There are many measures in AP Chemistry, especially relating to acids and bases and equilibrium. In this section, we will discuss the relationships between pH and pKa, two of the most important measures!
‘p’ Notation
'p' notation is actually fairly simple and seen throughout acid-base chemistry (we have pH, pOH, pKa, pKb, and many more!). 'p'-something is simply equal to the -log(something). For example, pH = -log(H+), and pOH = -log(OH-). Similarly, pKa = -log(Ka).
pKa and Acid Strength
An important use of pKa is in describing acid strength relative to other acids. For example, if one acid has a pKa of 3 and another has a pKa of 2, we know that the acid with a pKa of 2 is 10 times as acidic (note, however, that this does not mean that the pH is 10 times lower). Using 'p' notation gives us a logarithmic scale and not a linear one.
Like pH, where a lower pH corresponds to a higher [H+], a lower pKa implies a higher Ka. However, it is worth noting that a high pKa does not imply basicity. Another note is that, like pH and pOH, pKa + pKb = 14.
Predicting Predominant Forms: The pH vs pKa Rule
One of the most crucial concepts in acid-base chemistry is understanding which form of an acid or base predominates at a given pH. This relationship is fundamental and appears frequently on the AP exam!
Here's the key rule:
- When pH < pKa, the acid form (HA) predominates
- When pH > pKa, the base form (A⁻) predominates
- When pH = pKa, the concentrations of HA and A⁻ are equal
Let's understand why this works. Remember that Ka = [H⁺][A⁻]/[HA]. Rearranging this gives us [A⁻]/[HA] = Ka/[H⁺]. Taking the negative log of both sides leads to the Henderson-Hasselbalch equation, which shows that pH = pKa + log([A⁻]/[HA]).
When pH < pKa, the log term must be negative, meaning [A⁻]/[HA] < 1, so [HA] > [A⁻] - the acid form predominates!
When pH > pKa, the log term is positive, meaning [A⁻]/[HA] > 1, so [A⁻] > [HA] - the base form predominates!
Example: Acetic Acid (CH₃COOH, pKa = 4.74)
- At pH 3: Since 3 < 4.74, CH₃COOH (acid form) predominates
- At pH 6: Since 6 > 4.74, CH₃COO⁻ (base form) predominates
- At pH 4.74: Equal amounts of CH₃COOH and CH₃COO⁻
This concept is essential for understanding:
- Buffer behavior at different pH values
- Amino acid protonation states
- Drug absorption (many drugs are weak acids or bases)
- Indicator color changes
pH, pKa, and Buffers
pH and pKa are also related to buffers. As a reminder, a buffer is a mixture of an acid and its conjugate base and is important because it is resistant to changes in pH. However, a question arises: when is the buffer the strongest?
The Henderson-Hasselbalch Equation can be applied to find the pH of a buffer:
The strongest buffer occurs when the concentration of [A-] is equal to [HA]. In this case, pH = pKa + log(1) ⇒ pH = pKa. The relationship is vital, especially when looking at titration curves, because this same point occurs at the half-equivalence point, implying that you have the optimal buffer at the half-equivalence point.
Understanding the pH vs pKa relationship also helps predict buffer effectiveness:
- When pH is within ±1 of the pKa, you have a good buffer
- As pH moves further from pKa, one form begins to dominate and buffering capacity decreases
Practice Problem: Predicting Forms
For a weak base like ammonia (NH₃, Kb = 1.8 × 10⁻⁵), we first need the pKa of its conjugate acid NH₄⁺:
- pKb = 4.74, so pKa(NH₄⁺) = 14 - 4.74 = 9.26
At different pH values:
- pH 7: Since 7 < 9.26, NH₄⁺ (protonated form) predominates
- pH 11: Since 11 > 9.26, NH₃ (deprotonated form) predominates
- pH 9.26: Equal amounts of NH₄⁺ and NH₃ (optimal buffer!)
