Chapter 5 Flashcards by Audrey Booth

Physical Characteristics of Gases

assume the volume and shape of their
containers.
• much lower densities than liquids and
solids.
• most compressible state of matter.
• mix evenly and completely when
confined to the same container.
• differ from liquids and solids –
for a gas, the pressure, volume,
temperature and the amount are
related.

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Pressure

The force exerted per unit area

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Barometer

a device for measuring

the pressure of the atmosphere.

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Manometer

a device for measuring the pressure of a gas or

liquid in a vessel.

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Empirical Gas Laws

All gases behave quite simply
with respect to:
temperature
pressure
volume
molar amount

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At one
atmosphere
the volume of
the gas is

100 mL

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When pressure is doubled, the volume is halved to

50 mL

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When pressure
is tripled, the
volume
decreases to

33 mL

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As P (h) increases, V

decreases

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Boyle’s Law

The volume (V) of a sample of gas at
constant temperature varies inversely with the applied
pressure (P).

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As T increases

V increases (think of a cold balloon)

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absolute zero in degrees c

-273.15°C

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absolute zero definition

It is the temperature at which the volume of a gas is
hypothetically zero.
This is the basis of the absolute temperature scale, the
Kelvin scale (K)

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Charles’s Law

The volume of a sample of gas at
constant pressure is directly proportional to the absolute
temperature (K).

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Avogadro’s Law

Equal volumes of any two gases contain the same number of

molecules (at the same temperature and pressure).

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Equal volumes of any two gases at the same temperature

and pressure contain the same number of

molecules…

but will have different masses.

Standard Temperature and Pressure (STP)

exactly 0°
C
exactly 1 atm pressure
The reference condition for gases (chosen by convention)

At STP, 1 mole of an ideal gas occupies

22.414 L

the following are the

more commonly encountered R values

R = 0.08206 L atm K-1 mol-1
R = 62.36 L mmHg K-1 mol-1
R = 8.315 J K-1 mol-1

The volume of a gas varies with temperature and pressure -

therefore, the density of a gas

also varies with temperature and pressure.

Dalton found that in a mixture of unreactive gases

each
gas acts as if it were the only gas in the mixture as far as pressure is
concerned

Dalton’s Law of Partial Pressures

The sum of the partial pressures of all the different gases in a mixture is
equal to the total pressure of the mixture:
Ptotal = PA + PB + PC + . . .

Mole fraction

The composition of a gas mixture is often described in

terms of the component gases.

Kinetic-Molecular Theory (Kinetic Theory)

A theory, developed by physicists, that is based on the
assumption that a gas consists of molecules in constant
random motion.

The interpretation of a gas in terms of the kinetic molecular theory leads to

the ideal gas law.

Until mid nineteenth century: The popular view at that time (developed by Newton) was that gas pressure was due to the

mutual repulsions of the gas particles | that pushed them against the walls

Present explanation of gas pressure: According to kinetic theory, gas pressure is the result of the

bombardment of the container walls by constantly moving molecules (Maxwell Boltzmann)

Kinetic Molecular Theory of Gases: The Postulates (5)

Gases are composed of molecules whose sizes are negligible compared with the average distance between them. Molecules move randomly in straight lines in all directions and at various speeds. 4. When molecules collide with each other, the collisions are elastic. 5. The average kinetic energy of the molecules is proportional to the temperature of the gas (Kelvin).

An elastic collision is one in which

no kinetic energy is | lost (newton's craddle)

The frequency of collision is proportional to the

average speed, u, and the number of molecules, | N, and inversely proportional to the volume, V.

The average momentum (mu) depends on the

mass of the molecules, m, and | its average velocity, u

Molecular Speeds: | According to kinetic theory

molecular speeds vary over a wide range of values. • The distribution depends on temperature, so it increases as the temperature increases.

Root-Mean Square (rms) Molecular Speed, u:

urms is a type of average molecular speed, equal to the speed of a molecule that has the average molecular kinetic energy.

Gas diffusion

is the gradual mixing of molecules of one gas with molecules | of another by virtue of their kinetic properties.

Effusion:

The process by which a gas flows through a small hole in a | container. A pinprick in a balloon is one example of effusion.

Graham’s Law of Effusion

The rate of effusion of gas molecules through a particular hole is inversely proportional to the square root of the molecular mass of the gas (at constant temperature and pressure).

Real Gases: Deviations from Ideal Behavior

At high pressure, the volume of the gas molecule (Postulate 1) is not negligible. At high pressure, the intermolecular forces (Postulate 3) are not negligible

Real gases do not follow

PV = nRT perfectly. The van der Waals equation corrects for the non-ideal nature of real gases.

Van der Waals Equation

An equation that is similar to the ideal gas law, but which includes two constants, a and b, to account for deviations from ideal behavior