Chapter 16 Notes Chem 112

Acid-Base Definitions

  • Arrhenius Definitions

    • Acid: Substance that increases the concentration of hydrogen ions (H+) in water.

    • Base: Substance that increases the concentration of hydroxide ions (OH-) in water.

  • Brønsted–Lowry Definitions

    • Acid: Proton donor.

    • Base: Proton acceptor.

  • Lewis Definitions

    • Acid: Electron pair acceptor.

    • Base: Electron pair donor.

    • All Brønsted–Lowry acids/bases are also Lewis acids/bases, but not vice versa.

Water as a Proton Acceptor

  • In water, hydrogen ions (H+) don’t exist freely but form bonds, typically with water molecules.

Brønsted–Lowry Acid and Base

  • A Brønsted–Lowry acid must have a removable proton (H+) to donate.

  • A Brønsted–Lowry base must have a nonbonding pair of electrons to accept a proton.

Characteristics of Water

  • Water can act both as a Brønsted–Lowry base (accepts a proton) and an acid (donates a proton).

  • This property makes water amphiprotic.

Conjugate Acids and Bases

  • Conjugate acid-base pairs differ by a single proton (H+).

  • Reactions between acids and bases yield their respective conjugate bases and acids.

Writing Equations for Proton-Transfer Reactions

  • Example: Lithium oxide (Li2O) in water produces hydroxide ions, demonstrating a basic reaction.

Relative Strengths of Acids and Bases

  • Stronger acids donate protons more readily, leading to weaker conjugate bases.

  • Stronger bases accept protons better, leading to weaker conjugate acids.

Classifications of Acids & Bases

  • Strong Acids: Completely dissociate in water, leaving no undissociated molecules.

  • Weak Acids: Partially dissociate in water, existing in equilibrium with their undissociated forms.

  • Negligibly Acidic Substances: Contain hydrogen but do not show acidic behaviors in water.

Autoionization of Water

  • Water acts both as an acid and a base in an equilibrium known as autoionization, where pure water contains some molecules acting as acids and some as bases.

Ion Product Constant

  • The ion product constant for water (Kw) describes the equilibrium expression for water autoionization:

    • Kw = [H+][OH-]

  • At 25 °C, Kw = 1.0 x 10^-14.

pH Concept

  • pH is a logarithmic scale for hydrogen ion concentration:

    • Neutral water has a pH of 7.00

    • Acidic solutions have pH < 7.00

    • Basic solutions have pH > 7.00

Measuring pH

  • pH meters measure pH accurately through voltage changes; indicators provide quick but less accurate measurements.

Strong and Weak Acids

  • Strong acids: Completely dissociate in solution, include common acids (HCl, H2SO4).

  • Weak acids: Partially dissociate; examples include acetic acid and formic acid.

Polyprotic Acids

  • Polyprotic acids can donate more than one proton, with the first proton being removed more easily than successive protons.

  • The first dissociation governs the pH if the dissociation constants differ by a factor of three or more.

Acid-Base Properties of Salts

  • The acidic or basic nature of a salt solution is determined by the properties of its constituent ions.

    • Anions of strong acids are neutral, while anions of weak acids can create basic solutions.

    • Cations from strong bases are neutral; cations from weak bases can be acidic.

Factors Affecting Acid Strength

  1. Bond Polarization: Increased polarity of H–A bonds generally results in stronger acids.

  2. Bond Strength: Weaker bonds tend to break more easily, yielding stronger acids.

  3. Conjugate Base Stability: More stable conjugate bases result in stronger acids.

Binary and Oxyacids

  • Binary Acids: H with one other element; bond strength and polarity dictate acid strength.

  • Oxyacids: Contain H, O, and another nonmetal; acidity typically increases with the electronegativity of the nonmetal and the number of O atoms.