This section is a relatively simple and short one all about buffer capacity, which helps us gauge the effectiveness of buffers. For a quick reminder, buffers are important because they are resistant to changes in pH.
However, buffers are not infinitely resistant. Eventually, the buffer will weaken and succumb to the acid/base being added. This is why, despite there being buffers in your bloodstream, chugging hydrochloric acid is a very bad idea. Seriously, don't do it!
Buffer capacity helps us see how much acid/base one can add until there is a significant change in pH.
Describing Buffer Capacity
As said by the Henderson-Hasselbalch equation, the pH of a buffer is defined by the ratio of the concentrations of the conjugate base to the acid, or in math terms [A-]/[HA]. The capacity of a buffer is determined by the magnitudes of these concentrations.
What do we mean when we refer to magnitude? Essentially, the magnitude of each concentration describes how large the concentrations are. A concentration of 5M would have a higher magnitude than a 0.5M solution. The more concentrated the acid and conjugate base, the stronger the buffer is at reducing pH changes! There is more acid and conjugate base to be resistant to strong acids/strong bases in a similar volume.
Directional Buffer Capacity
Here's a crucial point about buffers: they don't necessarily resist added acid and base equally! The buffer's ability to neutralize added acid versus added base depends on the ratio of conjugate base to weak acid:
- When a buffer has more conjugate base than acid ([A⁻] > [HA]): It has greater capacity to neutralize added acid than added base
- When a buffer has more conjugate acid than base ([HA] > [A⁻]): It has greater capacity to neutralize added base than added acid
Why does this make sense? Think about what happens when you add acid or base:
- Added H⁺ reacts with A⁻: H⁺ + A⁻ → HA
- Added OH⁻ reacts with HA: OH⁻ + HA → A⁻ + H₂O
If you have lots of A⁻ but little HA, you can neutralize lots of added H⁺, but you'll quickly run out of HA to neutralize OH⁻!
Example: Directional Capacity
Consider two acetate buffers:
- Buffer A: 0.1 M CH₃COOH and 0.5 M CH₃COO⁻
- Buffer B: 0.5 M CH₃COOH and 0.1 M CH₃COO⁻
Buffer A has more conjugate base, so it better resists added acid. Buffer B has more weak acid, so it better resists added base.
Both might have similar total buffer capacity, but their directional capacities are opposite!
Example Problem: Identifying the Stronger Buffer
To practice applying buffer capacity, try thinking about two separate buffer systems, one with 5M acetic acid and 5M sodium acetate and another with 0.05M acetic acid and 0.05M sodium acetate. Because the ratios of the conjugate base to the acid are the same, both of these buffers will have the same pH (pH=4.74). However, our question asks us the following:
After HCl is added to each buffer system, the first one has a resulting pH of 4.74, and the second one has a resulting pH of 4.56. Which one has the better buffering capacity and why?
We can see that in the first buffer, the pH remained relatively unchanged, whereas, in the second, the pH dropped. Because the first system's pH remained constant, the first buffer is more effective at resisting pH changes and, therefore, has a better buffering capacity. The magnitude of the concentration of both the acid and the conjugate base is higher in the first buffer compared to the second, implying the first buffer will also have a stronger buffer capacity.
Practice Multiple Choice Questions
The College Board likes to use buffer capacity specifically on multiple choice questions because, unlike many other questions relating to buffers, questions about buffer capacity are very often qualitative. That means your answer will relate to some sort of non-numerical conclusion based on the information you are given.
They could ask you to identify the stronger buffer, like in the last question, or ask about a change in a system and how this may affect the buffer capacity.
Here is another example:
This problem gives us a BUNCH of information and can be really overwhelming at first. However, by breaking it down piece by piece, the problem is more approachable.
We know that a student is creating a buffer of acetic acid (CH3COOH, also stated as HC2H3O2) and sodium acetate since a weak acid mixes with its conjugate base. Now we can examine the numbers. The first set of numeric information indicates that the student wants to mix 250 mL of 0.100M acetic acid with 500 mL of 0.440M sodium acetate.
