Electrochemistry Cells and Batteries

Galvanic Cells Electrochemistry-- the study of the processes involved in converting chemical energy to electrical energy, and vice versa. Two concepts form the basis: Redox reactions involve the transfer of electrons from one reactant to another. An electric current is a flow of electrons in a circuit.

Galvanic cells Zinc metal reacts with a solution containing copper(II) ions, forming zinc(II) ions and metallic copper. The reaction is spontaneous (a spontaneous reaction is a reaction that occurs by itself; without an external energy source).

Galvanic Cells A galvanic cell, (or voltaic cell), converts chemical energy to electrical energy. For this to work we must prevent the reactants in a redox reaction from coming into direct contact with each other. Instead, electrons flow from one reactant to the other through an external circuit, which is a circuit outside the reaction vessel. This flow of electrons through the external circuit is an electric current.

Galvanic Cells: The Daniell Cell One half of the cell consists of a piece of zinc placed in a zinc sulfate solution. The other half of the cell consists of a piece of copper placed in a copper(II) sulfate solution. A porous barrier, separates these two half-cells. It stops the copper(II) ions from coming into direct contact with the zinc electrode.

Galvanic Cells: The Daniell Cell The pieces of metallic zinc and copper act as electrical conductors. The conductors that carry electrons into and out of a cell are named electrodes. The zinc sulfate and copper(II) sulfate act as electrolytes. Electrolytes are substances that conduct electricity when dissolved in water.

Galvanic Cells: The Daniell Cell The redox reaction takes place in a galvanic cell when an external circuit, such as a metal wire, connects the electrodes. The oxidation half-reaction occurs in one half-cell, and the reduction half-reaction occurs in the other half-cell. For the Daniell cell: Oxidation: Zn(s) → Zn2+(aq) + 2e− Reduction: Cu2+(aq) + 2e− → Cu(s)

Galvanic Cells: The Daniell Cell The electrode at which oxidation occurs is named the anode. Zinc atoms undergo oxidation at the zinc electrode. The zinc electrode is the anode of the Daniell cell. The electrode at which reduction occurs is named the cathode. Copper(II) ions undergo reduction at the copper electrode. Thus, the copper electrode is the cathode of the Daniell cell. Free electrons cannot travel through the solution. Instead, the external circuit conducts electrons from the anode to the cathode of a galvanic cell.

Galvanic Cells: The Daniell Cell At the anode of a galvanic cell, electrons are released by oxidation. The zinc anode of the Daniell cell, zinc metal atoms release electrons to become positive zinc ions. Making the anode of a galvanic cell is negatively charged. Relative to the anode, the cathode of a galvanic cell is positively charged. In galvanic cells, electrons flow through the external circuit from the negative electrode to the positive electrode.

Galvanic Cells: The Daniell Cell Each half-cell contains a solution of a neutral compound. These solutions are aqueous zinc sulfate and aqueous copper(II) sulfate. To maintain electrical neutrality in each half-cell, ions migrate through a porous barrier. Negative ions (anions) migrate toward the anode, and positive ions (cations) migrate toward the cathode. The barrier is sometimes a salt bridge, which contains an electrolyte solution. The open ends of the salt bridge are plugged with a porous material. The plugs allow ion migration to maintain electrical neutrality.

Galvanic Cell Notation The shorthand representation of a Daniell cell is as follows. Zn | Zn2+ || Cu2+ | Cu The phases or states may be included. Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) the anode is always shown on the left and the cathode on the right

Batteries A dry cell is a galvanic cell with the electrolyte contained in a paste thickened with starch. Dry cells are inexpensive. The cheapest AAA-, AA-, C-, and D-size 1.5-V batteries are dry cells. A battery is defined as a set of galvanic cells connected in series. The negative electrode of one cell is connected to the positive electrode of the next cell in the set. The voltage of a set of cells connected in series is the sum of the voltages of the individual cells. A 9-V battery containssix 1.5-V dry cells connected in series. http://www.youtube.com/watch?v=ksxSOwA933M&feature=kp

Cell Potentials The difference between the potential energy at the anode and the potential energy at the cathode is the electric potential, E, (cell voltage or cell potential) of a cell. The unit used to measure electric potential is called the volt, with symbol V. A cell potential can be measured using an electrical device called a voltmeter.

Standard Cell Potentials The half-cell potential for a reduction half-reaction is called a reduction potential. The numerical values of cell potentials and half-cell potentials depend on various conditions, so tables of standard reduction potentials are true when ions and molecules are in their standard states. Each half-cell reduction potential is given relative to the reduction potential of the standard hydrogen electrode, which has been assigned a value of zero.

Calculating Standard Cell Potentials Zn | Zn2+ (1 mol/L) || Cu2+ (1 mol/L) | Cu Method 1: E°cell = E°cathode − E°anode Cu2+(aq) + 2e− Cu(s) E° = 0.342 V Zn2+(aq) + 2e− Zn(s) E° = −0.762 V E°cell = E°cathode − E°anode = 0.342 V − (−0.762 V) = 0.342 V + 0.762 V = 1.104 V The standard cell potential for a Daniell cell is 1.104 V.

Then the oxidation half-reaction is: The half-cell potential for an oxidation half-reaction is called an oxidation potential. If the reduction half-reaction is as follows, Zn2+(aq) + 2e−  Zn(s) E° = −0.762 V Then the oxidation half-reaction is: Zn(s)  Zn2+(aq) + 2e− E°ox = +0.762 V

Method 2: E°cell = E°red + E°ox The standard cell potential can also be calculated as the sum of a standard reduction potential and a standard oxidation potential. Method 2: E°cell = E°red + E°ox Cu2+(aq) + 2e−  Cu(s) E°red = 0.342 V Zn(s) Zn2+(aq) + 2e− E°ox = +0.762 V E°cell = E°red + E°ox = 0.342 V + 0.762 V = 1.104 V

Calculate the standard cell potential for the galvanic cell in which the following reaction occurs. 2I−(aq) + Br2(l) → I2(s) + 2Br−(aq) Practice: page 773 #5-8

Electrolytic Cells Electrolytic cells- a type of cell that uses energy to move electrons from lower potential energy to higher potential energy. They convert electrical energy to chemical energy. They are non-spontaneous, and require energy to occur. The process that takes place in an electrolytic cell is called electrolysis.

Electrolytic Cells Electrolytic cells includes electrode, at least one electrolyte, and an external circuit. Electrolytic cells require an external source of electricity, sometimes called the external voltage.

Electrolytic Cells The external source of electricity forces electrons onto one electrode. This electrode becomes negative relative to the other. Na+ ions move toward the negative electrode, where they gain electrons and are reduced to the element sodium.

Electrolysis Electrolysis is the splitting of water. 2H2O  2H2 + O2 This is a non-spontaneous reaction that requires energy. Electrolytic cells can perform electrolysis.

Electroplating In the galvanic cell, the zinc anode gradually dissolves. The copper cathode grows as more copper is deposited onto it. In the electrolytic cell, the copper anode gradually dissolves. The zinc cathode grows as more zinc is deposited onto it. The process in which a metal is deposited, or plated, onto the cathode in an electrolytic cell is known as electroplating.