Unit 6: States of Matter
Honors Chemistry
I.
Gas State
Kinetic Molecular Theory for gases (ideal gases)- a model to
describe the observed behavior of gases.
1. Gas particle is negligibly small. The space between gas
particles is very large compared to the size of the atom
or molecule.
2. Average kinetic energy of a gas particle is proportional to
the kelvin temperature.
𝟏
𝟑

𝑲𝑬𝒂𝒗𝒈 = 𝒎𝒗𝟐 𝒐𝒓 𝑲𝑬𝒂𝒗𝒈 = 𝑹𝑻
𝟐
𝟐
J
𝑅 = 𝑔𝑎𝑠 𝑙𝑎𝑤 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 (8.31
)
mol ∙ K
Graham’s law of effusion
a. rate r of effusion (leaking) or diffusion (spreading out) is
inversely proportional to molar mass (MM)
b. The ratio of effusion of two different gases:
A.
𝒓𝒂𝒕𝒆𝒂
𝒓𝒂𝒕𝒆𝒃
3.
𝓜𝒃
𝓜𝒂
𝓜 = 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑔/𝑚𝑜𝑙
Gas particles collide with elastic collisions. No loss in
energy.
Attractive and repulsive forces between particles are
negligible.
4.
B.
=√
Properties of Gasestend to be small, non-polar molecules
form homogenous mixtures
takes the shape of the container and fills it.
particles are far apart and move randomly.
gas volumes changes significantly with pressure and
temperature changes.
6. relatively low densities.
1.
2.
3.
4.
5.
C.
1.
Units of Measure
Pressure- measure of the force of colliding molecules per
𝑓𝑜𝑟𝑐𝑒
unit area. 𝑝𝑟𝑒𝑠𝑠𝑢𝑟𝑒 =
𝑎𝑟𝑒𝑎


SI unit is the pascal (Pa), 1 Pa= 1N/m2
Other units: atmospheres, kilopascals, bars,
millimeters of Hg, inches of Hg, pounds/sq. in.

1 atm = 101.325 kPa = 760 mmHg = 1.01325 bar=
14.7 psi= 29.9 in. Hg

Pressure can be measured with a barometer or
manometer.
2. Volume- amount of space occupied by gas.

cm3, mL or L

1mL=1cm3, 1000 mL=1 L
3.
Temperature and Kinetic Energy- There is a direct
relationship between the kinetic energy and the
temperature of a molecule, see A2 above.

Kinetic Energy is a measure of the energy of motion;
𝟏
𝑲𝑬𝒂𝒗𝒈 = 𝒎𝒗𝟐 , if either the mass or velocity of a
𝟐
molecule is changed, then the KE is affected.

If KE is doubled, K temp. also doubles

Temperature is measured using the Kelvin
temperature scale. K = C + 273
1
4.

D.
1.
Amount of gas- moles are used, symbol is “n”
One mole of any gas at Standard Temperature and
Pressure, STP (273 K and 1 atm) will occupy 22.4 L.
This is called the molar volume of a gas.
Gas Laws
1
Boyles Law- 𝑉 ∝ T and n constant
𝑃
 𝑃1 𝑉1 = 𝑃2 𝑉2
Charles Law- 𝑉 ∝ 𝑇
2.
𝑉1

𝑃1

𝑇1
𝑇2
=
𝑃2
𝑇2
Avogadros’s Law- 𝑉 ∝ 𝑛 T and P constant

5.
=
Gay-Lussac’s law- 𝑃 ∝ 𝑇 V and n constant
3.
4.
𝑇1
P and n constant
𝑉2
𝑉1
𝑛1
=
𝑉2
𝑛2
Ideal Gas Law- One equation that combines the other gas
laws to describe the behavior of an ideal gas.
𝑷𝑽 = 𝒏𝑹𝑻

R= ideal gas constant (if P is atm, V is liters, T is
kelvin) 0.0821
6.
𝑎𝑡𝑚∙𝐿
𝑚𝑜𝑙𝑒∙𝐾
Rearrangement of the Ideal Gas Lawa.
𝑷𝟏 𝑽 𝟏
combined gas law
𝒏𝟏 𝑻𝟏
=
𝑷𝟐 𝑽 𝟐
𝒏𝟐 𝑻𝟐
b.
7.
density of a gas is proportional to it molar mass and
inversely proportional to temperature.
𝒎
𝑷∙𝓜
=𝒅=
𝑽
𝑹∙𝑻
Dalton’s Law of partial pressures- total pressure of a
mixture of gases is the sum of the partial pressures of
the individual gases.

