Redox and Electrochemistry

1. Redox and Electrochemistry

2. Redox Reactions • Reduction – Oxidation reactions • Involve the transfer of electrons from one substance to another + The oxidation numbers of the atoms will change…. one goes up (oxidation) and one goes down (reduction)

3. Oxidation Number (Oxidation State) • Used to keep track of the transfer of electrons • Number is assigned to every atom in a chemical formula, in accordance with certain rules • NOT an ionic charge, but is often the same as the ionic charge • Possible oxidation states are given on the periodic table (upper right hand corner)

4. Rules for assigning Oxidation Numbers • For a neutral compound, the sum of the oxidation states must be zero • The oxidation state of any atom in an uncombined element is zero • Element not in chemical combination with another element • Examples: Na, Mg, H2, Cl2

5. Rules for assigning Oxidation Numbers • The oxidation state of a monatomic ion is equal to its charge • Examples: Na+ = • In an ionic salt, the oxidation number of each ion is equal to its charge • Examples: CaCl2 • For a polyatomic ion, the sum of the oxidation states must equal the overall charge • Example: SO42-

6. Rules for assigning Oxidation Numbers • Metals of group 1 always have an oxidation number of +1 • Metals of groups 2 always have an oxidation number of +2 • Fluorine is always -1, other halogens are usually -1 • Aluminum is always +3

7. Rules for assigning Oxidation Numbers 10. Oxygen is usually -2 Exceptions: • When paired with F (OF2), oxygen will be +2 • Peroxides (H2O2), oxygen will be -1 11. Hydrogen is usually +1 Exceptions: • Metal hydrides (Group 1 or 2 metals paired with hydrogen), LiH, CaH2, hydrogen will be -1

8. Examples Assign an oxidation state to each element in the following: • H2SO4 • SO32- • K2CrO4 • CrCl3

9. Reduction • Reduction of charge by gaining electrons Na+ + e- Na O + 2e- → O2- • Oxidation • Increase in charge by loss of electrons Fe  Fe3+ + 3e- Cl-  Cl + e-

10. LEO the lion says GER Losing Electrons Oxidation Gaining Electrons Reduction

11. Conservation of Matter/Conservation of Charge • Mass must be conserved • Mass on both sides must be the same (balanced) • Charge must be conserved • Net charge on both sides must be the same (balanced) – add electrons to the higher side • Reduction and Oxidation reactions must occur together (REDOX reactions)

12. Half Reactions • Every Redox reaction consists of a reduction and oxidation reaction • Each reaction is called a ½ reaction • A separate equation can be written for each ½ reaction

13. Half Reactions • Net charge and mass must be the same on both sides of the equation • The number of electrons must balance out, electrons do not appear in the net equation • One ½ reaction is reduction and the other is oxidation

14. Spectator Ion • Does not change oxidation states in the reaction, same oxidation state on both sides of the equation • Not every species in an equation is oxidized or reduced, some are spectator ions

15. Assign oxidation states to each element in the reaction • Identify the 2 substances that are changing oxidation states • Write the half reactions • Balance the mass • Balance the charge (add electrons to the higher side)

16. Examples • H2 + Cl2 2HCl • Fe + ZnO  Zn + FeO

17. Reducing Agent • Substance which is oxidized • Serves as a source of electrons to make the reduction reaction occur • Good reducing agents are substances that lose (donate) electrons easily – elements with low ionization energies Examples: group 1 and 2 metals

18. Oxidizing Agent • Substance which is reduced • Accepts (gains electrons) • Good oxidizing agents are substances that gain electrons (highly electronegative elements) Examples: Group 17 elements

19. Activity SeriesReference Table J

20. Metals • The most reactive metals are listed at the top • A reaction will occur spontaneously if the metal is higher than the metal ion that it is trying to replace • Reactive metals lose electrons easily (low ionization energy) • Higher on the table = More likely to be oxidized

21. Examples Ba + ZnCl2→ Zn + BaCl2 • Ba will replace Zn because Ba is above Zn • Ba is more reactive than Zn • More reactive means that it loses electrons easier

22. Nonmetals • For the halogen nonmetals listed in Table J, the most reactive ones are at the top • For nonmetals, high reactivity means that they are likely to gain electrons (high electronegativity) • Higher on the table = More likely to be reduced Example: F2 will replace any other halogen (it is the most reactive)

23. Examples • Which metal is most reactive? a. Fe b. Zn Cu 2. Will Ba react with Mn2+? 3. Will Na+ react with Cr? 4. Will this reaction occur spontaneously? Mg + Co(NO3)2→ 5. If this reaction does occur, what products would be made?

24. Balancing Equations • Assign oxidation numbers to all substances in the equation • Write the oxidation and reduction ½ reactions • Balance (cancel out) the electrons in the ½ reaction • Balance the rest of the equation • Check

25. Examples • Fe + Cl2→ FeCl3 • Fe + Cu2+ Cu + Fe3+ • KMnO4 + HCl → KCl + MnCl2 + H2O + Cl2