79. Equivalent of metals from oxidation to oxides. Weigh accurately a small evaporating-dish or a large crucible, place in it about 1 g. of pure copper, and again weigh. Cover the dish with a watch glass and add dilute nitric acid, a little at a time, until the metal is all dissolved. Rinse the watch glass carefully into the dish, and place the latter on the rim of a somewhat smaller beaker two thirds full of water. By boiling the water in the beaker, evaporate the solution of copper nitrate until it is nearly dry. Place the dish on a pipestem triangle and very carefully heat it with a small flame, holding the burner in the hand and using every precaution to prevent loss by spattering. As the nitrate becomes dry, gradually increase the heat until red fumes cease to be given off, finally using the full heat of the burner. Allow the dish to cool and then weigh it. What is the product? From the increase in weight calculate the ratio by weight in which copper and oxygen combine. The experiment may be varied by the use of iron, tin, or zinc in place of copper.

80. Preparation and properties of nitrous acid. In a hemispherical iron dish heat 10 g. of potassium nitrate until it melts and just begins to evolve bubbles; then add 25 g. of lead. Continue the heating for about twenty minutes, stirring the mixture with an iron wire or file. Note the change in color. How do you account for it (R)? When the product becomes cool add 25 cc. of water and heat until the mass is disintegrated. Filter off the residue. What is its composition? What compound is present in the filtrate? Add to the filtrate a few drops of sulfuric acid. Account for the result. Write the equation for the reaction which takes place between sulfuric acid and potassium nitrate; between sulfuric acid and potassium nitrite.

81. Preparation and properties of some of the oxides of nitrogen. a. Nitrous oxide. Put 6 or 8 g. of ammonium nitrate in the hard-glass test tube used in the preparation of oxygen. Attach a delivery tube and heat gently, applying no more heat than is necessary to cause a slow evolution of gas. As soon as the gas is regularly evolved collect two or three bottles of

it over water. Notice the deposit of water on the sides of the test tube. What is the source of it (R)? Note the color, odor, and taste of the gas. Test it with a glowing splint. Account for the result. How can you distinguish it from oxygen? What is the common name of the gas, and for what is it used? Contrast the action of heat on ammonium nitrate and copper nitrate (R).

b. Nitric oxide and nitrogen dioxide. Put a few pieces of copper in your hydrogen generator (hood), just cover them with water, and add 2 or 3 cc. of nitric acid. Collect over water three bottles of the evolved gas, adding more nitric acid if necessary, and leaving the last one half filled in the pneumatic trough. Compare the color of the gas in the generator with that collected in the bottles and account for any difference. Write the equations for all the reactions involved. Uncover one of the bottles containing gas and account for the result. Test the gas in the second bottle with a burning splint. Which is the more stable, nitrous oxide or nitric oxide? Give reasons for your answer. To the third bottle standing over water add air in small portions at a time, transferring the air to the bottle with a test tube. After the addition of each portion of air, allow the red fumes to dissolve before adding another portion. Does the volume diminish indefinitely? Why? If pure oxygen had been added instead of air, would all of the gas in the bottle vanish? Can you suggest a modification of the experiment that could be used to determine the percentage of oxygen in the air?

CHAPTER X

EQUILIBRIUM

82. Velocity of reactions. a. To a dilute solution of sodium chloride add 1 or 2 cc. of a solution of silver nitrate. How rapidly does the reaction take place? How does it compare in rapidity with the reaction occasioned by the addition of a little barium chloride to dilute sulfuric acid? In general, reactions between freely ionized electrolytes are too rapid for measurement. Slow reactions occur when the reagents are very little ionized or when changes occur other than double decomposition between ions.

b. Dissolve a crystal of sodium phosphate no larger than a grain of wheat in 10 cc. of water and add 5 cc. of ammonium molybdate solution (side shelf). Does a precipitate at once form? Do you notice any change of color? Does this remain constant or increase? Set the tube aside and look at it from time to time during the laboratory period. Is the formation of a precipitate gradual or sudden? (The yellow solid has a very complex formula.)

83. Mass action. a. Make a concentrated solution of common salt by shaking 20 g. of salt with the least water that will serve to dissolve it. Add 5 cc. of concentrated sulfuric acid in small portions at a time, shaking gently after each addition (hood). The gas evolved has the formula HCl, the equation being

2 NaCl + H2SO1 = Na2SO1 + 2HC1

4

When all of the acid has been added, warm gently and set the solution aside to crystallize. (If no crystals appear by the time the solution has cooled, add a minute crystal of