Chapter 20: Electrochemistry

Section 1: Introduction to Electrochemistry Electrochemistry: deals with electricity-related applications of oxidation-reduction reactions. Oxidation-Reduction Reactions involve a transfer of electrons. If both substances are in contact with each other, there is also a transfer of energy as heat.

Energy as heat is given off when electrons are transferred directly from Zn atoms to Cu2+ ions. This causes the temperature of the aqueous CuSO4 solution to rise.

Electrochemical Cells If the substances are separated by a porous barrier or salt bridge, the transfer of energy as heat becomes a transfer of electrical energy. The barrier prevents the metal atoms from getting through but the ions can move through the barrier, which prevents a charge from building up on the electrodes. Allows for the movement of charge through ions.

Electrons can be transferred from one side to the other through an external connecting wire. Electric current moves in a closed loop path, or circuit, so the movement of electrons is balanced by the movement of ions in solution.

An electrode is a conductor used to establish electrical contact with a nonmetallic part of a circuits, such as an electrolyte. In the diagram, the Zn and Cu strips are electrodes. A single electrode immersed a solution of its own ions is a half-cell.

The Half-Cells Half-Cells can be represented by half-reactions. The Half-cell with the Zn electrode in ZnSO4 solution has the half-reaction: Zn(s)  Zn2+(aq) + 2e- Oxidation is occurring so this electrode is called the anode. The Half-cell with the Cu electrode in CuSO4 solution has the half-reaction: Cu2+(aq) + 2e-  Cu(s) Reduction is occurring so this electrode is called the cathode.

In Ch 19 we learned that both oxidation and reduction must occur in an electrochemical reaction. The two half-cells together make an electrochemical cell.

The Complete Cell Notation: Anode | Anode || Cathode | Cathode Electrode Solution Solution Electrode For the Zinc and Copper cell, the cell notation is: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) The overall electrochemical reaction is found by adding the anode half reaction to the cathode half reaction: 1st + 3rd  2nd + 4th Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) *If the change in charge is not equal, you may need to multiply the half-reaction to balance out.

Did You Know? In 1836, John Daniell developed the Daniell Cell, an electrochemical cell with Zn and Cu. This was the first battery to produce a constant electrical current over a long period of time, and therefore was important for the operation of the electric telegraph.

Homework None! 

Section 2: Voltaic Cells Voltaic Cells use spontaneous oxidation-reduction reactions to convert chemical energy into electrical energy. The most common application of voltaic cells is batteries. The three most common types of dry cells are the zinc-carbon battery, the alkaline battery, and the mercury battery.

2MnO2(s) + H2O(l) + 2e-  Mn2O3(s) + 2OH-(aq) Zinc-Carbon Dry Cells These consist of a zinc container (anode) filled with a moist paste of MnO2, NH4Cl, and ZnCl2. When the external circuit is closed, zinc atoms are oxidized at the negative electrode. Zn(s)  Zn2+(aq) + 2e- Electrons move across the circuit and reenter the cell through the carbon rod, or cathode. 2MnO2(s) + H2O(l) + 2e-  Mn2O3(s) + 2OH-(aq)

Alkaline Batteries These do not have a carbon rod which allows them to be smaller. This cell uses a paste of Zn metal and KOH instead of a solid metal anode. Zn(s) + 2OH-(aq)  Zn(OH)2(s) + 2e- The half-reaction at the cathode is the same as the zinc-carbon dry cell. 2MnO2(s) + H2O(l) + 2e-  Mn2O3(s) + 2OH-(aq)

Mercury Batteries Mercury Batteries have the same anode half-reaction as in the alkaline dry cell. Zn(s) + 2OH-(aq)  Zn(OH)2(s) + 2e- However, the cathode half-reaction is different HgO(s) + H2O(l) + 2e-  Hg(l) + 2OH-

Fuel Cells These are voltaic cells in which the reactants are being continuously supplied and the products are being continuously removed. So in principle, a fuel cell could keep changing chemical energy into electrical energy forever. The reaction below shows the type of fuel cell used in the United States Space Program. Cathode: O2(g) + 2H2O(l) + 4e-  4OH- Anode: 2H2(g) + 4OH-(aq)  4e- + 4H2O(l) Net Reaction: 2H2 + O2  2H2O

Fuel cells are very efficient and have very low emissions.

Corrosion and Its Prevention The metal most commonly affected by corrosion is iron. The formation of rust forms by the following reaction: 4Fe(s) + 3O2(g) + xH2O(l)  2Fe2O3●xH2O(s) The value of x will vary depending on the amount of water present and will affect the color of rust formed.

Preventing Corrosion The most common way is to coat steel with zinc in a process called galvanizing. Zinc is more easily oxidized than iron so it will react first, protecting the iron. This is called cathodic protection. Ex: The Alaskan Oil Pipeline

Electrical Potential In a voltaic cell, the oxidizing agent at the cathode pulls the electrons through the wire away from the reducing agent at the anode. This “pull” is called the electric potential and is expressed in volts (V). Current is the movement of the electrons and is expressed in units of amperes, or amps (A). Electrons flow from higher electric potential to lower electric potential.

