College Chemistry Quiz: Properties Of Buffers
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Properties Of BuffersQuestion 1 of 20

Which statement best explains why a mixture of HClHCl and CH3COONaCH_3COONa cannot function as an effective buffer system?

The concentrations of the two components are never equal in such mixtures
Hydrochloric acid completely neutralizes the acetate ion to form acetic acid
The pKa of the HCl/Cl⁻ system is too low for most biological applications
Sodium acetate is not sufficiently soluble in aqueous solutions
The ionic strength becomes too high when both compounds are dissolved
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College Chemistry Quiz

College Chemistry Quiz: Properties Of Buffers

Practice Properties Of Buffers in College Chemistry with focused quiz questions that help you check what you know, review explanations, and build confidence with test-style prompts.

What this quiz covers

This quiz focuses on Properties Of Buffers, giving you a quick way to practice the rules, question types, and explanations that matter most for College Chemistry.

How to use this quiz

Try each quiz question before looking at the correct answer. Use the explanations to review missed ideas, then come back to similar questions until the pattern feels familiar.

All questions

Question 1

Which statement best explains why a mixture of HClHCl and CH3COONaCH_3COONa cannot function as an effective buffer system?

  1. The concentrations of the two components are never equal in such mixtures
  2. Hydrochloric acid completely neutralizes the acetate ion to form acetic acid (correct answer)
  3. The pKa of the HCl/Cl⁻ system is too low for most biological applications
  4. Sodium acetate is not sufficiently soluble in aqueous solutions
  5. The ionic strength becomes too high when both compounds are dissolved
Explanation: When you encounter buffer questions, remember that effective buffers require a weak acid and its conjugate base (or weak base and its conjugate acid) to resist pH changes through equilibrium shifts. Let's examine what happens when you mix HClHCl and CH3COONaCH_3COONa. Since HClHCl is a strong acid, it completely dissociates and will react with the acetate ion (CH3COO−CH_3COO^-) from sodium acetate: HCl+CH3COO−→CH3COOH+Cl−HCl + CH_3COO^- \rightarrow CH_3COOH + Cl^-. This reaction goes essentially to completion, converting all the acetate ions to acetic acid molecules. You end up with acetic acid and chloride ions, not the weak acid/conjugate base pair needed for buffering. The strong acid has completely neutralized the base component, eliminating the buffer's ability to resist pH changes. Choice A is incorrect because buffer effectiveness doesn't require equal concentrations—buffers work best when the ratio is close to 1:1, but they function across a range of ratios. Choice C misses the point entirely; while HClHCl's extremely low pKa does make it unsuitable for biological buffers, this isn't why the HCl/CH3COONaHCl/CH_3COONa mixture fails—it fails because no buffer system forms at all. Choice D is wrong because sodium acetate is actually quite soluble in water. Study tip: When evaluating potential buffer systems, always check if you have a strong acid or base present. Strong acids/bases will completely neutralize their conjugate pairs, destroying the equilibrium necessary for buffering. Only weak acid/conjugate base pairs maintain the equilibrium needed for effective buffering.

Question 2

A buffer solution contains 0.15 M H2PO4−H_2PO_4^- and 0.25 M HPO42−HPO_4^{2-}. After adding a small amount of strong base, the buffer pH increases from 7.42 to 7.48. What can be concluded about the buffer's response?

  1. The buffer has reached its maximum capacity and will fail with further base addition
  2. The buffer is functioning normally with the dihydrogen phosphate consuming the added hydroxide (correct answer)
  3. The ratio of phosphate species has shifted to favor the dihydrogen phosphate form
  4. The buffer system has been completely converted to the hydrogen phosphate form
  5. The phosphate buffer is operating outside its effective pH range
Explanation: When you encounter buffer problems, focus on understanding how the conjugate acid-base pair responds to added acids or bases. Buffers work by having one species neutralize added base (the weak acid) and another neutralize added acid (the weak base). In this phosphate buffer system, H2PO4−H_2PO_4^- acts as the weak acid and HPO42−HPO_4^{2-} acts as the weak base. When strong base is added, the H2PO4−H_2PO_4^- donates a proton to neutralize the OH−OH^- ions: H2PO4−+OH−→HPO42−+H2OH_2PO_4^- + OH^- \rightarrow HPO_4^{2-} + H_2O. The small pH increase from 7.42 to 7.48 (only 0.06 units) indicates the buffer is effectively resisting dramatic pH change, which is exactly what a functioning buffer should do. Answer A is wrong because a 0.06 pH unit change shows the buffer is working well, not failing. A buffer at maximum capacity would show a much larger pH jump. Answer C incorrectly describes the species shift - when base is added, H2PO4−H_2PO_4^- is consumed and converted to HPO42−HPO_4^{2-}, so the ratio shifts to favor the hydrogen phosphate form, not the dihydrogen phosphate. Answer D is wrong because complete conversion would result in a massive pH change and buffer failure, not the small increase observed. Remember that effective buffers show small pH changes when small amounts of strong acid or base are added. Large pH changes indicate the buffer is overwhelmed or depleted.

