Unit 8: Acids and Bases

18 terms. Say what each one means, then open the card to check.

• pH and pOH

  pH is the negative base-10 log of the hydronium concentration and pOH is the same for hydroxide; at 25 °C they add up to 14.00.

  Example. A solution with $[\text{H}_{3}\text{O}^{+}] = 1.0 \times 10^{-3}$ M has pH 3.00 and pOH 11.00 at 25 °C.

• Kw, the water constant

  The equilibrium constant for water ionizing into hydronium and hydroxide, $K_w = [\text{H}_{3}\text{O}^{+}][\text{OH}^{-}]$, equal to $1.0 \times 10^{-14}$ at 25 °C and larger at higher temperatures.

  Example. In any water solution at 25 °C, if $[\text{OH}^{-}] = 1.0 \times 10^{-4}$ M, then $[\text{H}_{3}\text{O}^{+}] = 1.0 \times 10^{-10}$ M.

• Neutral solution

  A solution in which the hydronium and hydroxide concentrations are equal; its pH is 7.00 only at 25 °C.

  Example. Pure water near 0 °C has a pH above 7 but is still neutral, because $K_w$ is smaller in the cold and both ion concentrations fall together.

• Strong acid

  An acid that ionizes completely in water, so the hydronium concentration equals the acid concentration for a monoprotic strong acid.

  Example. A 0.020 M solution of nitric acid, $\text{HNO}_3$, has $[\text{H}_{3}\text{O}^{+}] = 0.020$ M and pH 1.70.

• Strong base

  A soluble metal hydroxide that dissociates completely; each formula unit gives as many hydroxide ions as its formula shows.

  Example. A 0.010 M solution of $\text{Ca(OH)}_2$ has $[\text{OH}^{-}] = 0.020$ M.

• Weak acid and Ka

  An acid that only partly ionizes in water; its acid dissociation constant $K_a = [\text{H}_{3}\text{O}^{+}][\text{A}^-]/[\text{HA}]$ is small, and a smaller $K_a$ means a weaker acid.

  Example. An acid with $K_a = 10^{-9}$ ionizes far less than one with $K_a = 10^{-3}$ at the same concentration.

• pKa

  The negative base-10 log of $K_a$. A lower $\text{p}K_a$ means a stronger acid.

  Example. An acid with $K_a = 1.0 \times 10^{-5}$ has $\text{p}K_a = 5.00$.

• Weak base and Kb

  A base that only partly reacts with water to make hydroxide; its base constant $K_b$ measures how far that reaction goes.

  Example. Methylamine reacts with water to make some methylammonium and hydroxide ions, with $K_b$ near $4 \times 10^{-4}$.

• Conjugate acid-base pair

  Two species that differ by exactly one hydrogen ion, such as an acid and the base it becomes after giving up that ion.

  Example. $\text{HCO}_3^-$ and $\text{CO}_3^{2-}$ are a conjugate pair; $\text{HCO}_3^-$ is the acid of the two.

• Ka times Kb equals Kw

  For a conjugate pair in water, $K_a K_b = K_w$, so the weaker the acid, the stronger its conjugate base.

  Example. If an acid has $K_a = 1.0 \times 10^{-6}$, its conjugate base has $K_b = 1.0 \times 10^{-8}$ at 25 °C.

• Percent ionization

  The share of a weak acid that has ionized: the equilibrium hydronium concentration divided by the starting acid concentration, times 100. It rises as the acid is diluted.

  Example. If a 0.050 M weak acid reaches $[\text{H}_{3}\text{O}^{+}] = 1.0 \times 10^{-3}$ M, it is 2.0 percent ionized.

• Buffer

  A solution of a weak acid and its conjugate base, both in large amounts, that changes pH only a little when a small amount of strong acid or base is added.

  Example. In blood, carbonic acid and hydrogen carbonate ions together hold the pH near 7.4 as acids from exercise enter the bloodstream.

• Henderson-Hasselbalch equation

  The buffer equation $\text{pH} = \text{p}K_a + \log(\frac{[\text{A}^-]}{[\text{HA}]})$, which gives the pH from the ratio of base to acid.

  Example. When the conjugate base is ten times the weak acid, the pH is one unit above the $\text{p}K_a$.

• Buffer capacity

  How much strong acid or base a buffer can absorb before its pH changes a lot; it grows with the amounts of weak acid and conjugate base present.

  Example. A buffer made from 1.0 M acid and 1.0 M conjugate base can absorb about ten times as much added acid, for the same pH change, as one made from 0.10 M of each.

• Equivalence point

  The point in a titration where the moles of added titrant exactly match the moles of acid or base in the sample, by the coefficients of the reaction.

  Example. Adding 0.0045 mol of NaOH to a sample that holds 0.0045 mol of a monoprotic acid reaches the equivalence point.

• Half-equivalence point

  The point where half of a weak acid has been turned into its conjugate base; the two are equal, so the pH equals the $\text{p}K_a$.

  Example. If the equivalence point is at 18.0 mL, the half-equivalence point is at 9.0 mL, and the pH read there is the $\text{p}K_a$.

• Acid-base indicator

  A weak acid whose two forms have different colors; it changes color over a pH range near its own $\text{p}K_a$, so it should be chosen to match the equivalence-point pH.

  Example. Bromocresol green changes from yellow to blue between about pH 3.8 and 5.4, so it suits a titration whose equivalence pH is near 4.6.

• Polyprotic acid

  An acid that can give up more than one hydrogen ion, one at a time, each step with its own $K_a$; its titration curve can show more than one equivalence point.

  Example. Phosphoric acid, $\text{H}_3\text{PO}_4$, loses its three hydrogen ions in three steps with $K_a$ values that fall sharply from one step to the next.