SAT Subject Test Chemistry
REVIEW OF MAJOR TOPICS
Gases and the Gas Laws
These skills are usually tested on the SAT Subject Test in Chemistry. You should be able to …
• Describe the physical and chemical properties of oxygen and hydrogen and the electronic makeup of their diatomic molecules.
• Explain how atmospheric pressure is measured, how to read the pressure in a manometer, and the units used to measure pressure.
• Read and explain a graphic distribution of the number of molecules versus kinetic energy at different temperatures.
• Know and use the following laws to solve gas problems: Graham’s, Charles’s, Boyle’s, Dalton’s, the Combined Gas Law, and the Ideal Gas Law.
This chapter will review and strengthen these skills. Be sure to do the Practice Exercises at the end of the chapter.
When we discuss gases today, the most pressing concern is the gases in our atmosphere. These are the gases that are held against Earth by the gravitational field. The principal constituents of the atmosphere of Earth today are nitrogen (78%) and oxygen (21%). The gases in the remaining 1% are argon (0.9%), carbon dioxide (0.03%), varying amounts of water vapor, and trace amounts of hydrogen, ozone, methane, carbon monoxide, helium, neon, krypton, and xenon. Oxides and other pollutants added to the atmosphere by factories and automobiles have become a major concern because of their damaging effects in the form of acid rain. In addition, a strong possibility exists that the steady increase in atmospheric carbon dioxide, mainly attributed to fossil fuel combustion over the past century, may affect Earth’s climate by causing a greenhouse effect, resulting in a steady rise in temperatures worldwide.
The major components of Earth’s atmosphere: 78% nitrogen, 21% oxygen
Studies of air samples show that up to 55 miles above sea level the composition of the atmosphere is substantially the same as at ground level; continuous stirring produced by atmospheric currents counteracts the tendency of the heavier gases to settle below the lighter ones. In the lower atmosphere, ozone is normally present in extremely low concentrations. The atmospheric layer 12 to 30 miles up contains more ozone that is produced by the action of ultraviolet radiation from the Sun. In this layer, however, the percentage of ozone is only 0.001 by volume. Human activity adds to the ozone concentration in the lower atmosphere where it can be a harmful pollutant.
The ozone layer became a subject of concern in the early 1970s when it was found that chemicals known as fluorocarbons, or chlorofluoromethanes, were rising into the atmosphere in large quantities because of their use as refrigerants and as propellants in aerosol dispensers. The concern centered on the possibility that these compounds, through the action of sunlight, could chemically attack and destroy stratospheric ozone, which protects Earth’s surface from excessive ultraviolet radiation. As a result, U.S. industries and the Environmental Protection Agency phased out the use of certain chlorocarbons and fluorocarbons as of the year 2000. There is still ongoing concern about both these environmental problems: the greenhouse effect and the deterioration of the ozone layer as it relates to possible global warming.
SOME REPRESENTATIVE GASES
Of the gases that occur in the atmosphere, the most important one to us is oxygen. Although it makes up only approximately 21% of the atmosphere, by volume, the oxygen found on Earth is equal in weight to all the other elements combined. About 50% of Earth’s crust (including the waters on Earth, and the air surrounding it) is oxygen. (Note Figure 15.)
Figure 15. Composition of Earth’s Crust
The composition of air varies slightly from place to place because air is a mixture of gases. The composition by volume is approximately as follows: nitrogen, 78%; oxygen, 21%; argon, 1%. There are also small amounts of carbon dioxide, water vapor, and trace gases.
PREPARATION OF OXYGEN. In 1774, an English scientist named Joseph Priestley discovered oxygen by heating mercuric oxide in an enclosed container with a magnifying glass. That mercuric oxide decomposes into oxygen and mercury can be expressed in an equation: 2HgO → 2Hg + O2. After his discovery, Priestley visited one of the greatest of all scientists, Antoine Lavoisier, in Paris. As early as 1773 Lavoisier had carried on experiments concerning burning, and they had caused him to doubt the phlogiston theory (that a substance called phlogiston was released when a substance burned; the theory went through several modifications before it was finally abandoned). By 1775, Lavoisier had demonstrated the true nature of burning and called the resulting gas “oxygen.”
Today oxygen is usually prepared in the lab by heating an easily decomposed oxygen compound such as potassium chlorate (KClO3). The equation for this reaction is:
2KClO3 + MnO2 → 2KCl + 3O2(g) + MnO2
A possible laboratory setup is shown in Figure 16.
Figure 16. A Possible Laboratory Preparation of Oxygen
In this preparation manganese dioxide (MnO2) is often used. This compound is not used up in the reaction and can be shown to have the same composition as it had before the reaction occurred. The only effect it has is that it lowers the temperature needed to decompose the KClO3, and thus speeds up the reaction. Substances that behave in this manner are referred to as catalysts. The mechanism by which a catalyst acts is not completely understood in all cases, but it is known that in some reactions the catalyst does change its structure temporarily. Its effect is shown graphically in the reaction graphs in Figure 17.
