Acids, Bases, and pH Concepts

Identifying Properties of Acids and Bases

• Aqueous solution:

  • pH High: Above 7 (neutral to basic: 7-14)

  • pH Low: Below 7 (acidic: 0-7)

• Common Properties:

  • Slippery to touch

  • Sour taste

  • Conduct electricity

  • Change the color of pH paper

  • Corrosive (can damage skin)

  • Neutralize before disposal

Understanding Strong Acids and Bases

• Note: The list of strong acids and bases may vary across different sources.

• No need to memorize for class but good to know for higher level chemistry (AP, SAT II).

Acidity and Basicity in Solutions

• Definitions:

  • Acidic:

  • Contains mostly water, [H_3O^+] > [OH^-]

  • Neutral:

  • Contains mostly water, [H_3O^+] = [OH^-]

  • Basic:

  • Contains mostly water, [H_3O^+] < [OH^-]

H2O+H2O ⇌ H3​O++OH−

Naming Acids

Binary Acids:

• Composition: Contains H and one other element.

• Naming Rules:

  1. Prefix “hydro” is used.

  2. Root of the anion is used in the name.

  3. Suffix “ic” is used.

  4. The word “acid” is last.

• Examples:

  • HCl
    ightarrow ext{Hydrochloric acid}

  • H_2S
    ightarrow ext{Hydrosulfuric acid}

Oxyacids:

• Composition: Contains H, O, and a third element (usually a nonmetal).

• Naming Rules:

  1. No prefix is used.

  2. Root of the anion is used.

  3. Change “ite” to “ous” and “ate” to “ic.”

  4. The word “acid” is last.

• Examples:

  • H2SO4
    ightarrow ext{Sulfuric acid}

  • H2SO3
    ightarrow ext{Sulfurous acid}

Acid-Base Reactions and Dissociation

• Acids: Produce H^+ ions in water, forming H_3O^+ (Arrhenius definition) or donate H^+ to bases (Bronsted-Lowry).

• Bases: Produce OH^- in solutions (Arrhenius definition) or accept H^+ from compounds (Bronsted-Lowry).

• Strong Acids & Bases: Complete dissociation, represented as K_a > 1.

• Weak Acids & Bases: Partial dissociation, K_a < 1, closer to neutral pH (7).

Neutralization Reactions

• General Formula:
  ext{acid + base}
  ightarrow ext{salt + water}

• Strong Acid + Strong Base (Example):

  • HNO3(aq) + NaOH(aq) ightarrow H2O(l) + NaNO_3(aq)

  • Net ionic:
    H^+(aq) + OH^-(aq)
    ightarrow H_2O(l)

• Example with Weak Acid:

  • LiOH(aq) + H2CO3(aq)
    ightarrow H2O(l) + Li2CO_3(aq)

  • Net ionic:
    OH^-(aq) + H2CO3(aq)
    ightarrow H2O(l) + CO3^{2-}(aq)

Conjugate Acids and Bases

• Conjugate Acid: Formed when a base gains an H^+.

• Conjugate Base: Formed when an acid loses an H^+.

• Examples:

  • H^+ + F^-
    ightleftharpoons HF ext{ (Conjugate acid)}

  • H2SO4
    ightleftharpoons HSO_4^- + H^+ ext{ (Conjugate base)}

Calculating pH and pOH

• Formulas:

  • pH = - ext{log}[H^+]

  • pOH = - ext{log}[OH^-]

  • [H^+] = 10^{-pH}

  • [OH^-] = 10^{-pOH}

  • Relationship: pOH + pH = 14.00

Example Calculations:

1. Tap Water with [H_3O^+] = 8.90 imes 10^{-3} M:

   • pH = - ext{log}(8.90 x 10^{-3}) = 2.05

2. Tap Water with pH = 7.9:

   • [H^+] = 10^{-7.9} = 1.3 imes 10^{-8} M

3. Solution with pOH = 3.00:

   • [OH^-] = 10^{-3.00} = 1.00 imes 10^{-3} M

   • pH = 14 - 3 = 11.00

4. Human blood (pH = 7.40):

   • [H_3O^+] = 10^{-7.40} = 3.98 imes 10^{-8} M

   • pOH = 14 - 7.40 = 6.60

   • [OH^-] = 10^{-6.60} = 2.51 imes 10^{-7} M

Additional Examples:

5. Calculate [H^+] from [OH^-] = 9.35 x 10^{-4} M:

   • pOH = - ext{log}(9.35 x 10^{-4}) = 3.03

   • pH = 14 - 3.03 = 10.97

   • [H^+] = 10^{-10.97} = 1.07 x 10^{-11} M

Summary of Acid-Base Concepts

• Acids are proton donors and produce H^+ in solution.

• Bases are proton acceptors, producing OH^- in solution.

• Understanding the relationship between pH, pOH, and ion concentrations is essential for determining solution properties and reactions.