Acids and Bases Chemistry

• CHE 101 explores the nature of acids and bases.

Concepts of Acid and Base

• S.A. Arrhenius proposed in 1884 that an acid is a substance that dissolves in water to produce hydrogen ions (H+) and corresponding negative ions.
• HCl, HNO3, and H2SO4 are typical acids that dissociate in water to yield hydrogen ions and their negative counterparts.
• A base produces hydroxyl ions (OH-) with corresponding positive ions when dissolved.
• KOH, NaOH, and Ca(OH)2 are examples of bases.

Limitations of Arrhenius Theory

• The theory is inapplicable to reactions outside aqueous solutions.
• Acids and bases are defined by ions in water.
• Free protons (H+) do not exist; non-solvated hydrogen ions cannot exist in solution.
• Hydroxonium ion (H3O+) characterizes an acid.
• The base definition is valid only for hydroxides, excluding compounds like ammonia gas and sodium carbonate, which can be neutralized without containing OH- ions.

Bronsted – Lowry Acid-Base Theory

• J.N. Bronsted and T.M. Lowry independently proposed that acids and bases bring about an improvement in the Arrhenius definition.
• It is based on proton transfer.
• An acid donates a proton when dissolved in a solvent (base).
• A base accepts a proton when dissolved in a solvent (acid).
• A conjugate acid is a corresponding base that gains a proton.
• A conjugate base is a corresponding acid that loses a proton.

Lewis Acid-Base Theory

• G.N. Lewis proposed the concept of acids and bases based on the sharing of an electron pair.
• An acid accepts a lone pair of electrons.
• This forms a co-ordinate covalent bond from the donor substance (base).
• BF3 is an example.
• A base donates a lone pair of electrons.
• This forms a co-ordinate covalent bond.
• NH3 is an example.

Water Equilibrium (Dissociation of water)

• Pure water is neutral
• It ionizes slightly to yield equal concentrations of hydroxonium and hydroxyl ions.
• Equilibrium exists between water molecules and the ions they form, which is self-ionization.
• The equilibrium constant of H2O is constant.
• The constant Kw, or [H3O+][OH-], refers to the Ionic Product of water or water dissociation equilibrium constant.
• If concentration is in moldm-³, then the unit of Kw is mol²dm-⁶.
• The numerical value depends on temperature, but it is 1.0 x 10-14 at 25°C for pure water.
• In acidic solutions, [H3O+] > 1.0 x 10-7 moldm-3.
• In neutral solutions, [H3O+] = [OH-] = 1.0 x 10-7 moldm-3.
• In basic solutions, [H3O+] < 1.0 x 10-7 moldm-3.

The pH Scale

• S.L. Sorenson suggested that the concentration of H3O+ ions should use the pH scale in 1909.
• The pH of a solution is the negative logarithm of the hydroxonium concentration in moldm-3.
• It measures the degree of acidity of a solution.
• pOH is the negative logarithm of the hydroxide ion concentration in moldm-3, measuring the degree of alkalinity.
• Lower pH indicates higher hydroxonium ion concentration and acidity.
• Acidic solutions have a pH value below 7.
• Basic solutions have a pH value above 7.
• A solution with pH of 7 is neutral.
• pH = - log [H3O+] = log 1/[H3O+]
• pOH = - log [OH-] = log 1/[OH-]
• At 25°C, Kw = 1.00 x 10-14
• pKw = - log Kw = - log (1.00 x 10-14).
• pKw = 14
• pKw = - log Kw = pH + pOH = 14
• pH + pOH = 14

Strength of Acids and Bases

• Strength is determined by the measurement of electrical conductivity.
• Strong acids/bases dissociate completely in solution with a stable conjugate base/acid.
• Weak acids/bases dissociate partially with an unstable conjugate base/acid.
• Relative strengths are determined by comparing dissociation constants.
• When an acid HA is dissolved in water, it will dissociate to form hydroxonium ion and an anion.
• A- + H2O НА + ОН
• The strength of a base (A-) also varies with concentration of the hydroxide ion [OH-] produced on dissociation.
• Weak acids have the concentration of water is constant.
• ais the acid equilibrium constant for dissociation of an acid and called the dissociation constant of an acid.
• Dissociation constants are unaffected by concentration changes just like pH but influenced only by temperature changes.

Relationship between Kₐ, K♭ and K𝑤

• A relationship exists between the water dissociation equilibrium constant, Kₐ, and the base dissociation equilibrium constant, K♭, of the conjugate base of the acid.
• Use HA as a weak acid and find the equilibrium expression.
• The equilibrium expression for conjugate base A- of the acid in water is: Kb = [HA] [OH-] / [A-]
• Strength is measured by pKa and pKb values; smaller values equal stronger acids or bases.
• • log Ka - log Kp = - log Kw = 14
• pKa + pKp = pKw = 14 (at 25°C).

Ostwald Dilution Law

• The degree of dissociation (α) of a weak electrolyte is proportional to the inverse square root of the concentration (C), or the square root of the dilution.
• The concentration of any one ionic species is given by the root of the product of the dissociation constant and the concentration of the electrolyte.
• α = ν (kd/C)
• [H3O+] = VKaC
• [OH-] = VkbC

Buffer Solutions

• A buffer resists changes in hydrogen ion concentration or pH upon the addition of acid or alkaline substances.
• Buffering Capacity (B) is the magnitude of a buffer action
• Buffering Capacity (B) is measured by the amount of strong base required to alter the pH by one unit.
• There are acidic and basic buffers.

Acidic Buffers

• Mixing a weak acid and its salt of a strong base prepares it.
• Examples are ethanoic acid and sodium ethanoate.
• A solution is acidic and resists changes in pH.
• pH = pKa - log ([Acid]/[Salt]) OR pH = pKa + log ([salt]/[Acid])

Basic Buffers

• Mixing a weak base and its salt of a strong acid prepares alkaline buffers.
• Aqueous ammonia and ammonium chloride are examples.
• A solution is basic and resists pH changes in a basic medium.
• • log [OH-] = - log Kb - log ([Base]/[Salt])
• pOH = pKb + log ([Salt]/[Base])

Explore acid-base theories including Arrhenius, Bronsted-Lowry, and Lewis definitions. Learn about conjugate acids and bases, water's ionic product, and pH scales. Understand acid and base strength measurements.