resulting solution is determined by which reagent (if any) is in excess.
True or False?
When a weak base is in excess after reaction with a strong acid, a buffer solution is formed.
True.
The excess weak base (B) and the conjugate acid salt (BH⁺) produced by the reaction form a buffer. Its pOH can be calculated using the Henderson–Hasselbalch equation: pOH = pK_b + log([BH⁺]/[B]).
Define a buffer solution in the context of acid-base reactions.
A buffer is a solution that resists significant changes in pH when small amounts of acid or base are added. It consists of a weak acid and its conjugate base (or a weak base and its conjugate acid) in comparable concentrations, typically formed when one reactant is in partial excess.
In a weak acid–weak base reaction, how do you predict which side of the equilibrium is favoured?
The equilibrium favours the side containing the weaker acid and weaker base.
1. If K > 1: products are the weaker acid/base — reaction favours products.
2. If K < 1: reactants are the weaker acid/base — reaction favours reactants.
Comparing K_a and K_b values for the species on each side determines the direction.
True or False?
When a strong base is in excess after reaction with a weak acid, the pH is determined using the Henderson–Hasselbalch equation.
False.
When the strong base is in excess (not the weak acid), the pH is determined from the moles of excess OH^- and the total solution volume: [H⁺] = K_w / [OH^-]. The Henderson–Hasselbalch equation applies only when both the weak acid and its conjugate base are present in significant amounts (buffer condition).
Why does a weak acid–strong base reaction produce a slightly basic solution at the equivalence point?
At the equivalence point, all the weak acid has been converted to its conjugate base A^-, which hydrolyzes water (A^- + H₂O ⇌ HA + OH^-), producing OH^- ions and raising the pH above 7. The solution is basic because the conjugate base has a greater tendency to accept protons than to donate them.
When excess OH^- ions remain after a strong acid–strong base reaction, the pH is calculated from the .......... of excess H⁺ or OH^- ions and the .......... of the combined solution.
When excess OH^- ions remain after a strong acid–strong base reaction, the pH is calculated from the moles of excess H⁺ or OH^- ions and the total volume of the combined solution.
Define the conjugate base of an acid.
The conjugate base is the species formed when an acid donates a proton (H⁺). For acid HA:
HA ⇌ A^- + H⁺
A^- is the conjugate base. The stronger the acid, the weaker (more stable) its conjugate base.
Why does acid strength increase down Group 17 for binary acids (HF → HI)?
Going down Group 17, the H–X bond becomes longer and weaker (bond enthalpy decreases: HF 567 kJ mol^-1, HI 299 kJ mol^-1), making the bond easier to break and proton release more favorable. HI therefore ionizes far more readily than HF and is a much stronger acid.
True or False?
A stronger acid has a more stable conjugate base.
True.
A more stable conjugate base is less likely to recapture a proton, so the equilibrium lies further to the right (more ionization). Greater stability of A^- therefore corresponds to greater acid strength of HA.
How does the inductive effect explain why trichloroacetic acid is a stronger acid than acetic acid?
The three electronegative Cl atoms in trichloroacetic acid (CCl₃COOH) withdraw electron density from the O–H bond via the inductive effect, weakening it and stabilizing the conjugate base (CCl₃COO^-) by dispersing its negative charge. This makes H⁺ far easier to donate than in acetic acid (pK_a 0.6 vs. 4.8).
True or False?
Strong bases such as NaOH have weak conjugate acids.
True.
Strong bases ionize completely, which means their conjugate acids have very little tendency to donate a proton. For example, the conjugate acid of OH^- is H₂O (pK_a ≈ 15.7), which is an extremely weak acid.
Define the inductive effect in the context of acid strength.
The inductive effect is the withdrawal of electron density through sigma bonds by electronegative atoms or groups. In oxy acids, electronegative substituents on the carbon attached to –COOH pull electron density away from the O–H bond, weakening it and increasing the ease of proton donation, thus increasing acid strength.
In comparing binary acids across a period, the .......... of the H–X bond is the dominant factor; in comparing them down a group, the .......... of the H–X bond is the dominant factor.
In comparing binary acids across a period, the polarity of the H–X bond is the dominant factor; in comparing them down a group, the strength (bond enthalpy) of the H–X bond is the dominant factor.
Define pK_a.
pK_a = –log(K_a), where K_a is the acid dissociation constant.
- A low pK_a indicates a strong (more ionized) acid.
- A high pK_a indicates a weak (less ionized) acid.
To recover K_a from pK_a: K_a = 10^-pKa.
State the Henderson–Hasselbalch equation and explain each term.
pH = pK_a + log([A^-] / [HA])
- pK_a: negative log of the acid dissociation constant
- [A^-]: molar concentration of the conjugate base
- [HA]: molar concentration of the weak acid
This equation links the pH of a buffer to the pK_a of the weak acid and the ratio of conjugate base to acid.
True or False?
At the half-equivalence point of a weak acid–strong base titration, pH = pK_a.
True.
At the half-equivalence point, exactly half the weak acid has been neutralized, so [HA] = [A^-]. Substituting into Henderson–Hasselbalch: pH = pK_a + log(1) = pK_a + 0 = pK_a.
How does pH differ from pK_a as measures of acidity?
pH measures the acidity of a specific solution and depends on both the acid identity and its concentration, whereas pK_a is a fixed property of the acid itself (independent of concentration) used to compare inherent acid strengths. A low pK_a means a strong acid; a low pH means an acidic solution at that moment.
True or False?
A buffer region appears on a strong acid–strong base titration curve.
False.
A buffer region only appears when a weak acid or weak base is involved (e.g. strong base–weak acid or strong acid–weak base titrations). In a strong acid–strong base titration, there is no weak acid/conjugate base pair to resist pH change, so no buffer region forms.
The Henderson–Hasselbalch equation states that pH = pK_a + log(.......... / .......... ), where the numerator is the concentration of the conjugate base and the denominator is the concentration of the weak acid.
The Henderson–Hasselbalch equation states that pH = pK_a + log([A^-] / [HA]), where the numerator is the concentration of the conjugate base and the denominator is the concentration of the weak acid.

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