2. Redox Reactions

3. Reactants: Zn + I 2 Product: Zn I 2

4. Redox Equations: At the conclusion of our time together, you should be able to: 1.Define redox 2.Figure out oxidation numbers for any element 3.Show the change in oxidation numbers in a reaction

5. What’s the Point ? Electrical production (batteries, fuel cells) REDOX reactions are important in … Purifying metals (e.g. Al, Na, Li) Producing gases (e.g. Cl 2, O 2, H 2 ) Electroplating metals Protecting metals from corrosion Balancing complex chemical equations Sensors and machines (e.g. pH meter) C 3 H 8 O + CrO 3 + H 2 SO 4  Cr 2 (SO 4 ) 3 + C 3 H 6 O + H 2 O

6. What is Redox? REDOX stands for REDuction/OXidation Oxidation is often thought of as a combination of a substance with oxygen (rusting, burning) Oxidation refers to a loss of electrons Reduction refers to a gain of electrons As a mnemonic remember LEO says GER Lose Electrons = Oxidation Gain Electrons = Reduction

7. Let’s See How You’re Doing?? Q- What is oxidation? What is reduction? Represent each as a chemical equation. oxidation = loss of e – … X  X + + e – reduction = gain of e – … X + e –  X – Q- Why are 2Na + Cl 2  2NaCl & 2H 2 + O 2  H 2 O considered redox reactions? Both involve the transfer of electrons. (Na has no charge, the atoms in diatomic molecules have no partial charge. After reaction the atoms have different shares of the electrons because of different EN values)

8. Let’s See How You’re Doing?? Q- Is it possible to oxidize a material without reducing something else? No. A lost e – is taken up by something else.

9. Determination of Oxidation and Reduction If oxidation # decreased; substance reduced If oxidation # decreased; substance reduced If oxidation # increased; substance oxidized If oxidation # increased; substance oxidized

10. Review of Oxidation Numbers We will see that there is a simple way to keep track of oxidation and reduction We will see that there is a simple way to keep track of oxidation and reduction This is done via “oxidation numbers” This is done via “oxidation numbers” An oxidation number is the charge an atom would have if electrons in its bonds belonged completely to the more electronegative atom An oxidation number is the charge an atom would have if electrons in its bonds belonged completely to the more electronegative atom e.g. in HCl, Cl has a higher EN. Thus, oxidation numbers are Cl = -1, H = +1 e.g. in HCl, Cl has a higher EN. Thus, oxidation numbers are Cl = -1, H = +1

11. 1.Any element, when not combined with atoms of a different element, has an oxidation # of zero. (O in O 2 is zero, Na by itself is zero) 2.Any simple monatomic ion (one-atom ion) has an oxidation number equal to its charge (Na + is +1, O 2– is –2) 3.The sum of the oxidation numbers of all of the atoms in a formula must equal the charge written for the formula. (if the oxidation number of O is –2, then in CO 3 2– the oxidation number of C is +4) Review of Oxidation Numbers

12. 4. In compounds, the oxidation # of IA metals is +1, IIA is +2, IIIA is +3, Zn & Cd is +2, Ag is +1. 5. In ionic compounds, the oxidation # of a nonmetal or polyatomic ion is equal to the charge of its associated ion. (MgCl 2, Mg is +2, therefore Cl is –1) 6. F is always –1, O is always –2 (unless combined with F), H is usually +1, except when it is bonded to metals in binary compounds. (ex. NaH, H oxidation # is –1 or when it’s in elemental form H 2, oxidation # is 0). Review of Oxidation Numbers

14. NaIO 3 Na = +1 O = -2 3x(-2) + 1 + ? = 0 Therefore, I = +5 IF 7 F = -1 7x(-1) + ? = 0 Therefore, I = +7 Oxidation numbers of all the elements in the following ?

15. Redox Equations: Let’s see if you can: 1.Define redox 2.Figure out oxidation numbers for any element 3.Show the change in oxidation numbers in a reaction

16. Rule (from earlier slide) TotalOx.# +1 6 -2 6 -6+5 3 +1 4 +2 -2 6 -14+12 3 +6 or rule 5 +1 6 +6 -2 6 -4 3 -2 5 +1 4 H N O 3 C2H6OC2H6OK 2 Cr 2 O 7 AgI +1 6 +2 -2 6 -8+5 3 H 2 PO 4 – Let’s Apply the Oxidation Rules

17. Corrosion – Deterioration of Metals by Electrochemical Process Determine and balance the reaction, determine what is oxidized and what is reduced.

18. Corrosion of Silver Ag + O 2  Ag 2 O 4 Ag + O 2  2 Ag 2 O - Each Ag loses 1e - - Each O gains 2e - Silver = oxidized Oxygen= reduced Silver = reducing agent Oxygen = oxidizing agent