Chart on page 664 lists the SEP values. Electrode Potentials Standard Electrode Potential: the potential of a half-cell under standard conditions measure relative to the standard hydrogen electrode. Chart on page 664 lists the SEP values. These are always written as reduction half-reactions. When they are changed to oxidation half-reactions, the sign is reversed. Zn2+ + 2e-  Zn E0 = -0.76 V Zn  Zn2+ + 2e- E0 = +0.76 V

Calculating Cell Potential Standard electrode potentials can be used to predict if a redox reaction will occur spontaneously. If E0cell is positive, it is spontaneous. E0cell = E0cathode - E0anode The half-reaction that has the more negative E0 value will be the anode. You do not need to switch the sign of the anode when doing these calculations because the subtraction in the formula takes care of that.

Example 1 Write the overall cell reaction, and calculate the cell potential for a voltaic cell consisting of the following half-cells: an iron (Fe) electrode in a solution of Fe(NO3)3 and a silver (Ag) electrode in a solution of AgNO3. Look up E0 for each half-reaction on pg 664. Fe3+(aq) + 3e-  Fe(s) E0 = -0.04 V Ag+(aq) + e-  Ag(s) E0 = +0.80 V Determine the cathode and anode. Anode is Fe (more negative) and Cathode is Ag

Example 1 Cont. 3. Determine the overall cell reaction. Electrons need to match so, multiple the Ag half-reaction by 3 and reverse the Fe reaction since we determined it is the anode and needs to be oxidation. Fe(s)  Fe3+(aq) + 3e- E0 = -0.04 V 3Ag+(aq) + 3e-  3Ag(s) E0 = +0.80 V *we do not multiply the E0 value by 3 Fe(s) + 3Ag+(aq)  3Ag(s) + Fe3+(aq) *electrons cancel out because they are both 3

Example 1 Cont. Calculate the cell potential E0cell = E0cathode - E0anode E0cell = +0.80 V – (-0.04 V) = +0.84 V Check your answer: The calculated value for E0cell is positive so that means it is spontaneous and a voltaic cell as the question stated.

Example 2: Determine the overall electrochemical equation and E0 value for H+/H2 and Fe2+/Fe3+. 2H+ + 2e-  H2 E0 = +0.00 V Fe3+ + e-  Fe2+ E0 = +0.77 V Anode is H2 (more negative) and Cathode is Fe2+ Multiply Fe half-reaction by 2 and reverse the H reaction. 2Fe3+ + 2e-  2Fe2+ E0 = +0.77 V H2  2H+ + 2e- E0 = +0.00 V H2 + 2Fe3+  2H+ + 2Fe2+ E0 = 0.77 V – 0.00 V = 0.77 V

Homework None!

Section 3: Electrolytic Cells If electrical energy is required to produce a redox reaction and bring about a chemical change in an electrochemical cell, it is an electrolytic cell. Example: In a voltaic cell consisting of zinc and copper, the electrons move from zinc to copper. When you attach a current, the electrons move from copper to zinc.

Important Differences The anode and cathode of an electrolytic cell are connected to a battery while a voltaic cell serves as a source of electrical energy. Electrolytic cells have nonspontaneous redox reactions occurring which is why they require an outside electrical energy source. Voltaic cells have spontaneous redox reactions occurring which produce electricity.

Electroplating An electrolytic process in which a metal ion is reduced and solid metal is deposited on a surface is called electroplating. Pennies consist of a zinc core (97.5%) electroplated with a layer of copper. Video of Electroplating

Rechargeable Cells A rechargeable cell combines the redox chemistry of both voltaic and electrolytic cells. When it operates as a battery it is like a voltaic cell, converting chemical energy into electrical energy. When the battery is being recharged, it is like an electrolytic cell, converting electrical energy into chemical energy.

Electrolysis Electroplating and recharging a battery are both examples of electrolysis. In electrolysis, electrical energy is used to force a nonspontaneous chemical reaction to occur. Electrolysis is used to purify many metals from the ores in which they are found chemically combined in the earth’s crust.

Electrolysis of Water Hydrogen gas and Oxygen gas will spontaneously combine to form water. To break apart water into H2 and O2 requires energy and is called the electrolysis of water.

Aluminum Production by Electrolysis Aluminum is the most abundant metal in the earth’s crust, however it is found as an oxide in an ore called bauxite. The Hall-Héroult process made the production of aluminum economically reasonable. However, it requires a lot of energy—nearly 5% of the national total. Recycling aluminum saves almost 95% of the cost of production and is the most economically worthwhile recycling program ever developed.

Homework Ch 20.3 pg 671 #1, 3 and pg 673 #21, 23, 24