edlted by
GEORGE L. GILBERT Denison University Granville. Ohio 43023
Safety in Oxygen Atmospheres (Don't worry, oxygen doesn't burn!)
Submitted bv:
Homer T. Knight Norbert lsenherg T h e University of Wisconsin-Parkside Kenosha, WI 53141 Checked by: Allen M. Scoffstall University of Colorado a t Colorado Springs Colorado Springs, CO 80907 On a flight from 1.m Angeles to Nrw York several years ago, a stewardess decided that one of the passengers needed oxygen. Before we could do anything or say anything, she started her procedure and she ignored some of our protests that she should ask people to extinguish their cigarettes. Her reply was simple: "Oh, don't worry, oxygen doesn't burn!" One of our co-nasseneers tried to correct her hut to no avail. Fortunatelv for us, the oxygen tank was empty. The trouble with safetv is that i t takes ex~erienceto make believers out of most of us. Accidents always happen to the other person; even when we have all the warning signs we chose to ignore them. No wonder that our students find i t so difficult to take some simple precautions. llsually we discuss the fact& which influence the rate of a chemical reaction, but we may not relate them tu everyday events. Surely, when we increase the concentration of our reagents, we increase the rate of the reaction. Things burn faster in oxygen; we know i t and we can show it-we use deflagration spoons containing sulfur or phosphoros or we burn some steel wool in oxygen. How many people "in real lifemgo around burning sulfur, phosphorus, or steel wool in oxveen? "Occasionally, we read about injuries to persons in oxygen atmosoheres: sometimes these accidents are fatal. Perhaps the best kno& accident involved three astronauts, ~ i r g i l ~ . Grissom, Edward H. White, and Roger B. Chafee, who died in their space vehicle at Cape Kennedy on January 27,1967, when a flash fire swept through their oxygen-filled space capsule.' There is a definite need for safety in oxygen atmospheres, including oxygen tents, masks, and other devices used in oxygen therapy.2J A practical way to illustrate the influence of concentration on the rate of a chemical reaction is to show a simulated oxygen tent to the class. The demonstration can be carried out as follows. Two olastic containers (sandwich haes or somewhat larger bags) a;e used in the demonstration (Ge Note I ). One container is filled withoxvaenand theother is filled withair.The containers are sealedtightly with aruhber band or a wire tie and are suspended from a suitable support on a ring stand. To avoid loss of oxygen through leakage i t is advisable to fill the containers shortly before class time. The instructor approaches the hag filled with air and uses a burning cigarette or a wwden splint to bum a small hole into the bag. When the
process is repeated with the oxygen-filled container, the plastic material bums rapidly and melts. This demonstration can he shown to a large class and the result is rather dramatic. One other precaution is necessary; a large pan or metal plate should he placed under the base of the ring stand so that the meltine olastic does not damaee the too of the demonstration table o i l h e floor. Note 1:The use of too large a container should be avoided t o prevent major fire.
a
Cobalt Complexes in Equilibrium Submitted by:
Charles E. Ophardt Elmhurst College Elmhurst, IL 60126
Checked by:
Wayne L. Smith Colby College Wateruille, ME 04901
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'Encyclopedia Britannica, No. 4 (1974),p. 747. ZGrenard,S., "Hazards of Respiration Therapy," Glenn Educational Medical Services, Sarasota, FL, 1973. SSeedor, M. M., "Therapy with Oqgen and Other Gases," Teachers College Press, Columbia University, New York, NY, 1966.
Preparation
Pour 5 ml of 0.2 M cobalt(I1) chloride mlution in each of two 15 X 150-mm test tubes. Have available the following: 10 ml 11.5 M HC1,5 ml distilled water, 5 mlO.l M AgN03, and 10 ml of acetone, and an empty 15 X 150 mm test tube? Demonstration I ) Into one test tube with robalt(ll) chloriden,lution,uk,wlyadd 10 ml uf 11.5 M HCI. Olmerve rulor change to blue 2) Pour half of the hlue solution inb, enother test tube. Add 5 ml of distilled water to one of the tuhes and observe color change back to pink. 3) To the blue tube add 5 mlO.l M AgNOs and observe precipitate and color change back to pink. 4) Into the second test tube. with cobalt(I1)chloride solution. slowlv ' with no mixing, add 10'ml of acetoneon top of the cobalt(1i) chloride solution. Observe the blue layer on top of a pink salution. Remarks This equilibrium illustrates the application of LeChatelier's principle to concentration effects. The equilibrium is [ C O ( H ~ O )+ ~ ]4CI+ ~ = [COCI~]-~ + 6Hz0 pink blue 1) In the first experiment the addition of excess chloride from HCI causes the equilibrium to shift right forming the blue tetrachlorocobalt(I1)complex. 2) The addition of excess water to the blue complex shifts the equilibrium left farming the pink heraaquocobalt(I1)complex. 3) Chloride is removed by precipitation with AgNOa. This causes the equilibrium to shift left forming the pink aquo complex. 4) The acetone hydrogen bonds to water and in effectdecreaseswater concentration on the top of the solution. The equilibrium shifts right to form the blue chloro complex. 'The demonstration may easily he scaled up to larger quantities hv multi~lvinp: . . .the amounts by 5 or 10. Demonstration &4 can be done in a large graduate cylinder. Volume 57, Number 8, June 1980 1 453
