Acids and bases are fundamental to chemistry, shaping reactions and properties of substances. They're defined by their ability to donate or accept protons, with theories evolving from Arrhenius to Brønsted-Lowry to Lewis, each broadening our understanding.The pH scale quantifies acidity and basicity, crucial in various applications. From neutralization reactions to buffer solutions, acids and bases play vital roles in biological systems, industrial processes, and environmental issues like ocean acidification.
Study Guides for Unit 8
8.0Unit 8 Overview: Acids and Bases3 min read
8.1Introduction to Acids and Bases3 min read
8.2pH and pOH of Strong Acids and Bases3 min read
8.3Weak Acid and Base Equilibria5 min read
8.4Acid-Base Reactions and Buffers 4 min read
8.5Acid-Base Titrations6 min read
8.6Molecular Structures of Acids and Bases2 min read
8.7pH and pKa3 min read
8.8Properties of Buffers3 min read
8.9Henderson-Hasselbalch Equation2 min read
8.10Buffer Capacity4 min read
Key Concepts
- Acids donate protons (H⁺) in aqueous solutions while bases accept protons
- Arrhenius theory defines acids as H⁺ donors and bases as OH⁻ donors
- Brønsted-Lowry theory expands the definition of acids and bases to include species without H⁺ or OH⁻
- Acids are proton donors and bases are proton acceptors
- Lewis theory further broadens the definition of acids and bases
- Acids are electron pair acceptors and bases are electron pair donors
- Conjugate acid-base pairs consist of a species and its corresponding acid or base after donating or accepting a proton
- Amphoteric substances can act as both acids and bases (water)
- Autoionization of water produces H⁺ and OH⁻ ions with a constant product Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C
Acid-Base Theories
- Arrhenius theory is the simplest and earliest acid-base theory
- Acids dissociate in water to produce H⁺ ions (HCl → H⁺ + Cl⁻)
- Bases dissociate in water to produce OH⁻ ions (NaOH → Na⁺ + OH⁻)
- Brønsted-Lowry theory is more comprehensive and includes species without H⁺ or OH⁻
- Acids donate protons to bases (HCl + H₂O ⇌ H₃O⁺ + Cl⁻)
- Bases accept protons from acids (NH₃ + H₂O ⇌ NH₄⁺ + OH⁻)
- Lewis theory is the most general and focuses on electron pair interactions
- Acids accept electron pairs (BF₃ + :NH₃ → F₃B←NH₃)
- Bases donate electron pairs (NH₃: + H⁺ → NH₄⁺)
- Conjugate acid-base pairs are related by the gain or loss of a proton
- A strong acid has a weak conjugate base and vice versa (HCl/Cl⁻, CH₃COOH/CH₃COO⁻)
Properties of Acids and Bases
- Acids have a sour taste (citric acid in lemons)
- Acids react with metals to produce hydrogen gas (Zn + 2HCl → ZnCl₂ + H₂)
- Acids change the color of pH indicators (litmus turns red in acidic solutions)
- Acids conduct electricity due to the presence of mobile ions (H⁺)
- Bases have a bitter taste (caffeine)
- Bases feel slippery due to the formation of soluble hydroxides (NaOH)
- Bases change the color of pH indicators (litmus turns blue in basic solutions)
- Bases conduct electricity due to the presence of mobile ions (OH⁻)
pH Scale and Calculations
- The pH scale measures the acidity or basicity of a solution
- pH = -log[H⁺] where [H⁺] is the molar concentration of hydrogen ions
- The pH scale ranges from 0 to 14 at 25°C
- Acidic solutions have pH < 7, neutral solutions have pH = 7, and basic solutions have pH > 7
- pOH is a measure of the hydroxide ion concentration
- pOH = -log[OH⁻] where [OH⁻] is the molar concentration of hydroxide ions
- The relationship between pH and pOH is pH + pOH = 14 at 25°C
- The concentration of [H⁺] and [OH⁻] can be calculated from pH and pOH using the inverse logarithm
- [H⁺] = 10⁻ᵖᴴ and [OH⁻] = 10⁻ᵖᴼᴴ
- The pH of a solution can be calculated from the concentration of a strong acid or base
- For a strong acid: pH = -log[HA] where [HA] is the molar concentration of the acid
