# Acids, Bases and pH

Chemistry I · Energy, Equilibrium and Electrochemistry · https://tryals.app/learn/chemistry-i/acids-bases-and-ph

## Measuring How Acidic

Three definitions of increasing generality coexist. **Arrhenius**: an acid releases $\mathrm{H^+}$ in water, a base releases $\mathrm{OH^-}$. **Brønsted-Lowry**: an acid donates a proton, a base accepts one, which introduces **conjugate pairs**, since every acid becomes its conjugate base on donating. **Lewis**: an acid accepts an electron pair, a base donates one, which covers reactions with no proton at all.

Water is **amphoteric**, acting as either, and undergoes autoprotolysis:

$$K_w = [\mathrm{H^+}][\mathrm{OH^-}] = 1.0 \times 10^{-14} \text{ at } 25\,^\circ\mathrm{C}$$

This single constant ties the scale together, giving $\mathrm{pH} + \mathrm{pOH} = 14$ and making neutrality pH 7, but only at 25 °C, since $K_w$ is itself temperature-dependent.

$$\mathrm{pH} = -\log[\mathrm{H^+}]$$

Because the scale is logarithmic, one pH unit is a **tenfold** change in concentration. A **strong** acid ionises essentially completely, so $[\mathrm{H^+}]$ equals its concentration. A **weak** acid ionises only partially, described by

$$K_a = \frac{[\mathrm{H^+}][\mathrm{A^-}]}{[\mathrm{HA}]}$$

For a weak acid of concentration $C$, assuming ionisation is small, $[\mathrm{H^+}] \approx \sqrt{K_aC}$.

Strength and concentration are independent axes. A 0.1 M solution of hydrochloric acid has pH 1; 0.1 M acetic acid, at the same concentration, has pH about 2.9, because only around 1 % of it ionises. Strength is about the *fraction* that ionises; concentration is about how much acid is present.

Conjugate strength runs inversely: the stronger the acid, the weaker its conjugate base. The chloride ion, conjugate of a very strong acid, is so weak a base that it has no effect on pH at all.

> **Common pitfall:** treating "weak" as a synonym for "dilute". They are independent. A concentrated weak acid can easily have a lower pH than a dilute strong one, 5 M acetic acid is more acidic than 0.001 M hydrochloric acid.

## Practice questions

7 of this lesson's 12 practice questions, with answers. The full set is in the app.

### 1. A 0.1 M solution of acetic acid has pH 2.9 while 0.1 M hydrochloric acid has pH 1. What does this show?

A. Acetic acid is more dilute than hydrochloric acid
B. Acetic acid ionises only partially at the same concentration
C. Acetic acid has a larger acid dissociation constant
D. The two acids release different numbers of protons per molecule

**Answer:** B. Acetic acid ionises only partially at the same concentration

**Why:** Both solutions contain the same amount of acid. Hydrochloric acid ionises completely, giving $[\mathrm{H^+}] = 0.1$ M; acetic acid ionises only about 1 %, so its hydrogen-ion concentration is far lower. Strength and concentration are independent.

Page: https://tryals.app/practice/chemistry-i/acids-bases-and-ph/a-0-1-m-solution-of-acetic-acid-has-ph-2-9-while-0-1-m-hydrochloric

### 2. Pure water undergoes autoprotolysis, and the equilibrium constant Kw increases with temperature. What follows regarding neutrality as water is heated above 25 °C?

A. The water turns weakly acidic because hydrogen ions outnumber hydroxides
B. Neutrality stays at pH 7 because autoprotolysis maintains equal ions
C. Neutral pH falls below 7 because hydrogen-ion concentration increases
D. The water turns weakly basic because hydroxide ions dissociate faster

**Answer:** C. Neutral pH falls below 7 because hydrogen-ion concentration increases

**Why:** Neutrality requires [H+] = [OH-], not a fixed pH of 7. As temperature rises, endothermic autoprotolysis yields higher [H+], lowering the neutral pH point while remaining strictly neutral.

Page: https://tryals.app/practice/chemistry-i/acids-bases-and-ph/pure-water-undergoes-autoprotolysis-and-the-equilibrium-constant-kw

### 3. A solution has $\mathrm{pOH} = 4.5$ at 25 °C. Compute its pH.

**Answer:** 9.5 (within ±0.05)

**Why:** $\mathrm{pH} = 14 - \mathrm{pOH} = 14 - 4.5 = 9.5$. The relation follows directly from $K_w = 10^{-14}$, and it holds only at 25 °C because $K_w$ varies with temperature.

Page: https://tryals.app/practice/chemistry-i/acids-bases-and-ph/a-solution-has-poh-4-5-at-25-c-compute-its-ph

### 4. A concentrated solution of a weak acid can have a lower pH than a dilute solution of a strong acid.

**Answer:** True

**Why:** True, 5 M acetic acid has a far lower pH than 0.001 M hydrochloric acid. Strength fixes the fraction that ionises; concentration fixes how much is there to ionise, and either can dominate.

Page: https://tryals.app/practice/chemistry-i/acids-bases-and-ph/a-concentrated-solution-of-a-weak-acid-can-have-a-lower-ph-than-a

### 5. Match each acid to its conjugate base.

**Answer:**

- $\mathrm{HCl}$ → $\mathrm{Cl^-}$
- $\mathrm{H_2O}$ → $\mathrm{OH^-}$
- $\mathrm{NH_4^+}$ → $\mathrm{NH_3}$
- $\mathrm{H_2CO_3}$ → $\mathrm{HCO_3^-}$

**Why:** Each conjugate base is the acid minus one proton, so the charge falls by exactly one. Note that water appears as an acid here, it is amphoteric and takes either role depending on its partner.

Page: https://tryals.app/practice/chemistry-i/acids-bases-and-ph/match-each-acid-to-its-conjugate-base

### 6. Sort each species by its behaviour in water.

**Answer:**

- Strong acid: Hydrochloric acid, Nitric acid
- Weak acid: Acetic acid, Carbonic acid
- Strong base: Sodium hydroxide, Potassium hydroxide

**Why:** Hydrochloric and nitric acids ionise essentially completely; acetic and carbonic acids only partially; the group 1 hydroxides dissociate completely as strong bases.

Page: https://tryals.app/practice/chemistry-i/acids-bases-and-ph/sort-each-species-by-its-behaviour-in-water

### 7. Why is the chloride ion unable to raise the pH of a solution?

A. It undergoes complete hydrolysis in water, producing hydronium ions instead
B. It is the conjugate of a very strong acid and so is a negligibly weak base
C. Its stable noble-gas electron configuration prevents it acting as a base
D. It behaves as a Lewis acid by accepting an electron pair from water molecules

**Answer:** B. It is the conjugate of a very strong acid and so is a negligibly weak base

**Why:** Hydrochloric acid is essentially completely ionised, which means the reverse reaction barely occurs, so chloride has almost no tendency to take a proton back. Its conjugate basicity is negligible and sodium chloride solutions are neutral.

Page: https://tryals.app/practice/chemistry-i/acids-bases-and-ph/why-is-the-chloride-ion-unable-to-raise-the-ph-of-a-solution