Premed · Premed · General Chemistry 2
Lecture 5: Acid-Base Chemistry: Definitions and Strength
General Chemistry II
Learning Objectives
By the end of this lecture, students will be able to:
- Define acids and bases according to Arrhenius, Bronsted-Lowry, and Lewis theories
- Identify conjugate acid-base pairs
- Distinguish between strong and weak acids and bases
- Relate acid/base strength to molecular structure
- Describe the autoionization of water and the ion-product constant Kw
- Classify salts as acidic, basic, or neutral based on their ions
Lecture Content
I. Arrhenius Definition
The Arrhenius definition, the earliest of the three major acid-base frameworks, defines an acid as a substance that produces H+ (or H3O+) ions in aqueous solution and a base as a substance that produces OH- ions in aqueous solution. For example, HCl dissociates in water to yield H+ and Cl-, while NaOH dissociates to give Na+ and OH-. Although straightforward, this definition has significant limitations: it applies only to aqueous solutions and cannot account for bases like ammonia (NH3) that do not contain hydroxide ions.
II. Bronsted-Lowry Definition
The Bronsted-Lowry definition is more general, defining an acid as a proton (H+) donor and a base as a proton acceptor. This framework extends beyond aqueous solutions to include reactions in non-aqueous solvents and even in the gas phase. Central to this definition is the concept of conjugate acid-base pairs. When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid. In the reaction HF + H2O <=> F- + H3O+, HF acts as the acid and F- is its conjugate base, while H2O acts as the base and H3O+ is its conjugate acid.
Some substances are amphoteric (or amphiprotic), meaning they can function as either an acid or a base depending on the reaction partner. Water is the classic example, capable of either donating or accepting a proton. Bicarbonate (HCO3-) is another important amphiprotic species.
III. Lewis Definition
The Lewis definition is the most expansive of the three. A Lewis acid is an electron pair acceptor, while a Lewis base is an electron pair donor. This definition encompasses all Bronsted-Lowry acids and bases while also including species that lack protons entirely. Lewis acids include metal cations such as Fe3+ and Al3+, molecules with incomplete octets like BF3 and AlCl3, and molecules with polar multiple bonds such as CO2 and SO3. Lewis bases include any species with available lone pairs, such as NH3, H2O, and OH-, as well as anions like Cl- and F-. The reaction of BF3 with NH3 to form the adduct F3B-NH3, in which BF3 accepts a lone pair from the nitrogen, is a prototypical Lewis acid-base reaction that falls outside the scope of both the Arrhenius and Bronsted-Lowry definitions.
<image>A Venn diagram showing the three acid-base definitions as nested circles. The outermost circle is labeled "Lewis acids/bases (electron pair acceptors/donors)" and includes examples like BF3 and AlCl3. The middle circle is labeled "Bronsted-Lowry acids/bases (proton donors/acceptors)" and includes examples like NH4+ and NH3. The innermost circle is labeled "Arrhenius acids/bases (produce H+/OH- in water)" and includes examples like HCl and NaOH. Arrows point to species that fit one definition but not the next inner one.</image>
IV. Autoionization of Water and Kw
Water undergoes autoionization, a process in which two water molecules react with each other: H2O + H2O <=> H3O+ + OH- (or, in simplified form, H2O <=> H+ + OH-). The equilibrium constant for this process is the ion-product constant of water, Kw = [H3O+][OH-] = 1.0 x 10^-14 at 25 degrees C. In pure water, the concentrations of H3O+ and OH- are equal, each at 1.0 x 10^-7 M.
The Kw expression applies to all aqueous solutions, not just pure water. If the concentration of H3O+ increases, the concentration of OH- must decrease proportionally to maintain the product Kw, and vice versa. This relationship provides the basis for classifying solutions: an acidic solution has [H3O+] > [OH-] (or [H3O+] > 1.0 x 10^-7 M), a basic solution has [OH-] > [H3O+], and a neutral solution has [H3O+] = [OH-] = 1.0 x 10^-7 M.
V. Strong Acids and Strong Bases
Strong acids ionize completely in aqueous solution, meaning that every molecule of acid donates its proton to water. The seven common strong acids to memorize are HCl (hydrochloric acid), HBr (hydrobromic acid), HI (hydroiodic acid), HNO3 (nitric acid), H2SO4 (sulfuric acid, first dissociation only), HClO4 (perchloric acid), and HClO3 (chloric acid).
