Oxidation Numbers, Redox Reactions, and Standard Charges

Pre-requisite to 'Topic 1: Inorganic Compounds 1 - Introduction to Transition Metals'

Reduction

Gain of an electron to an atom

Oxidized Species/Reductant

An atom that has oxidized - lost an electron

Reduced Species/Oxidant

An atom that has reduced - gained an electron

Oxidation Number

The charge and atom would have if the bonds were ionic

• It keeps track of electron transfers

Oxidation Number - Rules for Assigning

1. Free element: ON = 0

2. Monatomic ion: ON = ion charge

3. Polyatomic ion: sum of ONs = total charge

4. Neutral molecules: sum of ONS = 0

5. Oxygen: ON = -2, but is -1 in peroxides and is +2 with fluorine

6. Hydrogen: ON = +1 with non-metals and -1 with metals

7. Fluorine: ON = always -1

8. Chlorine, Bromine, Iodine: Usually -1, but takes a + number with oxygen/fluorine

9. Aluminum: ON = +3

10. Alkali metals: ON = +1

11. Alkaline earth metals: ON = +2

Monatomic Ion

Single atom with a positive (+) or negative (-) charge

• E.g: Na⁺, Cl^-, Al³⁺

Peroxide

Compounds that contain the peroxide ion: O₂^2-

• Each oxygen has an oxidation state of +1

Polyatomic Ions

Group of atoms covalently bonded but carry a charge as a whole

• E.g: OH^-, NO₃^-, NH₄⁺

Recognizing a Redox Reaction

Look for:

• An element reacting in its elemental state (as a pure substance)

• A monatomic ion changing charge

• An atom in a polyatomic ion changing oxidation states

Redox Reactions - Half Equations

Written as half equations for oxidation and reduction, where combined the electrons cancel

• e.g: 2Mg (s)  + O₂ (g) → 2MgO (s)

  • Oxidation half-equation: Mg → Mg²⁺ + 2e^- aka 2Mg → 2Mg²⁺ + 4e^-

  • Reduction half-equation: O₂ + 4e^- → 2O^2- (oxygen always splits into two separate oxide ions)

Oxidation Number - Tips

• Use the sum rule for polyatomic ions to solve for unknown ONs

• The overall charge isn’t only the charge of the ions/molecules, but also the charge multiplied by their molar coefficients

Redox Reactions - Steps and Principles for Balancing: in Acidic Conditions

1. Balance each half-reaction separately:

• Balance the amount of atoms by adding a molar coefficient

• Balance the oxygen by adding H₂O

• Balance the hydrogen by adding H⁺

• Balance the charges by adding electrons (e^-)

2. Ensure both half-equations have an equal amount of electrons, multiplying the whole equation by a coefficient if needed

3. Add them together, cancelling as needed

4. Ignore the spectator ions

Redox Reactions - Steps and Principles for Balancing: in Basic Conditions

1. Balance each half-reaction separately:

• Balance the amount of atoms by adding a molar coefficient

• Balance the oxygen by adding H₂O

• Balance the hydrogen by adding H⁺

• Balance the charges by adding electrons (e^-)

2. Ensure both half-equations have an equal amount of electrons, multiplying the whole equation by a coefficient if needed

3. Add them together, cancelling as needed

4. Ignore the spectator ions

5. Add OH^- equal to the amount of H⁺ on both sides to make H₂O, then cancel as needed

Common Charges of Monatomic Ions - Cations

Group	Names	Ions	Charge
1	Hydrogen, lithium, sodium, potassium	H+, Li+, Na+, K+	+1
2	Magnesium, calcium, strontium, barium	Mg2+, Ca2+, Sr2+, Ba2+	+2
13 (metals)	Aluminum	Al3+	+3
Transition Metals	Silver	Ag+	+1
	Zinc, nickel, manganese, lead	Zn2+, Ni2+, Mn2+, Pb2+	+2
	Copper(I), copper(II)	Cu+, Cu2+	+1/+2
	Iron(II), Iron(III)	Fe2+, Fe3+	+2/+3
	Cobalt(II), cobalt (III)	Co2+, Co3+	+2/+3
	Chromium	Cr3+	+3
	Tin(II), tin(IV)	Sn2+, Sn4+	+2/+4

Common Charges of Monatomic Ions - Anions

Group	Names	Ions	Charge
15	Nitride, phosphide	N3-, P3-	-3
16	Oxide, sulfide	O2-, S2-	-2
17	Fluoride, chloride, bromide, iodide	F-, Cl-, Br-, I-	-1

Common Charges of Polyatomic Ions - Anions

Names	Ions	Charges
Hydroxide	OH-	+1
Nitrate	NO3-	+1
Nitrite	NO2-	+1
Superoxide	O2-	+1
Bicarbonate	HCO3-	+1
Bisulfate	HSO4-	+1
Dihydrogen phosphate	H2PO4-	+1
Permanganate	MnO4-	+1
Acetate	C2H3O2- or CH3COO-	+1
Cyanide	CN-	+1
Thiocyanate	SCN-	+1
Perchlorate	ClO4-	+1
Chlorite	ClO2-	+1
Chlorate	ClO3-	+1
Hypochlorite	ClO-	+1
Peroxide	O22-	-2
Carbonate	CO32-	-2
Sulfate	SO42-	-2
Sulfite	SO32-	-2
Hydrogen phosphate	HPO42-	-2
Chromate	CrO42-	-2
Dichromate	CrO72-	-2
Oxalate	C2O42-	-2
Phosphate	PO43-	-3

Common Charges of Polyatomic Ions - Cations

Names	Ions	Charges
Ammonium	NH4+	+1
Hydronium	H3O+	+1
Mercury(I)	Hg2+	+2

Spectator Ions

Ions that do not take part in a chemical reaction — they are present before and after the reaction, but stay unchanged

Organic Compounds

Carbon-based molecules with C-H bonds, often together with oxygen, nitrogen, sulfur, phosphorus, or halogen,

Halogens

Elements in group 17 of the periodic table - non-metals including fluorine, chlorine, bromine, iodine, astatine, and tennessine

Inorganic Compounds

Carbon-based molecules (usually) without C-H bonds

s-block Elements

Section of the periodic table from groups 1-2 (alkali and alkaline earth metals)

p-block Elements

Section of the periodic table including some of the metallic elements in groups 13-16

Oxidation

Loss of an electron from an atom