Acids and Bases | Brilliant Math & Science Wiki

Overview of Acids and Bases

• General Characteristics:
  • Bases: Bitter taste, slippery feel. Examples include ammonia, antacids, baking soda, and toothpaste.
  • Acids: Sour taste, react with metals to release gas and cause corrosion. Examples include citric acid (oranges), acetic acid (vinegar), aspirin, and sulfuric acid (car batteries).
  • Biological Importance: DNA contains purines and pyrimidines (bases); dietary fats are acids.
• Dissociation and Strength:
  • Strong acids/bases: Dissociate completely in aqueous solutions.
  • Weak acids/bases: Reach an equilibrium where most molecules remain in their original form.
• Theoretical Definitions:
  • Arrhenius (late 1800s): Acids produce protons ($H^+$); bases produce hydroxide ions ($OH^-$). Limited to aqueous solutions.
  • Brønsted–Lowry (1920s): Acids donate a proton; bases accept a proton. Defined by conjugate acid-base pairs.
  • Lewis: Acids accept an electron pair; bases donate an electron pair. This is the most inclusive definition.
  • Usanovich: A broader theory covering electron/cation/anion transfer, including redox reactions.
• Water and Autoprotolysis:
  • Protons ($H^+$) in water exist as hydronium ions ($H_3O^+$).
  • Autoprotolysis: Water molecules constantly react with each other to form $H^+$ and $OH^-$. In pure water at room temperature, the concentration of each is approximately $10^{-7}$ moles per liter.
• pH and Measurement:
  • pH Scale: A logarithmic scale used to quantify hydrogen ion concentration ($pH = -\log[H^+]$).
  • pOH Scale: Used to measure hydroxide ion concentration.
  • Indicators: Dyes (liquid or paper) that change color based on hydronium concentration to measure pH levels (typically 1–14).
• Environmental Impact:
  • pH monitoring is critical for environmental science, particularly in river health and ecological stability, as most organisms have specific pH survival ranges.