(409) AP Chemistry 8.1 - Introduction to Acids and Bases

Introduction to Acids and Bases

  • Overview of Bronsted-Lowry Theory

Bronsted-Lowry Acid

  • Definition: A Bronsted acid is a substance that donates hydrogen ions (H+) to a solution.

  • Example: Hydrochloric acid (HCl) dissociates in water:

    • Reaction: HCl → H+ + Cl-

    • Result: Increases H+ concentration in solution.

Bronsted-Lowry Base

  • Definition: A Bronsted base is a substance that increases hydroxide ion (OH-) concentration in a solution.

  • Example: Sodium hydroxide (NaOH) dissociates in water:

    • Reaction: NaOH → Na+ + OH-

    • Result: Increases OH- concentration in solution.

  • Alternative Base Definition: Bases can also increase OH- concentration by removing H+ from water.

    • Example: Ammonia (NH3) reacts with water:

      • Reaction: NH3 + H2O → NH4+ + OH-

      • Result: H2O loses H+, forming NH4+ (conjugate acid) and OH- (base).

Acid-Base Reaction Dynamics

  • Acid-Base Reaction: Can be expressed as reactions with water.

    • HCl + H2O → H3O+ + Cl-

  • Remember: H+ ions in solution attach to water molecules, forming hydronium (H3O+).

Interchangeability of H+ and H3O+

  • You can use H+ and H3O+ interchangeably in equations.

  • Essential to remember that in solution, H+ typically exists as H3O+.

Conjugate Acid-Base Pairs

  • Key Definitions:

    • When a base gains a hydrogen ion, it becomes a conjugate acid.

    • When an acid loses a hydrogen ion, it becomes a conjugate base.

  • Example: In NH3 + H2O,

    • NH3 is a base → gains H+ → becomes NH4+ (conjugate acid).

    • H2O is an acid → loses H+ → becomes OH- (conjugate base).

  • Amphoteric Substances: Can act as either acid or base (e.g. water).

Strong Acids and Bases

Strong Acids

  • Strong acids completely dissociate into ions in solution.

    • Notable strong acids: HCl, HBr, HI, HClO3, HClO4, HNO3, H2SO4.

Strong Bases

  • Strong bases also fully dissociate in solution.

    • Notable strong bases: LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2.

pH and pOH Scales

  • Definition: "p" indicates the negative logarithm;

    • pH = -log[H+]

    • pOH = -log[OH-]

  • Understanding pH Values:

    • pH scale ranges from 0 to 14; 7 is neutral.

    • pH < 7 means acidic; pH > 7 means basic.

  • Examples of pH Calculation:

    • If [H+] = 0.1, then pH = 1.

    • If [H+] = 0.01, then pH = 2.

Relationships Between pH, pOH, and Ion Concentration

  • At 25 °C, the relation:

    • pH + pOH = 14.

  • Ability to convert between pH and pOH using this relation:

    • Example: If pH = 2, then pOH = 12.

Water at Equilibrium and Autoionization

  • Autoionization of Water:

    • Water can dissociate into H+ and OH- ions, but the extent is minimal.

  • Equilibrium Constant (Kw):

    • Kw = [H+][OH-] = 1 x 10^-14 at 25 °C.

Neutral Solutions

  • Definition: Neutral state occurs when [H+] = [OH-].

    • In pure water at 25 °C:

      • [H+] = [OH-] = 1 x 10^-7, resulting in pH = 7.

Temperature Dependence on Water Ionization

  • Changes in temperature affect Kw:

    • Higher temperature leads to higher Kw values and increased ion dissociation.

    • At higher temperatures, neutral pH can drop below 7.

Practical Calculations

Calculating Molarity from pH

  • Calculation Example: Given pH = 2.876,

    • Find [H3O+]:

      • [H3O+] = 10^-2.876 = 0.00133 M (Acidic solution).

Calculating OH- Concentration

  • From pH = 8.90:

    • Convert to pOH: pOH = 14 - 8.90 = 5.1.

    • [OH-] = 10^-5.1 = 7.9 x 10^-6 M (Basic solution).

Conversion Between H3O+ and OH-

  • Calculation using Kw:

    • Given [H3O+] = 3.76 x 10^-5 M, find [OH-]:

      • [OH-] = 1 x 10^-14 / [H3O+].

Conclusion

  • Be aware of how temperature affects the behavior of acids, bases, and neutral solutions.

  • Familiarity with strong acids, bases, pH, and pOH scales is crucial for understanding acid-base chemistry.

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