6.1 What is a redox reaction?

Oxidation → The loss of electrons from an atom, ion or molecule; the gain of oxygen by an element; or the removal of hydrogen from a compound.
Reduction → The gain of electrons by an atom, ion or molecule; the loss of oxygen; or the 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: $2Na + Cl_2 \rightarrow 2NaCl$ $2 N a + C l_{2} \to 2 N a Cl$
- Oxidation: $Na \rightarrow Na^+ + e^-$
- Reduction: $Cl_2 + 2e^- \rightarrow 2Cl^-$

Balancing half-equations

Assign oxidation states for each of the atoms on both sides of the reaction.
Balance the increase in oxidation state and the decrease in oxidation state.
Balance O by adding $H_2O$ .
Balance H by adding $H^+$ .
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 is.
The more negative the oxidation number, the more reduced a substance is.

🔥 Oxidation Number Rules

The oxidation number of an uncombined element is zero (e.g., $S_8$ , $Cl_2$ , and $Zn$ ).

In compounds, many atoms or ions have fixed oxidation numbers:

Group 1 elements are always $+1$ .

Group 2 elements are always $+2$ .

Fluorine is always $-1$ .

Hydrogen is $+1$ (except in metal hydrides such as $NaH$ , where it is $-1$ ).

Oxygen is $-2$ (except in peroxides, where it is $-1$ , and in $F_2O$ , where it is $+2$ ).

The oxidation number of an element in a monatomic ion is always the same as the charge.

The sum of oxidation numbers in a compound is zero.

The sum of oxidation numbers in an ion is equal to the charge on the ion.

In either a compound or an ion, the more electronegative element is given the negative oxidation number.

Oxidation numbers are usually whole numbers (they may 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.
$SO_2$ (sulfur dioxide) → S is $+4$ ; O is $-2$ .
$ICl_3$ (iodine trichloride) → I is $+3$ ; Cl is $-1$ .
$N_2H_4$ (hydrazine) → N is $-2$ ; H is $+1$ .

Compound ions

$NO_3^-$ (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.

🔥 Redox Example 1

$Cu + 2AgNO_3 \rightarrow Cu(NO_3)_2 + 2Ag$

The oxidation number of $Cu$ increases from $0$ to $+2$ (oxidation).

The oxidation number of $Ag$ decreases from $+1$ to $0$ (reduction).

🔥 Redox Example 2

$Sn + 4HNO_3 \rightarrow SnO_2 + 4NO_2 + 2H_2O$

The oxidation number of $Sn$ increases from $0$ to $+4$ (oxidation).

The oxidation number of $N$ decreases from $+5$ to $+4$ (reduction).

Nitric acid ( $HNO_3$ ) is acting as an oxidising agent.

Tin ( $Sn$ ) is acting as a reducing agent.

6.5 Oxidising agents and reducing agents

Oxidising agent (oxidant) → A substance which brings about oxidation by removing electrons from another atom or ion.
- It increases the oxidation number of another atom or ion.
- The oxidation number of the oxidising agent decreases.
- Examples: $O_2$ , $Cl_2$ , and $KMnO_4$ .
Reducing agent (reductant) → A substance which brings about reduction by donating electrons to another atom or ion.
- It decreases the oxidation number of another atom or ion.
- The oxidation number of the reducing agent increases.
- Examples: $H_2$ , $KI$ , 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.

🔥 Identifying Agents in a Reaction

$MnO_4^- + 5Fe^{2+} + 8H^+ \rightarrow Mn^{2+} + 5Fe^{3+} + 4H_2O$

$MnO_4^-$ is the oxidising agent (its oxidation state decreases from +7 to +2).

$Fe^{2+}$ is the reducing agent (its oxidation state increases from +2 to +3).

6.6 Naming compounds

Roman numerals 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 $+3$ .
This can also be used to distinguish between non-metal atoms in molecules and ions.
The oxidation number should be mentioned AFTER the ion it refers to.

Oxides of nitrogen

Several oxides of nitrogen exist, such as $N_2O$ , $NO$ , and $NO_2$ .
The oxidation number of N in $N_2O$ is $+1$ → nitrogen(I) oxide.
The oxidation number of N in $NO$ is $+2$ → nitrogen(II) oxide.
The oxidation number of N in $NO_2$ is $+4$ → nitrogen(IV) oxide.

Nitrate ions

The oxidation number of N in the $NO_2^-$ ion is $+3$ → $NaNO_2$ is sodium nitrate(III).
The oxidation number of N in the $NO_3^-$ ion is $+5$ → $NaNO_3$ is sodium nitrate(V).
Salts of common acids are usually named without including the oxidation number of the non-metal ion.
- $Mg(NO_3)_2$ is magnesium nitrate, not magnesium nitrate(V).
- $K_2SO_4$ is potassium sulfate, not potassium sulfate(VI).