chem test Solids
The test will cover properties of solids, focusing on their structure, types, and changes during different states. Key areas to review include:
Thermal energy: the sum of the kinetic
and potential energy of particles in an
object
Kinetic energy: (KE) energy in motion
Potential energy: (PE) stored energy
Kinetic-molecular theory helps us to
understand states of matter and thus their
behaviors and properties.
Kinetic-Molecular Theory
See Notes on Gas, Liquid, Solid slides to complete notes on this page.
All matter is made of small
particles. (atoms, ions, molecules,
etc.) (Notes will be finished on the Gas, Liquid and Solid slides.)
Those particles are in constant
random motion.
Thus, have KE.
This motion causes the particles
to collide with each other and the
container that they are in.
KE is transferred between
particles when they collide.
Elastic collision = when no KE
is lost overall in the collision.
Temperature
The average KE of the particles is
dependent on temperature.
Temperature: measure of the
average kinetic energy of particles
in an object.
KE = ½ mv2
The higher the temperature, the higher the
kinetic energy, the greater the motion of the
particles.
Standard unit for temperature = Kelvin = K
K = °C + 273.15
3 Main States of Matter
https://phet.colorado.edu/en/simulations/states-of-matter
Indefinite volume and shape
If confined, gas particles will spread out to
fill the container.
Particles have high kinetic energy and
thus collide with each other.
Gases
Energy is high enough to
overcome any forces that
hold the particles together
Thus no attractive or
repulsive forces.
Particles can move freely.
Particle diagram:
A Distinction
Ideal Gas: a hypothetical gas that
perfectly follows all of the
kinetic-molecular theory.
No attractive forces between particles.
Nonpolar monoatomic gases (like the
noble gases) are the closest to ideal
gases. (Next closest nonpolar diatomic
gases like H2 and N2)
Real Gas: a gas that doesn’t behave
entirely to the kinetic-molecular
theory.
Does exhibit attractive forces.
Polar gases are the least ideal and the
most real (Ex. H2O, NH3)
Kinetic-Molecular Theory
All matter is made of small particles.
Those particles are in constant random motion.
This motion causes the particles to collide with
each other and the container that they are in.
As it applies to gases
Gases consist of LOTS of TINY particles that are
REALLY far apart (more so than any other state of
matter).
Gases have the most kinetic energy and thus the most
rapid and random motion.
The higher the temperature, the faster they move.
There are no attractive or repulsive forces between
gases, thus particles move independently and have
elastic collisions.
Properties of (Ideal) Gases
Expansion: no definite volume or shape because
they fill whatever container they are in
Fluidity: no attractive forces between them so they
flow past each other, just like liquids
Low density: particles are so far apart that they are
way less dense as gases than when solids or liquids
Compressibility: can greatly
decrease the volume of a gas by
increasing pressure and pushing
the particles together
Properties of (Ideal) Gases
Gases naturally spread out to fill a space.
Diffusion: spontaneous mixing of particles caused
by random motion
Effusion: process where gas particles pass
through tiny openings
Rate is dependent on the velocity of particles.
Definite volume but indefinite
shape
Shape depends on the container it is in
Particles have less kinetic energy
than gases, but more than solids.
Liquids
Stronger intermolecular
forces (IMF) in liquids hold
the particles together more
than in gases
Attractive forces at work
Particles can flow or slide
past each other but have less
mobility than gases.
Particle diagram:
Summary Chart – a review of Intermolecular Forces
Type of IMF
Type of molecules
involved
Strength of
attraction
Dipole-Dipole
Hydrogen Bond
Permanent dipoles
in polar molecules
medium, ~5-20 kJ/mol
H atom in polar
molecule with a high EN
atom & unshared e- pair
medium-high,
~5-50 kJ/mol
London Dispersion
Temporary dipoles in
nonpolar molecules
and any polar
low, ~0.1-5 kJ/mol
Visual
representation
Kinetic-Molecular Theory
All matter is made of small particles.
Those particles are in constant random motion.
This motion causes the particles to collide with
each other and the container that they are in.
