Chapter 5 Flashcards
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.
Pressure
The force exerted per unit area
Barometer
a device for measuring
the pressure of the atmosphere.
Manometer
a device for measuring the pressure of a gas or
liquid in a vessel.
Empirical Gas Laws
All gases behave quite simply with respect to: temperature pressure volume molar amount
At one
atmosphere
the volume of
the gas is
100 mL
When pressure is doubled, the volume is halved to
50 mL
When pressure
is tripled, the
volume
decreases to
33 mL
As P (h) increases, V
decreases
Boyle’s Law
The volume (V) of a sample of gas at constant temperature varies inversely with the applied pressure (P).
As T increases
V increases (think of a cold balloon)
absolute zero in degrees c
-273.15°C
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)
Charles’s Law
The volume of a sample of gas at
constant pressure is directly proportional to the absolute
temperature (K).
Avogadro’s Law
Equal volumes of any two gases contain the same number of
molecules (at the same temperature and pressure).
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.
- When molecules collide with each other, the collisions are
elastic. - 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
