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Quizzes > Science > Chemistry

Buffer Solution Practice Problems Quiz - Challenge Yourself!

Ready to tackle buffer practice problems? Dive in and ace the buffer solution quiz!

Difficulty: Moderate
2-5mins
Learning OutcomesCheat Sheet
Paper art illustration for a quiz testing buffer solution practice problems skills on a coral background

Ready to master buffer solution practice problems? Dive into our free quiz that challenges your knowledge of acid-base buffer questions and pH buffer practice problems. Whether you're brushing up the Henderson-Hasselbalch equation or tackling real-world scenarios, this buffer solution quiz will test and strengthen your skills. Curious chemists and students can expand their practice with titration practice problems or reinforce learning through aqueous solution practice problems . Take the plunge now - tackle these buffer practice problems and see how you score today!

What is a buffer solution?
A solution that resists large pH changes upon addition of small amounts of acid or base
A solution that completely neutralizes any added acid
A solution that contains a strong acid and its salt
A solution whose pH remains exactly constant regardless of additions
A buffer solution contains a weak acid and its conjugate base (or a weak base and its conjugate acid) and resists pH change when small amounts of acid or base are added. It works by shifting the equilibrium of the acid–base pair to consume added H+ or OH– ions. This property is essential in many biological and analytical applications where pH stability is critical. Learn more.
Which equation correctly represents the Henderson–Hasselbalch relationship for a weak acid buffer?
pH = pKa + log([A–]/[HA])
pH = pKa – log([A–]/[HA])
pH = pKa + log([HA]/[A–])
pH = pKa – log([HA]/[A–])
The Henderson–Hasselbalch equation relates the pH of a buffer to the pKa of the weak acid and the ratio of its conjugate base to acid concentrations. It is derived by taking the negative logarithm of the acid dissociation equilibrium expression. This form is widely used for quick pH estimates in buffer preparation. See details.
Which of the following mixtures will act as a buffer?
Acetic acid and sodium acetate
Hydrochloric acid and sodium chloride
Sodium hydroxide and hydrochloric acid
Potassium nitrate and water
A buffer requires a weak acid and its conjugate base. Acetic acid (CH?COOH) with sodium acetate (CH?COONa) provides both species in solution and thus resists pH changes. Strong acid–salt or neutral salt solutions do not exhibit significant buffering action. More on buffer components.
Which acid/conjugate base pair is appropriate for a buffer at pH 4.75?
Acetic acid (pKa 4.76) and acetate
Formic acid (pKa 3.75) and formate
Benzoic acid (pKa 4.20) and benzoate
Phosphoric acid (pKa 2.15) and dihydrogen phosphate
A buffer is most effective when its pH is close to the acid’s pKa (±1 unit). Acetic acid has a pKa of 4.76, which matches the target pH of 4.75. Other acids have pKa values that are too far from 4.75 to provide optimal buffering. Buffer range details.
Calculate the pH of a buffer containing 0.10 M acetic acid (pKa 4.76) and 0.10 M sodium acetate.
4.76
5.00
4.50
7.00
When [A–] equals [HA], the Henderson–Hasselbalch equation simplifies to pH = pKa + log(1) = pKa. Thus the pH equals 4.76. This straightforward case illustrates how equal concentrations give pH equal to pKa. More examples.
What is the effective buffer pH range for a weak acid buffer?
pKa ± 1
pKa ± 0.1
pKa ± 2
pKa ± 3
A buffer effectively resists pH changes within approximately one pH unit above and below the acid’s pKa. Outside this range, the weaker component becomes too small to maintain pH stability. This guideline helps in selecting the appropriate buffer system. Buffer range guide.
Which best describes buffer capacity?
The amount of strong acid or base required to change the pH by one unit in one liter of buffer
The ratio of conjugate base to acid in a buffer
The total concentration of buffer components
The volume of buffer needed to neutralize strong acid
Buffer capacity quantifies how much strong acid or base a buffer can absorb before its pH changes by one unit (usually expressed per liter). It depends on the absolute concentrations of the buffering species. Higher concentrations yield higher capacity. Read more.
If the concentrations of acid [HA] and conjugate base [A–] are equal, what is the pH of the buffer?
Equal to the pKa of the acid
Neutral (pH 7)
Always less than 7
Always greater than 7
When [HA] = [A–], the Henderson–Hasselbalch equation reduces to pH = pKa + log(1) = pKa. This condition simplifies buffer calculations and is often used in lab preparations. It does not depend on absolute concentrations. Further reading.
What is the pH of a buffer with 0.20 M acetic acid (pKa 4.76) and 0.10 M sodium acetate?
4.46
4.76
5.06
3.76
Use Henderson–Hasselbalch: pH = 4.76 + log(0.10/0.20) = 4.76 + log(0.5) ? 4.76 – 0.301 = 4.46. This shows how changing the ratio shifts the pH. Calculation details.
If you add 0.01 mol HCl to 1 L of 0.10 M acetic acid/0.10 M acetate buffer, what is the new pH? (pKa 4.76)
4.67
4.76
4.56
4.90
HCl converts 0.01 mol acetate to acetic acid: [HA]=0.11 M, [A–]=0.09 M, so pH = 4.76 + log(0.09/0.11) ? 4.67. This illustrates buffer action against added acid. Buffer addition example.
To prepare a 1 L buffer at pH 7.40 using the H?PO??/HPO?²? system (pKa = 7.20), what ratio of [HPO?²?]/[H?PO??] is required?
1.58
0.63
1.00
2.00
Henderson–Hasselbalch: 7.40 = 7.20 + log(ratio) ? ratio = 10^(0.20) ? 1.58. This ratio of base to acid yields the desired pH. See the derivation.
What happens to the pH of a buffer solution upon dilution?
