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Lecture 7:
Gaseous and Liquid compounds
Course lecturer :
Jasmin Šutković
17th April 2014
Contents
International University of Sarajevo
1. Gaseous properties and compounds
2. Pressure, Volume, Temperature and Amount
relationship
3. The IDEAL GAS Law
4. Mixture of gases
4. Kinetic energy in gases
5. The REAL gas behavior
6. LIQUIDs and its kinetic energy
7. Intermolecular forces in liquids
8. Unique properties of LIQUIDS
9. Vapor pressure
10. CHANGE of states
11. Critical temperatures and pressures
1. Gaseous Elements and
Compounds
The three common phases (or states) of matter are gas, liquid, and solid

Gases
a. Have the lowest density of the three states of matter
b. Are highly compressible
c. Completely fill any container in which they are placed
d. Their intermolecular forces are weak
e. Molecules are constantly moving independently of the other
molecules present
 Liquids
a. Dense
b. Incompressible
c. Flow readily to adapt to the shape of the container
d. Sum of the intermolecular forces are between those of gases
and solids
Gaseous Elements and
Compounds
The state of a given substance depends strongly on conditions
 Solids
a. Dense
b. Rigid
c. Incompressible
d. Intermolecular forces are strong
e. Molecules locked in place
PSE location of G,L and S
elements
The figure of the periodic table shows the locations in the periodic table of those
elements that are commonly found in the gaseous, liquid, and solid states
Gaseous Elements and
Compounds
Many of the elements and compounds are typically found
as gases
• Gaseous substances include:
1. Many binary hydrides, such as the hydrogen halides
2. Hydrides of the Group-15 elements N, P, and As
3. Hydrides of the Group-14 elements C, Si, and Ge
4. Many of the simple covalent oxides of the nonmetals such
as
CO, CO2, NO, NO2, SO2, SO3, and ClO2
5. Many low-molecular-mass organic compounds
6. Most of the commonly used refrigerants
Gaseous Elements and
Compounds
• Gaseous substances contain covalent or polar
covalent bonds and are nonpolar or slightly polar
molecules.
• Boiling points of polar compounds are higher than those
of nonpolar compounds of similar molecular mass.
Gas Pressure
• At the macroscopic level, in order to describe the GAS
we require four quantities:
1. Temperature (expressed in K or Celsius (C))
2. Volume (expressed in liters (L) )
3. Amount (expressed in moles (n))
4. Pressure (given in atmospheres (atm))
• These variables are not independent — if the values of
any three of these quantities are known, the fourth can
be calculated.
Units of Pressure
Pressure is any force exerted from anything on any
surface with which it comes in contact.
Units of pressure are derived from the units used to
measure force and area
For scientific measurements, the S unit for pressure, is the newton
per square meter, N/m2, called the paschal (Pa):
1 pascal (Pa) = 1 newton/meter2 (N/m2)
Atmospheric Pressure
• Every point on Earth’s surface experiences a net
pressure called atmospheric pressure.
• Pressure is due to the weight of air above that surface in
the atmosphere of Earth (or that of another planet)
Air has weight.
Imagine a column ( some space) of air that is 2.5 square
meter in cross-section that extends vertically above
your head until the air runs out in space. At sea level
that column of air weighs around 7kg . If you climb
mount Everest there is much less air above you in that
column, so the weight of air above you will be much less
than at sea level.
Athmospheric pressure
example
This plastic bottle was sealed at approximately 4,300 m altitude, and
was crushed by the increase in atmospheric pressure —at 2,700 m
and 300 m— as it was brought down towards sea level.
Measurment of A.P
 Atmospheric pressure can be measured using a
barometer, a closed, inverted tube filled with mercury.

The height of the mercury column is proportional to the
atmospheric pressure, which is reported in units of millimeters of
mercury (mmHg), also called torr.
One (1) bar is atmospheric pressure at sea level.
1atm = 760mmHg = 760torr = 1.01325x10^5 Pa = 101.325
kPa = 14.7 psi
2. The Relationship between Pressure
and Volume
 Pressure on a gas increases, the volume of the gas
decreases ! (gas particles are forced closer together)
 Pressure on a gas gas decreases, the gas volume
increases ! (gas particles can now move farther apart)
 Boyle carried out some experiments that determined the
quantitative relationship between the pressure and
volume of a gas.
The Relationship between Pressure
and Volumen cont....
 Relationship between the two quantities is described by the
equation PV = constant (k).
 Dividing both sides by P gives an equation that illustrates the
inverse relationship between P and V:
• This relationship between pressure and volume is known
as Boyle’s law which states that at constant
temperature, the volume of a fixed amount of a gas is
inversely proportional to its pressure.
