Rucete ✏ SAT Chemistry In a Nutshell
11. Acids, Bases, and Salts
This chapter explains how acids, bases, and salts are defined, how they behave, and how to work with pH, indicators, titrations, and buffers. It also compares different acid–base theories and highlights key reactions involved in acid-base chemistry, including those found in environmental issues like acid rain.
Properties of Acids
- Conduct electricity in water; strength depends on degree of ionization (e.g., HCl = strong acid, HC₂H₃O₂ = weak acid)
- React with active metals to release hydrogen gas
- Change indicators:
- Litmus: red in acid
- Phenolphthalein: colorless in acid
- Neutralize bases to form salt and water (e.g., Mg(OH)₂ + H₂SO₄ → MgSO₄ + 2H₂O)
- React with carbonates to release CO₂ gas
- Arrhenius definition: Acids produce H⁺ in water (actually H₃O⁺) — Example: HX + H₂O → H₃O⁺ + X⁻
Properties of Bases
- Conduct electricity in water; strength depends on ionization
- Change indicators:
- Litmus: blue in base
- Phenolphthalein: pink in base
- Neutralize acids to form salt and water
- React with fats to form soap
- Aqueous bases feel slippery and can be caustic
- Arrhenius definition: Bases produce OH⁻ in water — Example: NaOH, KOH, Ca(OH)₂, NH₄OH
Brønsted–Lowry and Lewis Definitions
- Brønsted–Lowry theory:
- Acid = proton (H⁺) donor
- Base = proton (H⁺) acceptor
- Example: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
- Conjugate acid–base pairs:
- Differ by one H⁺
- Acid → conjugate base, Base → conjugate acid
- Lewis theory:
- Acid = electron pair acceptor
- Base = electron pair donor
pH, pOH, and Ion Concentration
- pH = −log[H⁺], pOH = −log[OH⁻]
- pH + pOH = 14 at 25°C
- [H⁺] = 10^−pH — Example: pH = 3 → [H⁺] = 1 × 10⁻³ M
- pH < 7 = acidic, pH = 7 = neutral, pH > 7 = basic
Indicators and Their Color Ranges
- Litmus: red in acid, blue in base
- Phenolphthalein: colorless in acid (pH < 8), pink in base (pH > 8)
- Methyl orange: red in acid (pH < 3.1), yellow in base (pH > 4.4)
Titration and Neutralization
- Titration: used to determine concentration of an unknown acid or base.
- Use standard solution to neutralize unknown.
- At equivalence point, moles of H⁺ = moles of OH⁻
- Formula: MaVa = MbVb (M = molarity, V = volume)
- Choose an indicator that changes color near the equivalence point.
Buffer Solutions
- A buffer resists changes in pH when small amounts of acid or base are added.
- Made from a weak acid and its conjugate base, or weak base and its conjugate acid.
- Examples: CH₃COOH and CH₃COONa; NH₃ and NH₄Cl
Salts and Hydrolysis
- Salts are ionic compounds formed from neutralization of acids and bases.
- Salts can be:
- Neutral (strong acid + strong base)
- Acidic (strong acid + weak base)
- Basic (weak acid + strong base)
- Hydrolysis: salt reacts with water to produce acidic or basic solution — e.g., NH₄Cl + H₂O → NH₃ + H₃O⁺
Amphoteric Substances
- Act as either an acid or base depending on conditions.
- Water is amphoteric: H₂O + H₂O ⇌ H₃O⁺ + OH⁻
- Other examples: HSO₄⁻, NH₃, amino acids
Acid Rain and Environmental Chemistry
- Caused by SO₂ and NOx gases forming H₂SO₄, HNO₃ in water
- Sources: burning fossil fuels (coal, oil, gas)
- Effects:
- Lowers pH of lakes and soil
- Harms aquatic life, trees, and buildings
- Prevention:
- Reducing emissions
- Using scrubbers in smokestacks
In a Nutshell
Acids and bases are defined by their ability to donate or accept protons or electron pairs. Their strength depends on ionization, and their behavior is measurable through pH and indicators. Titration allows concentration analysis, buffers maintain pH, and salts can influence acidity through hydrolysis. Understanding these principles is essential for acid–base chemistry and its real-world environmental impacts.