Question 3

A buffer solution has initial pH 8.3. After adding 0.01 mol of HClHCl to 1 L of buffer, the pH becomes 8.1. If instead 0.01 mol of NaOHNaOH had been added, what would be the approximate final pH?

  1. 8.4
  2. 8.5 (correct answer)
  3. 8.7
  4. 9.1
  5. 9.5
Explanation: Buffer problems test your understanding of how these solutions resist pH changes through equilibrium shifts. When you see a buffer question with symmetric additions of acid and base, look for the relationship between the magnitude of pH changes. The key insight is that a good buffer responds symmetrically to equal molar additions of strong acid and base. When 0.01 mol HClHCl was added, the pH dropped from 8.3 to 8.1 (a decrease of 0.2 units). The HClHCl converts some of the buffer's weak base component to its conjugate weak acid, shifting the equilibrium. If 0.01 mol NaOHNaOH were added instead, it would convert some of the buffer's weak acid component to its conjugate base, causing an equal but opposite shift. Since the buffer lost 0.2 pH units with acid addition, it should gain approximately 0.2 pH units with base addition: 8.3 + 0.2 = 8.5. Answer A (8.4) represents too small a change, suggesting the buffer is more resistant to base than acid, which isn't typical for equal molar additions. Answer C (8.7) shows twice the expected change, as if the buffer responds more dramatically to base. Answer D (9.1) represents an unrealistically large jump that would indicate buffer failure—the system would no longer be buffering effectively. Strategy tip: For buffer calculations involving symmetric acid/base additions, expect symmetric pH changes around the initial pH. If adding acid drops the pH by X units, adding the same amount of base should raise it by approximately X units.

Question 4

Which factor would most significantly reduce the effectiveness of an acetate buffer system in maintaining constant pH?

  1. Diluting the buffer solution by adding distilled water (correct answer)
  2. Increasing the temperature of the buffer solution by 10°C
  3. Adding a large excess of sodium acetate to the buffer
  4. Introducing a small amount of sodium chloride to the buffer
  5. Allowing the buffer solution to remain open to air for several hours
Explanation: Buffer effectiveness depends on maintaining adequate concentrations of both the weak acid and its conjugate base to resist pH changes. When evaluating threats to buffer capacity, consider what happens to these essential components. Diluting a buffer by adding distilled water (A) dramatically reduces the concentrations of both acetic acid and acetate ions proportionally. While the pH ratio remains initially unchanged according to the Henderson-Hasselbalb equation (pH=pKa+log⁡[A−][HA]pH = pK_a + \log\frac{[A^-]}{[HA]}), the buffer's capacity to neutralize added acids or bases plummets. With fewer buffer molecules available per unit volume, even small additions of strong acid or base will overwhelm the system and cause significant pH shifts. Choice B is incorrect because moderate temperature increases have minimal impact on buffer effectiveness. While pKapK_a values change slightly with temperature, a 10°C increase won't substantially compromise the buffer's ability to maintain pH. Choice C represents a common misconception. Adding excess sodium acetate shifts the buffer ratio but doesn't eliminate effectiveness. The system still contains both components needed for buffering, though it becomes better at neutralizing acids than bases. Choice D is wrong because sodium chloride is a neutral salt that doesn't react with buffer components. It may slightly affect ionic strength, but this has negligible impact on buffer capacity. Remember: buffer capacity depends on concentration, not just ratio. Dilution is the enemy of effective buffering because it reduces the absolute number of molecules available to consume added H⁺ or OH⁻ ions.

Question 5

Which combination would create the buffer system with the highest pH?

  1. 0.10 M CH3COOHCH_3COOH and 0.10 M CH3COONaCH_3COONa (pKa=4.74pK_a = 4.74)
  2. 0.10 M H2PO4−H_2PO_4^- and 0.10 M HPO42−HPO_4^{2-} (pKa=7.20pK_a = 7.20)
  3. 0.10 M HClOHClO and 0.10 M NaClONaClO (pKa=7.54pK_a = 7.54)
  4. 0.10 M NH4ClNH_4Cl and 0.10 M NH3NH_3 (pKb=4.74pK_b = 4.74) (correct answer)
  5. 0.10 M HFHF and 0.10 M NaFNaF (pKa=3.17pK_a = 3.17)
Explanation: When comparing buffer systems to find which has the highest pH, you need to use the Henderson-Hasselbalch equation and consider the relationship between pKₐ and pKᵦ values. For acid-base conjugate pairs with equal concentrations, the Henderson-Hasselbalch equation simplifies to pH = pKₐ when the ratio of base to acid is 1:1. Let's calculate the pH for each system: Option A gives pH = 4.74 (directly from the pKₐ). Option B gives pH = 7.20 (directly from the pKₐ). Option C gives pH = 7.54 (directly from the pKₐ). For option D, you have a base buffer system (NH₃/NH₄⁺), so you must first convert pKᵦ to pKₐ using the relationship pKₐ + pKᵦ = 14. Therefore, pKₐ = 14 - 4.74 = 9.26, giving pH = 9.26. Option D is correct because it produces the highest pH at 9.26. Option A is wrong because the acetate system has a much lower pKₐ (4.74), creating an acidic buffer. Option B is wrong because the phosphate system's pKₐ of 7.20 produces a neutral-to-slightly basic buffer, but not as high as option D. Option C is wrong because the hypochlorite system's pKₐ of 7.54 creates a basic buffer, but still lower than option D. Study tip: Remember that for equal-concentration buffers, pH equals pKₐ. When given pKᵦ values, always convert to pKₐ first using pKₐ + pKᵦ = 14. The highest pKₐ value will give the highest pH.

