Acids, Bases, and Salts Notes

Acids, Bases, and Salts

  • Acids have a lower pH.
  • Bases (alkaline) have a higher pH.
  • Neutral substances have a pH around 7.

Examples of Substances with Different pH Levels:

  • Hydrochloric acid
  • Vinegar
  • Baking Soda
  • Hand Soap
  • Ammonia
  • Bleach

Naming Rules for Acids

  • Binary Acids: Use the prefix "hydro-" followed by the nonmetal name with the suffix "-ic acid."
  • Ternary Acids: Refer to Table E for polyatomic ions.
    • If the polyatomic ion ends in "-ate", change the ending to "-ic."
    • If the polyatomic ion ends in "-ite", change the ending to "-ous."

Common Acids and Bases (Tables K and L)

  • Table K: Common Acids
    • HCl(aq)HCl(aq): hydrochloric acid
    • HNO3(aq)HNO3(aq): nitric acid
    • H2SO4(aq)H2SO4(aq): sulfuric acid
    • H3PO4(aq)H3PO4(aq): phosphoric acid
    • H2CO3(aq)H2CO3(aq) or CO2(aq)CO2(aq): carbonic acid
    • CH3COOH(aq)CH3COOH(aq) or HC2H3O2(aq)HC2H3O2(aq): ethanoic acid (acetic acid)
  • Table L: Common Bases
    • NaOH(aq)NaOH(aq): sodium hydroxide
    • KOH(aq)KOH(aq): potassium hydroxide
    • Ca(OH)2(aq)Ca(OH)2(aq): calcium hydroxide
    • NH3(aq)NH3(aq): aqueous ammonia

Properties of Acids

  • Taste sour.
  • Are electrolytes, meaning they produce ions in solution and conduct electricity.
    • Strong acids are strong electrolytes.
    • Weak acids are weak electrolytes.
  • React with active metals (above H2 in the activity series) to produce hydrogen gas (H2).
    • Example: Mg+2HCl(aq)→H2(g)+MgCl2(aq)Mg + 2HCl (aq) \rightarrow H2 (g) + MgCl2 (aq)

Properties of Bases

  • Strong bases consist of Group 1 or Group 2 metals with hydroxide ions.
    • Examples: NaOH(aq)NaOH (aq), LiOH(aq)LiOH (aq), Ca(OH)2(aq)Ca(OH)2 (aq)
  • Taste bitter and feel slippery.
  • Also electrolytes.
    • Strong bases are strong electrolytes.
    • Weak bases are weak electrolytes.
  • Do not react with active metals.

Indicators (Table M)

  • Indicators change color depending on the pH of the solution.
  • Below the indicator's range, the color corresponds to the left side of Table M.
  • Above the indicator's range, the color corresponds to the right side of Table M.

Electrolytes, Acids, Bases, and Salts

  • Electrolytes: Form ions in water and conduct electricity.
  • Salts: Ionic compounds formed from a metal and a nonmetal.
  • Acids:
    • pH < 7
    • H^+ > OH^-
  • Bases:
    • pH > 7
    • OH^- > H^+
  • Molecules (nonmetal to nonmetal) are generally not electrolytes because they do not produce ions in solution.

Water as an Acid/Base

  • Water molecules undergo self-ionization:
    • H2O(l)+H2O(l)⇌H3O+(aq)+OH−(aq)H 2O(l) + H2O (l) \rightleftharpoons H3O+ (aq) + OH- (aq)
  • In pure water, the equilibrium concentration of hydronium ions and hydroxide ions is:
    • [H3O+]=[OH−]=1×10−7M[H3O+] = [OH-] = 1 \times 10^{-7} M
  • The number of moles of H+H^+ ions equals the number of moles of OH−OH^- ions.

Ion-Product Constant for Water

  • Kw=[H+]×[OH−]=1.0×10−14K w= [H+] \times [OH-] = 1.0 \times 10^{-14}
  • The exponents of [H+][H^+] and [OH−][OH^-] concentrations always add up to 14.

The pH Scale

  • pH < 7: acidic
  • pH = 7: neutral
  • pH > 7: basic
  • pH=−log[H+]pH = -log[H^+]
  • pH scale is logarithmic, based on the hydrogen ion concentration in the aqueous solution.

pH Scale Changes

  • An increase of one unit on the pH scale represents a ten-fold (x10) decrease in the H+H^+ ion concentration.