- For a strong base: pOH = -log[MOH] where [MOH] is the molar concentration of the base
Strength of Acids and Bases
- The strength of an acid or base depends on its ability to ionize in aqueous solution
- Strong acids and bases completely ionize in water (HCl, H₂SO₄, NaOH)
- The concentration of H⁺ or OH⁻ equals the initial concentration of the acid or base
- Weak acids and bases partially ionize in water (CH₃COOH, NH₃)
- The concentration of H⁺ or OH⁻ is less than the initial concentration of the acid or base
- The acid dissociation constant (Ka) and base dissociation constant (Kb) measure the strength of weak acids and bases
- Ka = ([H⁺][A⁻])/[HA] and Kb = ([OH⁻][HB⁺])/[B]
- Larger Ka or Kb values indicate stronger acids or bases
- The relationship between Ka and Kb for a conjugate acid-base pair is Ka × Kb = Kw = 1.0 × 10⁻¹⁴ at 25°C
- The percent ionization of a weak acid or base can be calculated from its Ka or Kb value
- Percent ionization = (√(Ka/C)) × 100% where C is the initial concentration of the acid or base
Neutralization Reactions
- Neutralization reactions occur when an acid and a base react to form water and a salt
- HCl(aq) + NaOH(aq) → H₂O(l) + NaCl(aq)
- The net ionic equation for a strong acid-strong base neutralization is H⁺(aq) + OH⁻(aq) → H₂O(l)
- The equivalence point is reached when the moles of acid equal the moles of base
- At this point, the solution is neutral (pH = 7) if both the acid and base are strong
- Titration is a technique used to determine the concentration of an unknown acid or base solution
- A solution of known concentration (titrant) is added to the unknown solution (analyte) until the equivalence point is reached
- Indicators (phenolphthalein) or pH meters can be used to detect the endpoint of a titration
- The mole ratio of acid to base in a neutralization reaction depends on the balanced chemical equation
- For a monoprotic acid and a monobasic base, the mole ratio is 1:1 (HCl + NaOH)
- For a diprotic acid and a monobasic base, the mole ratio is 1:2 (H₂SO₄ + 2NaOH)
Buffers and Buffer Solutions
- A buffer is a solution that resists changes in pH when small amounts of acid or base are added
- Buffer solutions contain a weak acid and its conjugate base or a weak base and its conjugate acid
- Acetate buffer: CH₃COOH (weak acid) and CH₃COO⁻ (conjugate base)
- Ammonia buffer: NH₃ (weak base) and NH₄⁺ (conjugate acid)
- The Henderson-Hasselbalch equation relates the pH of a buffer solution to the pKa of the weak acid and the concentrations of the acid and its conjugate base
- pH = pKa + log([A⁻]/[HA]) where pKa = -logKa
- Buffer capacity is the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs
- Factors affecting buffer capacity include the concentrations of the acid and its conjugate base and the ratio of the two species
- Buffers play important roles in biological systems (maintaining blood pH) and in industrial processes (fermentation)
Real-World Applications
- Acid-base reactions are used in the production of fertilizers (ammonia), plastics (polyethylene terephthalate), and pharmaceuticals (aspirin)
- pH control is critical in water treatment, food processing, and soil management
- Lime (CaO) is added to acidic soils to increase pH and improve crop growth
- Buffers maintain the pH of bodily fluids within narrow ranges
- Carbonic acid-bicarbonate buffer system in blood (H₂CO₃/HCO₃⁻)
- Phosphate buffer system in intracellular fluid (H₂PO₄⁻/HPO₄²⁻)
- Acid-base titrations are used in environmental monitoring (water quality testing) and in the food industry (determining acidity of wines and vinegars)
- Antacids (Tums) neutralize excess stomach acid (HCl) to relieve indigestion and heartburn
- Acidic cleaning agents (toilet bowl cleaners) and basic cleaning agents (oven cleaners) are used in household maintenance
- Ocean acidification, caused by increasing atmospheric CO₂ levels, has negative impacts on marine ecosystems (coral bleaching)