Yikes! The student makes an error. They choose an acid with half the concentration and a conjugate base with half of the volume. For acetic acid, the volume remains at 250mL, but the molarity is now 0.0500M. Meanwhile, the molarity of the sodium acetate stays at 0.440M, but the volume falls from 500mL to 250mL. The problem wants to know the ramifications of these mistakes.
Before looking at the answer choices, what changed? In both sets (the weak acid and the conjugate base), the number of moles is halved. Therefore, we will have half the number of moles of each species in our buffer.
Again, before looking at any answer choices, how is the buffer impacted? As we discussed earlier, a lower number of moles of each reactant in the buffer creates less resistance to changes in pH. In other words, the buffer capacity lowers.
Which answer choice matches that concept?
If we read the answer choices, we see that what we just described fits answer A. You may be tempted to pick B, but remember that buffer capacity is determined by the number of moles of the weak acid and the conjugate base.
Vocabulary
The following words are mentioned explicitly in the College Board Course and Exam Description for this topic.
| Term | Definition |
|---|---|
| buffer capacity | The amount of acid or base that a buffer solution can neutralize while maintaining a relatively constant pH. |
| buffer components | The conjugate acid-base pair that makes up a buffer solution and determines its pH and capacity. |
| concentration ratio | The ratio of the concentration of the conjugate base to the concentration of the conjugate acid, [A-]/[HA], in a buffer solution. |
| conjugate acid | The species formed when a base accepts a proton; the acid form in an acid-base conjugate pair. |
| conjugate base | The species formed when an acid donates a proton; the base form in an acid-base conjugate pair. |
Frequently Asked Questions
What is buffer capacity and how is it different from just pH?
Buffer capacity is how much strong acid or base a buffer can neutralize before its pH changes significantly—it’s a measure of resistance to pH change, not the pH itself. pH tells you the current H+ concentration (and for a buffer you can get it from Henderson–Hasselbalch using pKa and the [conjugate base]/[conjugate acid] ratio). Buffer capacity depends on total buffer concentration and the relative amounts of conjugate acid and base: higher total concentration (same ratio) gives greater capacity (CED 8.10.A.1). Also, if a buffer has more conjugate acid than base it better resists added base; if it has more base it better resists added acid (CED 8.10.A.2). For AP style problems, expect to use Henderson–Hasselbalch and reason about moles added/neutralized and total buffer concentration. Want a quick review and practice? See the Topic 8.10 study guide (https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu) and try practice problems (https://library.fiveable.me/practice/ap-chemistry).
Why does adding more buffer components increase capacity but not change the pH?
Because pH depends on the ratio of conjugate base to conjugate acid, not their absolute amounts. The Henderson–Hasselbalch equation shows this directly: pH = pKa + log([A−]/[HA]). If you increase [A−] and [HA] by the same factor, the ratio (and therefore pH) stays the same. Buffer capacity, though, is about how many moles of strong acid or base the buffer can neutralize before the ratio (and thus pH) changes appreciably. Increasing both components (keeping the ratio constant) raises the total number of moles of HA and A−, so the buffer can absorb more added H+ or OH− before its ratio shifts—capacity increases. Practically: doubling both [HA] and [A−] doubles the amount of acid/base you can neutralize while keeping pH the same. This is exactly the CED point 8.10.A.1. For more review, see the Topic 8.10 study guide (https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu) and try practice problems (https://library.fiveable.me/practice/ap-chemistry).
I'm confused about buffer capacity - does it mean the buffer is stronger or something?
Buffer capacity basically tells you how much acid or base a buffer can "soak up" before its pH changes noticeably—so yes, a higher capacity means the buffer is stronger at resisting pH change. Two things matter (CED 8.10.A): - Total concentration of the buffer components: increasing both [acid] and [conjugate base] (keeping their ratio the same) raises capacity—more material to neutralize added H+ or OH– (8.10.A.1). - Ratio of acid to base: if you have more conjugate acid than base, the buffer resists added base better; more conjugate base than acid gives better protection against added acid (8.10.A.2). Maximum capacity around the point where [acid] ≈ [base] (pH ≈ pKa, half-equivalence). Use Henderson–Hasselbalch to see pH but remember capacity comes from total [buffer]. For a quick review, see the Topic 8.10 study guide (https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu) and try practice problems (https://library.fiveable.me/practice/ap-chemistry).
