Acids and Bases Notes

Acids & Bases

Acids

  • Aqueous solutions of acids have a sour taste.

  • Some acids react with active metals and release hydrogen gas, H_2.

    • Example: Ba + H2SO4 \rightarrow BaSO4 + H2

  • Acids turn litmus paper red.

  • Acids are electrolytes (conduct electricity when dissolved in a solution).

Bases

  • Aqueous solutions of bases taste bitter.

  • Dilute aqueous solutions of bases feel slippery.

  • Bases conduct electric current and are electrolytes.

  • Bases turn litmus paper blue.

  • Sometimes called alkaline solutions.

Acid Nomenclature

  • The acid name depends on the anion's normal ending:

    • -ide becomes hydro- -ic acid (e.g., HF: hydrogen monofluoride becomes hydrofluoric acid).

    • -ite becomes -ous acid (e.g., HNO_2: hydrogen nitrite becomes nitrous acid).

    • -ate becomes -ic acid (e.g., H3PO4: hydrogen phosphate becomes phosphoric acid).

  • For sulfur-containing oxyacids, the base name is sulfur instead of just S.

  • For phosphorus-containing oxyacids, the base name is phosphor instead of simply phosphe.

Binary Acids

  • Named as Hydro__ic Acid

    • Examples: HCl, H_2S, HBr, HI, HF

Oxy Acids

  • Polyatomic Ions

    • Examples: H3PO4 (Phosphoric Acid)

  • The "-ite" or "-ate" ending depends on the number or charge of Polyatomic ion

Acid Strengths

  • A strong acid or strong base is one that ionizes completely in aqueous solution (strong electrolyte).

  • A weak acid or weak base releases few hydrogen ions in aqueous solution (weak electrolyte).

  • Examples:

    • Strong acids: HCl, HBr, HI, HClO4, H2SO4, HNO3

    • Weak acids: HSO4, HF, CH3COOH, H2CO3, H_2S, HCN

  • H_3O^+: hydronium ion

    • HCl + H2O \rightarrow H3O^+ + Cl^-

    • HF + H2O \rightarrow H3O^+ + F^-

    • CH3COOH + H2O \rightarrow H3O^+ + CH3COO^-

    • H2CO3 + H2O \rightarrow H3O^+ + HCO_3^-

Strength of Bases

  • The strength of a base depends on the extent to which the base dissociates.

  • Strong bases are strong electrolytes.

  • Group 1 and 2 metal hydroxides are strong bases.

    • Examples: NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

    • Ca(OH)_2 \rightarrow Ca^{2+} + 2OH^-

Formula Practice

  • Hydroiodic acid: HI

  • Hydrochloric acid: HCl

  • Nitric acid: HNO_3

  • Carbonic acid: H2CO3

  • Acetic acid: CH_3COOH

  • Sulfuric acid: H2SO4

  • Iron (II) hydroxide: Fe(OH)_2

    • Roman numeral tells us charge

  • Sodium hydroxide: NaOH

  • Ammonium hydroxide: NH_4OH

  • Calcium hydroxide: Ca(OH)_2

Summary of Acid-Base Theories

Type

Acid

Base

Arrhenius

H^+ or H_3O^+ producer

OH^- producer

Brønsted-Lowry

Proton (H^+$) donor

Proton (H^+$) acceptor

Lewis

Electron-pair acceptor

Electron-pair donor

  • Main focus: Arrhenius, Brønsted-Lowry

Arrhenius Acids and Bases

  • An Arrhenius acid is a chemical compound that increases the concentration of hydrogen ions, H^+, in solution. [H^+]

  • An Arrhenius base is a substance that increases the concentration of hydroxide ions, OH^-, in solution. [OH^-]

  • What we typically think of when we think of acids & bases

Brønsted-Lowry Acids and Bases

  • A Brønsted-Lowry acid is a molecule or ion that is a proton donor (H^+$).

  • A Brønsted-Lowry base is a molecule or ion that is a proton acceptor.

  • Protons move from the acid to the base

  • Hydrogen chloride acts as a Brønsted-Lowry acid when it reacts with ammonia.

  • Ammonia accepts a proton from the hydrochloric acid and acts as a Brønsted-Lowry base.

    • Example: HCl + NH3 \rightleftharpoons NH4^+ + Cl^-

      • HCl (Acid), NH3 (Base), NH4^+ (Conjugate Acid), Cl^- (Conjugate Base)

Conjugate Acids and Bases

  • Conjugate Acid: A substance that accepted a proton AND can now act as a proton donor in the reverse reaction.

  • Conjugate Base: A substance that donated a proton AND can now act as a proton acceptor in the reverse reaction.

Brønsted-Lowry Practice

  • Label the following substances as the acid (A), base (B), conjugate acid (CA), or conjugate base (CB):

    • OH^- + HCN \rightleftharpoons H_2O + CN^-

      • B A CA CB

    • HCO3^- + OH^- \rightleftharpoons H2O + CO_3^{2-}

      • A B CA CB

  • Miss Neal recommends starting with identifying the acid first, then figure out the rest

Lewis Acids

  • A Lewis acid is an electron pair acceptor.

  • A Lewis base is an electron pair donor.

  • The Lewis definition is the broadest of the three acid definitions.

  • A bare proton (hydrogen ion) is a Lewis acid

    • Example: Lewis acid (acceptor) + Lewis base (donor) --> Lewis wide

Amphoteric Compounds

  • Any species that can react as either an acid or a base is described as amphoteric.

  • Example: water (H_2O)

    • Water can act as a base: H2SO4(aq) + H2O(l) \rightarrow H3O^+(aq) + HSO_4^-(aq)

      • Acid + Base --> Conjugate Acid + Conjugate Base

    • Water can act as an acid: NH3(g) + H2O(l) \rightarrow NH_4^+(aq) + OH^-(aq)

      • Base + Acid --> Conjugate Acid + Conjugate Base

  • H_2O can accept & lose H^+$$