Chem+120+S2025+Module+4revised

Evolution of Acid and Base Concepts

• Ancient perspectives:

  • Greeks and Romans viewed acids as sour substances, but lacked chemical understanding of acids and bases.

• Modern understanding initiated in 1887 with Svante Arrhenius:

  • Explored electricity conduction through water, laying groundwork for acid-base chemistry.

Electrodes and Metallic Bonding

• Electrode Definition:

  • A conductor (often metal) in contact with a non-metallic circuit (aqueous solutions).

  • Dipped into solutions to explore the interface of metals and electrolytes.

• Metallic Bonding:

  • Valence electrons in metals are loosely held and can move freely between cations.

  • Metals consist of cations surrounded by a "sea" of delocalized electrons, making them good conductors.

Arrhenius Experiment

• Arrhenius tested pure water and various chemicals:

  • Observed different conductivity levels in solutions, leading to classifications:

    • Non-electrolytes: Pure water (low conductance).

    • Weak Electrolytes: Dim light.

    • Strong Electrolytes: Bright light.

  • Conclusion: Electrical charge is carried by ions in solution.

Definitions of Acids and Bases (Arrhenius)

• Arrhenius Acid:

  • Increases [H+] concentration in solution upon dissolution in water.

  • Example reaction: HA ⇌ H+ + A−

  • Equilibrium constant, Ka, defined for strong acids as Ka >> 1 (e.g., HCl).

• Weak Acid:

  • Ionizes less than 5% in solution (e.g., HF).

Strong Acids to Memorize

• Hydrochloric acid (HCl), Hydrobromic acid (HBr), Hydroiodic acid (HI), Nitric acid (HNO3), Sulfuric acid (H2SO4), Perchloric acid (HClO4).

• Characteristics of strong acids:

  • Complete ionization in solution.

Weak Acids

• Example Weak Acid:

  • Hydrofluoric acid (HF): Ionizes weakly; Ka = 7.2 × 10−4.

  • Ionization Example: HF ⇌ H+ + F−

  • Calculate final concentrations based on initial conditions and Ka values.

Acid-Base Reactions

• Water can act as both an acid and a base (amphiprotic).

• Bronsted-Lowry Theory - 1923:

  • Coined definitions of acids and bases based on proton transfer.

  • Bronsted Acid: Donates protons (H+).

  • Bronsted Base: Accepts protons.

  • Conjugate Acid-Base Pairs: Identifies derived species post-transfer of protons.

pH and Acidity

• Use of logarithmic scales:

  • pH = -log[H+]

  • pOH = -log[OH−]

  • pKa and pKb allow for easier comparisons of acid/base strength.

  • Strong acids have lower pKa values.

Equations and Calculations

• Relationship Between pH and pOH:

  • pKw = pH + pOH = 14.

• Acid-Base Equilibrium:

  • Calculate concentrations and pH based on given solutions, particularly in regards to weak acids and bases.

Special Cases of Acids and Bases

• Polyprotic Acids:

  • Capable of donating multiple protons (e.g., H2SO4).

• **Lewis Acids and Bases:

  • Lewis Acid:** Electron pair acceptors, often metal ions.

  • Lewis Base: Electron pair donors, such as ligands.

Summary of Acid Strengths

• Trends in strength of binary acids and oxyacids:

  • Acid strength increases down groups (e.g., HCl < HBr < HI).

  • Oxyacid strength depends on the oxidation state of the central atom (e.g., HNO3 > HNO2).

Conclusion

• The definitions and understanding of acids and bases have evolved significantly from ancient times to modern chemistry, involving complex concepts such as electrical conductivity, ionization, and equilibrium.