GASES, LIQUIDS AND SOLIDS
Basic concepts:
 Matter (solids, liquids & gases) is made up of particles
 Solids hold their shape
 Liquids take the shape of their container but still hold together
 Gases occupy the maximum volume available to them
 Solids and liquids cannot be compressed but gases are easily
compressed
 Diffusion takes place rapidly in gases an liquids
 Solids are more dense than the liquid state of a substance (except for
water)
 As a substance changes phase (from solid to liquid, liquid to gas, gas
to liquid etc) the temperature remains constant even though energy is
being gained or lost)
 A substances melting point is always lower than its boiling point
Explanations:
 Particles in a solid are arranged in a regular pattern, very close
together. The particles have some kinetic energy, but are restricted to
vibrating and rotating and are not able to leave their position in the
lattice. Their energy is not sufficient to overcome the inter- particle
forces holding them together.
 Compressing and diffusing in solids is not normally possible because
the particles are already as close together as they can get.
 The particles in a liquid are close together because the inter-particle
force is holding them close, but they are free to move around. As the
temperature increases the particles move more rapidly and the liquid
expands.
 Diffusion is possible because the particles are free to move around.
Compression is very difficult because the particles are still very close
together.
 The particles in a gas have sufficient energy to overcome the attractive
forces holding them together. With no requirement to stay in any
position the gas particles are free to move in all direction and quickly
take up the volume available to them. Gas particles are widely
separated and so compression is easy. Diffusion is easy because
particles can move anywhere.
LRY 2013
Boiling, Vapour Pressure and Evaporation
The particles of a liquid (or solid, or gas) do not all have exactly the same
amount of energy at any given temperature. Some particles will have more
energy than others.
In a liquid of given temperature there will be a certain proportion of particles
which have sufficient energy to overcome the attractive forces between the
particles and, if they are on the surface of the liquid they can escape and become
a gas. Above the surface of the liquid substance there will always be vapour. The
pressure of this vapour is called the vapour pressure.
If the liquid is in a sealed beaker there will be an equilibrium set up between the
liquid and the gas phase of the substance. If the beaker is not sealed, particles of
gas can escape and never return. Eventually all the liquid will escape. We call
this evaporation.
Equilibrium
set up
Evaporation
Open beaker
Closed beaker
As the temperature rises, a greater proportion of particles in the liquid have
sufficient energy to evaporate and the vapour pressure rises. Eventually the
vapour pressure equals the atmospheric pressure and at this temperature we say
the liquid is boiling, because the particles within the body of the liquid (not just
those on the surface) can enter the gas phase.
Number of Particles
25oC
100oC
Maxwell-Boltzmann
Distribution
Energy needed
for evaporation
Energy
Boiling occurs when the vapour pressure equals atmospheric pressure.
The temperature at which this occurs is the ‘normal boiling point’ of the
liquid.
LRY 2013
Solids:
There are two types of solids:
 Crystals
 Amorphous solids
Crystals – particles (atoms, molecules or ions) are packed closely together and
are held in fixed positions. All true solids are crystalline. Particles in a crystal
vibrate about their positions but in general, are unable to move away from them.
Some, near the surface, may have sufficient energy to escape (vaporise or
sublime) and exert a vapour pressure. Vapour pressures of most solids at
ordinary temperature are very small.
Amorphous solids – do not form crystals. Particles not arranged in an orderly
fashion - usually formed when a substance has cooled fast and the particles
have not had time to arrange themselves in an orderly way. (E.g. plastic sulfur)
Forces between particles –
The attractive forces between particles are all electrical in origin. I.e. they all
depend upon the pull that something that is positively charged exerts on
something that is negatively charged.
 Forces between molecules (inter-molecular forces) known collectively
as Van der Waals forces, these are comparatively weak and include:
 The force between non-polar molecules (dispersion forces or
instantaneous dipole-dipole interactions)
 The force between polar molecules (permanent dipole-dipole
interactions)
 Hydrogen bonding (special case of permanent dipole forces)
 Force between atoms in molecules and polyatomic ions (intramolecular forces). Covalent Bonding.
 The force between ions – ionic bonding
 The force between metal atoms – metallic bonding. In metal crystals
the attraction is between the atoms themselves and the outer electrons
loosely held by the individual atoms, these electrons are ‘pooled’ and
held by the crystal as a whole.
LRY 2013
Intermolecular Forces
 Forces between polar and non-polar molecules.
Consider F2, Ar and HCl – all are gases, their molecules all contain 18
electrons and on cooling each condenses to a liquid and then freezes
to a solid.
