Unit 12 Redox and Electrochemistry Review of Terminology for Redox Reactions OXIDATION loss of electron(s) by a species; increase in oxidation number. REDUCTION gain of electron(s); decrease in oxidation number. OXIDIZING AGENT electron acceptor; species is reduced. REDUCING AGENT electron donor; species is oxidized. 1
Identifying Redox Reactions Double Replacement Reactions and Acid Base Reactions are NOT Redox reactions. Most other reactions are redox reactions. If the oxidation number of at least one element in the reaction is known to change then it is a redox reaction. Which of the following are Redox reactions? a. Fe(OH) 2 + H 2 S > FeS +2H 2 O b. SnS 2 + 2NH 4 HS > (NH 4 ) 2 SnS 3 + H 2 S c. H 2 S + Br 2 > S + 2HBr d. 3CuS + 8HNO 3 > 3Cu(NO 3 ) 2 + 3S + 2NO + 4H 2 O Balancing Equations Step 1- Write unbalanced complete ionic equation. Step 2: Write separate half reactions for oxidation and reduction. Write the number of electrons gained or lost in each process on the appropriate side. Step 3 - Balance atoms using H+ and H 2 O Steps 4 and 5 Multiply each half reaction to be sure the same number of electrons are gained or lost. Steps 6 and 7 - Add half reactions and add back in spectator ions 2
Tips on Balancing Equations Never add O 2, O atoms, or O 2- to balance oxygen. Never add H 2 or H atoms to balance hydrogen. Be sure to write the correct charges on all the ions. Check your work at the end to make sure mass and charge are balanced. PRACTICE! 3
Practice # 2 Using the half reaction method balance the following reaction: Aqueous sodium iodide plus aqueous sodium hypochlorite react to produce solid iodine and aqueous sodium chloride in basic solution. (OH) 4 2-4
CHEMICAL CHANGE --> ELECTRIC CURRENT To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire. Zn Zn 2+ ions wire electrons salt bridge Cu Cu 2+ ions This is accomplished in a. A group of such cells is called a. CHEMICAL CHANGE --> ELECTRIC CURRENT Named after the inventor (Alessandro Volta) are electrochemical cells that turn chemical energy into electrical energy via a spontaneous reaction. Examples include the batteries that power flashlights and calculators. Voltaic cells have a two in which either or occurs. The half-cells are connected by a salt bridge a tube containing a strong electrolyte, often potassium sulfate (K 2 SO 4 ). 5
Zn --> Zn 2+ + 2e- Oxidation Anode Negative Zn wire electrons salt bridge <--Anions Cations--> Cu 2+ + 2e- -->Cu Cu Reduction Cathode Positive Zn 2+ ions Cu 2+ ions Electrons travel allows anions and cations to move between electrode compartments. We can use that energy to do work if we make the electrons flow through a wire or an external circuit. We call such a setup a voltaic cell or battery. A Voltaic Cell 6
Reason for a Salt Bridge Once even one electron flows from the anode to the cathode in the external circuit, the charges in each beaker would not be balanced and the flow of electrons would stop (short out). Therefore, we use a salt bridge, usually a U-shaped tube that contains a salt solution, to keep the charges balanced. Cations (+) move toward the cathode. Anions (-) move toward the anode. Voltaic Cells As the electrons reach the cathode, cations in the cathode are attracted to the now negative cathode. The electrons are taken by the cation, and the neutral metal is deposited on the cathode. 7
Electromotive Force (emf) Water only spontaneously flows one way in a waterfall. Likewise, electrons only spontaneously flow one way in a redox reaction from higher to lower potential energy; from loser to taker or from anode to cathode. Quantitative Electrochemistry The reaction in an electrochemical cell is a result of a competition for electrons between the two half-cells. The of a voltaic cell is a measure of the cell s ability to produce an electric current. The tendency of a given half-reaction to occur as a reduction is called the. 8
Voltaic Cells - Current Direction Why do the electrons flow from Zn to Cu instead of the other way? Zn 2+ + 2e -----> Zn Reduction Potential = 0.763 Volts Cu 2+ + 2e -----> Cu Reduction Potential = 0.340 Volts Standard Hydrogen Electrode The reduction potential of other elements are referenced to a standard hydrogen electrode (SHE). By definition, the reduction potential for hydrogen is 0 V: 2 H + (aq, 1M) + 2 e H 2 (g, 1 atm) 9
Standard Reduction Potentials Reduction potentials for many electrodes have been measured and tabulated. Standard Electric Potential (E 0 ) Or Electromotive Force (EMF) The difference between the reduction potentials of the two halfcells is called the. (0r ). E 0 of a voltaic cell is a measure (in volts) of the cell s ability to produce an electric current. 10
How to Calculate a Cell Potential 1. Write the individual half-cell reactions with the correct number of electrons. Write them both as reduction half reactions. 2. For each reduction half-reaction look up the standard reduction potential in a table of standard reduction potentials. 3. For Galvanic cells the electron transfer between anode and cathode is spontaneous, so E cell has to be positive. The more positive reduction potential will denote the specie that is reduced. The other specie is oxidized and you will have to reverse the reaction and change the sign of the reduction potential. Example # 1 1. Determine the cell potential where iron(iii) reacts with nickel to become iron(ii) and nickel(ii). Fe 3+ (aq) + e ------> Fe 2+ (aq) Ni 2+ (aq) + 2e ------> Ni(s) E 0 Fe3+ = +0.77 V E 0 Ni2+ = 0.25 V Reduction occurs in the Iron half-cell since it is the more positive reduction potential 11
Example # 2 Determine whether the following reaction is spontaneous: 2Ag(s) + Zn 2+ (aq) ------> 2Ag + (aq) + Zn(s) Dry Cells A dry cell is a voltaic cell in which the electrolyte is a paste. Both dry cells and alkaline batteries are single electrochemical cells that produce about 1.5 V 12
Half-Reactions in a Dry Cell Oxidation (anode reaction) Zn(s) -------> Zn 2+ (aq) + 2e Reduction (cathode reaction) 2MnO 2 (s) +2NH 4 + (aq) +2e --> Mn 2 O 3 (s) +2NH 3 (aq) +H 2 O(l) In a normal dry cell the graphite only serves as an electrical conductor and does not undergo reduction. Lead Storage Batteries A battery is a group of cells connected together. The half-reactions for a lead storage battery are as follows: 13
Lead Storage Batteries A 12-V car battery consists of six voltaic cells connected together. One cell of a 12-V lead storage battery is illustrated here: Lead Storage Batteries When using the battery the sulfuric acid is slowly used up and a lead(ii) sulfate coating builds up on the plates. This overall net reaction is spontaneous. The reverse reaction is not. However, when a car is running, the generator sends electricity to battery and the current supplies the energy necessary to run the reaction in reverse. In theory this reversible reaction should lead to a battery that is good forever. However, in practice some of the lead(ii) sulfate falls off the electrodes to the bottom of the cell and are lost from the recharging process. 14