Electrochemistry

Electrochemistry Chapter 18

Electrochemistry • Electrochemistry • the branch of chemistry that studies the electricity-related application of oxidation-reduction reactions • Redox reactions involve the transfer of electrons from the oxidized substance to the reduced substance This picture shows the lab we did – we placed zinc into CuSO4 and the Cu+2 came out of solution (was reduced) and deposited copper metal (the black substance) onto the zinc metal. The zinc metal was oxidized and went into solution.

Electrochemistry • If we had taken the temperature of the CuSO4/Zinc solution during the reaction, we would have found that the temperature increased. • However, if we separate the two metals by a porous barrier and connect them by a wire, we can change that heat energy into electrical energy. This is called a “Daniel” cell after its inventor, John Frederic Daniell, a British chemist

Electrochemistry • Electrode: metal conductor used to establish electrical contact with the non-metallic (solution) part of the circuit. • Electrolyte: a substance that dissolved in water to give a solution that will conduct electricity (positive and negative ions) Half-cell: a single electrode in a solution of its own ions. One half-cell contains the oxidation reaction and one contains the reduction reaction.

Electrochemistry • Anode: the electrode where oxidation takes place • Cathode: the electrode where reduction takes place RED CAT and AN OX

Electrochemistry • Electrochemical Cell • A system of electrodes and electrolytes in which either chemical reaction produces electrical energy by the flow of electrons In the cell on the right: What are the electrodes made of?__________________________________________________________________________________________ What are the electrolytes?_______________________________________________________________________________

Electrochemistry • There are two kinds of electrochemical cells: 1. Voltaic 2. Electrolytic

Voltaic Cells • Uses redox reactions that are spontaneous (they happen by themselves without any outside energy input) The electrolyte solutions used are Copper Sulfate and Zinc Sulfate

Voltaic Cells • How do we know if the reactions will be spontaneous? We check Table J (Activity Series) to see if one metal can displace the other. According to Table J, can zinc displace copper in the cell to the right? Does this make the reaction spontaneous?

Voltaic Cells Remember: RED CAT and AN OX LEO the lion goes GER In this cell, cations (+) in solution are reduced (gain electrons) at the surface of the cathode (RED CAT) to become metal ions. In this cell, the reduction half reaction is: Cu+2(aq) + 2e-→ Cu(s)

Voltaic Cell RED CAT and AN OX / LEO the lion goes GER Electrons given up by the zinc at the anode travel through the wire to the cathode. However, without the salt bridge, the circuit would not be complete and would not work.

Voltaic Cell Anions (negatively charged ions) move toward the anode to replace the electrons that are flowing through the wire. Cations move toward the cathode since the positive charges have been lost during reduction.

Voltaic Cell • This occurs as the sulfate ions (SO4-) move through the salt bridge to the anode and the Na+2 ions move through to the cathode. The solution in the salt bridge is sodium sulfate

Voltaic Cells Now let’s look at it all together: • Zinc loses two electrons at the anode and forms Zn+2 (which goes into solution) • The electrons flow through the wire to the copper cathode where it gains two electrons to form solid copper metal. • Ions flow through the salt bridge to replenish the lost charges. • The flow of electrons through the wire generates electricity.

Electrolytic Cells If, according to Table J, the reaction is NOT spontaneous, we can force the reaction to happen by applying a power source, such as a battery, in the wire between half cells.

Electrolytic Cells • The process in which an electric current is used to produce a redox reaction is called electrolysis. • In electrolysis, the cathode and the anode SWITCH.

Electrolytic Cells • In this cell: • Electrons from the negative end of the battery flow through the wire to the electrode, which is NOW NEGATIVE and BECOMES THE CATHODE. • The electrode attached to the positive end of the battery loses electrons to it and it is NOW POSITIVEand BECOMES THE ANODE.

Electrolytic Cells • Let’s put it all together: • Electrons move from the battery to the cathode. • Positive ions (potassium ions) move to the cathode to gain electrons and become reduced (GER). • Negative ions (chlorine ions) move toward the anode and give up electrons to replace the ones lost to the battery. They become oxidized (LEO). • The electrons go through the wire back to the battery (positive end). K++ 1 e-→ K RED CAT at Cathode 2Cl-→ Cl2 + 2e- AN OX at Anode K+ Cl-

Electroplating An electrolytic cell is used for electroplating. - an electrolytic process by which a metal ion is reduced and solid metal is deposited on a surface

Electroplating To plate a silver spoon: make the spoon the cathode (GER). The silver ions from anode will gain an electron at the cathode and silver metal will deposit on the spoon. The electrolyte is Silver Nitrate which provides more silver ions (Ag+). AgNO3 solution +

Comparison of Cells • Both have: • Anode – where oxidation takes place (AN OX) • Cathode – where reduction takes place (RED CAT) Differences:

Electrochemistry