SAT Chemistry Gases and the Gas Laws - Some Representative Gases

SAT Chemistry Gases and the Gas Laws - 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.)
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.
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.
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 (03). 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 quanti-ties 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.

PREPARATION OF HYDROGEN. Although there is evidence of the preparation of hydrogen before 1766, Henry Cavendish was the first person to recognize this gas as a separate sub-stance. He observed that, whenever it burned, it produced water. Lavoisier named it hydro-gen, 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.

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 hydro-gen in sufficient quantities to collect in the laboratory. This general equation is:
Active metal + Dilute acid → Hydrogen + Salt of the acid
An example:
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.
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 1/14 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.

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 pres¬sure 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.

Measuring the Pressure of a Gas
Pressure is defined as force per unit area. With respect to the atmosphere, pressure is the result of the weight of a mixture of gases. This pressure, which is called atmospheric pressure, air pressure, or barometric pressure, is approximately equal to the weight of a kilogram mass on every square centimeter of surface exposed to it. This weight is about 10 newtons.
The pressure of the atmosphere varies with altitude. At higher altitudes, the weight of the overlying atmosphere is less, so the pressure is less. Air pressure also varies somewhat with weather conditions as low- and high-pressure areas move with weather fronts. On the average, however, the air pressure at sea level can support a column of mercury 760 millimeters in height. This average sea-level air pressure is known as normal atmospheric pressure, also called standard pressure.
The instrument most commonly used for measuring air pressure is the mercury barom­eter. The diagram below shows how it operates. Atmospheric pressure is exerted on the mercury in the dish, and this in turn holds the column of mercury up in the tube. This column at standard pressure will measure 760 millimeters above the level of the mercury in the dish below.
In gas-law problems pressure may be expressed in various units. One standard atmosphere (1 atm) is equal to 760 millimeters of mercury (760 mm Hg) or 760 torr, a unit named for Evangelista Torricelli. In the SI system, the unit of pressure is the pascal (Pa), named in honor of the scientist of the same name, and standard pressure is 101,325 pascals or 101.325 kilo- pascals (kPa). One pascal (Pa) is defined as the pressure exerted by the force of one newton (1 N) acting on an area of one square meter. In many cases, as in atmospheric pressure, it is more convenient to express pressure in kilopascals (kPa).

A device similar to the barometer can be used to measure the pressure of a gas in a con-fined container. This apparatus, called a manometer, is illustrated below. A manometer is basically a U-tube containing mercury or some other liquid. When both ends are open to the air, as in (1) in the diagram, the level of the liquid will be the same on both sides since the same pressure is being exerted on both ends of the tube. In (2) and (3), a vessel is connected to one end of the U-tube. Now the height of the mercury column serves as a means of reading the pressure inside the vessel if the atmospheric pressure is known. When the pressure inside the vessel is the same as the atmospheric pressure outside, the levels of liquid are the same. When the pressure inside is greater than outside, the column of liquid will be higher on the side that is exposed to the air, as in (2). Conversely, when the pressure inside the vessel is less than the outside atmospheric pressure, the additional pressure will force the liquid to a higher level on the side near the vessel, as in (3).
Kinetic-Molecular Theory
By indirect observations, the Kinetic-Molecular Theory has been arrived at to explain the forces between molecules and the energy the molecules possess. There are three basic assumptions to the Kinetic-Molecular Theory:
1. Matter in all its forms (solid, liquid, and gas) is composed of extremely small particles. In many cases these are called molecules. The space occupied by the gas particles themselves is ignored in comparison with the volume of the space in which they are contained.
2. The particles of matter are in constant motion. In solids, this motion is restricted to a small space. In liquids, the particles have a more random pattern but still are restricted to a kind of rolling over one another. In a gas, the particles are in continuous, random, straight-line motion.
3. When these particles collide with each other or with the walls of the container, there is no loss of energy.
Some Particular Properties of Gases
As the temperature of a gas is increased, its kinetic energy is increased, thereby increasing the random motion. At a particular temperature not all the particles have the same kinetic energy, but the temperature is a measure of the average kinetic energy of the particles. A graph of the various kinetic energies resembles a normal bell-shaped curve with the average found at the peak of the curve (see Figure 20).

When the temperature is lowered, the gas reaches a point at which the kinetic energy can no longer overcome the attractive forces between the particles (or molecules) and the gas condenses to a liquid. The temperature at which this condensation occurs is related to the type of substance the gas is composed of and the type of bonding in the molecules them­selves. This relationship of bond type to condensation point (or boiling point) is pointed out in Chapter 3, “Bonding.”
The random motion of gases in moving from one position to another is referred to as dif­fusion. You know that, if a bottle of perfume is opened in one comer of a room, the perfume, that is, its molecules, will move or diffuse to all parts of the room in time. The rate of diffusion is the rate of the mixing of gases.
Effusion is the term used to describe the passage of a gas through a tiny orifice into an evacuated chamber. The rate of effusion measures the speed at which the gas is transferred into the chamber.


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