REDOX Reactions

Redox reactions occur when electrons are transferred from one substance to another.

Occur in two opposite parts, which must both occur – Oxidation and Reduction
OXIDATION IS LOSS REDUCTION IS GAIN – OILRIG

Oxidation occurs when an oxidation agent causes reductant to lose electrons (oxidise)

Reduction occurs when reductive agent causes oxidant to gain electrons (reduce)

Reductant = reducing agent
Oxidant = oxidising agent

Conjugate redox pair is made up of two species that differ by a certain number of electrons.

 

Half Equations
·             Redox Reactions do not show electrons transferred.
·             Half equations consider the reactions separately, thus enabling electron transfer to be seen.
For complex equations:
1. Balance all atoms in the half equation except hydrogen and oxygen
2. Balance oxygen by adding water (oxygen atoms react to form water in acidic solution)
3. Balance hydrogen by adding H+ ions (present in acid solution)
4. Balance charges by adding electrons (e-)

EXAMPLE:


Oxidation of magnesium
Oxidants cause oxidations but are themselves reduced

Iron placed in copper sulfate solution:

As SO42- is not changed on either side, it is said to be a spectator ion, and can be removed from the equation.



If it is aqueous, this actually means they are floating together

Oxidation Numbers

·             Rules:
                               1.     All atoms are treated as ions for this, even if they are covalently bonded
                              2.     The oxidation number of an element in its free (uncombined) state is zero
o   E.g. Al(s) or Zn(s). This is also true for elements found in nature as diatomic (two-atom) elements
                               3.     The oxidation number of a monatomic ion is the same as its charge.
o   E.g. NaCl Na+ has ON of +1 Cl O.N. of -1
                               4.     The sum of all oxidation numbers in a neutral compound is zero. The sum of all oxidation numbers in a polyatomic (many-atom) ion is equal to the charge on the ion.
o   Allows the calculation of the oxidation number of an atom that may have multiple oxidation states, if the other atoms in the ion have known oxidation numbers.
§  Fe2O3 – overall oxidation number is 0.
§  oxidation number of oxygen (O) = -2 (total = -6)
§  so oxidation of iron (Fe) = +3
                               5.     The oxidation number of an alkali metal (IA family) in a compound is +1; the oxidation number of an alkaline earth metal (IIA family) in a compound is +2.
                               6.     The oxidation number of oxygen in a compound is usually   –2.
o   Ex. Oxygen in compounds called peroxides (for example, hydrogen peroxide), then the oxygen has an oxidation number of –1. If the oxygen is bonded to fluorine, the number is +1.
                               7.     The oxidation state of hydrogen in a compound is usually +1.
o   Ex. If the hydrogen is part of a binary metal hydride (compound of hydrogen and some metal), then the oxidation state of hydrogen is –1.
                               8.     The oxidation number of fluorine is always –1. Chlorine, bromine, and iodine usually have an oxidation number of –1, unless they’re in combination with an oxygen or fluorine.
                               9.     Roman numerals shows the oxidation number
o   Copper (II) Sulfate: Copper has oxidation number of +2
o    

REMEMBER:
·             Ionic charge written 2+, oxidation numbers written +2
·             Find O.N. of Oxygen and Hydrogen first, then the others
·             Cations+ are written before anions- in compounds
·             If an element is made of multiple atoms, divide charge by this.
·             Oxidation is an INCREASE in the Oxidation Number of an atom
·             Reduction is a DECREASE in the Oxidation Number of an atom


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