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Salts and Electrolysis

Subject: Chemistry
Topic: 5
Cambridge Code: 0620 / 0971 / 5070

Salt Preparation Methods​

Acid + Metal​

Condition: Metal must be below hydrogen in reactivity series

2HCl+Mg→MgCl2+H22HCl + Mg → MgCl_2 + H_2

Steps:

  1. Add metal to dilute acid
  2. Heat if necessary
  3. Filter to remove excess metal
  4. Crystallize by evaporation

Salts: Chlorides, sulfates (from dilute acids)

Acid + Base​

Condition: Soluble base or ammonia

2HCl+2NaOH→2NaCl+2H2O2HCl + 2NaOH → 2NaCl + 2H_2O

Steps:

  1. Add drops of acid to known base volume
  2. Use indicator to find equivalence point
  3. Repeat without indicator
  4. Evaporate and crystallize

Salts: Chlorides, sulfates, nitrates

Acid + Carbonate​

Condition: Generates CO₂ gas

2HCl+CaCO3→CaCl2+CO2+H2O2HCl + CaCO_3 → CaCl_2 + CO_2 + H_2O

Salts: Chlorides, sulfates (but not carbonates - would decompose)

Reactions with Copper​

Copper doesn't react with dilute acids (not reactive enough)

Uses:

  • Heat with concentrated sulfuric acid
  • Reaction with halogens
  • Displacement by more reactive metals

Precipitation Method​

Double displacement where product is insoluble

AgNO3+NaCl→AgCl(s)+NaNO3AgNO_3 + NaCl → AgCl(s) + NaNO_3

Steps:

  1. Mix solutions producing precipitate
  2. Filter to collect solid
  3. Wash with distilled water
  4. Dry in oven

Salts: AgCl, BaSO₄, CaCO₃

Salt Hydrolysis​

Hydrolysis - Salt reacts with water

Salts of Strong Acids and Strong Bases​

Properties:

  • Neutral solution
  • pH ≈ 7
  • No hydrolysis

Example: NaCl (from HCl + NaOH)

Salts of Weak Acids and Strong Bases​

Properties:

  • Alkaline solution
  • pH > 7
  • Anion hydrolyzes (weak acid restored)

CH3COO−+H2O⇌CH3COOH+OH−CH_3COO^- + H_2O ⇌ CH_3COOH + OH^-

Examples: NaCH₃COO, Na₂CO₃

Salts of Strong Acids and Weak Bases​

Properties:

  • Acidic solution
  • pH < 7
  • Cation hydrolyzes

NH4++H2O⇌NH3+H3O+NH_4^+ + H_2O ⇌ NH_3 + H_3O^+

Examples: NH₄Cl, NH₄NO₃

Salts of Weak Acids and Weak Bases​

Properties:

  • pH depends on relative strengths
  • Both hydrolyze
  • Usually weakly acidic or alkaline

Electrolysis​

Electrolysis - Chemical change using electrical energy

Key Components​

Electrode:

  • Anode - Positive electrode (oxidation)
  • Cathode - Negative electrode (reduction)

Electrolyte:

  • Liquid containing ions
  • Conducts current
  • Molten ionic compound or ionic solution

Electrodes:

  • Inert (platinum, graphite) - not consumed
  • Active (metals) - consumed (oxidized)

Electrolysis Equations​

At Electrodes​

Cathode (reduction - gain of electrons):

  • Cations gain electrons
  • Cation + e⁻ → Product

Anode (oxidation - loss of electrons):

  • Anions lose electrons
  • Anion - e⁻ → Product

Electrolysis of Molten Ionic Compounds​

Example: Molten NaCl

Cathode: 2Na++2e−→2Na2Na^+ + 2e^- → 2Na Anode: 2Cl−−2e−→Cl22Cl^- - 2e^- → Cl_2 Overall: 2NaCl→2Na+Cl22NaCl → 2Na + Cl_2

Electrolysis of Aqueous Solutions​

More complex: Competition between water and dissolved ions

Cathode (reduction priorities):

  1. Less reactive metals: Cu, Pb, Hg
  2. H₂O: 2H2O+2e−→H2+2OH−2H_2O + 2e^- → H_2 + 2OH^-
  3. More reactive metals: Na, K, Ca

Anode (oxidation priorities):

  1. Halides (Cl⁻, Br⁻, I⁻): Halogen produced
  2. Hydroxide (OH⁻): Oxygen produced
  3. Anions of weak acids

Example: NaCl(aq) with inert electrodes

  • Cathode: 2H2O+2e−→H2+2OH−2H_2O + 2e^- → H_2 + 2OH^-
  • Anode: 2Cl−−2e−→Cl22Cl^- - 2e^- → Cl_2
  • Overall: 2NaCl+2H2O→Cl2+H2+2NaOH2NaCl + 2H_2O → Cl_2 + H_2 + 2NaOH

Industrial Electrolysis​

Chlor-Alkali Process​

Purpose: Produce chlorine and sodium hydroxide

Reactant: Brine (NaCl solution)

Products:

  • Chlorine: Bleach, disinfectant
  • Hydrogen: Fuel
  • Sodium hydroxide: Detergent production

Aluminum Extraction​

Process: Hall-Héroult

  • Bauxite (Al₂O₃) dissolved in molten cryolite (Na₃AlF₆)
  • Electrolysis produces molten aluminum
  • Cathode: Al³⁺ + 3e⁻ → Al
  • Anode: 2O²⁻ - 4e⁻ → O₂

Copper Refining​

Purpose: Purify copper

Setup:

  • Cathode: Pure copper
  • Anode: Impure copper
  • Electrolyte: CuSO₄ solution

Cathode: Cu²⁺ + 2e⁻ → Cu (pure copper deposits) Anode: Cu - 2e⁻ → Cu²⁺

Faraday's Laws of Electrolysis​

First Law​

Charge transferred proportional to substance released

n=QFn = \frac{Q}{F}

where Q = charge (coulombs), F = Faraday (96,500 C/mol)

Second Law​

Same charge produces different amounts based on oxidation state

m1m2=M1A1M2A2\frac{m_1}{m_2} = \frac{M_1 A_1}{M_2 A_2}

Key Points​

  1. Salts prepared from acid + base/metal/carbonate
  2. Salt hydrolysis affects pH
  3. Electrolysis uses electrical energy for chemical change
  4. Cathode: Reduction (cations)
  5. Anode: Oxidation (anions)
  6. Different discharge priorities in aqueous solutions
  7. Industrial applications: Chlor-alkali, aluminum, copper

Practice Questions​

  1. Plan salt preparation from given reactants
  2. Predict pH of salt solutions
  3. Write electrolysis equations
  4. Calculate Faraday's law problems
  5. Explain industrial electrolysis
  6. Predict electrode products

Revision Tips​

  • Know salt preparation methods
  • Understand hydrolysis concept
  • Learn discharge priorities
  • Write electrolysis equations clearly
  • Practice Faraday calculations
  • Know industrial processes
  • Distinguish anode/cathode reactions