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: $Ba + H<em>2SO</em>4 \rightarrow BaSO<em>4 + H</em>2$

• 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., $H<em>3PO</em>4$: 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: $H<em>3PO</em>4$ (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, $HClO<em>4$, $H</em>2SO<em>4$, $HNO</em>3$

  • Weak acids: $HSO<em>4$, HF, $CH</em>3COOH$, $H<em>2CO</em>3$, $H_2S$, HCN

• $H_3O^+$: hydronium ion

  • $HCl + H<em>2O \rightarrow H</em>3O^+ + Cl^-$

  • $HF + H<em>2O \rightarrow H</em>3O^+ + F^-$

  • $CH<em>3COOH + H</em>2O \rightarrow H<em>3O^+ + CH</em>3COO^-$

  • $H<em>2CO</em>3 + H<em>2O \rightarrow H</em>3O^+ + HCO_3^-$

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: $H<em>2CO</em>3$

• Acetic acid: $CH_3COOH$

• Sulfuric acid: $H<em>2SO</em>4$

• 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

• 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^-$</p><ul><li><p>HCl (Acid),$NH3$(Base),$NH4^+$(Conjugate Acid),$Cl^-$(Conjugate Base)</p></li></ul></li></ul></li></ul><h5 id="9f540d14-694f-4099-b380-34170c649fe8" data-toc-id="9f540d14-694f-4099-b380-34170c649fe8" collapsed="false" seolevelmigrated="true">Conjugate Acids and Bases</h5><ul><li><p>Conjugate Acid: A substance that accepted a proton AND can now act as a proton donor in the reverse reaction.</p></li><li><p>Conjugate Base: A substance that donated a proton AND can now act as a proton acceptor in the reverse reaction.</p></li></ul><h5 id="df40b50e-3c09-481d-9390-538476cd6786" data-toc-id="df40b50e-3c09-481d-9390-538476cd6786" collapsed="false" seolevelmigrated="true">Brønsted-Lowry Practice</h5><ul><li><p>Label the following substances as the acid (A), base (B), conjugate acid (CA), or conjugate base (CB):</p><ul><li><p>$OH^- + HCN \rightleftharpoons H_2O + CN^-$</p><ul><li><p>B A CA CB</p></li></ul></li><li><p>$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$)</p><ul><li><p>Water can act as a base:$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