Acid-Base Indicators
Finally, we will discuss acid-base indicators. Acid-base indicators are a class of compounds that change color depending on the pH of the solution they are in. You may have used indicators in class during titrations to note when the equivalence point of a titration occurs. Some examples of acid-base indicators are bromothymol blue, phenolphthalein, and methyl red.
When choosing an acid-base indicator, you usually want to pick one in which your pH will end up in the effective range, which is the pKa plus or minus 1. While you will not need to memorize any indicators or their effective ranges on the exam, you may be asked to pick which one is the most effective for a certain experiment.
Let's see this concept with an FRQ from 2010:
We are given the following prompt:
In order for the indicator to be useful to us, we want it to change color at the equivalence point for this titration. Looking at the graph, we see that the pH at the equivalence point is 7. We also know this because it is a strong acid strong base titration. Therefore, we want to pick an indicator with a pH range closest to 7. This turns out to be methyl red, which is the correct answer.
Vocabulary
The following words are mentioned explicitly in the College Board Course and Exam Description for this topic.
| Term | Definition |
|---|---|
| acid-base indicator | A substance that exhibits different observable properties (such as color) depending on its protonation state, allowing it to signal pH changes in a 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. |
| deprotonated | The state of a molecule or ion after it has donated a proton (H+), decreasing its positive charge or increasing its negative charge. |
| equivalence point | The point in a titration where the analyte is completely consumed by the titrant in a quantitative reaction. |
| pKa | The negative logarithm of the acid dissociation constant (Ka); used to compare the relative strength of weak acids and predict protonation state at different pH values. |
| pKb | The negative logarithm of the base dissociation constant (Kb); used to compare the relative strength of weak bases. |
| predominant form | The chemical species (protonated or deprotonated) that exists in higher concentration in solution at a given pH. |
| protonated | The state of a molecule or ion after it has accepted a proton (H+), increasing its positive charge or decreasing its negative charge. |
| protonation state | The relative concentrations of the protonated (HA) and deprotonated (A-) forms of an acid or base in solution. |
| titration | An analytical procedure used to determine the amount of an analyte in solution by reacting it with a titrant of known concentration. |
| weak acid | An acid that only partially ionizes in solution, establishing an equilibrium between the molecular form (HA) and its conjugate base (A-). |
| weak base | A base that only partially dissociates in water; examples include ammonia and carboxylate ions. |
Frequently Asked Questions
What is pKa and how is it different from pH?
pKa is the negative log of an acid’s acid-dissociation constant (Ka): pKa = −log(Ka). It’s a property of a specific acid that tells you how easily that acid gives up a proton. pH, by contrast, is the negative log of the hydrogen-ion activity in a solution (pH = −log[H+])—it describes the actual acidity of the solution you're measuring. How they differ in practice (AP-relevant): - Compare pH and pKa to predict protonation (CED 8.7.A.1): if pH < pKa, the protonated form (HA) dominates; if pH > pKa, the deprotonated form (A−) dominates; if pH = pKa, [HA] = [A−]. - Use the Henderson–Hasselbalch equation: pH = pKa + log([A−]/[HA]) to calculate ratios, percent deprotonation, and buffer behavior (CED keywords). - Pick indicators whose pKa ≈ equivalence-point pH for titrations (CED 8.7.A.3). For more examples and quick practice, see the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG) and unit review (https://library.fiveable.me/ap-chemistry/unit-8). Practice problems are at (https://library.fiveable.me/practice/ap-chemistry).
Why does the form of a weak acid change when pH changes?
The form of a weak acid changes with pH because protonation depends on the H+ concentration. Compare the solution pH to the acid’s pKa (CED 8.7.A): when pH < pKa, [HA] > [A–] (the protonated acid predominates); when pH > pKa, [A–] > [HA] (the deprotonated base predominates). You can see this quantitatively with the Henderson–Hasselbalch equation: pH = pKa + log([A–]/[HA]). A one-unit change in pH shifts the ratio [A–]/[HA] by a factor of 10, so small pH shifts can strongly change which form is dominant. That’s why indicators change color near their pKa and why you choose an indicator whose pKa is near the titration equivalence pH (CED 8.7.A.2–3). For a focused review, check the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG) and practice more problems at (https://library.fiveable.me/practice/ap-chemistry).