Question 6

Consider a buffer made from 0.20 M CH3COOHCH_3COOH and 0.30 M CH3COONaCH_3COONa. If this buffer is compared to another buffer with 0.40 M CH3COOHCH_3COOH and 0.60 M CH3COONaCH_3COONa, which statement is correct?

  1. Both buffers have identical pH and identical buffer capacity
  2. Both buffers have identical pH but the second has greater buffer capacity (correct answer)
  3. The second buffer has higher pH and greater buffer capacity
  4. The first buffer has higher pH but lower buffer capacity
  5. The buffers have different pH values but identical buffer capacity
Explanation: When you encounter buffer problems, you need to analyze two key properties: pH (determined by the ratio of conjugate base to weak acid) and buffer capacity (determined by the absolute concentrations of both components). For pH calculation, use the Henderson-Hasselbalch equation: pH=pKa+log⁡[A−][HA]pH = pK_a + \log\frac{[A^-]}{[HA]}. In the first buffer, the ratio is [CH3COO−][CH3COOH]=0.300.20=1.5\frac{[CH_3COO^-]}{[CH_3COOH]} = \frac{0.30}{0.20} = 1.5. In the second buffer, this ratio is 0.600.40=1.5\frac{0.60}{0.40} = 1.5. Since both ratios are identical, both buffers have the same pH. Buffer capacity measures how much acid or base a buffer can neutralize before significant pH change occurs. This depends on the absolute amounts of both the weak acid and conjugate base present. The second buffer contains twice the concentration of both components (0.40 M and 0.60 M vs. 0.20 M and 0.30 M), giving it twice the buffer capacity. Looking at the answer choices: Choice A incorrectly states both properties are identical—while pH is the same, buffer capacities differ. Choice C wrongly claims the second buffer has higher pH, but identical ratios mean identical pH values. Choice D incorrectly suggests the first buffer has higher pH and also mischaracterizes the capacity relationship. Choice B correctly identifies that both buffers share the same pH while the second has greater buffer capacity. Remember: buffer pH depends only on the concentration ratio, while buffer capacity depends on absolute concentrations. Doubling both components doubles capacity without changing pH.

Question 7

Which statement best explains why the Henderson-Hasselbalch equation becomes less accurate for buffer calculations when the buffer components have very different concentrations (e.g., 100:1 ratio)?

  1. The equation assumes that water autoionization can be ignored
  2. Activity coefficients deviate significantly from unity at extreme concentration ratios
  3. The weak acid approximation breaks down when one component dominates (correct answer)
  4. Temperature effects become more pronounced at unequal concentrations
  5. The buffer no longer maintains a constant ionic strength
Explanation: When you encounter Henderson-Hasselbalch equation problems, focus on the underlying assumptions that make this equation work. The Henderson-Hasselbalch equation, pH=pKa+log⁡[A−][HA]pH = pK_a + \log\frac{[A^-]}{[HA]}, relies on several key approximations to simplify buffer calculations. The critical issue with extreme concentration ratios (like 100:1) is that the weak acid approximation fails. This approximation assumes that the initial concentrations of the buffer components remain essentially unchanged after equilibrium is established. When one component vastly outweighs the other, the small amount of acid-base reaction that occurs can represent a significant percentage change in the minor component's concentration. For example, if you have 0.01 M of the minor component and even a small equilibrium shift changes it by 0.001 M, that's a 10% change—far from the "negligible change" assumption. Option A is incorrect because water autoionization is typically negligible in buffer systems regardless of concentration ratios, except at very low buffer concentrations. Option B misidentifies the problem—while activity coefficients can deviate from unity, this isn't the primary issue with unequal concentration ratios in typical buffer calculations. Option D is wrong because temperature effects aren't specifically related to concentration ratios; temperature affects equilibrium constants equally regardless of component ratios. Remember this pattern: Henderson-Hasselbalch works best when buffer component concentrations are within about a 10:1 ratio. When ratios become more extreme, you need to solve the full equilibrium expression rather than relying on the simplified equation.

Question 8

Two buffer solutions are prepared: Buffer X contains 0.1 M CH3COOHCH_3COOH and 0.1 M CH3COONaCH_3COONa, while Buffer Y contains 0.1 M NH3NH_3 and 0.1 M NH4ClNH_4Cl. Both are at 25°C. If equal volumes of these buffers are mixed together, what is the most likely outcome?

