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Redox Year 12 Chemistry

Redox Year 12 Chemistry. What is Redox?. REDOX stands for REDuction /Oxidation Which species are oxidised ? 2Mg(s) + O 2 (g)  2MgO(s) Fe 2 O 3 (s) + 3CO(g)  2Fe(s) + 3CO 2 (g) Zn(s) + Cu 2+ ( aq )  Zn 2+ ( aq ) + Cu(s)

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Redox Year 12 Chemistry

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  1. RedoxYear 12 Chemistry

  2. What is Redox? • REDOX stands for REDuction/Oxidation • Which species are oxidised? 2Mg(s) + O2(g)  2MgO(s) Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g) Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) • Zn metal has been oxidised as it has lost electrons, and the Cu2+ has been reduced as it has gained electrons. • Zn  Zn2+ + 2e and Cu2++ 2e Cu • Oxidation refers to a loss of electrons • Reduction refers to a gain of electrons Redox reactions involve the transfer of electrons

  3. Definitions for Redox Reactions • OXIDATION—loss of electron(s) by a species; increase in oxidation number; increase in oxygen. • REDUCTION—gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen. • OXIDISING AGENT—accepts electrons (gets reduced) to facilitate oxidation of another species • REDUCING AGENT—donates electrons (gets oxidised) to facilitate reduction of another species.

  4. Oxidation and reduction always occur together • Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) • Zn  Zn2+ + 2e and Cu2++ 2e Cu • Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons. OIL RIG LEO GER xidation ose s lectrons oss xidation ain eduction lectrons s eduction ain

  5. Oxidant and Reductant Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) Zn  Zn2+ + 2eand Cu2+ + 2e Cu • The species that is oxidised, in this case Zn, is called the reducing agent or reductant. • The species that is reduced, in this case Cu2+, is the oxidising agent or oxidant. • Which is reduced /oxidised and which is the reductant /oxidant? • reduced oxidised • Fe3O4 + 4C  3Fe + 4CO • oxidant reductant

  6. Oxidation Number Rules The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. • Free elements and molecules have an oxidation number of zero. Na, Be, K, Pb, H2, O2, HCl, H2O = 0 • In monatomic ions, the oxidation number is equal to the charge on the ion. Li+ = +1; Fe3+ = +3; O2- = -2 • The oxidation number of oxygen is usually–2. • (In H2O2 and O22- it is –1.)

  7. Oxidation numbers of all the atoms in HCO3- and S in H2SO4? • The oxidation number of hydrogen is +1except when it is bonded to metals in hydrides (e.g. LiH). In these cases, its oxidation number is –1. • Fluorine is always –1. 6. The sum of the oxidation numbers of all the atoms in a molecule or polyatomic ion is equal to the charge. Remember charge on molecules = 0 HCO3- O = -2 H = +1 3x(-2) + 1 + ? = -1 C = +4 S = +6

  8. Recognising Oxidation and Reduction in a Redox reaction 2Mg (s) + O2 (g) 2MgO (s) • Write the ON for each atom • Increase in ON means oxidation • Decrease in ON means reduction 0 0 2+ 2- Explain which species is oxidised? Use ON to support your answer. S + O2 SO2 PCl3 + Cl2 PCl5 Zn(s) + 2HCl(aq) ZnCl2(s) + H2(g) NaCl + AgNO3 NaNO3 + AgCl ON of S increases from 0 to +4 therefore it is oxidised.

  9. Is it redox? Use ON to prove whether these reactions are redox • 2HCl + Ca(OH)2 CaCl2 + 2H2O • Cr2O72- + 2OH- 2CrO42- + H2O • Mg + Cl2 MgCl2 • Pb2+ + 2I- PbI2 • 2NH4NO3 2N2 + O2 + 4H2O Mg is oxidised as its ON increases from 0 in Mg to +2 in Mg2+ Cl- (not Chlorine) is reduced as its ON decreases from 0 in Cl2 to -1 in Cl-

  10. Test Yourself Q- Define oxidation and reduction and represent each as a chemical equation. A- oxidation = loss of e– … X X+ + e– reduction = gain of e– … X + e– X– Q- Why are 2Na+Cl22NaCl & 2H2+O22H2O considered redox reactions? A- Both involve the transfer of electrons (Na, Cl2 ,H2 and O2 have ON=0. After reaction ON are Na+ = 1, Cl- = -1,H+ = +1 and O2- = -2 Q- Is it possible to oxidise a material without reducing something else? A- No. A lost e– is taken up by something else.

  11. Test Yourself Q- Define oxidising and reducing agent. A- An oxidising agent causes oxidation by being reduced itself and a reducing agent causes reduction by being oxidised itself. Q- Explain using equations why Ca + Cl2 CaCl2 is a redox reaction. A- CaCl2 is an ionic compound made of positive calcium ion and negative chlorine ions Ca  Ca2+ + 2e–, Cl2 + 2e– 2Cl–. Thus Ca is losing electrons (oxidation) and Cl is gaining electrons (reduction).

