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The Gas Laws
Ideal Gases
You will definitely see some questions on gases and the
laws that govern them on the SAT II Chemistry exam. All of the gas
laws rely on some basic assumptions that are made about gases, and
together they constitute what it means for a gas to be in an ideal
state. In an ideal state
- All gas particles are in constant, random motion.
- All collisions between gas particles are perfectly elastic (meaning that the kinetic energy of the system is conserved).
- The volume of the gas molecules in a gas is negligible.
- Gases have no intermolecular attractive or repulsive forces.
- The average kinetic energy of the gas is directly proportional to its Kelvin temperature and is the same for all gases at a specified temperature.
Only four measurable properties are used to describe a
gas: its quantity, temperature, volume, and pressure. The quantity
(amount) of the gas is usually expressed in moles (n).
The temperature, T, of gases must
always be converted to the Kelvin temperature scale (the absolute
temperature scale). The volume, V,
of a gas is usually given in liters. Finally, the pressure, P,
of a gas is usually expressed in atmospheres. Gases are often discussed
in terms of standard temperature and pressure (STP),
which means 273K (or 0ºC) and 1 atm.
Example
Which of the following statements is not true of ideal
gases?
- The volume occupied by gas particles is only significant at very low pressures.
- Gas molecules occupy an insignificant volume compared to the volume of the container that holds them.
- The particles of a gas move in random straight line paths until a collision occurs.
- The collisions that occur between gas particles are considered elastic.
- At a given temperature, all gas molecules within a sample possess the same average kinetic energy.
Explanation
In this example, choice 1 is incorrect. Choices
2, 3, 4, and 5 all describe an ideal gas. Choice 1 makes an incorrect
assumption: it begins with a true statement about volume not being very
significant but then turns around and gives the incorrect scenario—if
the pressure is low, then gas particles undergo very few collisions,
so the volume is insignificant. The volume only becomes significant
if gas particles collide often, increasing the chances that intermolecular
forces will hold them together.
Measuring the Pressure of a Gas
Gas pressure is a gauge of the number and force of collisions
between gas particles and the walls of the container that holds
them. The SI unit for pressure is the pascal (Pa),
but other pressure terms include atmospheres (atms), millimeters
of mercury (mmHg), and torr. The following is
a list of all of the standard pressure in every unit for pressure.
Memorize these for the exam so you can convert units where necessary:
760 mmHg
760 torr
1.00 atm
101,325 Pa
101.325 kPa
The piece of lab equipment specifically designed to measure
the pressure of gases is known as the barometer. A barometer uses
the height of a column of mercury to measure gas pressure in millimeters
of mercury or torr (1 mmHg = 1 torr). The mercury is pushed up the tube
from the dish until the pressure at the bottom of the tube (due
to the mass of the mercury) is balanced by the atmospheric pressure.
When using a barometer, you calculate gas pressure with
the following equation:
Gas pressure = atmospheric pressure - h (height
of the mercury)
The open-tube manometer is another device
that can be used to measure pressure. The open-tube manometer is
used to measure the pressure of a gas in a container.
The pressure of the gas is given by h (the
difference in mercury levels) in units of torr or mmHg. Atmospheric
pressure pushes on the mercury from one direction, and the gas in the
container pushes from the other direction. In a manometer, since
the gas in the bulb is pushing more than the atmospheric pressure,
you add the atmospheric pressure to the height difference:
gas pressure = atmospheric pressure + h
There is one other possibility for a manometer question
that could appear on the SAT II Chemistry test: they could ask you
about a closed-tube manometer. Closed-tube manometers
look similar to regular manometers except that the end that’s open
to the atmospheric pressure in a regular manometer is sealed and
contains a vacuum. In these systems, the difference in mercury levels
(in mmHg) is equal to the pressure in torr.
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