CHEM 1412 Concept Review: Aqueous Ionic Equilibria

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Common Ion Effect: Whenever a weak electrolyte and a strong electrolyte containing a common ion are together in solution, the weak electrolyte ionizes less than it would if it were alone in solution. This is basically just an application of Le Chatlier’s Principle.

Buffers: Solutions that contain both a weak acid and its corresponding conjugate base. Such solutions resist changes in pH caused by the addition of small amounts of strong acid or strong base. The pH of buffers can be calculated easily by using the Henderson-Hasselbalch equation:

$LaTeX: pH=pK_a+\log\left(\frac{\left[A^-\right]}{\left[HA\right]}\right)$ $pH=pK_a+\log\left(\frac{\left[A^-\right]}{\left[HA\right]}\right)$ (Henderson-Hasselbalch equation)

Buffer Capacity: The amount of strong acid or base that a buffer can neutralize before the pH begins to change to a significant degree.

pH Range of a buffer: The pH range over which the buffer is able to act effectively at resisting pH changes. This is generally within of the pK_a of the weak acid which the buffer contains.

Effect of adding strong acid or base to a buffer

Acid-Base Titrations

pH titration curve: a graph of pH vs. mL of titrant added which allows one to determine the equivalence point of an acid-base titration. Such curves are useful for determining K_a’s of weak acids and for determining a suitable indicator for future titrations with a given species.

Equivalence Point: For an acid-base titration, the equivalence point is the point at which stoichiometrically equal amounts of acid and base have reacted. On the pH titration curve, this point is halfway up the sharp steep jump in the curve.

Indicator: A species that reacts with the titrant but less favorably than the analyte which changes color as it reacts. In acid-base titrations, this is typically a weak organic acid. For optimal results, you want to select an indicator whose pK_a is as close as possible to the equivalence point pH.

Polyprotic Acid Titrations: the pH titration curve for these titrations shows multiple equivalence points. Each equivalence point represents the completion of one “deprotonization” or the removal of one acidic hydrogen from the acid.

Strong Acid-Strong Base Titrations VS. Weak Acid-Strong Base Titrations

There are several distinct differences between Strong Acid-Strong Base (SA-SB) titration curves and Weak Acid-Strong Base (WA-SB) titration curves which can be seen clearly in the following graph:
CH 16 img 2.png

Difference #1: The WA-SB curve has a higher initial pH than the SA-SB curve.

Difference #2: The WA-SB curve has a much smaller jump in pH at the rapid rise portion of the curve near the equivalence point.

Difference #3: The pH at equivalence point for the SA-SB curve is 7.00, but the pH at the equivalence point for the WA-SB curve is above 7.00.

(For information on how to calculate the pH at different points in a titration, refer to the supplemental guide “Calculating pH at Different Points in Titrations” linked HERE.)

Solubility Equilibria

Solubility: The maximum amount of solute that can be dissolved in a given amount of solvent. This measurement is often expressed in grams per liter (g/L). When expressed as M (mol/L), it is called molar solubility.

Sparsely Soluble Salts: Ionic substances that dissolve only very slightly in water. Though these substances are often classified as “insoluble” in water, even substances classified as “insoluble” in a given solvent usually dissolve to a minute extent.

Solubility Product Constant (or K_sp): The equilibrium constant for the equilibrium between an ionic solid and its saturated solution. It is basically a measure of how much solid dissolves to make a saturated solution. The following chemical equation is an example of a K_sp equilibrium and it’s corresponding K_sp expression.

Example: $LaTeX: MX\left(s\right)\longleftrightarrow M^+\left(aq\right)+X^-\left(aq\right)$ $MX\left(s\right)\longleftrightarrow M^+\left(aq\right)+X^-\left(aq\right)$ $LaTeX: K_{sp}=\left[M^+\right]\left[X^-\right]$ $K_{sp}=\left[M^+\right]\left[X^-\right]$

(Recall that solids are not included in equilibrium expressions due to their essentially unchanging concentration.)

Several factors that can increase or decrease the solubility of a sparsely soluble salt include:

1. The Common Ion Effect: Adding a soluble salt that has an ion in common with an insoluble salt causes the insoluble salt to be LESS soluble. (In other words, adding common ions is adding products which shifts equilibrium to the left.)
2. The pH of the solution: Ions from the salt that react with either acid or base can shift equilibrium to the right
3. The formation of complex ions: The ability of Lewis bases to interact and form complexes with metal ions. By binding up the metal ions and effectively removing them from the equilibrium, the K_sp shifts hard to the right, leading to the dissolving of the salt.
4. Amphoterism: The capability of certain sparsely soluble salts to act as either an acid or a base, which allows them to dissolve in either.

K_sp, Q, and Precipitation

Using Le Chatlier’s Principle, we can predict whether a precipitate will form when we have a given concentrations of ions by calculating the reaction quotient (Q) and comparing it to the solubility product constant (K_sp):

If LaTeX: Q>K $Q>K$ , then there are too many ions in solution, equilibrium shifts to the left, and precipitation occurs until LaTeX: Q=K $Q=K$ .

If LaTeX: Q=K $Q=K$ , equilibrium exists (saturated solution).

If LaTeX: Q<K $Q<K$ , there are not enough ions in solution, equilibrium shifts to the right, and solid dissolves until LaTeX: Q=K $Q=K$ .