Notes on Acids and Bases

Acids and Bases Overview

  • Acids and bases are fundamental concepts in chemistry, play vital roles in various reactions and processes.

Definitions

  • Acid: A hydrogen-containing substance that is able to donate a proton (H+) to another substance.

  • Base: A molecule or ion that accepts a hydrogen ion (H+) from an acid.

Properties of Acids

  • Taste: Sour

  • Feel: Wet

  • pH Range: Less than 7 (0-6)

  • Litmus Test: Turns litmus paper red

  • Electrical Conductivity: Can conduct electricity

  • Reactivity with Metals: Reacts with metals to produce hydrogen gas (H2)

  • Ion Presence: Contains H+ ions and donates hydrogen ions in reactions.

Properties of Bases

  • Taste: Bitter

  • Feel: Slippery

  • pH Range: Greater than 7 (8-14)

  • Litmus Test: Turns litmus paper blue

  • Electrical Conductivity: Can conduct electricity

  • Ion Presence: Contains hydroxide ions (OH-) and accepts hydrogen ions in reactions.

Strong vs Weak Acids

  • Strong Acids:

    • Completely ionize in water

    • Produce maximum H+ ions

    • Good conductors of electricity

  • Weak Acids:

    • Partially ionize; most molecules remain intact

    • Example: Acetic acid (HC2H3O2)

    • Poor conductors of electricity

Strong vs Weak Bases

  • Strong Bases:

    • Dissociate completely in water into metal ions and hydroxide ions

    • Examples: Sodium hydroxide (NaOH), Lithium hydroxide (LiOH)

  • Weak Bases:

    • Partially ionize in water

    • Example: Ammonia (NH3)

Ionization in Water

  • Water molecules can transfer hydrogen ions between each other, producing:

    • Hydroxide Ion (OH-): Water molecule losing H+

    • Hydronium Ion (H3O+): Water molecule gaining H+

  • Balance exists between H+ and OH- determines solution's acidity, basicity, or neutrality.

    • Pure water (pH 7) has equal H+ and OH-.

Aqueous Solutions

  • Acidic Solutions: More H+ ions (pH 0-6)

  • Basic Solutions: More OH- ions (pH 8-14)

  • Neutral Solutions: Equal numbers of H+ and OH- ions (pH 7)

Naming Acids

  • All acids produce H+ when dissolved in water; formulas begin with H.

  • Naming Rules:

    • For -ate ions: Change -ate to -ic (e.g., H2SO4 is sulfuric acid)

    • For -ite ions: Change -ite to -ous (e.g., H2SO3 is sulfurous acid)

    • For halogen acids: Add prefix hydro and change ending to -ic (e.g., HCl is hydrochloric acid).

Common Acids and Bases

  • Common Acids:

    • HI: Hydroiodic acid

    • HBr: Hydrobromic acid

    • HF: Hydrofluoric acid

    • H2SO4: Sulfuric acid

    • H2CO3: Carbonic acid

    • H(C2H3O2): Acetic acid

    • H2SO3: Sulfurous acid

    • H3PO3: Phosphorus acid

    • H3PO4: Phosphoric acid

  • Common Bases:

    • NaOH: Sodium hydroxide

    • KOH: Potassium hydroxide

    • LiOH: Lithium hydroxide

    • Ca(OH)2: Calcium hydroxide

    • NaHCO3: Sodium bicarbonate

    • NH3: Ammonia

Bronsted-Lowry Model

  • Acid: Substance donating a hydrogen ion in a reaction.

  • Base: Substance accepting a hydrogen ion in a reaction.

  • Conjugate Acid-Base Pair: Related substances by the transfer of a hydrogen ion; an acid will have one more H+ than its conjugate base, and a base will have one less H+ than its conjugate acid.

Identifying Acids and Bases in Reactions

  • An acid has a conjugate base, and a base has a conjugate acid.

  • The conjugate base has fewer H+ ions compared to its conjugate acid.