Acid-Base Chemistry: Concepts and Equilibria

Acid-Base Equilibria and Concepts

• Unit focuses on acid-base chemistry.
• Includes Arrhenius, Bronsted-Lowry, and Lewis concepts.
• Explores water dissociation, weak acids/bases, and equilibrium calculations.
• Covers common ion effect, buffers, hydrolysis, titrations, and indicators.
• Relates buffering action to everyday life.
• Predicts solution acidity/basicity based on salt composition.
• Explains acid-base titration calculations.
• Writing chemical equations is important to show acid-base differences.

Arrhenius Concept

• Svante Arrhenius defined acids/bases based on their effect on water.
• Acids increase H⁺ concentration in aqueous solutions.
• Bases increase OH⁻ concentration in aqueous solutions.
• Strong acids completely ionize to form H₃O⁺ and an anion.
• HClO₄ is a strong acid that ionizes into H₃O⁺ and ClO₄⁻.
• Other examples of strong acids include HCl, HBr, HI, H₂SO₄, and HNO₃.
• Strong bases completely ionize to give OH⁻ and a cation (e.g., NaOH).
• Limitations exist in the Arrhenius theory.

Brønsted-Lowry Concept

• Introduced independently by Brønsted and Lowry in 1923.
• Acid-base reactions are viewed as proton-transfer processes.
• Acids are proton donors, while bases are proton acceptors.
• In the reaction of NH₃ with water, NH₃ acts as a base and water acts as an acid.
• Conjugate acid-base pairs are central to this concept.
• Includes conjugate acid and conjugate base pairs.
• The chloride ion (Cl⁻) is the conjugate base of HCl.
• H₃O⁺ is the conjugate acid of H₂O
• A conjugate base has one fewer H and one more negative charge than its acid.
• A conjugate acid has one more H and one less negative charge than its base.

Strengths of Conjugate Pairs

• Acid-base reaction direction depends on the relative strengths of acids and bases.
• Stronger acids/bases form weaker conjugate bases/acids.
• HCl (strong acid) has a weak conjugate base Cl⁻.
• CH₃COOH (weak acid) has a strong conjugate base CH₃COO⁻.

Auto-ionization and Amphiprotic Species

• Autoionization is a reaction between identical molecules to produce ions.
• Water can act as either an acid or a base, making it amphiprotic.
• Water's autoionization: 2H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq).
• Amphiprotic species donate or accept protons like HCO₃⁻.

Lewis Concept

• G.N. Lewis generalized acid/base concepts beyond proton transfer.
• Focuses on electron-pair donation.
• Lewis acid accepts an electron pair.
• Lewis base donates an electron pair.
• Boron trifluoride (BF₃) accepting electrons from ammonia (NH₃) is an example.
• Species with available empty orbitals (like B in BF₃) can act as Lewis acids.
• Species with lone-pair electrons (like N in NH₃) can act as Lewis bases.

Ionic Equilibria of Weak Acids and Bases

• Focuses on the ionization of water.
• Deriving the expression for the ion product of water (Kw) and use it in calculations.
• Water is a weak electrolyte.
• Describes how to calculate hydronium [H₃O⁺] or hydroxide [OH⁻] concentration.
• Describes writing expressions for the percent ionization of weak acids or bases.

Ionization of Water

• Pure water is a nonelectrolyte, but it undergoes slight self-ionization.
• The water molecule can act as an acid or a base (amphiprotic).
• The autoionization reaction is: 2H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq).
• The equilibrium constant for this process is: Kc = [H3O+][OH-] / [H2O]2
• Ion-product constant for water is Kw = [H3O+][OH-]
• The value of Kw is 1.0 x 10⁻¹⁴ at 25°C.
• Kw is temperature-dependent.
• In neutral solution [H3O+] = [OH-].
• In acidic solution [H3O+] > [OH–].
• In basic solution [OH-] > [H3O+].

pH and pOH scales

• The pH is defined as the negative logarithm of the hydronium ion concentration
• pH = -log[H3O+] or pH = -log[H+].
• The pOH is defined as the negative logarithm of the hydroxide ion concentration
• pOH = -log[OH-]
• pH + pOH = 14
• pH meters are used to determine the pH of a solution.

Strength of Acids and Bases

• Acid/base strength depends on hydrogen/hydroxide ion concentration.
• Includes pH and pOH, percent dissociation, Ka, and Kb.
• Strong acids dissociate completely in water.
• Strong acids yield higher [H3O+] than weak acids.
• Examples of strong acids: HCl, HNO₃, HClO₄, and H₂SO₄.
• Weak acids partially dissociate.