𝑃𝑡𝑜𝑡𝑎𝑙 = 𝑃1 + 𝑃2 + 𝑃3
= 𝑛𝑎
𝑅𝑇
𝑉
+ 𝑛𝑏
𝑅𝑇
𝑉
+ 𝑛𝑐
= (𝑛𝑎 + 𝑛𝑏 + 𝑛𝑐 )

II.
A.
B.
C.
= (𝑛𝑡𝑜𝑡𝑎𝑙 )
𝑃𝑎 = 𝑋𝑎 ∙ 𝑃𝑡𝑜𝑡𝑎𝑙
(𝑚𝑜𝑙𝑒 𝑓𝑟𝑎𝑐𝑡𝑖𝑜𝑛)
𝑅𝑇
𝑅𝑇
𝑅𝑇
𝑉
+. . .
𝑉
𝑉
𝑋𝑎 =
𝑛𝑎
𝑛𝑡𝑜𝑡𝑎𝑙
Liquid State- (11.1-11.6)
General Properties1. takes the shape of it’s container
2. does not expand to fill container
3. incompressible
4. flows
Comparison of Properties- all three states of matter are
dependent upon a balance between two factors:
1. The Kinetic Energies of the particles, (Temperature).

increase KE particles separate
2. The Energies of attraction between the particles,
(Intermolecular Forces).

Increase IMF and particles are more strongly
attracted
Intermolecular Forces - properties of liquids; vapor pressure,
boiling point are dependent upon the force of attraction
between the particles.
1. London Dispersion Forces (just dispersion)instantaneous dipole created by the random motion of
electrons throughout the molecule.
a. found in all molecules
b. strength dependent upon: polarizability of
molecule.

molar mass of molecule
2
2.
3.
4.
D.
Select Properties- based on IMF’s
1. Surface Tension- tendency of liquids to minimize their
surface area.

interior molecules attract with six “neighbors”,
surface molecules attract to only four.

surface tension decreases as IMF’s decrease.
2. Viscosity- resistance of a liquid to flow.

Greater the IMF’s, higher the viscosity
3. Capillary Action- ability of liquids to flow against gravity
up a narrow tube. Balance between cohesive and
adhesive forces.

cohesive – attraction of liquid molecules to each
other.

adhesive- attraction for liquid molecules and surface
of tube.
4. Vapor Pressure- In a closed container at constant
temperature vaporization and condensation are at
equilibrium: rates are equal 𝑔𝑎𝑠 ⟷ 𝑙𝑖𝑞𝑢𝑖𝑑
a. Pressure exerted by the gas at equilibrium is called
vapor pressure( Pvap.)
b. Effect of temperature- at increased temperatures
higher proportion of molecules have sufficient KE
to escape liquid state; therefore:
ℎ𝑖𝑔ℎ𝑒𝑟 𝑇 ⇒ ℎ𝑖𝑔ℎ𝑒𝑟 𝑃
c. At a given T all molecules have same KEave,
therefore molecules with weaker intermolecular
forces are held less tightly at surface; therefore:
𝑤𝑒𝑎𝑘𝑒𝑟 𝐼𝑀 𝑓𝑜𝑟𝑐𝑒𝑠 ⇒ ℎ𝑖𝑔ℎ𝑒𝑟 𝑃𝑣𝑎𝑝
d. high Pvap indicates volatility—tendency to
evaporate.
5.
III.
B.