A catalyst speeds up the rate of reaction by lowering the activation energy needed for the reaction. A catalyst is not consumed.
Figure 17. Effect of Catalyst on Reaction
Graphic representation of how a catalyst lowers the required activation energy
PROPERTIES OF OXYGEN. Oxygen is a gas under ordinary conditions of temperature and pressure, and it is a gas that is colorless, odorless, tasteless, and slightly heavier than air; all these physical properties are characteristic of this element. Oxygen is only slightly soluble in water, thus making it possible to collect the gas over water, as shown in Figure 16.
Although oxygen will support combustion, it will not burn. This is one of its chemical properties. The usual test for oxygen is to lower a glowing splint into the gas and see if the oxidation increases in its rate to reignite the splint. (Note: This is not the only gas that does this. N2O reacts the same.)
OZONE. Ozone is another form of oxygen and contains three atoms in its molecular structure (O3). Since ordinary oxygen and ozone differ in energy content and form, they have slightly different properties. They are called allotropic forms of oxygen. Ozone occurs in small quantities in the upper layers of Earth’s atmosphere, and can be formed in the lower atmosphere, where high-voltage electricity in lightning passes through the air. This formation of ozone also occurs around machinery using high voltage. The reaction can be shown by this equation:
3O2 + elec. → 2O3
Because of its higher energy content, ozone is more reactive chemically than oxygen.
The ozone layer prevents harmful wavelengths of ultraviolet (UV) light from passing through Earth’s atmosphere. UV rays have been linked to biological consequences such as skin cancer.
The ozone layer protects us from UV rays from the sun.
PREPARATION OF HYDROGEN. Although there is evidence of the preparation of hydrogen before 1766, Henry Cavandish was the first person to recognize this gas as a separate substance. He observed that, whenever it burned, it produced water. Lavoisier named it hydrogen, which means “water former.”
Electrolysis of water, which is the process of passing an electric current through water to cause it to decompose, is one method of obtaining hydrogen. This is a widely used commercial method, as well as a laboratory method.
Another method of producing hydrogen is to displace it from the water molecule by using a metal. To choose the metal you must be familiar with its activity with respect to hydrogen. The activities of the common metals are shown in Table 8.
Know the relative activity of metals.
As noted in Table 8, any of the first three metals will react with cold water; the reaction is as follows:
Very active metal + Water = Hydrogen + Metal hydroxide
Using sodium as an example:
2Na + 2HOH → H2(g) + 2NaOH
With the metals that react more slowly, a dilute acid reaction is needed to produce hydrogen in sufficient quantities to collect in the laboratory. This general equation is:
Active metal + Dilute acid → Hydrogen + Salt of the acid
Zn + dil. H2SO4 → H2(g) + ZnSO4
This equation shows the usual laboratory method of preparing hydrogen. Mossy zinc is used in a setup as shown in Figure 18. The acid is introduced down the thistle tube after the zinc is placed in the reacting bottle. In this sort of setup, you would not begin collecting the gas that bubbles out of the delivery tube for a few minutes so that the air in the system has a chance to be expelled and you can collect a rather pure volume of the gas generated.
Figure 18. Preparation of an Insoluble Gas by the Addition of Liquid to Other Reactant
In industry, hydrogen is produced by (1) the electrolysis of water, (2) passing steam over red-hot iron or through hot coke, or (3) by decomposing natural gas (mostly methane, CH4) with heat (CH4+ H2O → CO + 3H2).
PROPERTIES OF HYDROGEN. Hydrogen has the following important physical properties:
1. It is ordinarily a gas; colorless, odorless, tasteless when pure.
2. It weighs 0.9 gram per liter at 0°C and 1 atmosphere pressure. This is as dense as air.
3. It is slightly soluble in water.
4. It becomes a liquid at a temperature of −240°C and a pressure of 13 atmospheres.
5. It diffuses (moves from place to place in gases) more rapidly than any other gas. This property can be demonstrated as shown in Figure 19.
Figure 19. Diffusion of Hydrogen
Here the H2 in the beaker that is placed over the porous cup diffuses faster through the cup than the air can diffuse out. Consequently, there is a pressure buildup in the cup, which pushes the gas out through the water in the lower beaker.
The chemical properties of hydrogen are:
1. It burns in air or in oxygen, giving off large amounts of heat. Its high heat of combustion makes it a good fuel.
2. It does not support ordinary combustion.
3. It is a good reducing agent in that it withdraws oxygen from many hot metal oxides.