TEMPERATURE IS
CONSTANT
Temperature and Volumen
relationship
 Hot air rises and gases expand when heated.
 Aleksandar Charles carried out experiments to quantify the
relationship between the temperature and volume of a gas, he statet
his law :
At constant pressure, the volume of a given mass of an ideal gas
increases or decreases by the same factor as its temperature on the
absolute temperature scale.
where V is the volume of the gas; and T is the absolute temperature.
The law can also be usefully expressed as follows:
Example
The Relationship between
Amount of gases and Volumen
 Avogadro postulated that, at the same temperature and
pressure, equal volumes of gases contain the same
number of gaseous particles.
 Avogadro’s law describes the relationship between
volume and amount of gas: At constant temperature and
pressure, the volume of a sample of gas is directly
proportional to the number of moles of gas in the
sample.
 Stated mathematically:
Where:
V is the volume of the gas(es).
n is the amount of substance of the gas.
k is a proportionality constant.
In conclusion
• The relationships between the volume of a gas and its pressure,
temperature, and amount are summarized in the figure below.
• The volume increases with increasing temperature or amount but
decreases with increasing pressure.
3. Ideal Gas Law
 A law relating pressure, temperature, volume, and the amount of an
ideal gas.
Any set of relationships between a single quantity (such as V) and
several other variables (P, T, n) can be combined into a single
expression that describes all the relationships simultaneously.
R = 8.3145 J/mol·K (Joule /mol·Kelvin)
Deriving the Ideal Gas Law
 A particular set of conditions were chosen to use as a reference;
0ºC (273.15 K) and 1 atm pressure are referred to as standard
temperature and pressure (STP).
 The volume of 1 mol of an ideal gas at 0ºC and 1 atm pressure is
22.41L, called the standard molar volume of an ideal gas.
4. Mixture of gases Partial Pressures (P)
 If the volume, temperature, and number of moles of each gas in a
mixture is known, then the pressure exerted by each gas
individually, which is its partial pressure, can be calculated.
 Partial pressure is the pressure the gas would exert if it were the
only one present (at the same temperature and volume).
 This law is known as Dalton’s law of partial pressures and can be
written mathematically as
Pt = P1 + P2 + P3 - - - + Pi
where Pt is the total pressure and the other terms are the partial
pressures of the individual gases.
Example
Partial Pressures cont...
 For a mixture of two ideal gases, A and B, the expression for the
total pressure can be written as
• More generally, for a mixture of i components, the total pressure is
given by
The above equation makes it clear that, at constant temperature and
volume, the pressure exerted by a gas depends on only the total
number of moles of gas present, whether the gas is a single
chemical species or a mixture of gaseous species.
Mole Fractions of Gas Mixtures
 The composition of a gas mixture can be described by the mole
fractions of the gases present.
 Mole fraction ( ) of any component of a mixture is the ratio of the
number of moles of that component to the total number of moles of
all the species present in the mixture (nt)
 Mole fraction is a dimensionless quantity between 0 and 1.
 If A = 1, then the sample is pure A, not a mixture.
 If A = 0, then sample A is is not pure present in the mixture.
The sum of the mole fractions of all the components present must
equal 1.
4. Kinetic molecular theory of gases
The kinetic molecular theory of gases explains the laws
that describe the behavior of gases and it was developed
during the nineteenth century by Boltzmann, Clausius,
and Maxwell.
Kinetic molecular theory of gases provides a molecular
explanation for the observations that led to the
development of the ideal gas law !!
Several postulates ( rules ) exists but 2 of
them are essential for ideal gas behavior :
• Postulates 2 and 3 states that all gaseous particles behave
identically, regardless of the chemical nature of their component
molecules
 At at a given temperature, all gases have the same average kinetic
energy.
 The average kinetic energy of the molecules of a gas is
where 2 is the average of the squares of the speeds of the particles
and m is the mass of the object.
• The square root of 2 is the root mean square (rms) speed (rms)
•
To sum up ...
Diffusion of Gas
• Diffusion is the gradual mixing of gases due to the motion of their
component particles even in the absence of mechanical agitation
(movement) such as stirring.
 Result is a gas mixture with uniform composition - Graham’s law.
The ratio of the diffusion rates of two gases is the square root of the
inverse ratio of their molar masses.
If r is the diffusion rate and M is the molar mass, then
r1/r2 =  M2/M1
If M1  M2, then gas #1 will diffuse more rapidly than gas #2.
Effusion of Gas
• Effusion is the escape of a gas through a small (usually
microscopic) opening into an evacuated space.
• Rates of effusion of gases are inversely proportional to
the square root of their molar masses.
