Redox reactions Flashcards
What is oxidation
Oxidation is the process of electron loss:
Zn –> Zn2+ + 2e- It involves an increase in oxidation number
What is reduction
Reduction is the process of electron gain:
Cl2 + 2e- –> 2Cl- It involves a decrease in oxidation number
What are the Rules for assigning oxidation numbers
- All uncombined elements have an oxidation number of zero eg . Zn, Cl2, O2, Ar all have oxidation numbers of zero
- The oxidation numbers of the elements in a compound add up to zero In NaCl Na= +1 Cl= -1 Sum = +1 -1 = 0
- The oxidation number of a monoatomic ion is equal to the ionic charge e.g. Zn2+ = +2 Cl- = -1 4. In a polyatomic ion (CO3^2-) the sum of the individual oxidation numbers of the elements adds up to the charge on the ion e.g. in CO3
2- C = +4 and O = -2
sum = +4 + (3 x -2) = -2 - Several elements have invariable oxidation numbers in their common compounds e.g. Fe2O3=-6 because of 3x-2 and then 2/-6=+3 Fe=3
- Hydrogens are always +1 except in hydrides where it is -1e.g. HF and Na Hhydride-hydrogen bonded to metal
- Chlorine is always -1 except if in a compound with F and O- it would be a positive value
8.Fluorine alwasy -1 - Oxygen always -2 except it is -1 in peroxides and 2+ in OF2
Examples of stuff
Group 1 metals = +1
Group 2 metals = +2
Al = +3
H = +1 (except in metal hydrides where it is –1 eg NaH)
F = -1
Cl, Br, I = –1 except in compounds with oxygen and fluorine
O = -2 except in peroxides (H2O2
) where it is –1 and in compounds with fluorine.
Note
the oxidation number of Cl
in CaCl2 = -1 and not -2 because
there are two Cl’s
Always work out the oxidation
per one atom of the element.
What is the oxidation number of Fe in FeCl3
Using rule 5, Cl has an O.N. of –1
Using rule 2, the O.N. of the elements must add up to 0
Fe must have an O.N. of +3
in order to cancel out 3 x –1 = -3 of the Cl’s
What are reducing and oxidising agents
reducing agents are electron donors that are oxidised themselves
oxidising agents are electron acceptors and are reduced themselves
Note 2
When naming oxidising
and reducing agents
always refer to full name
of substance and not
just name of element or
ion
Rules for Writing half equations
- Work out oxidation numbers for element being oxidised/ reduced Zn –> Zn2+ Zn changes from 0 to +2
- Add electrons equal to the change in oxidation number
For reduction add e’s to reactants
For oxidation add e’s to products Zn –> Zn2+ + 2e-
Something to remember when Combining half-equations
To make a full redox equation combine a reduction half equation with a oxidation half equation
To combine two half equations there must be equal numbers of electrons in the two half equations so that the electrons cancel out
Examples of working out oxidation states: what is the oxidative states of H2SO4 and SO4-2
H2SO4=First look at the most electronegative element whihc is O then do 4x-2=-8 + 2x1(two hydrogens) so u get -6 that means S=+6
SO4^-2=4x-2=-8 and because its an ion u need to get -2 as its overall oxidattive state thingy so S=+6
RUles for balancoing half equations
- Write down the species before and after a reaction
- Balance any atoms apartf rom oxyegn and hydrgen
- Balance any oxygens with H2O if there are any
- Balance any Hydrogens with H+ ions if there are any
- Balance any charges with electrons (e-)
Example of combinign alf equations fe2+–>fe3+ and MnO4—>Mn^2+
Fe2+–>Fe3+e-
MnO4–>Mn2+
MnO4^-1 –> Mn2+4H2O
MnO4^-1 +8H^+ +5e- –>Mn2+ +4H2O
MnO4^- +8H^+ +5Fe^2+–>Mn2^+ +5Fe^3+ +4H2O
What does it mean by oxidation state
The number of electrons to balance the compound
If the electron is on the right side of the half equation it is oxidation or reduction
Oxidation and vice versa for reduction
