Acids and Bases: pH, Ka, Buffers, and Titrations
Acids and bases come together by connecting proton transfer, equilibrium, and concentration. You’ll identify acids and bases, calculate pH, compare strong vs. weak acids, use Ka and Kb, and work through buffers and titrations without mixing up the rules.
What you'll learn
15.1 Start Here: How to Identify the Acid, Base, and Conjugate Pair
A Brønsted-Lowry acid donates a proton (H+). A Brønsted-Lowry base accepts that proton.
That means every acid-base reaction is really a proton-transfer story. The acid loses H+ and becomes its conjugate base, while the base gains H+ and becomes its conjugate acid.
A species is amphiprotic if it can either donate or accept a proton, depending on the conditions.
Water is the most important amphiprotic species, and it undergoes autoionization:
The equilibrium constant for autoionization is the ion-product of water, Kw:
- Every conjugate acid-base pair is connected by exactly one proton.
- HCl donates H+ to become Cl-, so HCl/Cl- is a conjugate pair.
- Do not miss this: if two species differ by more than one proton, they are not a conjugate pair.
15.2 pH and pOH: Reading the Log Scale Correctly
Because hydronium and hydroxide concentrations are often tiny numbers, chemists use a logarithmic scale. The pH and pOH compress those wide ranges into values you can actually compare.
If you are confused, start with the neutral benchmark first, then decide whether the solution is acidic or basic before you calculate anything else.
| Condition | [H3O+] vs [OH-] | pH | Example |
|---|---|---|---|
| Acidic | [H3O+] > [OH-] | < 7 | Stomach acid (pH ≈ 1.5) |
| Neutral | [H3O+] = [OH-] | = 7 | Pure water |
| Basic | [H3O+] < [OH-] | > 7 | Bleach (pH ≈ 12) |
- Each unit change in pH represents a tenfold change in [H3O+].
- A solution of pH 3 has 100 times the hydronium concentration of a pH 5 solution, not 2 times.
- Notice the common mistake: the pH scale is logarithmic, not linear.
15.3 Strong vs. Weak Acids and Bases: Do Not Mix Up Strength and Concentration
A strong acid ionizes essentially completely in water. A weak acid ionizes only partially, so an equilibrium is established.
Strong and weak describe how much a substance ionizes. Concentrated and dilute describe how much of the substance is present.
Do not use strength and concentration as if they mean the same thing. A concentrated weak acid can still have a lower pH than a dilute strong acid.
The acid ionization constant Ka describes that equilibrium. The larger the Ka, the stronger the acid.
For bases, the analogous constant is Kb. A key relationship connects a conjugate pair:
| Strong Acids | Strong Bases |
|---|---|
| HCl, HBr, HI | NaOH, KOH, LiOH |
| HNO3, H2SO4, HClO4 | Ca(OH)2, Ba(OH)2, Sr(OH)2 |
Three patterns connect acid strength to structure — use these when comparing acids that are not on the strong acid list:
- Binary acids down the same group: strength increases going down because the H–X bond becomes easier to break.
- Binary acids across the same period: strength generally increases as the bonded atom becomes more electronegative.
- Oxyacids with the same central atom: more oxygen atoms usually mean a stronger acid because electron density is pulled away from the O–H bond more strongly.
15.4 pKa and pKb: The Strength Scale That Runs Backward
Just as pH is the negative log of a concentration, pKa and pKb are the negative logs of the ionization constants Ka and Kb.
Using a log scale makes these equilibrium constants easier to compare and easier to use in buffer calculations. It also creates a common mistake: the direction flips, and that is easy to lose track of.
- A stronger acid has a larger Ka → smaller pKa.
- A weaker acid has a smaller Ka → larger pKa.
The direction flips on the log scale. The same logic applies to bases and pKb.
| Acid | Ka | pKa | Relative strength |
|---|---|---|---|
| Hydrofluoric acid (HF) | 6.8 × 10-4 | 3.17 | Relatively strong weak acid |
| Acetic acid (CH3COOH) | 1.8 × 10-5 | 4.74 | Moderate |
| Carbonic acid (H2CO3) | 4.3 × 10-7 | 6.37 | Moderate–weak |
| Ammonium ion (NH4+) | 5.6 × 10-10 | 9.25 | Very weak acid |
| Hydrogen cyanide (HCN) | 4.9 × 10-10 | 9.31 | Very weak acid |
For a conjugate acid-base pair, pKa and pKb always sum to 14 at 25 °C — the same relationship as pH and pOH:
- For a buffer made from a weak acid and its conjugate base, pH = pKa when [HA] = [A-].
- At that point the Henderson-Hasselbalch log term equals zero.
- This is why buffers work best near the acid's pKa.
15.5 Polyprotic Acids: One Proton at a Time
A polyprotic acid can lose more than one proton, but it loses them one step at a time.
Each step has its own ionization constant. Later ionizations are usually much weaker because the next proton is being removed from a species that is already more negatively charged.
Because Ka1 ≫ Ka2, the first ionization usually controls the pH. For most introductory problems, you can calculate the pH using only Ka1 unless you are told to analyze later steps.
15.6 Buffers: Why pH Stays Nearly Steady
A buffer is a solution that resists changes in pH when small amounts of strong acid or strong base are added. Buffers contain an appreciable amount of both members of a weak conjugate acid-base pair. The acid component neutralizes added base, and the base component neutralizes added acid.
The pH of a buffer is calculated using the Henderson-Hasselbalch equation:
A buffer works best when [A-] ≈ [HA], meaning pH ≈ pKa. Adding too much acid or base exhausts one component and destroys the buffer. This limit is called the buffer capacity.
- Blood is buffered primarily by the carbonic acid / bicarbonate pair (H2CO3 / HCO3-, pKa ≈ 6.1) and maintains a pH of 7.35–7.45.
- Notice the key idea: a buffer only works if both parts of the conjugate pair are present in useful amounts.
15.7 Acid-Base Titrations: When Neutralization Does and Does Not Mean pH 7
A titration finds the concentration of an unknown acid or base by reacting it with a solution of known concentration.
The equivalence point is reached when the acid and base have reacted in the correct stoichiometric amounts. A titration curve plots pH vs. volume of titrant and helps you locate that point.
Here is the point that catches everyone off guard: the pH at the equivalence point is not always 7. It depends on what species are present after neutralization.
| Titration Type | pH at Equivalence Point | Why |
|---|---|---|
| Strong acid + Strong base | = 7 | Salt does not hydrolyze |
| Weak acid + Strong base | > 7 | Conjugate base hydrolyzes to give OH- |
| Strong acid + Weak base | < 7 | Conjugate acid hydrolyzes to give H3O+ |
An indicator is a weak acid or base that changes color near its own pKa. Choose an indicator whose color-change range overlaps with the steep part of the titration curve.
After Unit 15
This unit pulls together equilibrium, solutions, and concentration into one last major chemistry system. If you are confused, go back through Unit 14: Chemical Equilibrium, the Unit 15 Practice page, and the broader practice hub. For longer-term retention, pair this unit with Why Practice Tests Beat Rereading so pH, buffer, and titration setups keep showing up in retrieval practice.