How do I know if a buffer can handle more acid or more base being added?
Check the ratio and the total amount. If [HA] > [A−] the buffer can neutralize more added base (because there’s more acid to donate H+), and if [A−] > [HA] it can handle more added acid (because there’s more base to accept H+). A 1:1 ratio (HA = A−) gives equal capacity for acid and base; that’s also the maximum buffering around pH = pKa (half-equivalence point). Increasing the total concentration of HA + A− (while keeping the ratio the same) raises the overall buffer capacity—it just takes more added acid or base to change the pH (CED 8.10.A.1–2). Use the Henderson–Hasselbalch equation, pH = pKa + log([A−]/[HA]), to see how big a pH shift will be when small amounts are added. For more review and examples tied to the AP CED, check the Topic 8.10 study guide (https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu) and more practice problems at (https://library.fiveable.me/practice/ap-chemistry).
What's the relationship between conjugate acid/base concentrations and buffer capacity?
Buffer capacity depends on two things: the ratio of conjugate acid to conjugate base (which sets the buffer pH via Henderson–Hasselbalch) and the total concentrations of those components (which set how much acid or base the buffer can neutralize). If you keep the acid:base ratio constant, pH stays the same but increasing both concentrations (higher total buffer concentration) raises buffer capacity—it can neutralize more added H+ or OH– (CED 8.10.A.1). If the buffer has more conjugate acid than base, it resists added base better (more capacity to neutralize OH–); if it has more conjugate base, it resists added acid better (CED 8.10.A.2). Maximum capacity for resisting both directions is near the 1:1 ratio (half-equivalence). For practice and review see the Topic 8 study guide (https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu), the Unit 8 overview (https://library.fiveable.me/ap-chemistry/unit-8), and lots of practice problems (https://library.fiveable.me/practice/ap-chemistry).
Why does keeping the ratio constant matter when you increase buffer concentration?
Keeping the ratio constant matters because pH depends on the ratio [A–]/[HA] (Henderson–Hasselbalch), not on their absolute concentrations. If you double both HA and A–, the log([A–]/[HA]) term stays the same, so pH stays the same (this is exactly what CED 8.10.A.1 says). But total buffer concentration ( [HA]+[A–] ) controls buffer capacity: a higher total means more moles of acid and base are available to neutralize added H+ or OH– before the ratio—and therefore pH—shifts significantly. In short: ratio → pH; total concentration → how much acid/base the buffer can absorb. That’s why when you “strengthen” a buffer you keep the same ratio but raise concentrations to increase capacity (useful on the exam when explaining neutralization or calculating how many moles of titrant a buffer can resist). For a quick review, see the Topic 8.10 study guide (https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu). For more practice problems, check Fiveable’s AP Chem practice set (https://library.fiveable.me/practice/ap-chemistry).
Can someone explain why more conjugate acid means better at neutralizing base?
Think of a buffer as a pair: a weak acid (HA) and its conjugate base (A−). If you add a base (OH−), the conjugate acid HA neutralizes it: HA + OH− → A− + H2O. So the more HA you have, the more OH− moles the buffer can consume before the pH moves much—that’s greater buffer capacity for added base. Two quick CED points: capacity depends on the absolute concentrations (more total buffer concentration = more capacity) and on the ratio (if [HA] > [A−], capacity for added base > capacity for added acid). You can see this from neutralization stoichiometry and from the Henderson–Hasselbalch idea: pH changes less when there’s more HA available to react. For a short review, see the Topic 8.10 study guide (https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu). For extra practice, try problems at (https://library.fiveable.me/practice/ap-chemistry).