  1. A new buffer system forms with intermediate pH between the two original buffers (correct answer)
  2. The mixture loses all buffering capacity due to neutralization reactions
  3. The pH equals the average of the two original buffer pH values
  4. The stronger buffer dominates and maintains its original pH
  5. Multiple buffer systems coexist independently in the same solution
Explanation: When you encounter questions about mixing different buffer systems, think about how each buffer component will interact and whether new equilibria can establish. When these two buffers mix, you're combining an acetate buffer (weak acid CH3COOHCH_3COOH with its conjugate base CH3COO−CH_3COO^-) and an ammonia buffer (weak base NH3NH_3 with its conjugate acid NH4+NH_4^+). The key insight is that while some acid-base reactions occur between components, both buffer systems remain partially intact and continue functioning. The acetic acid will partially neutralize ammonia (CH3COOH+NH3→CH3COO−+NH4+CH_3COOH + NH_3 \rightarrow CH_3COO^- + NH_4^+), but since the concentrations are equal and both are weak, the reaction doesn't go to completion. You end up with a mixture containing all four species: CH3COOHCH_3COOH, CH3COO−CH_3COO^-, NH3NH_3, and NH4+NH_4^+. This creates a multi-component buffer system that can resist pH changes through multiple equilibria, with a pH somewhere between the original buffer pH values. Answer B is wrong because the neutralization isn't complete—weak acid/weak base reactions reach equilibrium, leaving active buffer components. Answer C incorrectly assumes simple averaging; buffer pH depends on equilibrium concentrations, not arithmetic means of starting pH values. Answer D is incorrect because neither buffer is inherently "stronger"—they have equal concentrations and both weak components remain active. Remember: when mixing buffers with weak components, partial reactions create hybrid systems. Complete neutralization only occurs with strong acid/strong base combinations or extreme concentration differences.

Question 9

A buffer solution is prepared by dissolving 0.15 mol NaH2PO4NaH_2PO_4 and 0.25 mol Na2HPO4Na_2HPO_4 in water to make 1.0 L of solution. What is the buffer's capacity for neutralizing added base? (Assume buffer failure occurs when 90% of the limiting component is consumed)

  1. 0.135 mol OH−OH^- (correct answer)
  2. 0.150 mol OH−OH^-
  3. 0.225 mol OH−OH^-
  4. 0.250 mol OH−OH^-
  5. 0.375 mol OH−OH^-
Explanation: When you encounter buffer capacity questions, you need to identify which component limits the buffer's ability to neutralize added acid or base. Buffer capacity depends on which component gets depleted first. This buffer contains two species: NaH2PO4NaH_2PO_4 (0.15 mol) acts as the weak acid component, and Na2HPO4Na_2HPO_4 (0.25 mol) acts as the conjugate base component. When base (OH−OH^-) is added, it reacts with the weak acid component according to: H2PO4−+OH−→HPO42−+H2OH_2PO_4^- + OH^- \rightarrow HPO_4^{2-} + H_2O Since you have less NaH2PO4NaH_2PO_4 (0.15 mol) than Na2HPO4Na_2HPO_4 (0.25 mol), the H2PO4−H_2PO_4^- will be consumed first, making it the limiting component for base neutralization. Buffer failure occurs when 90% of this limiting component is consumed, so the buffer can neutralize: 0.15 mol × 0.90 = 0.135 mol of OH−OH^-. Looking at the wrong answers: B (0.150 mol) represents 100% consumption of NaH2PO4NaH_2PO_4, ignoring the 90% buffer failure criterion. C (0.225 mol) incorrectly applies 90% to Na2HPO4Na_2HPO_4 (0.25 × 0.90), but this component doesn't limit base neutralization. D (0.250 mol) represents 100% of the Na2HPO4Na_2HPO_4, again using the wrong component. The correct answer is A (0.135 mol OH−OH^-). Study tip: Always identify the limiting component first by determining which buffer component will be consumed by the added species, then apply the buffer failure percentage to that limiting amount only.

Question 10

Which of the following combinations would produce the most effective buffer system for maintaining a pH of 7.4?

  1. HClHCl and NaClNaCl in a 1:1 molar ratio
  2. H2PO4−H_2PO_4^- and HPO42−HPO_4^{2-} with pKa=7.2pK_a = 7.2 (correct answer)
  3. CH3COOHCH_3COOH and CH3COONaCH_3COONa with pKa=4.8pK_a = 4.8
  4. NH3NH_3 and NH4ClNH_4Cl with pKb=4.7pK_b = 4.7
  5. HNO3HNO_3 and KNO3KNO_3 in a 2:1 molar ratio
Explanation: When evaluating buffer systems, you need to understand that the most effective buffer maintains a pH close to its pKapK_a value. The Henderson-Hasselbalch equation shows this relationship: pH=pKa+log⁡[A−][HA]pH = pK_a + \log\frac{[A^-]}{[HA]}. A buffer works best when the pH you want to maintain is within about 1 pH unit of the pKapK_a. Option B is correct because the H2PO4−/HPO42−H_2PO_4^-/HPO_4^{2-} system has a pKapK_a of 7.2, which is very close to the target pH of 7.4. This small difference (only 0.2 units) means the buffer will effectively resist pH changes around 7.4. This phosphate buffer system is actually what your body uses to maintain blood pH. Option A is wrong because HClHCl and NaClNaCl don't form a buffer at all - you need a weak acid and its conjugate base, but HClHCl is a strong acid that completely dissociates. Option C fails because acetic acid has a pKapK_a of 4.8, which is 2.6 units away from 7.4. This system would be effective around pH 4.8, not 7.4. Option D is incorrect because while NH3/NH4+NH_3/NH_4^+ can form a buffer, you need to convert the given pKbpK_b to pKapK_a using pKa+pKb=14pK_a + pK_b = 14. This gives pKa=9.3pK_a = 9.3, which is too far from 7.4. Study tip: Always check that the pKapK_a is within 1 pH unit of your target pH. If given pKbpK_b, convert it first using pKa=14−pKbpK_a = 14 - pK_b.

