6. Electrochemistry

6.1 What is a redox reaction?

• Oxidation: the loss of electrons from an atom, ion or molecule OR the gain of oxygen by an element OR removal of hydrogen from a compound
• Reduction: the gain of electrons by an atom, ion or molecule OR the loss of oxygen OR addition of hydrogen to a compound
• Redox reaction: a reaction in which oxidation and reduction take place at the same time

Redox and electron transfer

Half-equations

• Equations which show either oxidation or reduction only 
• Reaction: 2Na + Cl₂ → 2NaCl
  • Na → Na⁺ + e⁻
  • Cl₂ + 2e⁻ → 2Cl⁻

Balancing half-equations

1. Assign oxidation states for each of the atoms in both sides of the reaction 
2. Balance increase in oxidation state and decrease in oxidation state 
3. Balance O by adding H₂O
4. Balance H by adding H+
5. Balance the number of electrons in each of the half-equations (if the main equation is separated)

6.2 Oxidation numbers

What are oxidation numbers?

• A number assigned to atoms and ions in a compound that shows us its degree of oxidation 
  • Can be positive 
  • Can be negative 
  • Can be zero (for elements) 
  • + or – signs must be included when necessary 
• The more positive the oxidation number, the more oxidised a substance 
• The more negative the oxidation number, the more reduced the substance 

Oxidation number rules

1. The oxidation number of uncombined element is zero (e.g. S⁸, Cl2 and Zn)
2. In compounds many atoms or ions have fixed oxidation numbers 
   1. Group 1 elements are always +1
   2. Group 2 elements are always +2
   3. Fluorine is always -1
   4. Hydrogen is +1 (except in metal hydrides such as NaH, where it is -1)
   5. Oxygen is -2 (except in peroxides, where it is -1, and in F₂O, where it is +2)
3. The oxidation number of an element in a monatomic ion is always the same as the charge 
4. The sum of oxidation numbers in a compound is zero 
5. The sum of oxidation numbers in an ion is equal to the charge on the ion 
6. In either a compound or an ion, the more electronegative element is given the negative oxidation number
7. Oxidation numbers are usually whole numbers (appear to be fractions when the structure of a compound is mistaken)

6.3 Applying the oxidation number rules

Compounds of a metal with a non-metal 

• The metal always has the positive oxidation number 
• The non-metal always has the negative oxidation number 

Compounds of a non-metal with a non-metal

• The sign of the oxidation number depends on the electronegativity of each atom 
• SO2 (sulfur dioxide) → S is +4 ; O is -2
• ICl3 (iodine trichloride) → I is +3 ; Cl is -1
• N2H4 (hydrazine) → N is -2 ; H is +1

Compound ions 

• NO3– (nitrate ion) → N is +5 ; O is -2

6.4 Redox and oxidation number

• Oxidation: the increase in oxidation number 
• Reduction: the decrease in oxidation number 
• Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag
  
  • Oxidation number of Cu increases from 0 to +2 (oxidation)
  • Oxidation number of Ag decreases from +1 to 0 (reduction) 
• Sn + 4HNO₃ → SnO₂ + 4NO₂ + 2H₂O
  
  • Oxidation number of Sn increases from 0 to +4 (oxidation)
  • Oxidation number of N decreases from +5 to +4 (reduction)
• Nitric acid is acting as an oxidising agent 
• TIn is acting as a reducing agent 

6.5 Oxidising agents and reducing agents

• Oxidising agent (oxidant): a substance which brings about oxidation removing electrons from another atom or ion 

• • • Increases the oxidation number of another atom or ion 
    • The oxidation number of the oxidising agent decreases 
    • Examples: O₂, Cl₂ and KMnO₄

• Reducing agent (reductant): a substance which brings about reduction by donating electrons to another atom or ion 

• • Decreases the oxidation number of another atom or ion
  • The oxidation number of the reducing agent increases 
  • Examples: H₂, KI and reactive metals such as Al
• Many substances can act as either oxidising or reducing agents depending on the substances they are reacting with, and the conditions 
• MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O 
  • MnO₄⁻ is the oxidising agent 
  • 5Fe²⁺ is the reducing agent 

6.6 Naming compounds

• Roman numbers in brackets can be used while naming compounds to mention the oxidation state 
  • Iron(II) chloride → oxidation state of Fe is +2
  • Iron(III) chloride → oxidation state of Fe is +3
• Can also be used to distinguish between non-metal atoms in molecules and ions 
• Oxidation number should be mentioned AFTER the ion it refers to

Oxides of nitrogen

• Several oxides of nitrogen present → N₂O, NO and NO₂
• The oxidation number of N in N₂O is +1 → nitrogen(I) oxide 
• The oxidation number of N in NO is +2 → nitrogen(II) oxide
• The oxidation number of N in NO₂ is +4 → nitrogen(IV) oxide

Nitrate ions

• The oxidation number of N in the NO²⁻ ion is +3 → NaNO is sodium nitrate(III)
• The oxidation number of N in the NO³⁻ ion is +5 → NaNO₃ is sodium nitrate(V)
• Salts of common acids are usually named without including the oxidation number of the non-metal ion 
  • Mg(NO₃)₂ is magnesium nitrate, not magnesium nitrate(V)
  • K₂SO₄ is potassium sulfate, not potassium sulfate(VI)