Redox Reactions

What is it? Redox Reactions are chemical reactions that involve oxidation and reduction. Oxidation can be defined as a loss of electrons to another substance. Reduction can be defined as an acceptance of electrons from another substance. Redox reactions are those in which electrons are transferred from one reactant to another. Everyday redox reactions include: Photosynthesis Respiration Combustion of coal Production and use of fertilisers

Oxidation Is Loss of electrons Reduction Is Gain of electrons Key Terms If electrons are transferred, it is a redox reaction. A loss of electrons is called oxidation. A gain in electrons is called reduction. Reduction and oxidation happen simultaneously, hence the name “redox” An oxidising agent (oxidant) accepts electrons and thus gets reduced A reducing agent donates electrons and thus gets oxidised OIL RIG Oxidation Is Loss of electrons Reduction Is Gain of electrons

The oxidant is the species which causes oxidation and is itself reduced The reductant is the species which causes reduction and is itself oxidised

From this equation you can see that Na goes from an oxidation state of 0 to +1, it has donated an electron and has been oxidised. We can say that Na is the reducing agent (or reductant) as it has reduced the Cl Cl2 goes from 0 to -1, it has been reduced as it has gained an electron. It can also be called the oxidant (or oxidising agent) as it has oxidised the Na.

Oxidation/Reduction/Oxidant/Reductant Mg(s) + S (s) MgS (s) Fe(s) + 2Ag+ 2Ag(s) + Fe2+(aq) Cl2 (aq) + 2Br-(aq) Br2 (aq) + 2Cl-(aq) Fe3+(aq) + 3OH-(aq) Fe(OH)3 (s)

Oxidation Numbers - Rules There are a lot of rules: All atoms are treated as ions for this, even if they are covalently bonded Uncombined elements have an oxidation number of 0 Elements just bonded to identical atoms, eg O2 or H2 also have an oxidation number of 0. The oxidation number of a simple monatomic ion, eg Na+, is the same as its charge.

Oxidation Numbers - Rules 5) In compounds or compound ions, the overall oxidation number is just the ion charge SO42- - overall oxidation number is -2 Oxidation number of O = -2 (total -8) Oxidation number of S = +6 ***Within an ion, the most electronegative element has a negative oxidation number equal to its ionic charge***

Oxidation Numbers - Rules 6) The sum of the oxidation numbers in a neutral compound is 0 Fe2O3 – overall oxidation number is 0. oxidation number of O = -2 (total = -6) so oxidation of Fe= +3 7) If you see roman numerals, this is an oxidation number Copper (II) Sulphate: Copper has oxidation number of +2

Oxidation Numbers - Rules 8) The oxidation number of Hydrogen is +1 in its compounds with non-metals (eg HCl) The oxidation of Hydrogen is -1 in metal hybrides (eg NaH) 9) The oxidation number of Oxygen is usually -2 Exceptions: - peroxide compounds where O is -1 (eg H2O2) - compounds where it is bonded to Fluorine where O is +2

Assigning oxidation numbers to the atoms in the following substances ***Assign as many oxidation numbers as possible and then find the oxidation number of the unknown*** HBr Na2O CH4 Al2O3

Determine the oxidation number of S in each of the following: H2SO4 SO3 SO2 H2S

Has a redox reaction taken place?? Oxidation numbers are used to determine whether a REDOX reaction has taken place Oxidation is an INCREASE in the Oxidation Number of an ATOM Reduction is a DECREASE in the Oxidation Number of an atom ***Keep in mind that oxidation cannot happen without reduction***

Has a redox reaction taken place?? Assign oxidation numbers to all species present Determine whether a change in oxidation numbers has occurred 3. Has oxidation and reduction both taken place?

Identify the following equations as redox, state the substances that have been oxidised and reduced 2Fe(s) + 3Cl2 (g) 2FeCl3(s) NH3(g) + HCl(g) NH4(s) 2NO(g) + O2(g) 2NO2(g) P4O10(s) + 6H2O(l) 4H3PO4(aq)

Half Equations Half equations are a useful way of understanding the processes involved in a redox reaction. Although reduction and oxidation reactions occur simultaneously, it is possible to consider the two reactions separately. To do this we separate the conjugate pair of oxidant and reductant and we balance the equations by showing the electrons. Combining these half equations make up the ionic equation.

Half Equations When an iron nail is placed in a blue copper sulfate solution, the nail becomes coated with metallic copper and the blue colour of the solution fades. The full equation is: Fe(s) + CuSO4(aq) FeSO4(aq) + Cu(s) Fe(s) + Cu2+ + SO42-(aq) Fe2+ + SO42-(aq) + Cu(s) ***No change in SO42- oxidation number so can be disregarded ad spectator ions Fe(s) + Cu2+ Fe2+ + Cu(s)

Half Equations Oxidation Half-Equation: Fe(s) Fe2+ + 2e- Fe(s) + Cu2+ Fe2+ + Cu(s) Oxidation Half-Equation: Fe(s) Fe2+ + 2e- Reduction Half-Equation: Cu2+ (aq) + 2e- Cu(s)

2Ag+(aq) + Pb(s) 2Ag(s)+ Pb2+(aq) Half Equations The balanced ionic equation for the displacement of silver from an aqueous silver nitrate solution by metallic lead is: 2Ag+(aq) + Pb(s) 2Ag(s)+ Pb2+(aq) Write balance oxidation and reduction half-equations Which reactants accept electrons? Which reactant is oxidised? Which reactant is the reductant?

Chapter Review Questions 22, 23, 26, 27, 41