Question 11

What is the buffer capacity most directly related to in an aqueous buffer system?

  1. The difference between pKa and the solution pH
  2. The total concentration of buffer components present (correct answer)
  3. The ratio of weak acid to conjugate base concentrations
  4. The temperature at which the buffer is maintained
  5. The ionic strength of the solution containing the buffer
Explanation: Buffer capacity measures how much acid or base a buffer system can absorb before undergoing a significant pH change. Think of it as the buffer's "strength" or resistance to pH shifts when disturbed. Buffer capacity depends most directly on the total concentration of buffer components present (answer B). A buffer with higher concentrations of both the weak acid and its conjugate base can neutralize larger amounts of added acid or base. For example, a buffer containing 1.0 M acetic acid and 1.0 M acetate has much greater capacity than one with 0.1 M of each component, even though both have the same pH. Option A is incorrect because the difference between pKa and solution pH affects buffer efficiency (how well it resists pH change at a given addition), but not the total capacity. A buffer works best when pH equals pKa, but capacity still depends on concentration. Option C confuses buffer capacity with buffer pH. The ratio of weak acid to conjugate base determines the buffer's pH through the Henderson-Hasselbalck equation, but equal concentrations (1:1 ratio) don't guarantee high capacity if both concentrations are low. Option D is wrong because temperature has minimal direct effect on buffer capacity. While temperature can slightly affect equilibrium constants, it doesn't change the fundamental relationship between concentration and capacity. Remember: buffer capacity is about quantity, not quality. More concentrated buffers can handle more disturbance before failing, regardless of their pH or acid-to-base ratio.

Question 12

A research laboratory uses a HEPESHEPES buffer (pKa=7.5pK_a = 7.5) to maintain pH 7.4 in cell culture media. If the buffer concentration is 0.025 M total buffer components, what is the approximate ratio of the base form to acid form?

  1. 0.56:1
  2. 0.79:1 (correct answer)
  3. 1.0:1
  4. 1.26:1
  5. 1.78:1
Explanation: When you encounter buffer problems, you're dealing with the Henderson-Hasselbalch equation, which relates pH, pKa, and the ratio of conjugate base to acid forms. This is fundamental to understanding how biological systems maintain stable pH. To find the ratio of base form to acid form in this HEPES buffer, use the Henderson-Hasselbalch equation: pH=pKa+log⁡([A−][HA])pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right) Given pH = 7.4 and pKa = 7.5, substitute these values: 7.4=7.5+log⁡([A−][HA])7.4 = 7.5 + \log\left(\frac{[A^-]}{[HA]}\right) Solving for the ratio: 7.4−7.5=log⁡([A−][HA])7.4 - 7.5 = \log\left(\frac{[A^-]}{[HA]}\right) −0.1=log⁡([A−][HA])-0.1 = \log\left(\frac{[A^-]}{[HA]}\right) [A−][HA]=10−0.1=0.79\frac{[A^-]}{[HA]} = 10^{-0.1} = 0.79 This confirms answer B) 0.79:1 is correct. Answer A) 0.56:1 would result from calculation errors or using the wrong pH value. Answer C) 1.0:1 represents the ratio when pH equals pKa, which isn't the case here since pH < pKa. Answer D) 1.26:1 would be the reciprocal calculation error or using +0.1 instead of -0.1 in the logarithm. Remember: when pH < pKa, the acid form predominates, so the base-to-acid ratio will be less than 1. When pH > pKa, the base form predominates. The closer the pH to the pKa, the more effective the buffer system.

Question 13

A protein biochemist prepares a TrisTris buffer (tris(hydroxymethyl)aminomethane, pKa=8.1pK_a = 8.1) for enzyme studies. At pH 7.6, what percentage of the TrisTris molecules exist in the protonated (acid) form?

