1. 18.2-18.3 Balancing Redox Equations – Voltaic (Galvanic) Cells

2. Let’s Define Some Important Vocabulary Oxidation – the loss of electrons Reduction – the gain of electrons Redox (oxidation-reduction) reactions – a reaction in which there are changes in an atoms oxidation state. Great mnemonic device: LEO the lion goes GER. – Lose electrons oxidation … Gain electrons reduction.

3. Let’s Continue Balancing This Redox Reaction! +1-1 0 +2-1 0 HCl(aq) + Zn(s)  ZnCl 2 (aq) + H 2 (g) Next, we must identify the species being oxidized and reduced, and we must write the half reaction for each: Oxidation: H + (aq)  H 2 (g) Reduction: Zn(s)  Zn 2+ (aq) Now, we must balance mass. Oxidation: 2H + (aq)  H 2 (g) Reduction: Zn(s)  Zn 2+ (aq) Next, we must balance charge! Oxidation: 2H + (aq) + 2e -  H 2 (g) Reduction: Zn(s)  Zn 2+ (aq)+ 2e - Lastly, we must cancel electrons (they must be equal in both half reactions or we must make them equal) and add the two half-reactions together. 2H + (aq) + Zn(s)  Zn 2+ (aq) + H 2 (g)

4. Balancing Redox Reactions First we must identify each atoms oxidation number. As a reminder here are the rules for assigning oxidation numbers: – 1.) elements (even those that are polyatomic in nature) are assigned an oxidation number of zero. – 2.) monatomic ions have the same oxidation number as their charge. – 3.) Polyatomic ions: the sum of the oxidation numbers on the atoms that make up polyatomic ions, must sum up to the charge on the polyatomic ion. – 4.) Compounds – the sum of the oxidation numbers of a compound equals zero. Also as a reminder: Group 1 elements have an oxidation number of +1 Group 2 elements have an oxidation number of +2 Majority of the time, oxygen will have an oxidation number of -2 Majority of the time, halogens will have an oxidation number of -1 Let’s try assigning oxidation numbers to the following redox reaction: HCl(aq) + Zn(s)  ZnCl 2 (aq) + H 2 (g) +1-1 0 +2-1 0 HCl(aq) + Zn(s)  ZnCl 2 (aq) + H 2 (g)

5. Balancing Redox Reactions in Acid Solution When balancing redox reactions in acidic solution, the only difference is : 1.) when we are balancing mass, we must balance oxygen by adding water to the opposite side. 2.) When we balance hydrogen, we must add H +. Let’s Try a practice problem on the next slide 

6. Let’s Try a Practice Problem! Balance the following redox reaction in acidic solution: H + (aq) + Cr(s)  H 2 (g) + Cr 2+ (aq) 2H + (aq) + 2e -  H 2 (g) Cr(s)  Cr 2+ + 2e - 2H + (aq) + Cr(s)  H 2 (g) + Cr 2+ (aq)

7. Let’s Try Another!!! Balance the following redox reaction in acidic solution: Cu(s) + NO 3 - (aq)  Cu 2+ (aq) + NO 2 (g) Oxidation: Cu(s)  Cu 2+ (aq) + 2e - Reduction: 2(NO 3 - (aq) + 2H + (aq) + e -  NO 2 (g) + H 2 O(l)) 2NO 3 - (aq) + 4H + (aq) + 2e -  2NO 2 (g) + 2H 2 O(l) The overall reaction is: Cu(s) + 2NO 3 - (aq) + 4H + (aq)  Cu 2+ (aq) + 2NO 2 (g) + 2H 2 O(l)

8. Balancing Redox Reaction in Basic Solution When balancing a redox reaction in basic solution, the steps are the same as in an acidic solution, except in basic solution, after the hydrogen ions are used to balance the hydrogen atoms of the water molecules, we then need to add the same amount of hydroxide ions, as the number of hydrogen ions, to both sides of the equation. Since on the side with hydrogen ions the number of hydroxide ions are the same, it neutralizes the hydrogen ions. We can then combine the two as water!

9. Let’s Try a Practice Problem! Balance the following redox reaction in basic solution: ClO - (aq) + Cr(OH) 4 - (aq)  CrO 4 2- (aq) + Cl - (aq) +1-2 +3-2+1 +6-2 0 ClO - (aq) + Cr(OH) 4 - (aq)  CrO 4 2- (aq) + Cl - (aq) Oxidation: 2(Cr(OH) 4 - (aq) + 4OH - (aq)  CrO 4 2- (aq) + (4H + + 4OH - ) + 4H 2 O(l) + 3e - ) Reduction: 3(ClO - (aq) + 2H + (aq) + 2OH - (aq) 2H 2 O(l) + 2e -  Cl - (aq) + H 2 O(l) + 2OH - (aq) ) Now combine the two half reactions: 2Cr(OH) 4 - (aq) + 2OH - (aq) + 3ClO - (aq)  2CrO 4 2- (aq) + 5H 2 O(l) + 3Cl - (aq)

10. Voltaic Cells Electrical current – the flow of electrical charge. During a spontaneous redox reaction, such as placing a piece of solid zinc into a copper solution, the zinc atoms lose electrons, and the transfer to the copper ions. If a wire was used to connect zinc to copper, than the electrons would have to travel through the wire and could be used to do electrical work.

11. Electrochemical Cells Voltaic (Galvanic Cells) – an electrochemical cell that produces electrical current from a spontaneous chemical reaction. Electrolytic Cell - consumes electrical current to drive a non- spontaneous chemical reaction. In a voltaic cell, there are two half-cells. Each half-cell contains a metal placed into a salt solution. One half-cell is where oxidation occurs, and one is where reduction occurs. The metal (electrode) in both half-cells are connected with a wire. Electrons flow from the more negative electrode (the metal being oxidized) because the negative charge repels electrons, and they become attracted to the positive electrode (the metal being reduced).

12. Units Used In Electrochemical Cells

13. More on the Voltaic Cell You have already learned that electrons flow through a wire from one electrode to the other. The electrode where oxidation occurs (again in a voltaic cell, the negative electrode) is known as the anode. The electrode where reduction occurs is known as the cathode. A great mnemonic device, AN OX RED CAT. (anode oxidation, reduction cathode) Electrons always (in all electrochemical cells) flow from anode to cathode. A salt bridge, an inverted u-shaped tube containing a strong electrolyte (such as KNO 3 (aq)) with each end placed in the solution at both half-cells. This allows the flow of counter-ions without allowing the solutions to completely mix. The negative ions within the salt bridge neutralize the positively charged ions being accumulated in solution (and the positive ions within the salt bride neutralize the accumulated negative ions.) This completes the circuit.

14. Let’s Try a Practice Problem! In a voltaic cell, electrons flow (a)From the more negatively charged electrode to the more positively charged electrode. (b)From the more negatively charged electrode to the more positively charged electrode. (c)From lower potential energy to higher potential energy. (a) From the more negatively charged electrode to the more positively charged electrode.

15. Electrochemical Cell Notation Electrochemical cells can be represented by cell diagrams (also known as line notation). A cell diagram is a compact notation showing the what’s being oxidized and what’s being reduced in a cell, it is not a diagram of a cell. Example: Zn(s)│Zn 2+ (aq)││Cu 2+ (aq)│Cu(s) In a cell diagram, oxidation is on the left, and reduction is on the right. Single lines separate phases, and a double vertical line indicates a salt bridge. If the same state exists where one or both half-cells are in the same phase, commas are used to separate reactants, and or products.

16. 18.2-18.3 pgs. 904-905 #’s 38(a), 42(a), and 44(a) Read 18.4 pgs. 870-877