1.4: Acid Base Equilibrium

• Acid-Base Equilibria lecture notes provided by TJ Sotelo from the Institute of Chemistry, College of Science at the University of the Philippines Diliman

Module Overview

• Topics covered include acid-base theories, autoionization of water, pH scale, acid/base strengths, ionization, ionization constants, salt properties, common-ion effect, and buffer solutions.

Acid-Base Theories

• Arrhenius acids ionize in water to produce H+ ions, while Arrhenius bases ionize in water to produce OH- ions.
• Brønsted-Lowry acids donate protons (H+), and Brønsted-Lowry bases accept protons (H+). Solution needs to be aqueous.
• Lewis acids are electron pair acceptors, and Lewis bases are electron pair donors.
• H+ is a hydrogen ion or a proton and OH- is a hydroxide ion

Arrhenius Theory

• Acids produce H+ (or H3O+) in aqueous solution (e.g., HCl, HBr, HI, HNO3, HClO4, H2SO4).
• Bases produce OH- in aqueous solution (e.g., NaOH, LiOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2).
• H+ does not exist as free species in aqueous solution.
• H3O+ is the hydronium ion, a hydrated H+.
• Neutralization involves a reaction between an acid and a base to produce a salt and water in this theory.
• Arrhenius theory does not handle non-OH- bases well (e.g., ammonia).

Bronsted-Lowry Theory

• Acids are substances that donate a proton (H+), including any H-containing molecule or ion capable of releasing a proton (e.g., HCl, NH4+, HNO3, HClO4, H2SO4, H2O).
• Bases accept a proton (H+), including any molecule or ion capable of accepting a proton (e.g., OH-, NH3, F-, H2O).
• When acids ionize in water, they transfer a proton to water, creating a hydronium ion.
• When bases ionize in water, they accept a proton from water, creating a hydroxide ion.
• Acid-base reactions involve proton transfer.
• A conjugate base of an acid is formed when the acid donates a proton.
• A conjugate acid of a base is formed when the base accepts a proton.
• For conjugate acid-base pairs, the stronger the acid, the weaker its conjugate base, and vice versa.
• Amphoteric substances can act as either an acid or a base.
• Amphiprotic substances can either donate or accept H+.

Lewis Theory

• Acids accept an electron pair.
• Bases donate an electron pair.
• It includes molecules/ions capable of accepting a proton.
• Lewis acid-base reactions include many reactions that do not involve Brønsted acids.

Strengths of Acids and Bases

• Strong acids and bases undergo complete ionization in water.
• Weak acids and bases undergo partial ionization in water.
• Strong acids: HCl, HBr, HI, HClO4, HNO3, H2SO4
• Strong bases: LiOH, NaOH, KOH, RbOH, CsOH, Mg(OH)2, Ca(OH)2, Sr(OH)2, Ba(OH)2
• Binary acid strength (H-X) increases down a group and, in general, increases across a period with increasing electronegativity of X.
• In aqueous solution, the strength of hydrohalic acids follows the order HF < HCl = HBr = HI due to the leveling effect.
• For oxoacids (H-O-X) with different X from the same group, acid strength increases with increasing electronegativity of X.
• For oxoacids having the same X but with different oxidation numbers, acid strength increases with increasing oxidation number of X.
• Organic acids contain an acidic hydrogen and carboxylic acids are the most common type.
• Alkyl chain length has little effect on acid strength.
• Attaching an EN atom increases acidity by making the acidic H easier to release
• Polyprotic acids release successive H atoms with increasing difficulty.
• Strong bases undergo complete ionization in water to give OH-.
• Weak bases do not completely ionize in water.

Autoionization of Water

• Pure water dissociates to form equal concentrations of H3O+ and OH-.
• The ion-product constant, Kw, is the product of the molar concentrations of H3O+ and OH- at a particular temperature.
• Kw = [H3O+][OH-] = 1.0 x 10-14 at 25°C
• In pure water at 25°C, [H+] = [OH-] = 1.0 x 10-7 M
• If you add either an acid or a base to water, they are no longer equal, but the Kw expression still holds, meaning that if one goes up, the other goes down to keep produce equal to Kw.

pH Measurement

• Can be changed by dissolving substances in water.
• Acidity or basicity is dictated by respective [H3O+] and [OH-].
• pH = -log[H3O+].
• pKw = pH + pOH = 14.00
• Acidic solution: [H3O+] > [OH-], pH < 7
• Neutral solution: [H3O+] = [OH-], pH = 7
• Basic solution: [H3O+] < [OH-], pH > 7
• The higher the [H3O+], the lower the pH, and equal to the amount of significant digits in the hydrogen-ion concentration.

Ionization Constants

• Weak acids and bases have ionization constants (Ka and Kb, respectively) that describe the extent of their ionization in water.
• Ka = [H3O+][A-]/[HA], Kb = [OH-][NH4+]/[NH3]
• ↑ Ka, ↑ [H3O+] eq, ↑ Acid strength, ↑ Kb,↑ [OH-]eq, ↑ Base strength
• For a monoprotic acid (% ionization = [H3O+]eq / [HA]o x 100%)
• Values are commonly listed as Ka, relationships determine conjugates.

Calculating pH

• Strong acids: dissociates in water, pH is calculated directly (i.e.: pH = -log[H+])
• pH can be calculated with: pH = 14 - (-log[OH-])
• Weak acids: dissociation is determined by its Ka value, use an ICE table
• Weak bases: dissociation is determined by its Kb value, use an ICE table
• [HA]eq must be greater than 100Ka
• If [H3O+] from the acid is less or equal to 1.00 * 10^-5 M, the pH calculation should include the H2O autoionization.

Salt Solutions

• Salts derived from a strong base and a strong acid form neutral solutions (pH = 7).
• Ions derived from strong acids/bases do not hydrolyze in water.
• Salts derived from a strong base and a weak acid form basic solutions (pH > 7)
• Na+ does not hydrolyze, but CH3COO- will.
• Salts derived from a weak base and a strong acid form acidic solutions (pH < 7).
• Cl- does not hydrolyze, but NH4+ will.
• Solutions must be neutral, or relative Ka of cation and Kb of anion must be considered.

Predicting Reaction Direction

• To determine the reaction, must compare the following components: HF(aq) + NH3(aq) ↔ F-(aq) + NH4+(aq) and its pKa.
• It is based on the principle that stronger acid/base > weaker acid/base

Buffer Solutions

• Two-component systems (weak acid/conjugate base or weak base/conjugate acid) can resists changes in pH upon addition of small amounts of acid or base.
• Components cannot self neutralize
• The Henderson-Hasselbalch equation relates the pH of a buffer solution to the pKa of the weak acid and the ratio of the concentrations of the conjugate base and acid:
• pH = pKa + log([A-]/[HA])
• General steps to prepare solutions are select a weak acid with close pKa value, calculate the conjugate ratio for pH, and calculate concentrations of conjugate base and acid.
• Prepare with components or titrate with with a base or acid.
• Capacity is the amount of acid/base neutralized and range is the pH effectively added.
• The HHE relates the pH of a buffer for different concentrations of conjugate acid and base.

Summary of Key Concepts

• Three acid-base theories are differentiated and explained and are: the Arrhenius Theory, Brønsted Lowry theory and Lewis acid-base theory.
• Solutions of salts maybe acidic or basic because of hydrolysis of ions, which the constant equals to Ka for a cation or Kb for an anion, which KaKb= Kw.
• Two important determining relative acid strengths factors are polarity of the bond of H, atomic radius, and electronegativity all relating to bond strength.