shape
Dipole-Dipole Forces- attraction between oppositely
charged ends of polar molecules.
Ion-Dipole Forces- attraction between an ion and a polar
molecule.
Hydrogen Bonding- especially strong type of dipoledipole. Only for hydrogen atoms bonded to O, F or N of
one molecule and O, F and N of another molecule.
Boiling Point- Temperature at which the vapor pressure
of a liquid equals the external pressure, acting on the
surface of the liquid.
a. At this temperature, the KE of the molecules in the
interior of the liquid are sufficient to break their
IMF’s and enter the gas phase.
b. Stronger the IMF ⟹lower the vapor pressure⟹
higher the boiling point.
c. Normal Boiling Point- boiling point at 1 atm.
Phase Changes
Heating curve- graph of the temperature of the system versus
the amount of heat added or removed.
a. Kinetic energy- absorbed (KE increasing/endothermic) or
released (KE decreasing/exothermic) :
𝑞 = 𝑚 ∙ 𝐶𝑝 ∙ ∆𝑇
𝑚 = 𝑚𝑎𝑠𝑠 𝑖𝑛 𝑔𝑟𝑎𝑚𝑠
∆𝑇 = 𝑇𝑓 − 𝑇𝑖
𝐶𝑝 = 𝑠𝑝𝑒𝑐𝑖𝑓𝑖𝑐 ℎ𝑒𝑎𝑡 𝑐𝑎𝑝𝑎𝑐𝑖𝑡𝑦 𝑜𝑓 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒
𝐽

𝐻2 𝑂 𝑠𝑜𝑙𝑖𝑑 = 2.09

𝐻2 𝑂 𝑙𝑖𝑞𝑢𝑖𝑑 = 4.18

𝐻2 𝑂 𝑔𝑎𝑠 = 2.01
𝑔∙°𝐶
𝐽
𝑔∙°𝐶
𝐽
𝑔∙°𝐶
3
b.
C.
Phase change- bond breaking=endothermic, bond
formation= exothermic

𝑣𝑎𝑝𝑜𝑟𝑖𝑧𝑎𝑡𝑖𝑜𝑛, 𝑙𝑖𝑞𝑢𝑖𝑑 𝑡𝑜 𝑔𝑎𝑠
∆𝐻𝑣𝑎𝑝 = 40.7𝑘𝐽/𝑚𝑜𝑙

𝑐𝑜𝑛𝑑𝑒𝑛𝑠𝑎𝑡𝑖𝑜𝑛, 𝑔𝑎𝑠 𝑡𝑜 𝑙𝑖𝑞𝑢𝑖𝑑
∆𝐻𝑐𝑜𝑛𝑑 = −40.7 𝑘𝐽/𝑚𝑜𝑙

𝑓𝑢𝑠𝑖𝑜𝑛, 𝑠𝑜𝑙𝑖𝑑 𝑡𝑜 𝑙𝑖𝑞𝑢𝑖𝑑
∆𝐻𝑓𝑢𝑠𝑖𝑜𝑛 = 6.01 𝑘𝐽/𝑚𝑜𝑙

𝑠𝑜𝑙𝑖𝑑𝑖𝑓𝑖𝑐𝑎𝑡𝑖𝑜𝑛, 𝑙𝑖𝑞𝑢𝑖𝑑 𝑡𝑜 𝑠𝑜𝑙𝑖𝑑
∆𝐻𝑠𝑜𝑙𝑖𝑑 = −6.01 𝑘𝐽/𝑚𝑜𝑙

sublimation, solid to gas
Phase Diagram-pressure and temperature relationship.
1. point A: triple point (three phases at equilibrium)

above triple point: melting and vaporization

below triple point: sublimation (deposition)
2. line A-C: equilibrium vapor pressure curve for liquid (nbp:
normal boiling point occurs at 1 atm pressure)
3. C: critical point, where there is no distinction between
liquid and vapor (no liquid-vapor surface)
4. line A-D: equilibrium vapor-pressure curve for solid
(sp: sublimation point depends on pressure)
5. line A-B: melting point of solid at various pressures
(mp: normal melting point occurs at 1 atm)
a. negative slope when liquid is the densest phase
(melting point decrease with pressure)
b. positive slope (more common) when solid is the
densest phase (melting point increases with
pressure)
4
5
6