Chapter 8 - Acids, Bases, and pH

Chapter 8.1: Explaining the Properties of Acids and Bases

• The Arrhenius Theory of Acids and Bases:
  • An acid is a substance that dissociates in water to form H+
  • A base is a substance that dissociates in water to form OH−
• Hydronium ion, H3O+(aq): hydrated proton
  • Makes hydrogen bonds with water
• Brønsted-Lowry theory:
  • Acid is a substance a proton can be removed (proton-donors)
  • The base is a substance that accepts a proton (proton-acceptors)
• Conjugate acid-base pair: substances that differ with a proton
  • Conjugate acid: the proton is added to the base
  • Conjugate base: the proton is removed from the acid
• Amphoteric: a substance that can act as a proton donor in one reaction and a base in another
• Strong acids:
  • Hydrohalic acids: (HCl, HBr,) have hydrogen-bonded to atoms
  • Oxoacids: number of oxygen atoms by two or more
  • Increase in strength with more oxygen atoms
• Monoprotic acids: single hydrogen atoms that dissociate in water
• Polyprotic acids: more than one hydrogen atom that dissociates
• Strong bases**:**
  • Oxides and hydroxide of alkali metals
  • Alkaline earth metal oxides and hydroxides below beryllium
• Strong acids and bases (strong electrolytes dissociate in water
  • Can’t find a concentration of weak acids and bases

Chapter 8.2: The Equilibrium of Weak Acids and Bases

• Equilibrium constant Kc for dissociation of water:
  • Kc = ( [H3O+][OH−] ) / [H2O]2
  • Kc [H2O]2 = product of concentration of hydronium ions and hydroxide ions
• Ion product constant for water, Kw: product of [H3O+][OH−]
  • Equal to 1.0 x 10-14 mol/L
• Acids solution: mol/L
  • [H3O+] > 1.0 x 10-7 mol/L
  • [OH−] < 1.0 x 10-7 mol/L
• Neutral solution:
  • [H3O+] = 1.0 x 10-7 mol/L
  • [OH−] = 1.0 x 10-7 mol/L
• Basic solution:
  • [H3O+] < 1.0 x 10-7 mol/L
  • [OH−] > 1.0 x 10-7 mol/L
• pH: exponential power of hydrogen/ hydronium ions, in moles per litre
  • pH = -log[H3O+]
• pOH: power of hydroxide ions of a solution from the [OH−]
  • pOH = −log[OH−]
• Kw = [H3O+][OH−] = 1.0 x 10-14 mol/L at 25 ̊C
• ∴ pH + pOH = 14
• [H3O] = 10−pH
• [OH−] = 10−pOH
• Acid dissociation constant, Ka: acid ionization constant, measures the strength of the acid
  • Kc[H2O] = Ka = ( [H3O+][A−] ) / [HA]
• Percent dissociation: fraction of acid molecules that dissociate compared with the initial concentration of the acid, expressed as a percent depends on Ka and initial concentration of weak acid
• Polyprotic acids pH calculated for the first dissociation are used in the second dissociation and again for as many steps required

Chapter 8.3: Bases and Buffers

• Base dissociation constant, Kb: base ionization constant, measures strength of base Kb = ( [[HB+][OH−] ) / [B]
• Buffer solution: a solution that contains a weak acid/conjugate base mixture or a weak base/conjugate acid mixture made by:Using weak acid and its salting weak base and its salt Characteristics: Unchanging pHBuffer capacity: the amount of acid or base that can be added before the considerable change occurs to the pH

Chapter 8.4: Acid-Base Titration Curves

• Acid-base titration curve: a graph of the pH of an acid (or base) against the volume of an added base
• Titrations: analytical procedures, usually to find:
  • Equivalence point: point in a titration when the acid and base that are present completely react with each other
  • Can help calculate unknown concentrations
  • Indicators can be used to find endpoints close to the equivalence point
  • Strong acid with strong acid
  • Equivalence of pH of 7
  • Weak acid and strong base
  • Equivalence of pH of above 7
  • Weak base with strong acid
  • Equivalence of pH below 7

\