  12. 2I- I2 + 2e- 2I- I2 I- I2 2I- I2 MnO4- Mn2+ MnO4- Mn2+ Mg Mg2+ + 2e- Cl2 + 2e- 2Cl- Cl2 Cl- Balancing Half Equations 1. Write half equation by identifying reactant and product 2. Balance atoms that are not O or H 3. Balance O by adding H2O and H by adding H+ MnO4- + 8H+ Mn2+ + 4H2O 4. Balance charge by adding e- to the most positive side MnO4- + 8H+ +5e- Mn2+ + 4H2O Balance the following, are they oxidation or reduction? Oxidation half-reaction (lose e-) Reduction half-reaction (gain e-)

  13. Fe2+ + Cr2O72- Fe3+ + Cr3+ +2 +3 Fe2+ Fe3+ +6 +3 Cr2O72- Cr3+ Cr2O72- 2Cr3+ Balancing Redox Equations The oxidation of Fe2+ to Fe3+ by Cr2O72- in acid solution? • Write the unbalanced equation for the reaction in ionic form. • Separate the equation into two half-reactions. Oxidation: Reduction: • Balance the atoms other than O and H in each half-reaction.

  14. Balancing Redox Equations Fe2+ Fe3+ + 1e- 6Fe2+ 6Fe3+ + 6e- 6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O 6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O 14H+ + Cr2O72- 2Cr3+ + 7H2O Cr2O72- 2Cr3+ + 7H2O • For reactions in acid, add H2O to balance O atoms and H+ to balance H atoms. • Add electrons to the most positive side of each half-reaction to balance the charges on the half-reaction. • If necessary, equalise the number of electrons in the two half-reactions by multiplying the half-reactions by appropriate coefficients.

  15. 14H+ + Cr2O72- + 6Fe2+ 6Fe3+ + 2Cr3+ + 7H2O 6Fe2+ 6Fe3+ + 6e- 6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O Balancing Redox Equations • Add the two half-reactions together and balance the final equation by inspection. The number of electrons on both sides must cancel. You should also cancel like species. Oxidation: Reduction: • Combine H+ and OH- to make water. Balance the following Fe2+ +MnO4- Fe3+ +Mn2+ 2H2O + 2e- H2 + 2OH- 2H20O2 + 4H+ + 4e-

  16. Halogens as Oxidants • Fluorine is so powerful an oxidant that it oxidises water to oxygen. 2F2 + 2H2O 4HF + O2 A halogen higher in the Group can oxidise the ions of one lower down. • Chlorine reacts with bromide and iodide Cl2 + 2Br- 2Cl- + Br2 Cl2 + 2I- 2Cl- + I2 • Bromine reacts with iodide Br2 + 2I- 2Br- + I2 • Iodine does not react with either of the other halide ions.

  17. Test Yourself • Write a balanced equation for Chlorine reacting with Magnesium and use oxidation numbers to explain which element is oxidised. Cl2 + Mg Mg2+ + 2Cl- • Mg changes its oxidation number from 0 to +2, an increase in oxidation number means it is oxidised. Chlorine decreases its oxidation number from 0 to -1 so it is reduced. 0 0 +2 -1

  18. Test Yourself • Write a balanced equation for Bromine reacting with Potassium Iodide and use oxidation numbers to explain which element is the oxidant. Br2 + KI KBr + I2 • Bromine changes its oxidation number from 0 to -1, a decrease in oxidation number means it is reduced. Iodine increases its oxidation number from -1 to 0 so it is oxidised. An oxidant gets reduced to help something else become oxidised therefore Br2 is the oxidant. 0 +1 -1 +1 -1 0

  19. Halogens can Oxidise Water • Halogens are not very soluble in water • They do react with water Cl2(s) + H2O(l) HCl(aq) + HOCl(aq) • Hypochlorous acid (HOCl) and hypochlorite (OCl-) are the main components of free active chlorine used in disinfectants and swimming pools.

  20. Metals Reacting with Metal Ions Metals will reduce metal ions if the metal is higher on the reactivity series than the ion. E.g. Magnesium in copper sulfate solution Mg Mg2+ + 2e- Cu2+ +2e- Cu But if the reaction doesn’t work water is getting reduced E.g. Sodium in Zinc sulfate 2Na 2Na+ + 2e- 2H20 + 2e- H2 + 2OH-

  21. Common Reductants Any metal will displace a less reactive one from solution K>Na>Li>Ca>Mg>Al>Zn>Fe>Sn>Pb>Cu>Ag

  22. Common Oxidants Cl2 + Mg MgCl2 Cl2 + 2e- 2Cl- and Mg Mg2+ + 2e-

  23. Source of electricity Electrode carbon or platinum Cell – + – + – + Electrolysis– used to separate ions e- • “Cells” are containers of liquid with electrodes: Molten or aqueous ions (Electrolyte) Cations Reduced Cathode is negative Anions Oxidised Anode is positive • In “electrolytic cells”, electricity is used to force chemicals to undergo a redox reaction

  24. – + – + – + The electrolytic cell • Electric current forces charges on electrodes Na+ Na+ Cl– Cl– • Na+ is attracted to cathode, Cl– to anode • Na+ takes up an electron: Na+(l) + e– Na • Cl– gives up an electron: 2Cl–(l)  Cl2 + 2e– • Electricity flows until ions are used up • Pure Na is deposited, Cl2 gas is produced

  25. Reactivity Series • Whether you get the metal or hydrogen during electrolysis depends on the position of the metal in the reactivity series: E.g. copper chloride solution Anode chlorine Cathode copper sodium chloride solution Anode chlorine Cathode hydrogen Hydrogen will be produced unless the metal is lower on the reactivity series.

  26. Electrolysis of Water

  27. Test Yourself

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