• Heavy molecules effuse through a porous material more
slowly than light molecules.
The Behavior of Real Gases
• Postulates of the kinetic molecular theory of gases ignore
both the volume occupied by the molecules of a gas and
all interactions between molecules, whether attractive or
repulsive.
• In reality, all gases have nonzero molecular volumes and
the molecules of real gases interact with one another in
ways that depend on the structure of the molecules and
differ for each gaseous substance.
The Behavior of Real Gases
• Real gases behave differently from ideal gases at high pressures
and low temperatures.
A ) In an ideal gas the interactions between molecules are not accounted
B ) in an ideal gas the actual volume taken up by the molecules of gas is not taken
into account.
So…
Molecules of an ideal gas are assumed to have zero volume; volume
available to them for motion is the same as the volume of the
container.
P,V and T Relationships
in Real Gases
Molecules of a real gas have small but measurable volumes.
 At low pressures, gaseous molecules are far
apart
 As pressure increases, intermolecular
distances become smaller
 Total volume occupied by gas is greater than
the volume predicted by the ideal gas law, so
the experimentally measured value of
PV/nRT is greater than the value predicted by
the ideal gas law.
V.W constants of some gases
Liquefaction of Gases
Liquefaction of gases is the condensation of gases into a
liquid form.
• Both the theory and the ideal gas law predict that
gases compressed to very high pressures and cooled to
very low temperatures should still behave like gases.
However, as gases are compressed and cooled, they
condense to form liquids.
6. LIQUIDS
The Kinetic Molecular Description
 Molecules in liquids are very close together, with
essentially no empty space between them ( In solids
they are more apart )
 Molecules in liquids are in constant motion, and their
kinetic energy (and hence their speed) depends on their
temperature.
 Properties of liquids can be explained using a modified
version of the kinetic molecular theory.
The Kinetic Molecular
Description of Liquids
Some characteristic properties of liquids are:
1. Density
– Molecules of a liquid are packed relatively close together.
– As a result, liquids are much denser than gases
– Densities of liquids measured in units of grams per cubic
centimeterr (g/cm3) or grams per milliliter (g/mL).
The Kinetic Molecular
Description of Liquids cont...
2. Molecular order
– Liquids exhibit short-range order because strong
intermolecular attractive forces cause the molecules to
pack together tightly.
– Because of the high kinetic energy of the molecules, they
move rapidly with respect to one another.
– Arrangement of the molecules in a liquid is not completely
random.
– Molecules in liquids are ordered because of strong
intermolecular attractive forces.
The Kinetic Molecular
Description of Liquids cont...
3. Compressibility
– Liquids cannot be readily compressed because
they have so little empty space between the
component molecules.
4. Thermal expansion
– Intermolecular forces in liquids are strong enough
to keep them from expanding significantly
when heated.
– Volumes of liquids are somewhat fixed.
The Kinetic Molecular
Description of Liquids cont...
5. Fluidity
– Liquids can flow, adjusting to the shape
of the container, because their
molecules are free to move.
6. Diffusion
– Molecules in liquids diffuse because they
are in constant motion.
Fluidity
and
Diffusion
Intermolecular forces in liquids
 intermolecular interactions, weaker than the
intramolecular interactions that hold the atoms together
in molecules and polyatomic ions.
 Transitions between the solid, liquid, and gaseous states
are due to changes in intermolecular interactions but do
not affect intramolecular attractions.
 Intermolecular forces are electrostatic; they arise from
the interaction between positively and negatively
charged species.
Intermolecular Forces cont...
There are three major types of intermolecular interactions:
1. Dipole-dipole interactions
2. London dispersion forces
3. Hydrogen bonds
The first two are described collectively as van der Waals
forces
Intermolecular Forces cont...
Dipole-dipole interactions
– There are two types of dipole-dipole
interactions:
1. Attractive — molecular orientations in
which the positive end of one dipole is near the
negative end of another (and vice versa)
2. Repulsive — molecular orientations that
juxtapose the positive or negative ends
of
the dipoles on adjacent molecules
–The attractive intermolecular interactions are
more stable than the repulsive intermolecular
interactions
Intermolecular Forces cont...
London dispersion forces
The London dispersion force is the weakest
intermolecular force.
The London dispersion force is a temporary attractive
force that results when the electrons in two adjacent
atoms occupy positions that make the atoms form
temporary dipoles.
Intermolecular Forces cont...
Hydrogen bonds
– Molecules that contain hydrogen atoms bonded to
electronegative atoms such as O, N, F, and to a lesser extent
Cl and S, tend to exhibit strong intermolecular interactions and
have high boiling points.
