2. Chemical Reactions Non-Redox Reactions (Double Replacement) Redox Reactions Precipitation 2 solutions  Solid ppt Neutralization Acid + Base  salt + H 2 O Synthesis A + B  AB Decomposition AB  A + B Single Replacement A + BC  B + AC Combustion C x H y + O 2  CO 2 + H 2 O

3. Redox Reactions

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? – Page 27 Involves a transfer of electrons. Redox reactions have a change in charge. REDOX stands for REDuction/OXidation 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. Determination of Oxidation and Reduction If oxidation # increased; substance oxidized (reducing agent) If oxidation # decreased; substance reduced (oxidizing agent)

8. Review of Oxidation Numbers – Page 28 1. Write down complete reaction. 2. Assign an oxidation number to each element. We will see that there is a simple way to keep track of oxidation and reduction

9. Oxidation Rules Any element by itself is zero. Oxidation # for ionic compounds are the same as the charge. Group 1 metals = +1, Group 2 metals = +2, F = -1. O is -2, H = +1, and Cl = -1 in compounds. The sum of oxidation numbers in a compound = zero, sum of oxidation numbers in an ion = the charge. An ion by itself has that oxidation number. Na + = 1 +

10. The Remainder of Redox Step 3: Determine which element is being reduced. This element’s oxidation # will decrease. Connect them and mark # of electrons gained. Step 4: Determine which element is being oxidized. This element’s oxidation # will increase. Connect them and mark # of electrons lost.

11. Example 0 0 +1 -1 2K + F 2  2KF -4 +1 0 +4 -2 +1 -2 CH 4 +O 2  CO 2 + H 2 O

12. Examples 1&2 – Page 29 Al + O 2  Al 2 O 3 FeO  Fe + O 2

13. Example 3 Mg + PbCl 4  Pb + MgCl 2

14. 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

15. 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).

16. Example 4 Ba(NO 3 ) 2 + Na 3 PO 4  Ba 3 (PO 4 ) 2 + NaNO 3