chem2 ai 2

Solubility

The maximum amount of solute that can dissolve in a given solvent at a specific temperature.

Saturated Solution

A solution that contains the maximum amount of dissolved solute in equilibrium with undissolved solute.

Unsaturated Solution

A solution that contains less solute than it can hold at a given temperature.

Supersaturated Solution

A solution that contains more solute than theoretically possible at equilibrium; unstable.

Ksp

The solubility product constant; equilibrium constant for the dissolution of an ionic compound.

Common Ion Effect

A decrease in solubility of an ionic compound when a common ion is added to the solution.

Selective Precipitation

A technique for separating ions by gradually adding a reagent that precipitates one ion before others.

Ion Product (Qsp)

The product of ion concentrations in a solution; compared with Ksp to predict precipitation.

If Qsp > Ksp

A precipitate will form.

If Qsp < Ksp

No precipitate will form.

If Qsp = Ksp

The solution is saturated and at equilibrium.

Solubility vs. Ksp

For compounds with the same stoichiometry, smaller Ksp means lower solubility.

Effect of pH on Solubility

Acidic solutions can increase solubility of salts containing basic anions.

Example of pH Affecting Solubility

CaCO3 dissolves more in acid because H+ reacts with CO3²⁻.

Molar Solubility

The number of moles of solute that dissolve per liter of solution.

Spectator Ion

An ion that remains unchanged on both sides of a chemical equation.

Net Ionic Equation

An equation showing only species that actually change during the reaction.

Complex Ion Formation

A metal ion bonded to ligands; increases solubility of some salts.

Example of Complex Ion Formation

AgCl dissolves in NH3 because [Ag(NH3)2]+ forms.

Ligand

A molecule or ion that donates an electron pair to a metal ion to form a complex.

Chelation

Binding of a metal ion by multiple donor atoms in a single ligand.

Bronsted-Lowry Acid

Proton (H+) donor.

Bronsted-Lowry Base

Proton (H+) acceptor.

Arrhenius Acid

Substance that produces H+ in water.

Arrhenius Base

Substance that produces OH− in water.

Lewis Acid

Electron pair acceptor.

Lewis Base

Electron pair donor.

Amphoteric Substance

A substance that can act as either an acid or a base.

Example of Amphoteric Substance

Water, H2O, can donate or accept a proton.

Conjugate Acid

The species formed when a base gains a proton.

Conjugate Base

The species formed when an acid loses a proton.

Strong Acid

Completely dissociates in water.

Weak Acid

Partially dissociates in water.

Strong Base

Completely dissociates in water.

Weak Base

Partially dissociates in water.

List of Strong Acids

HCl, HBr, HI, HNO3, HClO4, H2SO4.

List of Strong Bases

Group 1 and 2 hydroxides (NaOH, KOH, Ba(OH)2, etc.).

Kw

The ion product of water; Kw = [H+][OH−] = 1.0 × 10−14 at 25°C.

pH

−log[H+]

pOH

−log[OH−]

Relationship between pH and pOH

pH + pOH = 14 at 25°C.

Neutral Solution

[H+] = [OH−], pH = 7 at 25°C.

Acidic Solution

[H+] > [OH−], pH < 7.

Basic Solution

[H+] < [OH−], pH > 7.

Ka

Acid dissociation constant, measuring acid strength.

Kb

Base dissociation constant, measuring base strength.

Relationship between Ka and Kb

Ka × Kb = Kw.

pKa

−log(Ka)

pKb

−log(Kb)

Strong Acid Ka Value

Very large (>1)

Weak Acid Ka Value

Small (<1)

Henderson-Hasselbalch Equation

pH = pKa + log([A−]/[HA])

Buffer

A solution that resists pH changes when small amounts of acid or base are added.

Buffer Components

A weak acid and its conjugate base, or a weak base and its conjugate acid.

Example of Buffer System

CH3COOH and CH3COONa.

Buffer Capacity

The amount of acid or base a buffer can neutralize before pH changes significantly.

Best Buffer pH

The pKa value of the weak acid (pH ≈ pKa).

How to Prepare a Buffer

Mix a weak acid with its conjugate base, or partially neutralize the weak acid with strong base.

Titration

A process to determine the concentration of an unknown solution using a known one.

Titrant

The solution of known concentration added during a titration.

Analyte

The solution of unknown concentration being titrated.

Equivalence Point

Point at which moles of acid = moles of base.

End Point

The point where the indicator changes color.

Indicator

A dye that changes color at (or near) the equivalence point.

Strong Acid + Strong Base Titration

Equivalence point pH = 7.

Weak Acid + Strong Base Titration

Equivalence point pH > 7.

Weak Base + Strong Acid Titration

Equivalence point pH < 7.

Polyprotic Acid

An acid that can donate more than one proton.

Diprotic Acid

Can donate two protons, e.g., H2CO3.

Triprotic Acid

Can donate three protons, e.g., H3PO4.

Ka1 vs. Ka2

Ka1 > Ka2 > Ka3 (first proton is easiest to remove).

pKa1, pKa2, pKa3

Corresponding negative log values of Ka1, Ka2, Ka3.

Example of Polyprotic Acid Behavior

At intermediate pH, both H2A and HA− exist in equilibrium.

Fractional (Alpha) Diagram

Shows the fraction of each acid form (H2A, HA−, A2−) vs. pH.

Dominant Species at Low pH

Fully protonated acid form.

Dominant Species at High pH

Fully deprotonated base form.

Hydrolysis of Salts

Reaction of anion or cation with water to produce H+ or OH−.

Example of Acidic Salt

NH4Cl (produces H+ from NH4+).

Example of Basic Salt

NaCN (produces OH− from CN−).

Neutral Salt Example

NaCl (no hydrolysis).

Common Ion Effect on Weak Acid

Adding a salt containing the conjugate base decreases ionization of the weak acid.

Example of Common Ion Effect

Adding NaCH3COO to acetic acid decreases [H+].

Solubility and Common Ion Effect

Decreases solubility of a salt with a common ion.

pH and Weak Acid Ionization

As pH increases, weak acid ionizes more.

pKa and Acid Strength

Smaller pKa = stronger acid.

pKb and Base Strength

Smaller pKb = stronger base.

Conjugate Strength Relationship

Stronger acid → weaker conjugate base.

Percent Ionization

([H+]/[HA]_initial) × 100%.

When to Use ICE Table

For equilibrium problems involving weak acids or bases.

When to Use Henderson-Hasselbalch

For buffer or partially neutralized solutions.

Ka Expression for HA ⇌ H+ + A−

Ka = [H+][A−]/[HA].

Kb Expression for B + H2O ⇌ BH+ + OH−

Kb = [BH+][OH−]/[B].

Polyprotic Ka2 Expression

Ka2 = [H+][A2−]/[HA−].

Solubility from Ksp Example

For AB2 → A2+ + 2B−, Ksp = 4s³.

Predicting Precipitation

Compare Qsp and Ksp values.

Effect of Temperature on Solubility

Most solids become more soluble at higher temperature.

Titration Curve

A plot of pH versus volume of titrant added.

Buffer Region in Titration

The flat region where pH changes slowly (before equivalence).

Half-Equivalence Point

The point where [HA] = [A−]; pH = pKa.

Post-Equivalence Region

Excess titrant controls pH.