Chapter
12 – Liquids and Solids
1. The three states of matter:
(there are actually 2 others: Plasma and the Bose-Einstein Condensate but
we will consider only the basic three which we commonly find on Earth)
Solid: A substance with a definite volume and shape
– most highly organized
Liquid: A substance which flows and takes the shape
of its container – less organized than a solid.
Gas: A substance which takes the shape of and
fills its container – less organized than a liquid.
Kinetic
energy: Energy of an object because
of its motion
2. Changing States of
matter: See table 12-2 p. 372 and
diagram 12-15 on p. 382
Intermolecular
forces: The force of attraction
between two or more molecules. Gases
have the weakest/lowest intermolecular forces, solids have the
strongest/largest.
Condensation: This is the process of going from a gas to a
liquid. If the temperature is lowered
enough, the average kinetic energy of a particle is too low to remain as a gas,
therefore, it is attracted to other gas particles and condenses into a
liquid. If intermolecular forces are
large, then the condensation will happen at a higher temperature. Example:
Water has higher intermolecular forces than alcohol. Water condenses from a vapor to a liquid at
100 degrees while alcohol might do it at 80 degrees. This is because it takes more energy to break water apart from
liquid to gas.
Boiling
or Vaporization or Evaporation: (p.
378) This is the process of going
from a liquid to a gas. It is the exact
opposite of condensation. The higher
the intermolecular forces, the higher the evaporation temperature. Boiling temp. and Condensation temp. are
exactly the same.
You can alter the boiling/vaporization/evaporation point of a liquid by altering the pressure on the liquid. Polar compounds will have a higher boiling point than non-polar compounds.
Freezing
(p. 379-380): This is the process
of a liquid becoming a solid. The
higher the intermolecular forces, the higher the freezing point. For Example: The metal pipes in the room have very strong intermolecular
forces. Their freezing point is very
high – several hundred degrees. You can
alter the freezing point by altering the outside pressure. Polar compounds will have a higher freezing
point than non-polar compounds.
Look at table 12-1 on p. 370 and you will
see that covalent molecules (which have lower intermolecular forces than ionic
compounds) have lower melting and boiling points.
Melting
or Fusion (p. 379-380): This is the
process of a solid becoming a liquid.
It happens at the exact same temperature as the freezing point.
Sublimation
(p. 380): This is the process of a
solid becoming a gas and NEVER entering the liquid state. The reverse of this is called
Deposition.
Note: Refer to the phase diagrams on
p. 381 (water) and 390 (CO2) to help you see where
these phases of matter changes occur as a function of pressure and temperature.
3. Let’s take a trip down memory lane…….:
Remember from chapter 6:
Intramolecular forces are the forces which
hold individual atoms together to make a molecule. These forces are determined by the atoms’ individual
electronegativity rating (see p. 143). Intermolecular forces, on the other
hand, are the forces between molecules holding them together. For example, water molecules clinging
together in a glass of water. You will
need to be able to distinguish between inter and intra molecular forces.
4. Equilibrium (p. 372-375): Equilibrium is a dynamic (ever changing) condition
in which two opposing changes occur at equal rates in a closed system. Example: Stand on one leg and close both of your eyes
– you will lose your equilibrium of balance.
Example: Look at the
Sparkletts water bottle in class.
Notice that it is in equilibrium between liquid and gas phase.
Le Chatelier’s Principle: When a system at equilibrium is disturbed by
application of a stress, it attains a new equilibrium position which minimizes
the stress. (p. 374) See table 12-3 on
p. 375 and understand what changes in conditions will cause a shift.
5. Vapor Pressure: (page 376 figure 12-12) This is the pressure of the vapor above a
liquid at a given temperature.
Example: As water is heated to
100oC, its vapor pressure approaches that of the air (1 atm). As Ethanol is heated to 80oC, its
vapor pressure approaches that of the air.
This figure shows that water must have the highest INTERMOLECULAR forces
among these three molecules. It takes a
lot of heat to make water reach 1 atm at its surface. When vapor pressure reaches atmospheric pressure (1 atm or
101.325 kPa), the liquid boils.
Strong Intermolecular forces make for LOW
vapor pressures. Water, for example has
strong intermolecular forces. Its vapor
pressure is very low. This means that
not much of the water is in vapor form above the liquid. Water wants to remain a liquid and therefore
there is not much vapor above a swimming pool (for example). You must heat the water up a lot in order to
increase the vapor pressure to 1 atm.
Bottom
line: If a liquid has a low vapor
pressure, it is going to have a high boiling point. If a liquid has a high vapor pressure, it is going to have a low
boiling point – because its liquid is readily available to converting to a
gas. (Make sure you read p. 374-377 to
re-enforce this idea)
Volatile liquids: These have
low intermolecular forces and very high vapor pressures. They are ready to convert to a gas at a
moments notice. Gasoline would be an
example of a volatile liquid with a high vapor pressure and a low boiling
point. Water is not volatile.
How to manipulate boiling point: Boiling point is the temperature at which the vapor pressure
equals the external pressure acting on a liquid. If you reduce the external pressure on the liquid, then the vapor
pressure will reach that external pressure at a lower temperature. The same is true for water.
How to manipulate freezing point: Reduce the outside pressure and the freezing point will
decrease. Without outside pressure
pushing the molecules together, the molecules will have to be COLDER to fuse
together to make a solid. EXCEPT FOR
WATER! With water, as pressure increases,
the freezing point will decrease (look at phase diagram on p. 381) This is why, in the mountains, where there
is less pressure, it can snow more easily.
It doesn’t have to be as cold in the mountains (as it does as in San
Clemente) to make snow fall.
Compare the phase diagram of water with that of Carbon dioxide
on p. 390
6. Water’s properties
Density of water based on temperature: See page 385 figure 12-17. As water freezes to form ice, the hydrogen bonding which occurs
allows for spaces to open up between the molecules of water and therefore water
is less dense as a solid than as a liquid.
As water freezes, it expands. This is why bottles blow up in the freezer
and pipes may burst in the winter in cold climates.
Water is at its most dense right before it freezes. At 4 degrees Celsius, water is at its most dense and it sinks to the bottom of a lake (for example). But at 0 degrees, it is at its least dense and therefore it rises. (Realize that 4 degree water is a liquid but 0 degree water is a solid)
Water’s heat capacity:
Because of the strong hydrogen bonding, water has a high heat
capacity. It takes a lot of heat to
warm up water. And water can store a
lot of heat too. Substances with low
heat capacity (such as metals) heat up quickly and then cool down quickly too.
Study Guide
You must be able to define
the following terms and know how they are related.
1. The four states of matter: (look in the
glossary in the back of the book)
Solid
Liquid
Gas
2.
Changing States of matter:
Intermolecular forces
Condensation
Vaporization or Evaporation
Boiling
Freezing
Melting or Fusion
Sublimation
Phase diagram of water
Phase diagram of CO2.
3. Intermolecular Forces holding molecules
together: (Review Chapter 6)
Electronegativity
The van der Waals forces:
Dispersion
forces
Dipole-Dipole
forces
Hydrogen
Bonding
4. Equilibrium
Le Chatelier’s Principle
5. Vapor Pressure
Partial
pressure
Volatile liquids
How to manipulate boiling point
How to manipulate freezing point
6. Water’s properties
Density of water based on temperature
Water’s heat capacity
7. Common Temperatures of substances at 1 atm of pressure
oC oF
Human Body 37 98.6
Freezing Water 0 32
Boiling Water 100 212
Dry Ice (CO2) -78 -94
Liquid Nitrogen (N2) -196 -321
Room Temperature 25
77