Reduction-Oxidation Reactions

1. Reduction-Oxidation Reactions (Redox)

2. Redox-defined Oxidation state (oxidation number) • A system for tracking electrons in redox reactions • An atom in a pure element has no charge and is assigned an oxidation state of zero

3. Redox-defined • REDuction-OXidation • chemical reactions in which atoms have their oxidation number (oxidation state) changed. • This can be either a simple or a complex process

4. Redox-defined • REDuction-OXidation • Oxidation is the loss of electrons or an increase in oxidation state • Reduction is the gain of electrons or a decrease in oxidation state • “OIL RIG” (oxidation is loss, reduction is gain)

5. Redox-defined • Electrons will go to the molecule, atom, or ion with the greater electronegativity • The oxidant(oxidizing agent) removes electrons from another substance i.e. it oxidizes other substances, and is thus itself reduced • The reductant(reducing agent) transfers electrons to another substance i.e. it reduces others, and is thus itself oxidized

6. Redox-defined

7. Redox-defined

8. Redox-examples

9. Trends in oxidation state

11. Oxidation-Reduction Between Nonmetals

12. Oxidation-Reduction Between Nonmetals

13. Rules for Assigning Oxidation Numbers • The oxidation number of a free element is always 0. • The atoms in He and N2, for example, have oxidation numbers of 0. • The oxidation number of a monatomic ion equals the charge of the ion. • For example, the oxidation number of Na+ is +1; the oxidation number of N3- is -3.

14. Rules for Assigning Oxidation Numbers • The usual oxidation number of hydrogen is +1. • The oxidation number of hydrogen is -1 in compounds containing elements that are less electronegative than hydrogen • The oxidation number of oxygen in compounds is usually -2. • Exceptions include OF2, since F is more electronegative than O, and BaO2, due to the structure of the peroxide ion, which is [O-O]2-.

15. Rules for Assigning Oxidation Numbers • The oxidation number of a Group IA element in a compound is +1. • The oxidation number of a Group IIA element in a compound is +2. • The oxidation number of a Group VIIA (17) element in a compound is -1, except when that element is combined with one having a higher electronegativity. • The oxidation number of Cl is -1 in HCl, but the oxidation number of Cl is +1 in HOCl.

16. Rules for Assigning Oxidation Numbers • The sum of the oxidation numbers of all of the atoms in a neutral compound is 0. • The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion. • For example, the sum of the oxidation numbers for SO42- is -2.

17. Balancing Oxidation-Reduction Reactions by the Half-reaction Method • Consider the reaction of aluminum with oxygen to produce aluminum oxide 4Al(s) + 3O2(g) 2Al2O3 • This reaction can be separated into a half-reaction for the substance oxidized 4Al(s) 4Al3+(s)+ 12e- • And a half-reaction for the substance reduced 3O2(g) + 12e- 6O2-(s)

18. Balancing Oxidation-Reduction Reactions by the Half-reaction Method 4Al(s) + 3O2(g) 2Al2O3 4Al(s) 4Al3+(s)+ 12e- 3O2(g) + 12e- 6O2-(s) The number of electrons lost (oxidized) must equal the number of electrons gained (reduced)

19. Electrochemistry • A galvanic cell, or voltaic cell, named after Luigi Galvani, or Alessandro Volta respectively, is an electrochemical cell that derives electrical energy from spontaneous redox reactions taking place within the cell. • It generally consists of two different metals connected by a salt bridge, or individual half-cells separated by a porous membrane.

20. Electrochemistry

21. Electrochemistry • A galvanic cell consists of two half-cells, • the electrode of one half-cell is composed of metal A, and the electrode of the other half-cell is composed of metal B; the redox reactions for the two separate half-cells are thus: • An+ + ne- A • Bm+ + me- B

22. Electrochemistry • these two metals can react with each other: • the metal atoms of one half-cell are able to induce reduction of the metal cations of the other half-cell

23. Electrochemistry • The standard electrical potential of a cell can be determined by use of a standard potential table for the two half cells. • identify the two metals reacting in the cell. Then look up the standard electrode potential, in volts, for each of the two half reactions.

24. Electrochemistry

25. Electrochemistry • the reaction of zinc metal and copper ion • To figure the potential or voltage for a redox reaction, add the voltages for the two half-reactions. One of the reactions will have to be “flipped” to make it an oxidation reaction. • Cu2+ + 2e- → Cu 0.34v • Zn2+ + 2e- → Zn -0.76v • Zn → Zn2+ + 2e- 0.76v • Zn + Cu2+ → Zn2+ + Cu 1.10v

26. Electrochemistry • What is the potential difference (voltage) created in a redox reaction between zinc metal and iron (II) ions? • Zn2+ + 2e- → Zn -0.76v • Fe2+ + 2e- → Fe -0.45v • Zn + Fe2+ → Zn2+ + Fe 0.31v

27. Electrochemistry • What is the potential difference (voltage) created in a redox reaction between liquid bromine and iodine ions • Br2 + 2I- → 2Br- + I2 0.53v

28. Relative Strength of Oxidizing and Reducing Agents Reducing Agents Oxidizing Agents S Li Li+ W T K K+ E R Ca Ca2+ A O Na Na+ K N Mg Mg2+ E G Al Al3+ R E Zn Zn2+ R Cr Cr3+ Fe Fe2+ Ni Ni2+ Sn Sn2+ Pb Pb2+ H2 H3O+ H2S S Cu Cu2+ I– I2 MnO42– MnO4– Fe2+ Fe3+ Hg Hg22+ Ag Ag+ NO2– NO3– S Br– Br2 T W Mn2+ MnO2 R E SO2 H2SO4 (conc.) O A Cr3+ Cr2O72– N K Cl– Cl2 G E Mn2+ MnO4– E R F– F2 R

29. Redox-examples • H2 + F2 → 2 HF (synthesis) • Fe + CuSO4 → FeSO4 + Cu (replacement) • 4 Fe + 3 O2 → 2 Fe2O3 (synthesis, “rusting”) • CH4 + 2 O2 → CO2 + 2 H2O (replacement, “combustion”)

30. Redox –applications • Cleaning products, fertilizer • Electrochemical cells (batteries) • Electroplating • Production of compact discs • Cellular respiration • Photosynthesis • etc…