It remains essentially constant
It increases significantly
It decreases significantly
It fluctuates unpredictably
Dilution changes both [HA] and [A–] equally, so their ratio remains constant and the pH stays nearly the same. Only ionic strength and activity coefficients might cause minor shifts. This is why buffers are effective even when diluted. More on dilution effects.
At what pH is a buffer’s capacity highest?
Equal to the pKa of the acid
Two units above the pKa
At pH 7 for all buffers
When [A–] = 0
Buffer capacity peaks when [HA] = [A–], which corresponds to pH = pKa. Here the solution has maximal ability to neutralize added acid or base. Deviations from this ratio reduce capacity. Buffer capacity.
Which weak acid is most suitable for a buffer at pH 5.20?
Acetic acid (pKa 4.76)
Formic acid (pKa 3.75)
Phosphoric acid (pKa 2.15)
Benzoic acid (pKa 4.20)
You choose an acid with pKa closest to the desired pH. Acetic acid’s pKa of 4.76 is within ±1 of pH 5.20, making it the best choice. Other candidates are too far from 5.20 to buffer effectively. Buffer selection guide.
Calculate the pH of a buffer containing 0.10 M NH? (pKb = 4.75) and 0.05 M NH?Cl.
9.55
9.25
8.95
10.00
First find pKa: 14.00 – 4.75 = 9.25. Then Henderson–Hasselbalch: pH = 9.25 + log(0.10/0.05) ? 9.25 + 0.301 = 9.55. This shows use of conjugate acid–base pair for a base buffer. Calculation steps.
For citric acid (pKa1 = 3.13, pKa2 = 4.76, pKa3 = 6.40), which species predominates at pH 2?
H?A (fully protonated form)
H?A?
HA²?
A³?
At pH 2, which is below the first pKa, all acidic protons remain bound, so the fully protonated form H?A dominates. As pH rises past each pKa, different deprotonated species appear. Polyprotic acid behavior.
What is the pH after adding 0.05 mol NaOH to 1 L of a buffer containing 0.10 mol HA and 0.10 mol A– (pKa 4.76)?
5.24
5.00
4.76
5.76
Adding 0.05 mol NaOH converts that amount of HA to A–: new [HA]=0.05 mol, [A–]=0.15 mol. Then pH = 4.76 + log(0.15/0.05) ? 4.76 + 0.477 = 5.24. This uses stoichiometric adjustment in Henderson–Hasselbalch. Stoichiometry in buffers.
Which statement about activity coefficients in buffer solutions is correct?
They generally decrease with increasing ionic strength
They are constant regardless of concentration
They increase pH directly
They only affect neutral species
As ionic strength rises, electrostatic interactions are screened and activity coefficients of ions decrease. This alters the effective concentration for equilibrium calculations. Ignoring this can lead to pH prediction errors. Activity coefficient theory.
For a diprotic acid H?A with pKa? and pKa?, which species predominates at pH between pKa? and pKa??
HA?
H?A
A²?
H?A?
Between pKa? and pKa?, the monoanionic form HA? is most abundant because the first proton has dissociated but the second has not. Understanding this helps in predicting buffer composition for polyprotic systems. Polyprotic acid speciation.
What is the pH at the half-equivalence point in the titration of a weak acid with a strong base?
Equal to the pKa of the acid
Neutral (pH 7)
pH = pKb of the conjugate base
Undefined
At the half-equivalence point half of the acid has been neutralized, so [HA] = [A–]. By Henderson–Hasselbalch, pH = pKa. This is a key feature of titration curves. Half-equivalence point.
How does increasing ionic strength affect the apparent pKa of an acid in buffer solution?
It decreases the apparent pKa slightly
It increases the apparent pKa slightly
It has no effect
It reverses the acid–base equilibrium
Higher ionic strength lowers activity coefficients of ions, altering the equilibrium constant expression and making the acid appear stronger, thus decreasing the apparent pKa. This effect is described by Debye–Hückel theory. Debye–Hückel details.
Why does buffer capacity vary with pH?
Because the ratio of conjugate base to acid changes with pH
Because buffer concentration is pH-dependent
Because water dissociation dominates at all pH values
Because only strong acids have capacity
Buffer capacity depends on both the absolute concentrations and the ratio [A–]/[HA]. As pH shifts away from pKa, one component becomes less abundant and capacity decreases. Maximum capacity occurs at pH = pKa. Buffer capacity explained.
Which buffer system exhibits the smallest temperature dependence of pH?
Phosphate buffer system
Tris buffer
Acetate buffer
Borate buffer
The phosphate buffer system has a small ?pKa/?T, making its pH change minimal with temperature fluctuations. Other buffers like Tris exhibit larger temperature coefficients, altering pH more upon heating or cooling. Phosphate buffer details.
When mixing volumes V? of acid HA at concentration C? and V? of conjugate base A– at concentration C?, which expression gives the final pH using Henderson–Hasselbalch?
pH = pKa + log((C?·V?)/(C?·V?))
pH = pKa + log((C?·V?)/(C?·V?))
pH = pKa + log((C? + C?)/(V? + V?))
pH = pKa + log((V?)/(V?))
Total moles of A– and HA are C?·V? and C?·V? respectively. The Henderson–Hasselbalch equation uses the ratio of these moles (or resulting concentrations). This general form handles unequal volumes. Volume mixing.
Which modification of the Henderson–Hasselbalch equation accounts for activity coefficients?
pH = pKa + log((?A– [A–])/(?HA [HA]))
pH = pKa + log([A–]/[HA]) + I
pH = pKa – log((?HA [HA])/(?A– [A–]))
pH = pKa + log([A–] · ?A–)
In non-ideal solutions, each species’ activity equals its concentration times its activity coefficient (?). The correct form replaces concentrations with activities (?·[ ]). This yields more accurate pH predictions at higher ionic strengths. Activity corrections.
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Study Outcomes