Percent Ionization

• % ionization is the proportion of ionized molecules on a percentage basis.
• The formula is: % ionization = (Ionized acid concentration at equilibrium / Initial concentration of acid) x 100%

Acid Ionization Constant (Ka)

• Ka is a measure of acid strength.
• HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq)
• The constant is expressed as: Ka = [H3O+][A-]/[HA].

Base Dissociation Constant (Kb)

• Equilibria of weak bases are treated like those of weak acids.
• NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq)
• The constant is expressed as: Kb = [NH4+][OH-]/[NH3]

Common Ion Effect

• Shift in ionic equilibrium caused by adding a solute with a common ion.
• Acetic acid ionization is repressed by adding HCl, reducing ionization degree.
• The common-ion effect occurs when a given ion is added to an equilibrium mixture that already contains that ion.

Buffer Solutions

• Resist pH changes upon adding acid/base or dilution.
• Consist of a weak acid/base and its conjugate.
• Need a large concentration of acid to react with added OH-.
• Need a large concentration of base to react with added H⁺.

How Buffers Work

• A buffer solution must contain large concentration of acid to react with added hydroxide ions.
• A buffer system must contain large concentration of base to react with added hydronium ions.
• Buffer molecules must not react with each other.
• Adding CH3COOH and CH3COONa to water creates a simple buffer solution.
• CH3COOH(aq) + H2O(l) ⇌ CH3COO-(aq)+ H3O+(aq)
• The degree of ionization of acetic acid is decreased by the addition of a strong acid.

Henderson-Hasselbalch Equation

• Relates pH to concentrations of acid and conjugate base.
• The formula is: pH = pKa + log ([conjugate base]/[weak acid]).

Hydrolysis of Salts

• Reaction where water breaks chemical bonds in a substance.
• Greek words hydro meaning water and lysis meaning break or to unbind.
• Involves water molecules attaching to two parts of a molecule.
• Salt hydrolysis is the interaction of salt ions with water.

Salts of Strong Acids and Bases

• The reaction between a strong acid (say, HCl) and a strong base (say, NaOH) can be represented by NaOH(aq) + HCl(aq) → NaCl(aq) + H O(l) or in terms of the net ionic equation H+(aq)+OH-(aq) → H2O(l)
• Anions from strong acids are weak conjugate bases and do not undergo hydrolysis
• Cations from strong bases are weak conjugated acids and do not undergo hydrolysis.
• The solution is neutral (pH = 7)

Salts of Weak Acids and Strong Bases

• Neutralizing a weak acid with a strong base yields a salt.
• Acetate ions undergo hydrolysis.
• CH3COO-(aq)+H2O(aq) → CH3COOH(aq)+OH-(aq)
• Solution becomes basic due to excess OH⁻.

Salts of Strong Acids and Weak Bases

• Neutralizing a weak base with a strong acid produces a salt.
• Ammonium chloride (NH4Cl) results .
• The solution contains ammonium and chloride ions
• The ammonium ions, the conjugate acid of ammonia, reacts with water and increases the hydronium ion concentration: NH4+ (aq) +H2O(l) → H3O+ (aq) + NH3(aq)

Acid-Base Indicators

• Indicate solution acidity, basicity, or neutrality.
• Indicators are weak organic acids (HIn) or bases (In⁻).
• Indicator color depends on pH.
• HIn(aq) + H₂O(l) ⇌ H₃O⁺(aq) + In⁻(aq)
• Acidic conditions favor HIn form.
• Basic conditions favor In⁻ form.

Equivalents of Acids and Bases

• An equivalent is the amount of a substance that reacts with an arbitrary amount of another substance in a given chemical reaction (typically one mole).
• Equivalent = the amount of substance needed to react with or supply one mole of hydrogen ions (H+) in an acid-base reaction. Equivalent of Sulfuric acid (H₂SO₄) = 2.
• Normality is the number of equivalents of a solute/liter of solution.
• Normality = Number of equivalents of solutes/Liters of solution
• If enough acid is added to neutralize a given volume of base, the equation below holds:
• where NA and VA refer to the normality and volume of the acid solution, respectively, and NB and VB refer to the normality and volume of the base solution, respectively.

Acid-Base Titrations

• An acid-base titration is a procedure for determining the amount of acid (or base) in a solution.
• Done by determining the volume of base (or acid) of known concentration that will completely react with it.
• In a titration, one of the solutions to be neutralized is placed in a flask or beaker, together with an acid base indicator
• The other solution, the base, will be added from a burette and is called titrant.
• The titrant is added to the acid (titrand) drop by drop, up to the equivalence point.
• The equivalence point is the point at which the amount of titrant added is just enough to completely neutralize the analyte solution
• During an acid base titration, moles of base are equal to moles of base
• The point in a titration at which the indicator changes color is called the end poin