Polyprotic Acids, Salt Solutions, Buffers, and Titrations: Key Concepts in Acid-Base Chemistry

Successive deprotonations

For inorganic polyprotic acids, the ionization constant ($K_a$) for each successive loss of a proton is generally $10^4$ to $10^6$ times smaller than the previous step.

Factors impacting acid strength

Acid strength correlates with the sum of the bond-dissociation enthalpy and the electron-attachment enthalpy of the conjugate base.

Ranking acidity of hydrohalic acids

For hydrohalic acids, bond strength dominates; HF is the weakest acid because it has the strongest bond, and HI is the strongest because it has the weakest bond.

Ranking acidity of oxoacids

Acidity increases as the number of oxygen atoms bonded to the central element increases (e.g., $HNO_3 > HNO_2$).

Ranking acidity of carboxylic acids

Substituting hydrogen atoms with more electronegative elements increases acidity by stabilizing the negative charge on the conjugate base.

Acidity of hydrated metal cations

Hydrated metal cations become more acidic with increasing positive charge (e.g., 3+ is more acidic than 2+).

Classifying salt solutions

A salt's pH is determined by the acid-base properties of its constituent ions.

Neutral anions

Anions that are conjugate bases of strong acids (like $Cl^-$) and alkali/alkaline earth metal cations (like $Na^+$) are neutral.

Basic anions

Anions that are conjugate bases of weak acids (like $CH_3CO_2^-$) are basic.

Acidic cations

Ammonium ions and transition metal cations with 2+ or 3+ charges are acidic.

Calculating the pH of salt solutions

Write the balanced chemical equation for the ion's reaction with water, set up an ICE table, and use the appropriate $K_a$ or $K_b$ expression.

pH impact on solubility

Insoluble salts containing basic anions will dissolve to a greater extent in acidic solutions.

The Common Ion Effect

Adding a common ion to a weak acid or base solution shifts the equilibrium backward, per Le Chatelier's principle.

How buffers resist drastic pH changes

Buffers consist of a weak acid and its conjugate base.

Calculating buffer pH

You can calculate the pH using an ICE table or the Henderson-Hasselbalch equation: $pH = pK_a + ext{log}rac{[ ext{conjugate base}]}{[ ext{acid}]}$.

Preparing a buffer

Select a weak acid whose $pK_a$ is as close to the desired pH as possible.

Buffer capacities

A buffer must have sufficient concentrations of reagents to react with reasonable quantities of added acid or base.

Identifying titration curves

The equivalence point pH for strong acid with strong base is exactly 7.00.

Calculating pH during a titration

Use stoichiometry to determine what species remain.

Selecting indicators

An acid-base indicator is a weak acid whose color changes depending on the ratio of its acidic form ($HInd$) to its basic form ($Ind^-$).

Calculating molar solubility and $K_{sp}$

The solubility product constant ($K_{sp}$) represents the equilibrium of an insoluble salt.

Ranking compounds by solubility

You can directly compare the solubilities of different salts using their $K_{sp}$ values only if the salts have the exact same cation-to-anion ratio.

Predicting precipitates using $Q_{sp}$

To determine if a precipitate will form, calculate the reaction quotient ($Q$) using the given ion concentrations and compare it to $K_{sp}$.