– The large difference in electronegativity results in a large
partial positive charge on hydrogen and a corresponding large
partial negative charge on the O, N, or F atom.
Example
8. Unique Properties of Liquids
There are three unique properties of liquids that
depend intimately on the nature of
intermolecular interactions
1. Surface tension
2. Capillary action
3. Viscosity
Surface Tension
 Is the energy required to increase the surface area of a
liquid by a specific amount!
 Measured as energy per unit area, such as joules per
square meter (J/m2)
 The stronger the intermolecular forces, the higher the
surface tension!
 Surfactants are molecules such as soap and detergents
that reduce the surface tension of polar liquids like water
by disrupting the intermolecular attractions between
adjacent molecules
Surface tension example...
(a) A paper clip can “float” on
water because of surface
tension.
(b) Surface tension
also allows insects such as
this water strider to “walk on
water.”
Capillary Action
Is the ability of a liquid to flow in narrow spaces without the assistance
of, and in opposition to, external forces like gravity.
The height to which the liquid rises depends on the diameter of the tube
(the smaller the diameter, the higher the liquid rises) and the
temperature of the liquid, but not on the angle at which the tube
enters the liquid.
Viscosity
• Viscosity () is the resistance of a liquid to flow.
– The higher the viscosity, the slower the liquid flows through the tube and the
slower the steel balls fall.
• Viscosity is expressed in units of the poise (mPa•s); the higher the
number, the higher the viscosity.
• Liquids that have strong intermolecular forces have high viscosities.
9. Vapor Pressure
 When the liquid is heated, its molecules obtain sufficient
kinetic energy to overcome the forces holding them in
the liquid and they escape into the gaseous phase.
The result of this phenomenon Vapor!
Boiling Points
 As the temperature of a liquid increases, the vapor
pressure of the liquid increases until it equals the
external pressure, or the atmospheric pressure.
 Bubbles of vapor begin to form throughout the liquid, and
the liquid begins to boil.
 Temperature at which a liquid boils at exactly 1 atm
pressure is the normal boiling point of the liquid.
10. Changes of State
 Changes of state are examples of phase changes, or phase
transitions
 Any of the three forms of matter (gas, liquid, solid) is converted
to either of the other two
 Six most common phase changes:
1. fusion (melting)
2. freezing
3. vaporization
4. condensation
5. sublimation
6. deposition
solid
liquid
liquid
gas
solid
gas
→ liquid
→ solid
→ gas
→ liquid
→ gas
→ solid
Energy Changes that
results in phase Changes
 Phase changes are always accompanied by a change in the energy
of the system – we have to types of transition phases ;
1.
2.
Any transition from a more-ordered to a less-ordered state (solid to
liquid, liquid to gas, or solid to gas) requires an input of energy — it is
ENDOTHERMIC.
Any transition from a less-ordered to a more-ordered state (liquid to
solid, gas to liquid, or gas to solid) releases energy — it is
EXOTHERMIC.
 The conversion of a solid to a liquid is called melting or fusion —
energy required to melt 1 mol of the substance is its enthalpy of
fusion, ΔHfus.
 Energy change required to vaporize 1 mol of a substance is the
enthalpy of vaporization, ΔHvap.
Energy Changes that
results in phase Changes
 Direct conversion of a solid to a gas, without an intervening liquid
phase, is called sublimation.
 Amount of energy required to sublime 1 mol of a pure solid is the
enthalpy of sublimation, ΔHsub.
 Enthalpy of sublimation of a substance is the sum of its enthalpies of
fusion and vaporization (an application of Hess’s law):
ΔHsub = ΔHfus + ΔHvap.
Fusion, vaporization, and sublimation are endothermic
processes; they occur only with the absorption of heat.
11. Critical Temperature and Pressure
 Critical temperature (Tc)
The critical temperature of a substance is the temperature at and above
which vapor of the substance cannot be liquefied, no matter how much
pressure is applied.
 Critical pressure (Pc)
The critical pressure of a substance is the pressure required to liquefy a gas at
its critical temperature
 Critical point — combination of critical temperature and
critical pressure
Readings for this lecture
Book chapter 10 and 11 – pages : 655-720,
731-782
Follow this lecture if you decide to read the
book , what is mentioned in the powerpoint
lectures ,this sould be also read in the
book for more details.
Reminder for Quiz 2 and
Homework 3
Submit by 24th April Homework 3
6th May Quiz 2
( Lecture 7, 8, 9 and 10)
 Homework 3 question:
1. why can a gas be compressed into a smaller volume?
2. why can you not squeeze a liquid into a smaller volume?
3. why can you not stir a solid?
4. what makes a gas exert pressure?
5. why does a gas completely fill its container?
6. why does a solid expand when it is heated?