Page 221 - Theory and Problems of BEGINNING CHEMISTRY

210 OXIDATION AND REDUCTION [CHAP. 14

If you electrolyze a concentrated solution of NaCl instead, H 2 is produced at the cathode and Cl 2 is produced at
the anode:
electricity
−
2 e + 2H 2 O −−−→ H 2 + 2OH −
NaCl
electricity
−
2Cl −−−→ Cl 2 + 2 e −
NaCl
It is obvious that the reaction conditions are very important to the products.
Electrolysis is used in a wide variety of ways. Three examples follow: (1) Electrolysis cells are used
to produce very active elements in their elemental form. The aluminum industry is based on the electrolytic
reduction of aluminum oxide, for example. (2) Electrolysis may be used to electroplate objects. A thin layer
of metal, such as silver, can be deposited on other metals, such as steel, by electrodeposition (Fig. 14-2).
(3) Electrolysis is also used to purify metals, such as copper. Copper is thus made suitable to conduct electricity.
The anode is made out of the impure material; the cathode is made from a thin piece of pure copper. Under
carefully controlled conditions, copper goes into solution at the anode, but less active metals, notably silver and
gold, fall to the bottom of the container. The copper ion deposits on the cathode, but more active metals stay
in solution. Thus very pure copper is produced. The pure copper turns out to be less expensive than the impure
copper, which is not too surprising when you think about it. (Which would you expect to be more expensive,
pure copper or a copper-silver-gold mixture?)

Battery
Pure silver
Spoon

Ag +

NO 3 −

Fig. 14-2. Electroplating

Galvanic Cells
When you place a piece of zinc metal into a solution of CuSO 4 , you expect a chemical reaction because the
more active zinc displaces the less active copper from its compound (Sec. 8.3). This is an oxidation-reduction
reaction, involving transfer of electrons from zinc to copper.
Zn −→ Zn 2+ + 2 e −

Cu 2+ + 2 e −→ Cu
−
It is possible to carry out these same half-reactions in different places if we connect them suitably. We must
2+
deliver the electrons from Zn to Cu , and we must have a complete circuit. The apparatus is shown in Fig. 14-3.
A galvanic cell with this particular combination of reactants is called a Daniell cell. The pieces of zinc and
copper serve as electrodes, at which chemical reaction takes place. It is at the electrodes that the electron current
is changed to an ion current or vice versa. The salt bridge is necessary to complete the circuit. If it were not there,
the buildup of charge in each beaker (positive in the left, negative in the right) would stop the reaction extremely
quickly (less than 1 s). The same chemical reactions are taking place in this apparatus as would take place if we
dipped zinc metal in CuSO 4 solution, but the zinc half-reaction is taking place in the left beaker and the copper
half-reaction is taking place in the right beaker. Electrons flow from left to right in the wire, and they could be
made to do electrical work, such as lighting a small bulb. To keep the beakers from acquiring a charge, cations