I'm confused about when to use HA vs A- in acid base problems - can someone explain?
Think of HA (protonated acid) vs A– (deprotonated base) as a pH “tug-of-war” decided by pKa. Rule from the CED (8.7.A): - If pH < pKa → solution is more acidic, so HA predominates. - If pH > pKa → solution is more basic, so A– predominates. - If pH = pKa → [HA] ≈ [A–] (half protonated). Use the Henderson–Hasselbalch equation to quantify this: pH = pKa + log([A–]/[HA]). Rearranged, it tells you the ratio [A–]/[HA] for any pH. Example: pH = pKa + 1 → [A–]/[HA] = 10, so ~91% is A–. This is exactly what AP expects you to explain (CED 8.7.A, indicator/titration selection, buffer region). For practice, see the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG) and try problems from the Unit 8 page (https://library.fiveable.me/ap-chemistry/unit-8) or the practice bank (https://library.fiveable.me/practice/ap-chemistry).
How do I predict whether an acid will be protonated or deprotonated at a given pH?
Compare the solution pH to the acid’s pKa. If pH < pKa the protonated form (HA) predominates; if pH > pKa the deprotonated form (A−) predominates—that’s the CED rule (8.7.A.1). You can quantify this with the Henderson–Hasselbalch equation: pH = pKa + log([A−]/[HA]). For example, when pH = pKa, [A−] = [HA] (50:50). If pH = pKa + 1, [A−] ≈ 10×[HA] (about 91% deprotonated). If pH = pKa − 1, [HA] ≈ 10×[A−] (about 91% protonated). Use this for weak acids/bases and picking indicators (choose an indicator with pKa near the equivalence pH—CED 8.7.A.3). For a quick refresher, see the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG) and practice problems at (https://library.fiveable.me/practice/ap-chemistry).
What's the relationship between pKa and the strength of an acid?
pKa tells you how easily an acid gives up a proton: lower pKa = stronger acid. pKa is just −log(Ka), so a big Ka (more dissociation) gives a small pKa. Use the Henderson–Hasselbalch form, pH = pKa + log([A−]/[HA]), to see the protonation state: when pH < pKa the protonated form (HA) predominates; when pH > pKa the deprotonated form (A−) predominates (CED 8.7.A.1). For titrations and indicators, pick an indicator whose pKa is near the equivalence-point pH so the color change matches the endpoint (CED 8.7.A.3). Want practice? Review the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG) and more Unit 8 resources (https://library.fiveable.me/ap-chemistry/unit-8) or try practice problems (https://library.fiveable.me/practice/ap-chemistry).
Why do acid-base indicators change color at different pH values?
Indicators change color because the molecule of the indicator exists in two different forms (protonated HA and deprotonated A–) that have different colors. Which form dominates depends on pH relative to the indicator’s pKa (CED 8.7.A.1–A.2). The Henderson–Hasselbalch equation, pH = pKa + log([A–]/[HA]), shows that when pH ≈ pKa the two forms are about equal, so you see the “transition” color. Most indicators have a visible color change over about pKa ±1 unit (the transition range) because that’s where the ratio [A–]/[HA] changes enough to shift color. For titrations you pick an indicator whose pKa (transition range) is close to the equivalence-point pH so the endpoint matches the actual equivalence point (CED 8.7.A.3). For a quick review, see the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG) and practice more problems at Fiveable (https://library.fiveable.me/practice/ap-chemistry).
How do I choose the right indicator for a titration experiment?
