BATTERIES AND ELECTROLYTIC CELLS Practical Electrochemistry
How Batteries Work One of the most practical applications of spontaneous redox reactions is making batteries. In a battery, a spontaneous electron transfer from a reducer to an oxidizer is forced to occur through an external wire. As the electrons crowd through the wire, they give up their energy to do work.
Spontaneous Electron Transfer Zinc atoms will give up their 2 valence electrons to copper ions easily. 2+ Zn Cu 0 (s) Cu 2+
Remote Transfer If a barrier is placed between Zn and Cu 2+, the transfer can still take place, but through a wire. Zn
Electrochemical Cells Electrochemical cells are those in which spontaneous electron transfer takes place through a wire. They are also called voltaic or galvanic cells. You probably know them best by their practical name: batteries. We think of batteries as electrical, but they are really chemical devices that produce electricity.
Batteries Batteries are really quite simple devices that contain 2 half-cells with electrodes and solutions. Each half-cell is usually made of a metal electrode and a solution containing a salt of the same metal. There is an external wire connecting the 2 electrodes. The electrodes are the negative anode and the positive cathode.
Batteries An electrical power using device (resistor) is placed in the wire, and as the electrons crowd through the device, they give up their energy to make it operate: ex, a light bulb in a flashlight. The circuit must however be complete for electrons to flow, so a salt bridge or porous barrier between half-cells is also necessary.
Basic Battery Zinc metal atoms with 2 electrons Copper atoms with 2 electrons Zn 2+ aqueous ions 2+ Cl- 2+ Hungry Cu 2+ ions Cl- Zn Cu Anions move across the 2+ Cl- Cl- barrier to Cl- Cl- complete the circuit. Anode half-cell (-) Cathode half-cell (+) Zn Zn 2+ + 2e - Cu 2+ + 2 e - Cu
Disclaimer Unlike the animations, electrons do not really zip through a wire from one end to the other. It s actually more like a tube full of ping-pong balls. If you push a new ball into one end, they all move down a little, and the one on the other end pops out. Electrons in a real electric circuit actually move quite slowly.
Basic Battery The anode is slowly dissolving away as the atoms become soluble ions. The cations of the cathode compartment are becoming metal atoms, so the cathode is getting more massive. When either the anode is completely gone or the cathode cations are gone, the battery is dead.
Why batteries die on the shelf 2+ 2+ Cl- Cl- Cl- Cl- - Cl- + Zn Cu Cl- All the energy is released as heat and no work is done. However, the barrier must be porous for the battery to function normally.
Rechargeable Batteries Zn 2+ 2+ Generator: electron pump Cl- 2+ Cl- Cl- Cl- Cu Cl- Cl- Cl- Cl- With a rechargeable battery, a generator pushes the electrons backwards in the circuit. Even though the electrons don t want to go in this direction, the generator forces them. The electrons are restored to the anode and the battery can be used again.
Electromotive Force (emf) What is electromotive force? Emf is symbolized E. It is the pressure with which electrons flow through a circuit. Emf is measured in volts. So, how does one go about measuring the voltage produced by a battery? Or, are all batteries the same except for size?
Half-Cell Potentials Years ago, scientists agreed on a standard reference half-cell, against which to measure all other halfreactions. The reference they chose was a halfcell with H + in the solution, and a platinum electrode. The reaction occurring in the cell is: 2H + (aq) + 2 e - H 2 (g) Hydrogen ions are reduced. (The Pt is an inert electrode.)
Half-Cell Potentials The hydrogen half-cell was assigned a voltage of E = 0.0Volts. Various reactions were used for the other half-cell and the voltages recorded. If an element is easier to reduce than hydrogen, the voltage is a positive value. For example, when the hydrogen half cell is hooked up to a copper half-cell, the voltage is +0.34V.
Half-Cell Potentials So, a copper/hydrogen battery has an overall voltage (E cell ) of +0.34 volts. If the element is more difficult to reduce than hydrogen, the recorded voltage is negative. Ex. Cr 3+ + 3e - Cr E = -0.74V All such half-cell voltages are recorded in a table of Standard Electrode Reduction Potentials.
Standard Reduction Potentials Reduction Half-Reaction E F 2 + 2e - 2 F - +2.87V Au 3+ + 3e - Au +1.42V Ag + + e - Ag +0.80V I 2 + 2e - 2I - +0.54V Cu 2+ + 2e - Cu +0.34V 2 H + + 2 e - H 2 (g) 0.0 0V Pb 2+ + 2e - Pb -0.13V Ni 2+ + 2e - Ni -0.28V Cd 2+ + 2e - Cd -0.41V Zn 2+ + 2e - Zn -0.76V Li + + e - Li -3.04V
Battery Voltages In order for an electrochemical cell to work as a battery, the overall cell voltage must be positive. Remember that a reduction cannot occur without an oxidation. The table we just saw shows only reductions. So, one of the half-cell reactions will need to be turned around to become an oxidation.
Battery Voltages Guess what happens to the symbol on the voltage if we turn the equation around. The +/- symbol changes! Let s examine our Cu/Zn battery again with a look at the voltage it generates.
Cu/Zn Battery Here are the 2 half-reactions from the table: Cu 2+ + 2e - Cu +0.34V Zn 2+ + 2e - Zn -0.76V One of the reactions will have to be turned around to show an oxidation. Which one will it be to leave a positive voltage? Right! The Zn reaction must be reversed.