Practice Problems: pH and H+H^+ concentration

  • pH 6 to pH 4: [H+][H^+] increases by a factor of 100.
  • pH 5 to pH 6: [H+][H^+] decreases by a factor of 10.
  • pH 2 to pH 4: [H+][H^+] decreases by a factor of 100.

Neutralization Reactions

  • Occur when an acid and a base react to form a salt and water.
    • Acid + Base → Salt + Water
    • Example: HCl(aq)+NaOH(aq)→NaCl(aq)+H2O(l)HCl (aq) + NaOH (aq) \rightarrow NaCl (aq) + H2O (l)
      • Acid: HCl
      • Base: NaOH
      • Salt: NaCl
      • Water: H2O
  • Net ionic reaction: H+(aq)+OH−(aq)→H2O(l)H^+ (aq) + OH^- (aq) \rightarrow H2O (l)

Titrations

  • Used to determine the concentration of an acid or base.
  • The equivalence point is reached when the moles of H+H^+ equal the moles of OH−OH^-.
  • Titration Equation:
    • MaVa=MbVbM a V a = MbVb
      • MaM a = Molarity of the acid
      • VaV a = Volume of the acid
      • MbMb = Molarity of the base
      • VbVb = Volume of the base
  • Any unit of volume can be used as long as consistency is maintained.

Titration Setup

  • Burette containing titrant of known concentration.
  • Titration flask containing the solution of unknown concentration.
  • Clamp stand to hold the burette.

Indicators for Titrations

  • Phenolphthalein (phth) is commonly used to approximate the endpoint.
  • Colorless in acidic solutions and pink in basic solutions.
  • The endpoint is reached when a pale pink color persists.
  • The endpoint, or equivalence point, occurs when moles H+H^+ = moles OH−OH^- (pH = 7).

Arrhenius Theory of Acids and Bases

  • Arrhenius Acid: A substance that, when dissolved in water, produces H+H^+ ions as the only positive ions in solution.
    • Example: HCl(aq)→H+(aq)+Cl−(aq)HCl (aq) \rightarrow H^+ (aq)+ Cl^- (aq)
  • Arrhenius Base: A substance that, when dissolved in water, produces OH−OH^- ions as the only negative ions in solution.
    • Example: NaOH(aq)→Na+(aq)+OH−(aq)NaOH (aq) \rightarrow Na^+ (aq) + OH^- (aq)

Bronsted-Lowry Theory of Acids and Bases

  • Based on the proton (H+H^+
  • Bronsted Acid: A proton donor; must release an H+(aq)H^+ (aq) ion.
    • Example: HCl(aq)HCl (aq)
  • Bronsted Base: A proton acceptor; does not need to have the OH−(aq)OH^-(aq) ion but must have a lone pair of electrons.
    • Example: NH3(aq)NH3 (aq)

Reversible Acid-Base Reactions

  • Acid-base reactions can be reversible.
    • H2SO4(aq)+NH3(aq)⇌HSO4−(aq)+NH4+(aq)H 2SO4(aq) + NH3(aq) \rightleftharpoons HSO4 -(aq) + NH4 +(aq)
      • H2SO4(aq)H 2SO4(aq) (B.A.)
      • NH3(aq)NH3(aq) (B.B.)
      • HSO4−(aq)HSO4 -(aq) (B.B.)
      • NH4+(aq)NH4 +(aq) (B.A.)
  • A conjugate acid-base pair consists of two substances related by the loss or gain of a single hydrogen ion (H+H^+
  • Strong acids have weak conjugate bases.
  • Strong bases have weak conjugate acids.

Amphoteric Substances

  • A substance that can act as both an acid and a base.
  • Example: HCO3−(aq)+H2O(l)⇌CO32−(aq)+H3O+(aq)HCO3 - (aq) + H2O (l) \rightleftharpoons CO3 2- (aq) + H3O+(aq)
    • Here, HCO3−HCO3 - acts as an acid.
  • Example: HCO3−(aq)+H2O(l)⇌H2CO3(aq)+OH−(aq)HCO3 - (aq) + H2O (l) \rightleftharpoons H2CO3 (aq) + OH- (aq)
    • Here, HCO3−HCO3 - acts as a base.
  • Water (H2O) is a common amphoteric substance.