  1. 24%
  2. 32%
  3. 54%
  4. 68%
  5. 76% (correct answer)
Explanation: This question tests your understanding of the Henderson-Hasselbalch equation and buffer chemistry. When you see a buffer problem asking about the ratio of acid to base forms at a specific pH, you need to determine which form predominates and calculate the exact percentage. Using the Henderson-Hasselbalch equation: pH=pKa+log⁡[A−][HA]pH = pK_a + \log\frac{[A^-]}{[HA]} Substituting the given values: 7.6=8.1+log⁡[base][acid]7.6 = 8.1 + \log\frac{[base]}{[acid]} Solving: log⁡[base][acid]=7.6−8.1=−0.5\log\frac{[base]}{[acid]} = 7.6 - 8.1 = -0.5 Therefore: [base][acid]=10−0.5=0.316\frac{[base]}{[acid]} = 10^{-0.5} = 0.316 This means for every 1 part acid, there are 0.316 parts base. The fraction of protonated form is: 11+0.316=11.316=0.76=76%\frac{1}{1 + 0.316} = \frac{1}{1.316} = 0.76 = 76\% Since the correct answer E isn't shown but should be 76%, let's examine why the given options are wrong: A) 24% incorrectly calculates the base form percentage instead of the acid form. B) 32% appears to use an incorrect logarithmic conversion. C) 54% might result from calculation errors in the Henderson-Hasselbalch equation. D) 68% is close but represents a systematic error, possibly from rounding mistakes or formula misapplication. Study tip: Remember that when pH < pKa, the acid form predominates. Always double-check whether you're calculating the acid or base percentage, as this is a common source of errors on buffer problems.

Question 14

When a buffer solution is diluted with distilled water, which property changes most significantly?

  1. The pH of the buffer solution
  2. The ratio of conjugate acid to base concentrations
  3. The buffer capacity per unit volume (correct answer)
  4. The pKa value of the weak acid component
  5. The temperature coefficient of the buffer system
Explanation: When you encounter questions about buffer dilution, focus on understanding what happens to each buffer property when you add water to the system. Buffer capacity per unit volume changes most dramatically during dilution. Buffer capacity depends on the absolute concentrations of the conjugate acid-base pair - when you dilute a buffer, you're decreasing these concentrations proportionally. Since buffer capacity is typically expressed per unit volume, halving the concentrations by doubling the volume cuts the buffer capacity in half. This represents a significant, measurable change in the buffer's ability to resist pH changes. Let's examine why the other options don't change significantly: Choice A is incorrect because pH remains essentially constant during dilution. The Henderson-Hasselbalch equation shows that pH depends on the ratio of conjugate base to acid concentrations, not their absolute values. Choice B is wrong for the same reason - dilution decreases both the acid and base concentrations proportionally, so their ratio stays the same. Choice D is incorrect because pKapK_a is an intrinsic property of the weak acid that doesn't depend on concentration or dilution. The key insight is distinguishing between intensive properties (like pH, concentration ratios, and pKapK_a) that don't change with dilution, and extensive properties (like buffer capacity) that do change. Remember that buffer capacity is directly proportional to the concentrations of the buffer components - dilute the buffer, and you proportionally reduce its capacity to neutralize added acid or base.

Question 15

A carbonic acid buffer system (H2CO3/HCO3−H_2CO_3/HCO_3^-) maintains blood pH at 7.40. If the pKapK_a is 6.37 and [H2CO3]=1.2×10−3[H_2CO_3] = 1.2 \times 10^{-3} M, what is the concentration of bicarbonate ion?

  1. 1.1×10−41.1 \times 10^{-4} M
  2. 2.4×10−32.4 \times 10^{-3} M
  3. 1.3×10−21.3 \times 10^{-2} M (correct answer)
  4. 2.6×10−22.6 \times 10^{-2} M
  5. 5.2×10−25.2 \times 10^{-2} M
Explanation: Buffer problems require you to connect pH, pKa, and the ratio of conjugate acid-base pairs using the Henderson-Hasselbalch equation: pH=pKa+log⁡[A−][HA]pH = pK_a + \log\frac{[A^-]}{[HA]}. For this carbonic acid system, you need to find the bicarbonate concentration that maintains blood's precise pH of 7.40. Substituting the given values: 7.40=6.37+log⁡[HCO3−][H2CO3]7.40 = 6.37 + \log\frac{[HCO_3^-]}{[H_2CO_3]} Rearranging: 7.40−6.37=log⁡[HCO3−]1.2×10−37.40 - 6.37 = \log\frac{[HCO_3^-]}{1.2 \times 10^{-3}} 1.03=log⁡[HCO3−]1.2×10−31.03 = \log\frac{[HCO_3^-]}{1.2 \times 10^{-3}} Taking the antilog: 101.03=[HCO3−]1.2×10−310^{1.03} = \frac{[HCO_3^-]}{1.2 \times 10^{-3}} Since 101.03=10.710^{1.03} = 10.7: [HCO3−]=10.7×1.2×10−3=1.3×10−2[HCO_3^-] = 10.7 \times 1.2 \times 10^{-3} = 1.3 \times 10^{-2} M This confirms answer C is correct. Answer A (1.1×10−41.1 \times 10^{-4} M) results from incorrectly using 10−1.0310^{-1.03} instead of 101.0310^{1.03} - a sign error in the Henderson-Hasselbalch rearrangement. Answer B (2.4×10−32.4 \times 10^{-3} M) comes from using 100.310^{0.3} instead of 101.0310^{1.03}, likely from miscalculating 7.40−6.37=0.37.40 - 6.37 = 0.3. Answer D (2.6×10−22.6 \times 10^{-2} M) suggests doubling the correct answer, possibly from calculation errors. Remember: blood buffers maintain a bicarbonate-to-carbonic acid ratio of about 10:1, which should seem reasonable when checking your final answer against physiological norms.

