CHEM30S
Mr. Gluchowski
Name:
Physical Properties of Matter Unit Review
Section 1: The Kinetic Molecular Theory
1. Outline the 4 postulates of the Kinetic Molecular Theory.
All matter is made up of small particles
Particles (molecules) are in constant motion – Gases have straight line motion, solids and
liquids experience vibrational motion
All collisions between particles and the containers are completely elastic
There are no attractive forces experienced between gas molecules. Liquid and solids
experience intermolecular forces (IMF)
2. What do we mean when we say elastic collisions? How does this apply to the motion of gases
and pressure?
Elastic collisions mean that no energy is lost after a collision. The energy going in, is equal to
the energy coming out (E1=E2). If there was a difference in energy coming out (greater than
initial, or less than initial) this would cause the pressure to either increase or decrease,
respectively.
3. What type of relationship do temperature and average kinetic energy have?
Temperature is directly proportional to KE. This means that as temperature goes up, KE goes
up, and the opposite is also true.
4. What is the relationship between the motion of molecules, the temperature of the sample of
matter, and the average kinetic energy of the sample of molecules?
As molecules move faster they gain more energy, and this is related to the temperature, since
it is a measure of movement. HOWEVER, it is important to remember that KE won’t be the
same between two different substances at the same temperature. Do not forget about
differences in intermolecular forces between different substances.
Section 2: Properties of the Three States of Matter
Define: Diffusion: the ability for particles to move away from each other, from a high to low
concentration, Density: How close molecules/particles are to each other, Compressibility: how easily
particles can be pushed together.
1. Describe the following properties of a solid and then explain why the solid has these
properties.
a. Density: high density, particles are close together, no room for movement.
b. Compressibility: low due to how close particles in solids are.
c. Diffusion: stronger IMFs hold on to particles, do not allow much movement, cannot
easily move away from each other.
o Describe ionic, polar covalent and non-polar covalent bonds. Give an example of each.
Ionic bonds are between metals and non-metals where one element is more
electronegative (electron hungry) than the other element and causes the electrons to be
pulled to one side, ionic bonds are polar, ex: NaCl (table salt). Polar covalent bonds are
where electrons are unequally shared, because one element wants the electron more,
example of this would be water. Non-polar covalent bonds are where electrons are
shared equally over all bonds, an example of this would be a diamond, or sand.
Duke of Marlborough 2018
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CHEM30S
Mr. Gluchowski
Name:
o Why can metallic solid conduct electricity?
Metallic solids have delocalized bonding (meaning that they are all over the place)
electrons are free to move through the valence shells of all the atoms in the solid. Think
of it as a “sea of electrons”
2. Describe the following properties of a liquid and then explain why the liquid has these
properties.
a. Density moderate density, there is a bit of room between molecules in a liquid, but IMF’s
are still keeping the molecules close together.
b. Compressibility considered to be very low, this characteristic is affected by the bonds
between the molecules in a liquid.
c. Diffusion particles have more movement and can easily move away from each other.
Bonds start to play into this as well.
3. Describe the following properties of a gas and then explain why the gas has these properties.
a. Density low density, due to a lot of space between the particles
b. Compressibility high compressibility due to space between particles
c. Diffusion high diffusion, the particles have more kinetic energy so they can easily “float”
away from the sample
Section 3: Phase Changes
Use arrows and the following words to make an appropriate concept diagram:
Solid
Deposition
liquid
Vaporization(evaporation)
Gas
Freezing (solidification)
Sublimation
Condensation (liquidification)
a) Indicate where heat is added/removed
b) Indicate where kinetic energy increases/decreases, intermolecular forces increase/decrease and
molecular motion increases/decreases.
Changing Phases – Solid to Liquid
1. Explain what happens to the molecules, temperature, IMFs and kinetic energy during the
process of melting. When melting, molecules start to speed up as the temperature goes up, KE
Duke of Marlborough 2018
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CHEM30S
Mr. Gluchowski
Name:
of the molecules also increases. As more energy is put in IMFs start to weaken, and particles
have more movment.
