Redox Flashcards
When a redox reaction occurs…
… e-s are transferred and therefore oxidation numbers change
Oxidation
- increase in oxidation number
* occurs in formation of metal ions
Reduction
- decrease in oxidation number
* occurs in formation of non-metal ions
Oxidation number
Shows the number of electrons which an atom has lost or gained as a result of forming a compound
Oxidation number determining rules
1) the oxidation number of an uncombined element is 0
2) the sum of oxidation numbers in a neutral compound is always 0
3) the oxidation number of a monatomic ion is equal to the charge
4) the sum of oxidation number in polyatomic ions always are up to the charge on the ion
Monoatomic ion charges
- all group 1 elements are +1
- al group 2 elements are +2
- hydrogen is always +1 (except in hydrides)
- fluorine is always -1
- chlorine is always -1 (except when combined with fluorine or oxygen), as are bromine and iodine
- oxygen is always 2- (except in peroxides [1-], superoxides and when combined with fluorine)
You must use the oxidation number when naming these substances
- NaClO: sodium chlorate (I)
- NaClO3: sodium chlorate (V)
- K2SO4: potassium sulphate (VI)
- K2SO3: potassium sulphate (IV)
- NaNO3: sodium nitrate (V)
- NaNO2: sodium nitrate (III)
-ate
The quantity of oxygen will be the number of combined oxygen atoms to give the anion the smallest negative charge
Oxidising agent
- causes another material to become oxidised
- is reduced in the reaction (takes up electrons)
- electron acceptor
Reducing agent
- caused another material to become reduced
- is oxidised in the reaction (gives away electrons)
- electron donor
Half equations
The charges on both sides must be equal
Reduction half equations
Electrons on the left because they are being gained
Oxidation half equations
Electrons are on the right because they are being lost
In acidic (ordinary) conditions for half equations
add H2O and H+ to balance equations
In alkaline conditions for half equations
Use H2O and OH- to balance half equations