Question 16

A buffer solution contains 0.10 M CH3COOHCH_3COOH and 0.10 M CH3COONaCH_3COONa. When 0.020 mol of HClHCl is added to 1.0 L of this buffer, what is the change in pH? (KaK_a for acetic acid = 1.8×10−51.8 \times 10^{-5})

  1. Decreases by 0.18 (correct answer)
  2. Decreases by 0.35
  3. Decreases by 0.52
  4. Increases by 0.18
  5. No change in pH
Explanation: When you encounter a buffer problem involving the addition of acid or base, you're dealing with the Henderson-Hasselbalch equation and how the buffer components react to neutralize the added species. First, calculate the initial pH using pH=pKa+log⁡[A−][HA]pH = pK_a + \log\frac{[A^-]}{[HA]}. Since pKa=−log⁡(1.8×10−5)=4.74pK_a = -\log(1.8 \times 10^{-5}) = 4.74 and both concentrations are 0.10 M, the initial pH is 4.74+log⁡(1)=4.744.74 + \log(1) = 4.74. When 0.020 mol HCl is added, it reacts completely with the acetate ion: CH3COO−+HCl→CH3COOH+Cl−CH_3COO^- + HCl \rightarrow CH_3COOH + Cl^-. This consumes 0.020 mol of acetate and produces 0.020 mol of additional acetic acid. New concentrations become: [CH3COOH]=0.10+0.020=0.12[CH_3COOH] = 0.10 + 0.020 = 0.12 M and [CH3COO−]=0.10−0.020=0.08[CH_3COO^-] = 0.10 - 0.020 = 0.08 M. The new pH is 4.74+log⁡0.080.12=4.74+log⁡(0.667)=4.74−0.18=4.564.74 + \log\frac{0.08}{0.12} = 4.74 + \log(0.667) = 4.74 - 0.18 = 4.56. The change in pH is 4.56−4.74=−0.184.56 - 4.74 = -0.18, confirming answer A. Answer B (-0.35) likely results from calculation errors in the logarithm. Answer C (-0.52) suggests confusion about stoichiometry or using incorrect concentrations. Answer D (+0.18) represents the correct magnitude but wrong direction—a common error when students forget that adding acid decreases pH. Remember: in buffer calculations, always account for the complete reaction between added strong acid/base and the appropriate buffer component before applying Henderson-Hasselbalch.

Question 17

A student measures the pH of several buffer solutions and finds: Buffer A (pKa=4.2pK_a = 4.2) has pH 4.8, Buffer B (pKa=7.1pK_a = 7.1) has pH 7.1, and Buffer C (pKa=9.3pK_a = 9.3) has pH 8.7. Which buffer would be most resistant to pH change upon addition of a small amount of strong base?

  1. Buffer A, because it has the lowest pKa value
  2. Buffer B, because its pH equals its pKa value (correct answer)
  3. Buffer C, because it has the highest pH value
  4. All buffers would show equal resistance to pH change
  5. Cannot be determined without knowing the concentrations of buffer components
Explanation: Buffer resistance to pH change depends on how close the buffer's pH is to its pKapK_a value. A buffer works most effectively when the concentrations of its weak acid and conjugate base are roughly equal, which occurs when pH = pKapK_a. This relationship comes from the Henderson-Hasselbalch equation: pH=pKa+log⁡[A−][HA]pH = pK_a + \log\frac{[A^-]}{[HA]}. When pH equals pKapK_a, the log term equals zero, meaning the acid and base forms are present in equal concentrations. Buffer B is most resistant to pH change because its pH (7.1) exactly equals its pKapK_a (7.1). This means it has optimal buffering capacity with equal amounts of weak acid and conjugate base available to neutralize added strong base. Choice A is incorrect because having the lowest pKapK_a doesn't determine buffering effectiveness. Buffer A's pH (4.8) is 0.6 units above its pKapK_a (4.2), meaning it has more conjugate base than weak acid, making it less effective against added base. Choice C is wrong because the absolute pH value doesn't determine buffer resistance. Buffer C's pH (8.7) is 0.6 units below its pKapK_a (9.3), indicating excess weak acid but insufficient conjugate base to effectively neutralize added strong base. Choice D is incorrect because buffers operating at different distances from their pKapK_a values have different capacities. Remember: A buffer's effectiveness peaks when pH = pKapK_a. The further the pH drifts from the pKapK_a, the less resistant the buffer becomes to pH changes.

Question 18

A buffer solution contains carbonic acid and bicarbonate ion. When carbon dioxide gas is bubbled through the solution, what happens to the buffer's pH?

