Acids and Bases

Acids and Bases: Definitions

  • Hydronium Ion: Formed when an acid reacts with water.
  • Arrhenius definition of an acid: Produces hydronium ions (H3O+H_3O^+) when dissolved in water.
  • Arrhenius definition of a base: Produces hydroxide ions (OH−OH^-) when dissolved in water.
  • Brønsted-Lowry acid: Donates a hydrogen ion (H+H^+) to another molecule or ion (proton donor).
  • Brønsted-Lowry base: Accepts a proton (H+H^+.
  • Acid-base reaction: A proton is transferred.
  • Monoprotic acids: Acids with one proton (e.g., HCl).
  • Diprotic acids: Acids with two protons to donate (e.g., H<em>2SO</em>4H<em>2SO</em>4).
  • Triprotic acids: Acids with three protons to donate (e.g., H<em>3PO</em>4H<em>3PO</em>4).
  • Acidic H atoms: Bonded to electronegative atoms.
  • Conjugate acid-base pairs: Chemical species that differ by one H+H^+.

Acid and Base Strength

  • Strong acids: Give up a proton easily and are essentially 100% dissociated in water.
  • Weak acids: Give up a proton with difficulty and are substantially less than 100% dissociated in water.
  • Strong bases: Grab and hold a proton tightly.
  • Weak bases: Have little affinity for a proton.
  • Inverse relationship between acid strength and base strength: The stronger the acid, the weaker its conjugate base.
  • Acid-base proton-transfer equilibrium: Favors reaction of the stronger acid with the stronger base, forming the weaker acid and base.

Acid Dissociation Constants

  • Acid dissociation constant (KaK_a):
    • HA(aq)+H<em>2O(l)⇌H</em>3O+(aq)+A−(aq)HA(aq) + H<em>2O(l) \rightleftharpoons H</em>3O^+(aq) + A^-(aq)
    • K=[H<em>3O+][A−][HA][H</em>2O]K = \frac{[H<em>3O^+][A^-]}{[HA][H</em>2O]}
    • K<em>a=[H</em>3O+][A−][HA]K<em>a = \frac{[H</em>3O^+][A^-]}{[HA]}
  • Strong acids: KaK_a values much greater than 1.
  • Weak acids: KaK_a values much less than 1.
  • Polyprotic acids: Donation of each successive H+H^+ is more difficult, so KaK_a values become successively lower.

Water as Both an Acid and a Base

  • Amphoteric: A substance that can react as either an acid or a base.
  • Water acts as both an acid and a base.
  • Ion-product constant for water (KwK_w): The product of the molar concentrations of hydronium and hydroxide ions in water or any aqueous solution.
    • K<em>w=[H</em>3O+][OH−]=1.0×10−14K<em>w = [H</em>3O^+][OH^-]= 1.0 \times 10^{-14}
  • Acidic, neutral, or basic solutions are identified according to their ion concentrations.

Measuring Acidity in an Aqueous Solution: The pH Scale

  • pH: A measure of the acid strength of a solution.
  • pH is a number, usually between 0 and 14, that indicates the H3O+H_3O^+ concentration of the solution.
  • pH=−log⁡[H3O+]pH = -\log[H_3O^+]
  • [H3O+]=10−pH[H_3O^+] = 10^{-pH}
  • A change of 1 pH unit means a tenfold change in [H3O+][H_3O^+].
  • Acid-base indicator: A dye that changes color depending on the pH of the solution.

Working with pH

  • Converting from pH to [H3O+][H_3O^+] requires finding the antilogarithm of the negative pH.
  • Significant figures: An antilogarithm contains the same number of significant figures as the original number has to the right of the decimal point

Acid and Base Equivalents

  • One equivalent of any acid neutralizes one equivalent of any base.
  • Equivalent of acid: The amount of an acid that contains 1 mole of H+H^+ ions.
  • Equivalent of base: The amount of base that contains 1 mole of OH−OH^- ions.
  • Normality (N) of an acid or base solution: The number of equivalents of acid or base per liter of solution.
  • For monoprotic acids, molarity and normality are the same.
  • Normality = Molarity × number of H+H^+ or OH−OH^- ions produced per formula unit.

Some Common Acid-Base Reactions

  • Acids with hydroxide ion: One equivalent of an acid reacts with 1 Eq of a metal hydroxide to yield water and a salt in a neutralization reaction.
  • Acids with bicarbonate and carbonate ion: The bicarbonate ion and carbonate ion react with acid by accepting H+H^+ to yield carbonic acid, which decomposes to carbon dioxide gas and water.
  • Acids with Ammonia: Acids react with ammonia to yield water-soluble ammonium salts.

Acidity and Basicity of Salt Solutions

  • Salt solutions can be neutral, acidic, or basic, depending on the ions present.
  • Salts are classified according to the acid and base from which they are formed in a neutralization reaction.
  • Some ions react with water to produce H3O+H_3O^+ and some ions react with water to produce OH−OH^-.

Buffer Solutions

  • Buffer: A combination of substances that act together to prevent a drastic change in pH; usually a weak acid and its conjugate base.
  • Henderson-Hasselbalch equation: pH=pKa+log⁡[A−][HA]pH = pK_a + \log \frac{[A^-]}{[HA]}

Titration

  • Titration: A procedure for determining the total acid or base concentration of a solution.
  • Total acid concentration gives the sum of dissociated plus undissociated acid and can be determined by carrying out a titration.
  • When the titration involves a neutralization reaction in which one mole of acid reacts with one mole of base:
    M<em>acidV</em>acid=M<em>baseV</em>baseM<em>{acid}V</em>{acid} = M<em>{base}V</em>{base}
  • When the coefficients for the acid and base in the balanced neutralization reaction are not the same, then we can use equivalents of acid and base instead of moles, and normality instead of molarity:
    N<em>acidV</em>acid=N<em>baseV</em>baseN<em>{acid}V</em>{acid} = N<em>{base}V</em>{base}