Redox Reactions

of the degree of oxidation (loss of electrons) of an

atom in a chemical compound.
The oxidation state, which may be positive, negative or zero, is the hypothetical charge that an atom would have if all bonds to atoms of different elements were 100% ionic, with no covalent component. This is never exactly true for real bonds.

the reason for different properties.
Fe(OH)2
Fe(OH)3
iron is in oxidation state

+2

iron is in oxidation state
+3

forms diatomic molecules like chlorine (Cl2)—has an oxidation state of

zero.
For monatomic ions, the oxidation state is the same as the charge of the ion. E.g., the chloride anion (Cl−) has an oxidation state of −1, whereas the lithium cation (Li+) has an oxidation state of +1.
The sum of oxidation states for all atoms in a polyatomic ion is equal to the charge of the ion. Thus, the oxidation state of one element can be calculated from the oxidation states of the other elements.
The sum of the oxidation states of all atoms in a neutral molecule must be zero.

of −1.
Halogens other than fluorine have an oxidation state of

−1 except when they are bonded to oxygen, to nitrogen, or to another halogen that is more electronegative.
Hydrogen has an oxidation state of +1 except when bonded to more electropositive elements such as sodium, aluminium, and boron.
In compounds, oxygen typically has an oxidation state of −2.
Alkali metals have an oxidation state of +1 in virtually all of their compounds.
Alkaline earth metals have an oxidation state of +2 in virtually all of their compounds.

electrons.


Any reaction involving the transfer of electrons is an oxidation-reduction

(or redox) reaction
Oxidation cannot take place without reduction.
During a redox reaction, the oxidation numbers of reactants will change.

(cause them to lose electrons) are known as oxidizing agents,

oxidants, or oxidizers. The oxidizing agent removes electrons from another substance, and is thus itself reduced. And, because it "accepts" electrons, the oxidizing agent is also called an electron acceptor.
Oxygen is the quintessential oxidizer.
Substances with elements in high oxidation states (e.g., concentrated sulfuric acid H2SO4, potassium permanganate KMnO4, potassium dichromate(VI) K2Cr2O7, manganese(IV) oxide MnO2)
or else highly electronegative elements (O2, F2, Cl2, Br2) that can gain extra electrons by oxidizing another substance.

and which is oxidized?
Aluminium (Al) removes the oxygen atoms from

the iron(III) oxide (Fe2O3). The heat needed to start the reaction is usually provided by a magnesium fuse.
2Al + Fe2O3  Al2O3 + 2Fe
The iron(III) ions are reduced and the aluminium ions are oxidized.
This reaction produces much heat. It is used in incendiary weapons and in underwater welding.

(cause them to gain electrons) are said to be reductive

or reducing and are known as reducing agents, reductants, or reducers. They transfer electrons to another substance, and are thus oxidized. And, because it "donates" electrons, the reducing agent is also called an electron donor.
Electropositive elemental metals, such as Li, Na, Mg, Fe, Zn, and Al. These metals donate or give away electrons readily.
C, CO, H2

of each type on the left side of the arrow

must equal the number of atoms of each type to the right of the arrow.

The total charges on the left side of the arrow must equal the total charges to the right of the arrow.

The electrons lost (during oxidation) must equal the electrons gained (during reduction).

O2 = 2 H2O
2. Decomposition
General equation: AB → A +

B
2Ag2O = 4Ag + O2­ (heating)
(NH4)2Cr2O7 = N2­ + Cr2O3 + 4H2O­ (heating)
Decomposition reactions are the reverse of combination reactions, meaning they are the breakdown of a chemical compound into its component elements.

Mn+2SO4 + K2SO4 + H2O
Mn+7 + 5e- = Mn+2

reduction 2
2Cl- - 2e- = Cl20 oxidation 10 5

2 KMnO4 + 10KCl + 8H2SO4 = 5Cl2 + 2MnSO4 + 6K2SO4 + 8H2O

AB + CA

Cl2 + 2 NaBr = 2 NaCl +

Br2

2AgNO3 + Cu = Cu(NO3)2 + 2Ag

reactions
always involve oxygen
and an organic fuel.

reduced. These are known as disproportionation reactions.
2 H2O2(aq) →

2 H2O(l) + O2(g)
In H2O2, oxygen has an oxidation state of -1.
In H2O, its oxidation state is -2,
and it has been reduced.
In O2 its oxidation state is 0,
and it has been oxidized.
Oxygen has been both oxidized and reduced in the reaction, making this a disproportionation reaction.

= NO + H2O
CuO + NH3 = Cu + N2

+ H2O
NO2 + H2O = HNO3 + NO
NH4NO3 = N2O + H2O
K2Cr2O7 + KI + H2SO4 = Cr2(SO4)3 + K2SO4 + I2 + H2O
KMnO4 + HCl = Cl2↑ + MnCl2 + KCl + H2O