Strong bases dissociate completely as well. These include the Group 1A hydroxides (LiOH, NaOH, KOH, RbOH, CsOH) and the heavier Group 2A hydroxides (Ca(OH)2, Sr(OH)2, Ba(OH)2). For a strong monoprotic acid, [H3O+] equals the initial acid concentration. For a strong monohydroxide base, [OH-] equals the initial base concentration, while for a dihydroxide base, [OH-] equals twice the base concentration.
VI. Weak Acids and Weak Bases
Weak acids ionize only partially in aqueous solution, establishing an equilibrium that lies to the left: HA + H2O <=> H3O+ + A-. The extent of ionization is quantified by the acid dissociation constant, Ka = [H3O+][A-] / [HA]. Similarly, weak bases react only partially with water: B + H2O <=> BH+ + OH-, with the base dissociation constant Kb = [BH+][OH-] / [B]. A larger Ka indicates a stronger weak acid, and a larger Kb indicates a stronger weak base.
A fundamental relationship connects conjugate acid-base pairs: Ka x Kb = Kw = 1.0 x 10^-14 at 25 degrees C. This means that a strong acid necessarily has a weak conjugate base, and vice versa. The stronger an acid, the weaker its conjugate base must be.
<image>A horizontal arrow diagram showing the relationship between acid strength and conjugate base strength. The top arrow spans left to right and is labeled "Increasing acid strength" with strong acids (HCl, HNO3) on the right and weak acids (CH3COOH, HF, HCN) in the middle. Below it, a reversed arrow spans right to left labeled "Increasing conjugate base strength" showing that Cl- and NO3- are very weak bases while CH3COO-, F-, and CN- are stronger conjugate bases. Ka values are listed beneath each acid, and Kb values beneath each conjugate base, demonstrating Ka x Kb = Kw.</image>
VII. Molecular Structure and Acid Strength
The strength of an acid is determined by how easily it releases a proton, which in turn depends on molecular structure. For binary acids of the form HX, acid strength increases going down a group in the periodic table because bond strength decreases (making it easier to lose H+): HF < HCl < HBr < HI. Across a period, acid strength increases with increasing polarity of the H-X bond: CH4 < NH3 < H2O < HF.
For oxyacids, which have the general structure H-O-Y, two structural factors govern acidity. When comparing oxyacids with the same number of oxygen atoms, increasing electronegativity of the central atom Y increases acid strength: HOCl > HOBr > HOI. When comparing oxyacids of the same central atom, increasing the number of terminal oxygen atoms (those not bonded to H) increases acid strength dramatically: HClO4 > HClO3 > HClO2 > HClO. The additional electronegative oxygen atoms withdraw electron density from the O-H bond, polarizing it and making it easier for the proton to leave. In carboxylic acids (R-COOH), electron-withdrawing substituent groups attached to R similarly increase acidity.
<image>A structural diagram comparing oxyacid strength for the chlorine oxyacid series. Four molecular structures are shown: HClO (hypochlorous acid, Ka = 3.0 x 10^-8), HClO2 (chlorous acid, Ka = 1.1 x 10^-2), HClO3 (chloric acid, strong), and HClO4 (perchloric acid, strong). Each structure shows the Cl atom bonded to increasing numbers of terminal O atoms (0, 1, 2, 3 respectively). Red arrows indicate electron-withdrawing effects of the additional oxygen atoms on the O-H bond. An arrow beneath all four indicates "Increasing acid strength" from left to right.</image>
VIII. Acid-Base Properties of Salts
Salts are ionic compounds formed from acid-base neutralization reactions, and the pH of a salt solution depends on whether its constituent ions hydrolyze (react with water). The conjugate base of a strong acid (such as Cl-, NO3-, or Br-) does not hydrolyze and is neutral. Similarly, the conjugate acid of a strong base (such as Na+, K+, or Ca2+) does not hydrolyze. The conjugate base of a weak acid (such as CH3COO-, F-, or CN-) hydrolyzes to produce OH-, making the solution basic. The conjugate acid of a weak base (such as NH4+) hydrolyzes to produce H3O+, making the solution acidic. Small, highly charged metal cations like Al3+ and Fe3+ also hydrolyze to produce acidic solutions by polarizing their coordinated water molecules.
Predicting the pH of salt solutions follows from these principles. NaCl is neutral because it comes from a strong acid and a strong base. Sodium acetate (NaCH3COO) is basic because the acetate ion is the conjugate base of a weak acid. Ammonium chloride (NH4Cl) is acidic because the ammonium ion is the conjugate acid of a weak base. For salts like ammonium acetate (NH4CH3COO), where both ions hydrolyze, you must compare Ka of the cation with Kb of the anion to determine whether the solution is acidic, basic, or neutral.
