Honors Chemistry Unit 10 Review: Acids, Bases, and pH

Acids and bases have properties you can spot in the lab. Acids taste sour and turn blue litmus red, and many acids react with active metals such as zinc or magnesium to release hydrogen gas. Bases taste bitter, feel slippery, and turn red litmus blue. These descriptions come from everyday examples like lemon juice and soap, so never taste or touch lab chemicals.

An acid reacting with a metal hydroxide usually forms a salt and water, as in HCl + NaOH → NaCl + H₂O. Acids also react with carbonates and bicarbonates to produce carbon dioxide, which is why vinegar and baking soda fizz.

The biggest idea in this topic is strength. A strong acid ionizes completely in water, so every HCl molecule becomes $\text{H}^+$ and $\text{Cl}^-$. A weak acid like acetic acid (CH₃COOH) only partly ionizes, and in vinegar less than 1% of its molecules have given up a proton at any moment. Strong acids are strong electrolytes, and weak acids are weak electrolytes.

Strength is not the same as concentration. A 0.001 M HCl solution is dilute but still strong, and 6 M acetic acid is concentrated but still weak.

Memorize the common strong acids: HCl, HBr, HI, HNO₃, HClO₃, HClO₄, and H₂SO₄, which is strong for its first proton only. The common strong bases are the Group 1 hydroxides (LiOH, NaOH, KOH) and heavier Group 2 hydroxides like Ca(OH)₂ and Ba(OH)₂.

pH is a shortcut for writing tiny hydrogen ion concentrations. The definition is $\text{pH} = -\log[\text{H}^+]$, so a solution with $[\text{H}^+] = 1 \times 10^{-3}$ M has pH 3. pOH works the same way for hydroxide, $\text{pOH} = -\log[\text{OH}^-]$.

Water always contains a little of both ions because it self-ionizes. At 25°C, $K_w = [\text{H}^+][\text{OH}^-] = 1.0 \times 10^{-14}$, which gives $\text{pH} + \text{pOH} = 14$. In pure water, both concentrations are $1.0 \times 10^{-7}$ M, so the pH is 7.

The scale is logarithmic. Dropping 1 pH unit means 10 times more $\text{H}^+$, and dropping 2 units means 100 times more. Low pH means acidic, but low pOH means basic. Most solutions fall between 0 and 14, though very concentrated acids and bases can go past those limits. To work backward from pH, use $[\text{H}^+] = 10^{-\text{pH}}$.

Worked example. Find the pH of 0.0025 M NaOH. NaOH is a strong base, so $[\text{OH}^-] = 0.0025$ M. Next, $\text{pOH} = -\log(0.0025) = 2.60$, so $\text{pH} = 14.00 - 2.60 = 11.40$. You can check this with $K_w$. $[\text{H}^+] = (1.0 \times 10^{-14}) / 0.0025 = 4.0 \times 10^{-12}$ M, and $-\log(4.0 \times 10^{-12}) = 11.40$. A basic answer makes sense for a hydroxide solution.

A titration finds an unknown concentration by reacting the unknown solution with a solution of known concentration. You add the titrant from a buret to the analyte until the reaction is exactly complete. That moment is the equivalence point, where the acid and base have reacted in the exact stoichiometric ratio shown by the balanced equation. An indicator such as phenolphthalein changes color near that point, going from colorless to pink around pH 8.2-10. The color change is called the endpoint.

Worked example. It takes 18.6 mL of 0.100 M NaOH to neutralize 25.0 mL of HCl. Moles of NaOH are $0.0186 \text{ L} \times 0.100 \text{ mol/L} = 1.86 \times 10^{-3}$ mol. HCl and NaOH react 1:1, so the HCl sample also contains $1.86 \times 10^{-3}$ mol. Its concentration is $1.86 \times 10^{-3} \text{ mol} / 0.0250 \text{ L} = 0.0744$ M.

The pH at equivalence depends on what you mixed. At 25°C, a strong acid titrated with a strong base has an equivalence-point pH of 7. A weak acid with a strong base lands above 7 because the conjugate base left behind is slightly basic. A strong acid with a weak base lands below 7.

A buffer resists pH change when you add a small amount of acid or base. It contains a weak acid and its conjugate base, such as acetic acid and sodium acetate. Added $\text{OH}^-$ reacts with the weak acid, and added $\text{H}^+$ reacts with the conjugate base, so neither ion builds up much. Your blood uses the $\text{H}_2\text{CO}_3/\text{HCO}_3^-$ buffer to stay near pH 7.4.

Chemists use three definitions of acids and bases, and each one covers more substances than the last.

An Arrhenius acid increases $\text{H}_3\text{O}^+$ in water, and an Arrhenius base increases $\text{OH}^-$ in water. HCl and NaOH fit this definition. NH₃ also acts as an Arrhenius base because it reacts with water and raises $[\text{OH}^-]$, while the Brønsted-Lowry definition explains this behavior more directly as proton acceptance.

In Brønsted-Lowry terms, an acid is a proton ($\text{H}^+$) donor and a base is a proton acceptor. In $\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-$, ammonia takes a proton from water, so ammonia is the base and water is the acid.

Every Brønsted-Lowry reaction has two conjugate pairs whose members differ by exactly one $\text{H}^+$, here NH₃/NH₄⁺ and H₂O/OH⁻. In each pair, the member with the extra H is the acid. Water is amphiprotic, so it acts as an acid with ammonia but as a base when HCl dissolves to form $\text{H}_3\text{O}^+$.

A strong acid like HCl has a conjugate base (Cl⁻) that barely accepts protons. Acetic acid's conjugate base, acetate, is a weak base, though it is strong enough to make sodium acetate solutions slightly basic.

The Lewis definition is the broadest. A Lewis acid accepts an electron pair, and a Lewis base donates one. In $\text{BF}_3 + \text{NH}_3 \rightarrow \text{F}_3\text{BNH}_3$, the lone pair on nitrogen bonds to boron, and no proton moves at all.