What happens to buffer capacity when you have unequal amounts of conjugate acid and base?
If the conjugate acid and base aren’t present in equal amounts, the buffer still resists pH change but it’s stronger in the direction for which it has more reagent. That is exactly what CED 8.10.A says: - If [HA] > [A–], the buffer has a greater capacity to neutralize added base (because there’s more HA to react with OH–). - If [A–] > [HA], it has a greater capacity to neutralize added acid (there’s more A– to react with H+). Remember: buffer pH depends on the ratio [A–]/[HA] (Henderson–Hasselbalch), while capacity depends on the absolute amounts (total buffer concentration) and the ratio. The maximum (most balanced) resistance to both acid and base additions occurs when [HA] ≈ [A–] (the half-equivalence point). For more practice or a quick review of Topic 8.10, see the Fiveable study guide (https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu) and try problems at (https://library.fiveable.me/practice/ap-chemistry).
How do I calculate or predict buffer capacity for AP exam questions?
Buffer capacity is basically how many moles of strong acid or base a buffer can neutralize without a big pH change. For AP problems, do two things: 1) Predict qualitatively from the CED rules: capacity increases as the total concentration of buffer components increases (same ratio, higher concentration → stronger buffering). A buffer with more conjugate acid than base resists added base better; more base than acid resists added acid better (CED 8.10.A). 2) Calculate quantitatively by counting moles: capacity ≈ the moles of the component that will neutralize added reagent. Example: 0.100 L of 0.10 M HA + 0.10 M A− has 0.010 mol A− available to neutralize added strong acid. If you add 0.004 mol H+, A− → HA uses 0.004 mol; remaining A− = 0.006 mol. Use Henderson–Hasselbalch (pH = pKa + log[A−]/[HA]) to find new pH if needed. On the AP exam you may be asked to explain the relationship or do neutralization math—practice both (see the Topic 8.10 study guide: https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu and more practice: https://library.fiveable.me/practice/ap-chemistry).
I don't understand why buffer capacity changes but pH stays the same - isn't that contradictory?
Not contradictory—pH depends on the ratio [conjugate acid]/[conjugate base] (Henderson–Hasselbalch: pH = pKa + log([A–]/[HA])), while buffer capacity depends on the total amount of those components (how many moles are available to neutralize added acid or base). If you keep the ratio constant but increase both concentrations (say 0.10 M HA + 0.10 M A– → 0.20 M HA + 0.20 M A–), pH stays the same because the ratio is unchanged, but buffer capacity doubles because there are twice as many moles to react with added H+ or OH–. Also, if a buffer has more conjugate acid than base it resists added base better (and vice versa)—that’s buffer capacity vs. directionality (CED 8.10.A.1–2). For more examples and AP-style practice, see the Topic 8.10 study guide (https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu) and Unit 8 resources (https://library.fiveable.me/ap-chemistry/unit-8). For extra practice problems, check the AP practice bank (https://library.fiveable.me/practice/ap-chemistry).
What's the difference between buffer capacity and buffer range?
Buffer capacity = how much strong acid or base a buffer can neutralize before its pH changes significantly. It depends on the total concentrations of the conjugate acid and base (higher total [HA]+[A−] → higher capacity) and on their ratio. A buffer has maximum capacity against both acid and base when [HA] ≈ [A−] (at the half-equivalence/ pH ≈ pKa) but if [HA] > [A−] it resists added base better, and vice versa (CED 8.10.A; Henderson–Hasselbalch). Buffer range (or effective buffering pH range) = the pH window where the buffer actually works well, roughly pKa ± 1. Outside that window the conjugate pair is too imbalanced to keep pH steady. Short rule: capacity = “how much” it can absorb (controlled by total concentration and ratio); range = “where” it works (set by pKa). For more AP-aligned review, see the Topic 8.10 study guide (https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu) and extra practice (https://library.fiveable.me/practice/ap-chemistry).
Why would you want a buffer with higher capacity instead of just any buffer?