Pick an indicator whose transition range (its effective color-change pH) overlaps the pH at the equivalence point of your titration. On the AP CED this is exactly why indicators are chosen with pKa ≈ equivalence-point pH (8.7.A.3). Practically: first predict the equivalence pH from the acid/base strengths (strong/strong → ≈7; weak acid + strong base → >7; weak base + strong acid → <7). Then use a titration curve or quick pH estimate to get the equivalence pH and choose an indicator whose transition range contains that value (e.g., methyl orange ~3.1–4.4 for acidic endpoints; bromothymol blue ~6.0–7.6 for near-neutral; phenolphthalein ~8.3–10.0 for basic endpoints). Always check the actual titration curve if possible and pick the indicator that gives the sharpest color change near the steep region. For a refresher and practice problems, see the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG) and unit resources (https://library.fiveable.me/ap-chemistry/unit-8).
What happens when solution pH is less than the pKa of an acid?
If solution pH < pKa, the protonated (acid) form predominates. Use the Henderson–Hasselbalch equation: pH = pKa + log([A–]/[HA]). So when pH < pKa the log term is negative → [A–]/[HA] < 1, meaning [HA] > [A–]. Quantitatively, 1 pH unit below the pKa gives [A–]/[HA] = 0.10 (about 90% HA); 2 units below gives ≈99% HA. This is exactly what the CED expects you to explain for 8.7.A: compare pH and pKa to predict protonation state. Practically, that’s why acid–base indicators (which have different colors for HA vs A–) change near their pKa and why you pick an indicator with pKa near the equivalence-point pH for titrations. For a quick refresher, see the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG) and try practice problems (https://library.fiveable.me/practice/ap-chemistry).
Can someone explain in simple terms what it means when pH equals pKa?
When pH = pKa it means the acid and its conjugate base are present in equal amounts. Using the Henderson–Hasselbalch equation: pH = pKa + log([A–]/[HA]). If pH = pKa then log([A–]/[HA]) = 0, so [A–] = [HA]—the acid is 50% deprotonated. Practically: this is the center of a buffer region (maximum buffer capacity) and the point where a weak acid/base pair resists pH change best. For indicators, the color-change midpoint (about 50% of each form) corresponds to pKa, so you pick an indicator whose pKa is near the equivalence-point pH for accurate titrations (CED 8.7.A). Want more practice or review on this topic? Check the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG), the Unit 8 overview (https://library.fiveable.me/ap-chemistry/unit-8), or >1000 practice questions (https://library.fiveable.me/practice/ap-chemistry).
Why is it important that an indicator's pKa matches the equivalence point pH?
You want the indicator’s pKa near the equivalence-point pH because an indicator changes color where its conjugate acid and base are about equal. By Henderson–Hasselbalch, pH = pKa + log([A−]/[HA])—when pH = pKa the indicator is 50% protonated and 50% deprotonated, so the color is mid-transition. Most indicators show a visible change over about pKa ±1, so if the indicator’s pKa matches the equivalence pH the steep part of the titration curve crosses the indicator’s transition range and you get an accurate end point. If pKa is far from the equivalence pH, the color change happens too early or too late and the titration result is wrong. This idea ties directly to CED 8.7.A (protonation state, pKa, indicator behavior) and is important for AP titration questions. For a quick review check the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG) and more practice problems (https://library.fiveable.me/practice/ap-chemistry).
I don't understand how to compare pH and pKa to determine the predominant form - help?
Quick rule you can always use: compare the solution pH to the acid’s pKa. - If pH < pKa → protonated (HA) predominates. - If pH > pKa → deprotonated (A–) predominates. - If pH = pKa → about 50:50 HA and A– (buffer midpoint). A handy quantitative check is the Henderson–Hasselbalch equation: pH = pKa + log([A–]/[HA]). So if pH = pKa + 1, log([A–]/[HA]) = 1 → [A–]/[HA] = 10 (base form ~90%, acid ~10%). If pH = pKa – 1, acid form ~90%. That ±1 pH-unit idea also explains indicator ranges and picking an indicator with pKa near the equivalence pH for titrations (CED 8.7.A). This is an AP-tested idea (use it in free-response and titration questions). For more practice and a short study guide, see the Topic 8.7 page (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG). For broader review and extra problems, check Unit 8 (https://library.fiveable.me/ap-chemistry/unit-8) and practice sets (https://library.fiveable.me/practice/ap-chemistry).