Cu/Zn Battery Cu 2+ + 2e - Cu Zn Zn 2+ + 2e - +0.34V +0.76V Notice that the voltage became positive. To calculate battery voltage, simply add the 2 equations together. Cu 2+ + 2e - Cu +0.34V Zn Zn 2+ + 2e - +0.76V Cu 2+ + Zn Zn 2+ + Cu +1.10V
Battery Electrodes In an earlier slide, we said that a battery has an anode and a cathode. How can you tell which is which? Do you know the legend of Paul Bunyan? Paul, the giant lumberjack had a well known pet: Babe, the blue ox. But Paul had another, less famous pet: Fritz, the red cat.
Battery Electrodes So, when you think of batteries, think of Paul Bunyan. He had 2 pets: an ox and a red cat. The anode is where oxidation happens, and reduction is at the cathode. The cathode is the + electrode, and the anode is the electrode. Electron flow is always from anode to cathode.
Try This Example A battery is composed of a silver half-cell and a cadmium half-cell. Write the reduction and oxidation half reactions. Draw the battery. Calculate the voltage. Label the anode and cathode. Draw the direction of electron flow.
Silver-Cadmium Battery e - flow: Cd to Ag Cd is oxidized: anode Ag is reduced: cathode Ag Cathode + Cd Ag + Cd 2+ NO 3 - Cl - Write half-reactions and voltages. Ag + + e - Ag E = 0.80V Cd Cd 2+ Cd + 2e 2+ - + 2eCd - Anode - E = +0.41V -0.41V Total Voltage = 1.21V Put in the porous barrier. Add the electrodes. Put cations in the solution. We also need some anions. Connect the external circuit. Reverse the Cd equation to be an oxidation.
Alternate Battery Diagrams Sometimes a simple battery diagram is shown as 2 separate half-cells. Instead of a porous barrier, there is a salt bridge. The bridge is usually some sort of tube that is filled with salt solution. It allows transfer of ions from one half-cell to the other.
Electrolysis Electrolysis involves using electric current and energy to force a nonspontaneous redox reaction to go. It is the opposite of an electrochemical cell. So, the overall cell voltage (E ) is negative rather than positive. This is the minimum voltage needed to make the cell operate.
Electrolysis The rechargeable battery we saw earlier is an example of an electrolytic reaction. We use a direct current generator for force the electron flow to go in reverse. Electrolysis is also used extensively for electroplating one metal onto another.
Electrolytic Cell Example An easy example of an electrolysis reaction would be the decomposition of NaCl. First the NaCl must be melted (around 1500 C). Then inert metal electrodes are placed in the liquid and the current is turned on. An electron is removed from Cl - and transferred to Na +.
Electrolytic Cell Example + Na + ion Na atom At the anode, Cl - ions lose electrons and are oxidized to Cl 2 molecules. - Cl - ion electrons Cl atom Electrons still flow from anode to cathode. At the cathode, Na + ions are reduced to Na atoms. + - + - A - + C
Electrolytic Cell The key differences between the electrolytic cell and the electrochemical cell are: 1. Generator rather than a resistor in the circuit. 2. Lack of a porous barrier or salt bridge. 3. Polarities of the electrodes are reversed. The anode is positive and the cathode is negative. + - + - A - + C
Aluminum Production Another very important industrial electrolysis reaction is the Hall- Heroult process for making aluminum from bauxite ore. Al 2 O 3 (l) + 3 C(s) 2 Al(l) + 3 CO(g) Al 3+ + 3 e - Al metal (reduction) C(s) + O CO (oxidation of C) The production of aluminum metal takes about 1% of all the electricity used in the U. S. each year.
Aluminum Production Although aluminum is relatively inexpensive, (about $1.50-$2.00 per pound), it is MUCH cheaper to melt and recycle aluminum than it is to smelt new metal. In the early 1800 s, when only a small amount of metal could be produced, Al cost $250,000 per pound ($545/gram) while silver was only $17/gram!
Calculating Voltage Voltage for electrolytic cells is calculated the same way that it is for batteries, except that the value will be negative rather than positive for this non-spontaneous cell. 2Cl - Cl 2 + 2 e - E = -1.36V 2Na + + 2 e - 2Na E = -2.71V Total voltage: E cell = -4.07V 4.07V is the minimum voltage needed to decompose NaCl(l).
Uses for Electrolysis Electrolysis of seawater is used to make Cl 2, H 2 and NaOH, three very useful and important materials. Electrolysis is also used extensively for electroplating one metal onto another (sometimes called anodizing. ) For example silver flatware is really silver plated flatware.
Electroplating Silverware is not solid silver. It is made of some cheaper, harder metal, like nickel. Then the nickel piece is coated with silver in an electrolysis cell. A bar of solid silver acts as the anode, and the nickel piece is the cathode. The cell simply transfers silver atoms from anode to cathode.
Electroplating Silver bar anode: Ag atoms change to ions and dissolve into the solution. e - The generator pulls electrons from the anode and pushes them into the cathode. e - e - e - e - e - e - e - e - Knife and threek cathode: Ag + ions change back to atoms and stick to, or plate the metal. Ag + Ag anode Ag AgNO 3 (aq)
Schematic Symbols Here are some symbols commonly used in electronic schematics: Generator Resistor Battery Voltmeter V Ammeter A Switch