You want a higher-capacity buffer when you need the solution to resist larger amounts of added acid or base without changing pH. Buffer capacity depends on the total concentration of the conjugate acid/base pair (not just their ratio): if you keep the ratio constant but increase both components (for example, 0.10 M HA / 0.10 M A– → 1.0 M HA / 1.0 M A–), the pH stays the same (Henderson–Hasselbalch) but the buffer can neutralize ~10× more strong acid or base. Also, the relative amounts set directionality: more conjugate acid gives greater capacity against added base; more conjugate base gives greater capacity against added acid (CED 8.10.A.1–2). For AP exam problems, watch for total buffer concentration and which component is in excess. Review examples and practice problems in the Topic 8.10 study guide (https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu) and more practice at the unit page (https://library.fiveable.me/ap-chemistry/unit-8) or the practice bank (https://library.fiveable.me/practice/ap-chemistry).
How does the Henderson-Hasselbalch equation relate to buffer capacity?
Henderson–Hasselbalch (pH = pKa + log([A−]/[HA])) connects directly to buffer capacity. The log ratio term shows how the relative concentrations of conjugate base [A−] and acid [HA] set the buffer pH—if [A−]=[HA], pH = pKa (the half-equivalence point), and the buffer resists pH change best near that pKa (maximum usable buffer range). Buffer capacity, however, depends on total amounts: increasing both [A−] and [HA] by the same factor keeps pH (the ratio) the same but increases how much acid or base the buffer can neutralize (higher capacity). Also, if a buffer has more HA than A− it can neutralize more added base; more A− than HA gives greater capacity against added acid (CED 8.10.A). For AP review, make sure you can explain both the ratio (pH control) and total concentration (capacity)—see the Topic 8.10 study guide for examples (https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu). For extra practice, try problems at (https://library.fiveable.me/practice/ap-chemistry).
If I have 0.1 M buffer vs 0.5 M buffer with same ratio, which neutralizes more acid and why?
The 0.5 M buffer will neutralize more added acid than the 0.1 M buffer if both have the same conjugate acid:base ratio. Buffer capacity depends on the total concentrations (the “total buffer concentration” or common-ion amounts), not the pH ratio itself—increasing both [acid] and [base] by the same factor keeps pH (via Henderson–Hasselbalch) the same but raises how many moles of H+ (or OH–) the buffer can neutralize (CED 8.10.A.1). Practically: a 0.5 M acetate/acetate– buffer contains five times more moles of each component per liter than a 0.1 M buffer, so it can consume about five times more strong acid before the ratio (and pH) shifts significantly. For more practice and AP-aligned review of buffer capacity, see the Topic 8.10 study guide (https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu) and Unit 8 resources (https://library.fiveable.me/ap-chemistry/unit-8).
What are some real world examples where buffer capacity actually matters?
Buffer capacity matters everywhere you need pH to stay nearly constant when small amounts of acid/base are added. Real examples: - Blood: the bicarbonate/CO2 buffer keeps pH ~7.35–7.45; the buffer’s total concentration (≈10–30 mM bicarbonate range) sets how much acid it can neutralize before pH shifts dangerously. - Saliva and teeth: saliva buffers protect enamel; low buffer capacity + acid (soda) → cavities. - Pharmaceuticals: drug formulations use buffers so medication stays active and safe (small pH changes can ruin drugs). - Antacids: they’re buffers or neutralizers—capacity determines how long they relieve heartburn. - Brewing/winemaking and fermentation: yeast activity and flavor depend on stable pH; brewers adjust buffer strength. - Pools and wastewater: buffering prevents big swings when people add acids/bases or when effluent varies. Link to the AP Topic for review and how buffer capacity depends on concentrations and the acid/base ratio (Henderson–Hasselbalch, pKa): https://library.fiveable.me/ap-chemistry/unit-8/buffer-capacity/study-guide/TRzbdif6DAyXpITT77wu. For more practice, check the unit overview (https://library.fiveable.me/ap-chemistry/unit-8) and thousands of problems (https://library.fiveable.me/practice/ap-chemistry).