edited by GEORGEL. GILBERT Denison University Granvilie. Ohio 43023
Electrodepositionof Nickel on Copper SUBMITTEDBY
Joseph Manlkowl and Dana LevlneZ New Jersey lnrtitute ot Technology Newark, N J 07102 CHECKED BY
Steve Rowley Carleton College Northfield, M N 55057
Michael Faraday discovered that an electric current flowing through an aqueous solution of a metal salt will deposit the metal on the electrode serving as the cathode. Two familiar examples of electrodeposition are chromium-plated automobile bumpers and gold-plated jewelry. The former acts as protection against corrosion and the latter improves the appearance of less costly materials. In the classroom one can demonstrate electroplating by the simple, fast, inexpensive, and visually interesting reaction between nickel ions and copper metal. The half reactions are anode cathode
Cu 2e
+ Ni2+
-
+2
+ cU2+
Ni
ex= -0.34 V eed = -0.25 V
T o carry out the experiment, one needs a 6-V battery, a tall, l-L beaker, 1L of 1MNiC12.6H20 in 0.1 M HCI, and two strips of copper metal, 6 X 1 X %, in.. to act as electrodes. I t is important the copper cathode-have a clean, nonoxidized surface (steel wool or a metal cleaner such as Noxon works well) because dirt or oxides will interfere with the plating. The copper strips, attached with alligator clips to the banana leads of the battery, are immersed in the acidic nickel ion solution. Within 60 s, nickel will evenly coat the copper cathode (see figure). The nickel-plated metal can he bent into a bracelet and given to a member of the class. Prior to ~erforminethe exneriment. one mieht " show that ~-~~~~ no reaction will occur with& electricity. After the plating demonstration, one can discuss oxidation ootentials and the relative positions of nickel and copper in the activity series. An ammeter in the circuit will measure the current flow: however, it is not possible to determine accurately the mass of nickel plated out in a given amount of time because hvdro.gen ions are reduced a t t h e cathode as well as nickel ions. A variable resistor in the circuit will decrease the amount of hydrogen gas produced, but it will also slow the electrodeposition of the nickel. Other variations of this demonstration might be to use a solar-powered cell to power the reaction or common metallic objects to serve as the cathode, such as brass keys or nails." ~
~
~
The overall reaction is
-
Cu + Ni2+
Cu2++ Ni
EL, = -0.59 V
' Present address
Lnwers ly of Callforma Berkeley. CA AJhor lo whom correspondence snoulo oe aodressed 'D x G , rlultscn R J Chem € d m 1978. 55. 259
The Electrolysis of Water: An Improved Demonstration Procedure SUSM~-0
BY
Stephen Heideman Iowa State University Arnes. IA 50011
CHECKED BY
George Wollaston Clarion University Clarion, PA 16214
Nickel ions, inasalutionaf 1 MNiClp3H10 in 0.1 MHCI, plating onthecathode 01 copper metal. The anode is also copper. Electric current supplied by a 6 V baltery.
The use of dilute sulfuric acid as the electrolvte in the usual demonstration of the electrolysis of water] does not allow the students to observe the accomnanvine OH chanees a t the electrodes. A simple modification df ti;; procedEre allows all the chemical reactions that occur in this demonstration to be observed and used to elucidate each half-cell reaction as well as the net cell reaction. The procedure described here uses a 1M Na2S04solution, to which hromocresol green indicator has been added, as the electrolyte. The pH of this solution is adjusted to near 4.5, the approximate pK, of bromocresol green. As the electrolvsis proceeds the color of the solution at each electrode changes indicating changes in the pH of'the solution near each electrode. Volume 63 Number 9
September 1986
809
Equipment
12-Vde power supply electrolysis apparatus 2 250-mL beakers 1600-mLbeaker white paper or poster board (to use as a background) Reagents
400 mL of l M NazS04 20 mL bromocresol green indicator solution (Prepare by dissolving 0.1 zof the indicator in 15 mL of 0.01 MNaOH, then dilute to 250 m ~ w i t hwater.) 1 MCHICOOH I MH2S04 l M NaOH Demonstration
Add 20 mL of the bromocresol green indicator to 400 mL of 1M NalSOa. Adiust the DH of the solution to anoroximately p ~ b . by 5 adding t h e i M CH3COOH dropwigwhile stirrina. The solution will have a clear. areen color a t this DH. Fill the electrolysis unit with the sdLtion and divide'the remainina solution eauallv between two beakers. Connect the dc poker supply i o start the electrolysis and allow the experiment to run for several minutes to produce sufficient volumes of the gases to be easily seen by the students and to dlow the indicator to change color. Although the time required varies from apparatus to apparatus, as a rule i t is necessary to allow a minimum of 15 to 20 min of electrolysis time to produce sufficient volumes of the gases for the students to observe easily the 2 1 volume ratio of hydrogen to oxygen.
into flame. Add some 1 M H2S04 to one of the beakers containing the indicator to demonstrate that the acid form of bromocresol green has a yellow color. At the cathode: The solution is a deep blue color with two unit volumes of gas present above it. If the gas is collected and brought near a flame a "pop" is heard. Add some 1 M NaOH to the second beaker of indicator to show that the base form of bromocresol green has a blue color. Drain all the solution from the electrolysis unit into a 600mL beaker, allowing the students to observe the indicator return to its initial green color. Conclusions
Oxvaen aas and hvdronium ions a r e ~ r o d u c e da t the anode w h i ~ ~ h y d r o g egasand n hydroxide ions are produced a t the cathode. Two moles of hvdroaen . - pas - are produced at the cathode for each mole of oxygen gas produced at the anode. The acid and base, produced a t the anode and cathode, respectively, just neutralize each other. From these observations and conclusions the following reactions may he written.
--
At the anode 6 HzO(1) 4 H30t(aq) + Odg) + 4 eAt the cathode 4 HzO(l)+ 4 e4 OH- + 2 Hdg) Net cell reaction 2 HzO(l) 2 Hdg) + Oz(g) The observations made in this demonstration lead directly to the writing of the appropriate equations for the reactions at each electrode. The combination of these two electrode reactions yields the net cell reaction while giving the student a more complete explanation of the chemical reactions involved in the electrolysis of water.
Obse~atlons
Atthe anode: The solution has turned light yellow in color, and one unit volume of gas is present above it. If the gas is collected and a glowing splint thrust into it, the splint bursts
810
Journal of Chemical Education
'
Alyea, H. N. "Tested Demonstrations in Chemistry", 6th ad.; Journal of Chemical Education: Easton, PA, 1965.