2. Explain what happens to the molecules, temperature, IMFs and kinetic energy during the
process of freezing. When freezing, molecules are losing energy and bonds start to reform. As
energy is taken away the movement of the particles slows down, along with kinetic energy.
3. What relationship do the freezing point and melting point of a substance have? They appear at
the same time on a graph, they are just the reverse of each other.
Changing Phases – Liquid to Gas
1. Explain what happens to the molecules, temperature, IMFs, and kinetic energy during the
process of condensation. During condensation gas particles lose energy when they interact
with lower energy surfaces or lower energy particles. As energy (KE) goes down IMFs start to
get stronger,
2. Explain what happens to the molecules, temperature, IMFs, and kinetic energy during the
process of evaporation. As more energy is added to the system, molecules at the surface start
to escape from their IMFs which at this point are now very week.
3. Define dynamic equilibrium. Provide an example of dynamic equilibrium using evaporation and
condensation dynamic equilibrium where the forward reaction equals the reverse reaction. An
example of this is vapour pressure in a closed system.
4. What part do collisions play in changing phases? Collisions can sometimes stop other
molecules from moving on to the next level.
5. What makes a liquid more volatile than another? Volatility is how easily something can
evaporate, or move away from the surface. Lower IMFs result in particles with a lot of KE
Section 4: Vapour Pressure
1. Define pressure in terms of molecules (what are they doing). Pressure is equal to the number
of collisions between particles and the container.
2. What is standard atmospheric pressure? This is the pressure that we all feel, as the
atompshere is always pushing down on us
3. What is vapour? Escaped partilces of a liquid
4. Define vapour pressure the pressure exerted by a liquid (often in a closed system)
Section 5: Boiling
1. Define boiling a liquid or substance boils when the vapour pressure of the boiling bubbles is
equal. As the pressure in the bubbles increases more and more can escape leaving their
vapour behin.
2. Describe what is happening on the molecular level when a substance boils. As a substance
boils, the liquid molecules start to move faster, and overcome their IMFs and they are able to
be es
3. Define the term “normal boiling point” the boiling point of a substance at normal pressure
4. What is the difference between boiling and evaporation? Evaporation happens when no extra
energy added (remember that graph, some particles have enough energy to escape), while
boiling requires energy to be added to the system
5. Using the data below, graph the curves and then find the vapour pressures and boiling points
of the substances below.
Duke of Marlborough 2018
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CHEM30S
Mr. Gluchowski
Name:
6. How do intermolecular forces influence the temperature at which something melts? The
stronger the forces between and within molecules will cause the melting point to increase.
Circle the Correct Answer
1. What does raising the temperature do to molecules? (increase or decrease their kinetic
energy)
2. What happens to the intermolecular forces between molecules if their kinetic energy
decreases? (get weaker or get stronger)
3. Molecules with weaker intermolecular forces will move (faster or slower).
4. The stronger the intermolecular forces that exist between solid molecules, the (more or less)
energy is required to cause it to melt.
5. As the boiling point of a (solid or liquid) increases, (more or less) kinetic energy is required to
(strengthen or weaken) the intermolecular forces.
Table 1: Vapour Pressures of Liquids (remember 760mmHg or 1 atm is normal)
Temperature(⁰C)
Water (mmHg)
0
10
20
30
40
50
60
70
80
90
100
5
9
18
32
55
93
149
234
355
526
760
a.
b.
c.
d.
e.
Ethyl Alcohol
(mmHg)
12
24
43
79
135
222
253
543
813
1187
1693
Dimethylether
(mmHg)
175
300
450
700
1000
Ethylene glycol
(mmHg)
2
4
12
24
60
What is the normal boiling point of Ethyl Alcohol? Between 70-80C
At what temperature is the vapour pressure of Dimethyl ether 650 mmHg? 20 – 30C
At What temperature does water vapour exert a pressure of 900mmHg? ++100C
What pressure does Ethylene glycol have to be at for it to boil at 70⁰C? 4C
List the substances in increasing order (weakest to strongest) of intermolecular forces.
Ethylene glycol, Ethyl alcohol, Dimethyleher, water.
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