Substance
Fluorine
Argon
Hydrogen chloride
Fluorine
F
F
Boiling point (K)
85
87
188
Argon
Ar
Non-polar molecule
Non-polar monatomic
molecule
Melting point (K)
53
84
159
Hydrogen chloride
H
Cl
Polar molecule – charge
separation due to uneven
distribution of electrons
The polar nature of the HCl molecule explains the force that acts between the
molecules. The partially positive end on one HCl molecule attracts the partially
negative end of another etc. Because the dipole is a property of the HCl molecule
and is there all the time it is called a permanent dipole.
In the case of fluorine and Argon, both non-polar molecules, it is harder to see
where the force comes from. Since the electrons are constantly moving their
distribution does not remain symmetrical. At one moment in time there could be
more negative charge on one side of the nucleus that the other. This produces
and instantaneous dipole (there one minute and gone the next). While it is
there the dipole is able to distort the electrons on neighbouring molecules and
create or induce dipoles on neighbouring molecules. These are called induced
dipoles. This leads to a force of attraction between the molecules.
The dipoles between non-polar molecules are much weaker than those between
polar molecules with the same number of electrons.
 Hydrogen bonding
Water is the one compound whose properties cannot be explained by
the above ideas on bonding. According to its number of electrons,
shape and polarity, water should be a gas. CO2 has 22 electrons and is
a gas. HCl is polar, has 18 electrons and is a gas. SO2 is a gas with 32
electrons. Water has only 10 electrons, yet it is a liquid up to 100 oC.
LRY 2013
 Other compounds that behave unusually, like water, are NH3 and HF.
Fluorine is the most electronegative element there is and it is closely
followed by oxygen and nitrogen. When hydrogen is bonded to these
atoms, the electrons are shared unequally, with hydrogen getting the
least share. The F, O and N end of the molecules all have a ‘large’ charge while the hydrogen end has a ‘large’ + charge. The positively
charged hydrogen is attracted to the lone-pair of electrons on the
oxygen, fluorine or nitrogen of the neighbouring molecule. This linkage
of the two molecules through the hydrogen is called a hydrogen bond.
The hydrogen bond has about 10% the strength of a covalent bond.
Hydrogen bonding accounts for the high melting and boiling points of
water, ammonia and hydrogen fluoride.
Boiling Point vs No of electrons
100
Boiling Point (oC)
50
0
0
10
20
30
40
50
-50
-100
-150
-200
No. of electrons per molecule
Hydrogen bonds also explain why water is less
dense than ice. In ice the water molecules are held
in an hexagonal pattern which is more open and
therefore less dense than the random structure of
water. Hydrogen bonds are also important in natural
fibres such as paper and wool. They hold protein
molecules together and link the stands of DNA
together.
LRY 2013
60
group 14
group 15
group 16
group 17
Heating and cooling curves
When a pure substance is heated it goes through a series of periods of constant
temperature and temperature increases.
Temperature
F
D
E
B
C
A
Time
 From point A to B temperature increases, particles are increasing in
kinetic energy, solid particles vibrate more and take up more room,
solid expands.
 At point B solid starts to melt. Added energy is used to overcome
attractive forces between molecules. Potential energy increase.
Temperature remains constant, so kinetic energy of particles remains
constant. At point C solid has melted.
 From C to D temperature increases so kinetic energy of particles is
increasing.
 At D liquid has reached boiling point. The added energy is used to
overcome attractive forces between particles and allow particles to
escape into gas phase. By E all liquid has turned to vapour.
 After point E the added energy goes into increasing the kinetic energy
of particles in the gas phase. (provided container is closed)
Note: Temperature is a measure of the average kinetic energy of particles. Two
substances at the same temperature, means that their particles have the same
average kinetic energy.
Molar enthalpy of fusion and vaporisation
The energy required to change one mole of substance from a solid to a liquid at
the melting point is called the molar enthalpy of fusion (fusH). It is measured in
kJmol-1.
H2O(solid at 0oC)
H2O (liquid at 0oC) fusH = 6.0 kJmol-1
The molar enthalpy of fusion provides a measure of the strength of the force
holding the particles together in the solid phase
LRY 2013
The energy required to change one mole of substance from a liquid to a vapour
at the boiling point is called the molar enthalpy of vaporisation ((vapH). it is
measured in kJmol-1.
H2O(liquid at 100oC)
H2O (vapour at 100oC) vapH = 41.0 kJmol-1
The molar enthalpy of vaporisation provides a measure of the strength of the
force between particles in the liquid phase.
 vapH is always greater than fusH (it takes more energy to completely
overcome the attractive forces between the particles to allow them to
go anywhere as in a gas).
 The greater the fusH, the higher the melting point – particles held
more firmly.
 The greater the vapH the higher the boiling point.
LRY 2013