What's the difference between the protonated and deprotonated forms of an indicator?
The protonated form of an indicator is the version that still has the acidic proton (written HA); the deprotonated form has lost that proton (A−). They’re the same conjugate acid–base pair but different in charge and electronic structure, so they absorb light differently and show different colors (that’s why indicators change color). Which form dominates depends on pH vs pKa: if pH < pKa, the protonated (HA) form predominates; if pH > pKa, the deprotonated (A−) form predominates (use Henderson–Hasselbalch to quantify ratios). Practically, indicators have a transition range of about pKa ±1, so you pick one whose pKa is near the titration’s equivalence-point pH for an accurate end point (CED 8.7.A.1–3). For a quick review and practice questions on this topic, see the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG), the Unit 8 overview (https://library.fiveable.me/ap-chemistry/unit-8), and practice problems (https://library.fiveable.me/practice/ap-chemistry).
How do indicators actually work in titrations and why do we need them?
Indicators are weak acids or bases that change color depending on their protonation state—HA (one color) vs A− (another). In a titration the solution’s pH shifts as you add titrant; the indicator’s color change happens around its pKa (roughly pKa ±1 is the visible transition range). You use an indicator because the equivalence point (where stoichiometric amounts of acid and base have reacted) is a single pH, but you need a visible signal when the solution crosses that pH. For accuracy, pick an indicator whose pKa (and thus transition range) is close to the expected equivalence-point pH for your titration—e.g., phenolphthalein (pKa ≈ 9.7, changes ~8–10) for strong base/weak acid titrations. Remember the equivalence pH depends on whether the conjugate species are weak or strong, so calculate the titration curve or estimate equivalence pH before choosing an indicator (CED 8.7.A.1–A.3). For more review, see the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG) and Unit 8 resources (https://library.fiveable.me/ap-chemistry/unit-8). For extra practice, check the AP problem set (https://library.fiveable.me/practice/ap-chemistry).
When solution pH is greater than acid pKa, which form is more concentrated?
If pH > pKa, the deprotonated (base) form A– is more concentrated than the protonated acid form HA. Use the Henderson–Hasselbalch equation: pH = pKa + log([A–]/[HA]). When pH exceeds pKa, log([A–]/[HA]) is positive, so [A–] > [HA] (for example, if pH = pKa + 1, then [A–] = 10·[HA]). This is exactly the CED essential knowledge (8.7.A.1). This idea matters on the exam for predicting protonation states, buffer behavior, and choosing indicators (pick an indicator whose pKa is near the equivalence pH). For a quick review, see the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG) and try practice problems (https://library.fiveable.me/practice/ap-chemistry).
Why do we care about the protonation state of acids and bases in solution?
Because whether a molecule is protonated (HA) or deprotonated (A–) changes its charge, shape, solubility, reactivity and even color—so knowing the protonation state tells you how it will behave in solution. AP CED learning objective 8.7.A says you predict the predominant form by comparing pH to pKa: if pH < pKa, HA dominates; if pH > pKa, A– dominates. That matters for buffers (control pH, buffer capacity near pKa), for choosing indicators (pick an indicator with pKa near the equivalence-point pH), for titration end-point accuracy, and for percent deprotonation calculations using Henderson–Hasselbalch. In real labs and biology, protonation alters solubility and binding (affects absorption, transport, enzyme activity). For more on pH, pKa, and practice problems tied to the AP framework, check the Topic 8.7 study guide (https://library.fiveable.me/ap-chemistry/unit-8/ph-pka/study-guide/yKYJCv37E6gkI01yeRwG) and the AP Chem practice question bank (https://library.fiveable.me/practice/ap-chemistry).