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