  1. pH increases because CO₂ acts as a Lewis base
  2. pH decreases because CO₂ forms additional H₂CO₃ (correct answer)
  3. pH remains constant because CO₂ is already part of the buffer system
  4. pH increases because CO₂ displaces H⁺ ions from solution
  5. pH becomes undefined due to gas-liquid equilibrium effects
Explanation: When you encounter buffer questions involving gas addition, focus on how the gas interacts chemically with the existing equilibrium system. This carbonic acid/bicarbonate buffer operates through the equilibrium: H2CO3⇌H++HCO3−\text{H}_2\text{CO}_3 \rightleftharpoons \text{H}^+ + \text{HCO}_3^- When CO₂ gas bubbles through the solution, it doesn't just sit there—it reacts with water to form more carbonic acid: CO2+H2O⇌H2CO3\text{CO}_2 + \text{H}_2\text{O} \rightleftharpoons \text{H}_2\text{CO}_3 This additional H₂CO₃ then dissociates to release more H⁺ ions, shifting the buffer equilibrium toward the acidic side and lowering the pH. Answer B correctly identifies this process. Answer A is wrong because CO₂ acts as a Lewis acid (electron pair acceptor), not a Lewis base, when it forms bonds with water. Answer C misses the key point—while CO₂ is related to the buffer components, adding more of it disrupts the established equilibrium by increasing the concentration of the weak acid component. Answer D incorrectly suggests CO₂ removes H⁺ ions, when it actually generates them through carbonic acid formation. Remember that buffers resist pH change but aren't immune to it. When you add a component that can form more of either the weak acid or weak base in the buffer pair, you'll shift the equilibrium and change the pH. Always trace the chemical pathway: what does the added substance become, and how does that affect H⁺ concentration?

Question 19

In preparing a formic acid buffer (HCOOH/HCOO−HCOOH/HCOO^-, pKa=3.75pK_a = 3.75), a chemist wants to achieve maximum buffer capacity. Which molar ratio of HCOOHHCOOH to HCOONaHCOONa should be used?

  1. 10:1
  2. 3:1
  3. 1:1 (correct answer)
  4. 1:3
  5. 1:10
Explanation: Buffer capacity is maximized when a buffer can resist pH changes in both directions equally well. This occurs when the buffer has equal concentrations of the weak acid and its conjugate base, making the molar ratio 1:1. Using the Henderson-Hasselbalch equation: pH=pKa+log⁡[A−][HA]pH = pK_a + \log\frac{[A^-]}{[HA]}. When [HCOO−]=[HCOOH][HCOO^-] = [HCOOH], the ratio equals 1, so log⁡(1)=0\log(1) = 0, and pH=pKa=3.75pH = pK_a = 3.75. At this point, the buffer is equally effective at neutralizing added acid or base because there are equal amounts of both buffering species available. Choice A (10:1) creates a buffer heavily weighted toward the acid form, making it effective against added base but poor against added acid. Choice B (3:1) still favors the acid form significantly, reducing overall buffer capacity. Choice D (1:3) favors the conjugate base form, making it good against added acid but weak against added base. The key insight is that maximum buffer capacity occurs at the buffer's pKapK_a, where pH=pKapH = pK_a. This happens when the weak acid and conjugate base concentrations are equal. Moving away from this 1:1 ratio in either direction reduces the buffer's ability to handle additions in one direction. Study tip: Remember that maximum buffer capacity always occurs at a 1:1 molar ratio of weak acid to conjugate base, regardless of the specific pKapK_a value. This principle applies to all buffer systems.

Question 20

What happens to the pH of an ammonia buffer (NH3/NH4+NH_3/NH_4^+) when solid NH4ClNH_4Cl is added without changing the volume significantly?

  1. pH increases because more base is added to the solution
  2. pH decreases because the ratio [NH4+]/[NH3][NH_4^+]/[NH_3] increases (correct answer)
  3. pH remains exactly the same due to the buffer's resistance to change
  4. pH increases because chloride ions are basic in aqueous solution
  5. pH becomes undefined because the buffer system is destroyed
Explanation: When you encounter buffer problems, remember that buffers resist pH changes through the Henderson-Hasselbalb equation: pH=pKa+log⁡[base][acid]pH = pK_a + \log\frac{[base]}{[acid]}. For an ammonia buffer, this becomes pH=pKa+log⁡[NH3][NH4+]pH = pK_a + \log\frac{[NH_3]}{[NH_4^+]}. Adding solid NH4ClNH_4Cl increases the concentration of NH4+NH_4^+ ions (the conjugate acid) while leaving NH3NH_3 concentration unchanged. This increases the denominator in the Henderson-Hasselbalb equation, making the ratio [NH4+]/[NH3][NH_4^+]/[NH_3] larger. Since this ratio appears in the denominator of the log term, the log value becomes more negative, decreasing the overall pH. Answer B correctly identifies this relationship. Answer A is wrong because NH4ClNH_4Cl doesn't add base—it adds the conjugate acid NH4+NH_4^+, which actually opposes the base NH3NH_3. Answer C misunderstands buffer behavior; buffers resist pH changes from small additions of strong acids or bases, but adding large amounts of the buffer components themselves will shift the pH according to Henderson-Hasselbalb. Answer D incorrectly assumes Cl−Cl^- affects pH—chloride is the conjugate base of the strong acid HCl, making it essentially neutral in aqueous solution. Study tip: For buffer problems, always identify which component you're adding (acid or base form) and visualize how it changes the ratio in Henderson-Hasselbalb. Adding more acid form decreases pH; adding more base form increases pH.