Unit 1: Acid-Base Equilibria

Acid-base equilibria are fundamental to a vast array of chemical and biological processes, spanning from industrial manufacturing to the delicate balance of physiological systems. This unit explores how substances behave as acids and bases through various theoretical lenses, including the Arrhenius, Brønsted-Lowry, and Lewis models. By mastering quantitative tools such as pH, dissociation constants, and titration techniques, students can predict and control the behavior of complex chemical solutions.+4

1.1 Acid-Base Concepts

• Arrhenius acid: A substance that increases the concentration of $H^{+}$ (proton ion) in aqueous solution.
• Arrhenius base: A substance that increases the concentration of $OH^{-}$ (a hydroxide ion) in aqueous solution.
• Strong acid: A substance that completely ionizes in aqueous solution to produce $H_{3}O^{+}$ and an anion.
• Strong base: A substance that completely ionizes in aqueous solution to give $OH^{-}$ and a cation.
• Group IA and Group IIA elements: These form the principal strong bases (except for Beryllium).
• Brønsted-Lowry acid: A species that acts as a proton donor.
• Brønsted-Lowry base: A species that acts as a proton acceptor.
• Conjugate acid-base pairs: A pair consisting of an acid and its conjugate base, or a base and its conjugate acid.
• Conjugate base: The species that remains after one proton has been removed from an acid.
• Conjugate acid: The species resulting from the addition of a proton to a base.
• Amphiprotic species: A substance (like water) that possesses the ability to act as either an acid or a base.
• Lewis acid: A substance that can accept an electron pair to form a covalent bond.
• Lewis base: A substance that can donate an electron pair to form a covalent bond.

1.2 Ionic Equilibria of Weak Acids and Bases

• Molecular auto-ionization (self-ionization): A reaction between two identical neutral molecules to produce an anion and a cation.
• Weak electrolyte: A substance, such as water, that shows very small electrical conduction due to limited self-ionization.
• Ionic product of water ($K_w$): The product of $[H^{+}]$ and $[OH^{-}]$, which always equals $1.0 \times 10^{-14}$ at $25^{\circ}C$.
• pH scale: A quantitative measure of acidity defined as the negative of the logarithm of the molar hydronium-ion concentration.
• pOH: The negative logarithm of the hydroxide ion concentration of a solution.
• Acid ionization constant ($K_a$): A quantitative measure of the strength of an acid in solution, representing the equilibrium constant for acid dissociation.
• Base dissociation constant ($K_b$): The equilibrium constant for the ionization of a weak base in water.
• Percent ionization (percent dissociation): The ratio of the ionized concentration to the initial concentration of the acid or base, expressed as a percentage.

Steps to Calculate pH and pOH:

1. Determine the molar concentration of hydronium ions ($[H_{3}O^{+}]$) or hydroxide ions ($[OH^{-}]$).
2. For pH, calculate the negative log: $pH = -\log[H_{3}O^{+}]$.
3. For pOH, calculate the negative log: $pOH = -\log[OH^{-}]$.
4. Verify the relationship $pH + pOH = 14$.

1.3 Common Ion Effect and Buffer Solution

• Common-ion effect: The repression of the ionization of a weak electrolyte caused by the addition of a strong electrolyte that provides a shared ion.
• Le Châtelier’s principle: Used to explain how the equilibrium composition shifts to the left when a common ion is added, decreasing the degree of ionization.
• Buffer solution: A solution that resists changes in pH when small amounts of acid or base are added.
• Buffer capacity: The ability of a buffer to maintain pH; a concentrated buffer undergoes smaller changes than a dilute buffer.
• Henderson-Hasselbalch equation: A formula for calculating buffer pH: $pH = pK_a + \log\left(\frac{[Base]}{[Acid]}\right)$.

1.4 Hydrolysis of Salts

• Salt hydrolysis: The interaction of the anions or cations of a salt (or both) with water, often involving the breakdown of chemical bonds.
• Hydrolysis of Salts of Strong Acids and Strong Bases: These do not undergo hydrolysis; the resulting solution remains neutral (pH 7).
• Hydrolysis of Salts of Weak Acids and Strong Bases: The anion interacts with water to yield $OH^{-}$, producing a basic solution.
• Hydrolysis of Salts of Strong Acids and Weak Bases: The cation undergoes hydrolysis to yield a strong conjugate acid, resulting in an acidic solution.
• Hydrolysis of Salts of Weak Acids and Weak Bases: The solution’s nature (acidic, basic, or neutral) depends on the relative values of the ions.

1.5 Acid-Base Indicators and Titrations

• Acid-base indicators: Weak organic acids (HIn) or bases ($In^{-}$) that signal whether a solution is acidic, basic, or neutral by changing color.
• Normality (N): A measure of the reacting power of a solution, defined as the number of equivalents of solute per liter of solution.
• Acid-base titration: A laboratory procedure used to determine the concentration of an unknown analyte by reacting it with a titrant of known concentration.
• Equivalence point: The point in a titration where the amount of titrant added is exactly enough to completely neutralize the analyte.
• End point: The point at which the indicator changes color; it must be chosen to match the equivalence point.
• Titration curve: A plot of the solution pH against the volume of titrant added during the procedure.

Steps for Acid-Base Titration:

1. Place a measured volume of the analyte (the solution to be neutralized) in a flask.
2. Add a few drops of a suitable acid-base indicator.
3. Fill a burette with the titrant (the base or acid of known concentration).
4. Add the titrant to the analyte rapidly at first, then drop by drop as the end point nears.
5. Stop the addition at the exact moment the indicator changes color (end point).
6. Use the volume of titrant added to calculate the unknown concentration.

Key Terminology

• Acid-base indicators: Weak organic substances that change color based on pH.
• Acid-base titration: Procedure to determine solution concentration via neutralization.
• Acid ionization constant ($K_a$): Quantitative measure of acid strength.
• Amphiprotic species: Substances that can both donate and accept protons.
• Analyte: The solution of unknown concentration in a titration.
• Arrhenius acid-base concept: Theory defining acids by $H^{+}$ and bases by $OH^{-}$ production.
• Autoionization (Self-ionization): Reaction where like molecules react to produce ions.
• Base dissociation constant ($K_b$): Equilibrium constant for weak base ionization.
• Brønsted-Lowry concept: Theory defining acids as proton donors and bases as acceptors.
• Buffer solution: A solution that resists changes in pH.
• Common ion effect: Suppression of ionization by adding a shared ion.
• Conjugate acid: Species formed when a base accepts a proton.
• Conjugate base: Species remaining after an acid donates a proton.
• End point: The point where an indicator signals titration completion.
• Equivalence point: The point where stoichiometric amounts of acid and base react.
• Equivalents of acids and bases: Measure of chemical reacting power.
• Hydrolysis of salts: Reaction of salt ions with water to affect pH.
• Ionic product of water ($K_w$): Constant value ($1.0 \times 10^{-14}$) for water ionization.
• Lewis’s concept: Theory defining acids as electron-pair acceptors and bases as donors.
• Normality (N): Concentration expressed in equivalents per liter.
• Percent ionization: The extent of ionization expressed as a percentage.
• pH scale: Measure of hydronium ion concentration.
• pOH: Measure of hydroxide ion concentration.
• Strong acid/base: Substances that ionize completely in solution.
• Titrant: The solution of known concentration in a titration.
• Titration curve: A graph of pH versus titrant volume.
• Weak electrolyte: A substance that conducts electricity poorly due to low ionization.