Acids and bases are everywhere, from the food we eat to the cleaning products we use. This section dives into how oxides behave as acids or bases in water, and how we measure their strength.
We'll explore the pH scale, buffer solutions, and acid-base indicators. Understanding these concepts helps us make sense of chemical reactions in everyday life and in the lab.
Acid-Base Theories
Brønsted-Lowry and Lewis Acid-Base Concepts
- Brønsted-Lowry acids donate protons (H+ ions) in chemical reactions
- Brønsted-Lowry bases accept protons in chemical reactions
- Lewis acids accept electron pairs forming coordinate covalent bonds
- Lewis bases donate electron pairs forming coordinate covalent bonds
- Brønsted-Lowry theory applies to proton transfer reactions (HCl + H2O → H3O+ + Cl-)
- Lewis theory encompasses broader range of reactions including those without proton transfer (BF3 + NH3 → F3B:NH3)
Conjugate Acid-Base Pairs and Amphoteric Behavior
- Conjugate acid-base pairs consist of two species differing by one proton
- Conjugate acid forms when a base accepts a proton
- Conjugate base forms when an acid donates a proton
- Strength of conjugate acid inversely relates to strength of corresponding base
- Amphoteric oxides exhibit both acidic and basic properties depending on reaction conditions
- Aluminum oxide (Al2O3) reacts as an acid with strong bases and as a base with strong acids
Oxides and Their Properties
Acidic and Basic Oxides
- Acidic oxides form acids when dissolved in water
- Nonmetal oxides typically produce acidic solutions (CO2, SO2, N2O5)
- Basic oxides form bases when dissolved in water
- Metal oxides typically produce basic solutions (Na2O, CaO, MgO)
- Oxide acidity increases with increasing oxidation state of central atom
- Oxide basicity increases with increasing ionic character of metal-oxygen bond
Hydrolysis and Neutralization Reactions
- Hydrolysis involves reaction of oxides with water to form acids or bases
- Acidic oxide hydrolysis: SO3 + H2O → H2SO4
- Basic oxide hydrolysis: CaO + H2O → Ca(OH)2
- Neutralization reactions occur between acidic and basic oxides
- Neutralization produces salt and water: CaO + SO3 → CaSO4
- Amphoteric oxides undergo both acid and base neutralization reactions
Aqueous Solutions and pH
pH Scale and Buffer Solutions
- pH scale measures hydrogen ion concentration in aqueous solutions
- pH ranges from 0 (most acidic) to 14 (most basic)
- pH calculated as negative logarithm of hydrogen ion concentration: pH = -log[H+]
- Buffer solutions resist changes in pH when small amounts of acid or base added
- Buffers consist of weak acid and its conjugate base (CH3COOH/CH3COO-)
- Buffer capacity depends on concentration and ratio of acid to conjugate base
Acid-Base Indicators and Titrations
- Acid-base indicators change color at specific pH values
- Phenolphthalein changes from colorless (acidic) to pink (basic) around pH 8.2
- Methyl orange changes from red (acidic) to yellow (basic) around pH 4.4
- Acid-base titrations determine unknown concentration of acid or base
- Titration curve shows pH change as titrant added to analyte
- Equivalence point occurs when moles of acid equal moles of base
- Endpoint detected visually using indicator color change
Acid-Base Equilibrium Constants
Acid and Base Dissociation Constants
- Acid dissociation constant (Ka) measures acid strength in aqueous solution
- Ka expressed as equilibrium constant for acid dissociation: HA ⇌ H+ + A-
- Ka = [H+][A-] / [HA]
- Base dissociation constant (Kb) measures base strength in aqueous solution
- Kb expressed as equilibrium constant for base dissociation: B + H2O ⇌ BH+ + OH-
- Kb = [BH+][OH-] / [B]
pKa, pKb, and Relationship to pH
- pKa defined as negative logarithm of Ka: pKa = -log Ka
- pKb defined as negative logarithm of Kb: pKb = -log Kb
- Lower pKa indicates stronger acid, higher pKa indicates weaker acid
- pKa + pKb = 14 for conjugate acid-base pair in water at 25°C
- pH of weak acid solution related to pKa: pH = pKa + log([A-] / [HA])
- Henderson-Hasselbalch equation used to calculate pH of buffer solutions