  1. Apply the Henderson - Hasselbalch Equation -

    Calculate the pH of buffer solutions by linking acid and conjugate base concentrations using the Henderson - Hasselbalch equation.

  2. Analyze Buffer Capacity -

    Determine how changes in concentration, volume, or added strong acids/bases affect a buffer's ability to resist pH shifts.

  3. Calculate pH After Dilution -

    Perform pH buffer practice problems involving dilution and mixing of buffer components to predict final pH values.

  4. Predict pH Changes on Acid/Base Addition -

    Evaluate real-world acid - base buffer questions by estimating pH response when titrating buffers with strong acids or bases.

  5. Differentiate Between Buffer Types -

    Classify common buffer systems (e.g., acetate, phosphate) and select appropriate buffers for specific pH ranges.

  6. Interpret Buffer Solution Quiz Results -

    Review detailed explanations to identify strengths and weaknesses in your buffer solution practice problem-solving skills.

Cheat Sheet

  1. Henderson - Hasselbalch Foundation -

    When tackling buffer solution practice problems, the Henderson - Hasselbalch equation (pH = pKa + log([A - ]/[HA])) is your go-to tool for converting concentrations of conjugate acid - base pairs into pH values. For example, mixing 0.1 M acetic acid (pKa 4.76) with 0.1 M sodium acetate yields pH 4.76. Remember the mnemonic "HA over A minus keeps pH in line."

  2. Buffer Capacity and Effectiveness -

    Buffer capacity measures how well a system resists pH changes when acids or bases are added, and it peaks when [A - ] ≈ [HA]. Higher concentrations of both components boost capacity, so a 0.2 M acetate buffer handles more base than a 0.05 M solution. Think "concentration counts" to recall that more solutes mean stronger defense against pH swings.

  3. Choosing Conjugate Acid - Base Pairs -

    Select a weak acid or base whose pKa is within ±1 unit of your target pH for optimal buffering in acid-base buffer questions. Common examples include acetic acid/acetate for pH 4.5 - 6.5 and phosphate buffers for pH 6.8 - 7.4 in biochemical assays. Matching pKa to desired pH simplifies buffer solution quiz scenarios and improves accuracy.

  4. Predicting pH Changes on Additions -

    In real-world scenarios, adding a known volume of strong acid or base shifts the ratio [A - ]/[HA], and you recalculate pH with Henderson - Hasselbalch. For instance, adding 0.005 mol HCl to a 0.1 M acetate buffer reduces [A - ] and gives a new pH via updated log ratio. This stepwise calculation is a staple of buffer practice problems and sharpens your problem-solving flow.

  5. Practical Buffer Preparation -

    To prepare a buffer solution for a specific pH, calculate required amounts of acid and conjugate base using the Henderson - Hasselbalch equation, then dilute to volume. Always verify pH with a calibrated meter and adjust with tiny additions of acid or base if necessary. This hands-on approach cements theory from buffer solution practice problems into lab-ready skills.

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