As it applies to liquids
Liquids are made of particles that can be tiny or a bit larger
but that are closer together due to the effects of IMF.
Liquid particles have less kinetic energy than gases, and
thus are in motion, but at a lower rate.
Liquids form at lower temperatures than gases.
Liquid particles have attraction due to IMF so they slide
past each other more than collide like gases.
Fluidity: particles flow past each other and take the
shape of their container
Relatively high density: most substances are 100s
of times more dense as liquids than gases
As solids they get 10% more dense, but that’s it (with the
exception of water which is less dense as a solid)
Incompressibility: unlike gases, they can’t really be
compressed together much more
Properties of Liquids
Ability to diffuse: just like gases,
but slower
Surface tension: force that
pulls adjacent parts of a
liquid’s surface together,
minimizing the surface area
This happens because of
IMF!
The higher the force of the
attraction, the greater the
surface tension.
Capillary action: attraction
of the surface of a liquid to
the surface of a solid,
against the pull of gravity
Properties of Liquids
Definite volume and definite
shape
Particles have lowest kinetic
energy, thus move the least.
Solids
Strongest intermolecular
forces (IMF), holding the
particles tightly together in
a rigid structure.
Particle motion is limited
due to rigidity, thus they
vibrate in place.
Particle diagram:
Crystalline solid: made of crystals;
particles arranged in an orderly
geometric pattern
Ex. Salt
The repeating coordinated formation =
lattice
Amorphous solid: particles are
arranged more randomly
Ex. Glass, plastics, candles, cotton candy
Sometimes classified as “supercooled
liquids” because they retain certain liquid
properties (like fluidity) even when they
appear to be solids.
2 types
Solids
Kinetic-Molecular Theory
All matter is made of small particles.
Those particles are in constant random motion.
This motion causes the particles to collide with
each other and the container that they are in.
As it applies to solids
These particles are very close together in solids.
Solids have the lowest kinetic energy and thus the most
limited movement.
Solids form at the lowest temperatures.
Because solid particles have the strongest attraction due
to IMF, they vibrate within their rigid structure next to each
other.
High density: most substances are the MOST
dense in their solid state.
Exception = water
Incompressibility: solids are virtually entirely
incompressible.
Properties of Solids
Low ability to diffuse: can
occur but if it does, it is very
slow.
2 other States of Matter
Plasma: matter
composed of positive
ions and electrons
with extremely high
kinetic energy
Most common form of
matter in the universe
and the least common
on Earth
Makes up stars, neon
lights, auroras
Changing States
When particles gain or lose thermal energy,
they can undergo a state change.
This is a physical change because the identity of
the matter is still the same.
Ex. When ice melts it is still water (H2O).
Changing States
When particles gain or lose thermal energy,
they can undergo a state change.
Heat of fusion: (aka, the molar enthalpy of
fusion) the amount of energy, as heat, needed to
turn 1 mole of solid into a liquid at its melting
point.
When heat is added to a solid, it has more kinetic
energy so the particles vibrate faster and start moving
farther apart as they transition to the liquid state.
Heat of vaporization: (aka, the molar enthalpy of
vaporization) the amount of energy, as heat,
needed to turn 1 mole of liquid into a gas at its
boiling point.
A diagram to summarize how matter undergoes phase changes
= adding
energy
= removing
energy
MELTING
VAPORIZATIO
N
DEPOSITION
SUBLIMATION
FREEZING
CONDENSATION
Solid
Liquid
Gas
Changing States
Melting point = temperature at which a solid becomes
a liquid due to the KE of particles overcoming the
attractive IMF that hold their order together.
This is a definite point in crystalline solids but not amorphous
solids.
Freezing: (solidification) the phase/state change
from liquid to solid by the removal of energy in
the form of heat.
Freezing point = temperature at which a liquid turns
into a crystalline solid.
This is the same temperature as a melting point, just the
process is going in the opposite direction.
Solid Liquid
Melting: the phase/state
change of a solid to a liquid
from adding energy in the
form of heat.
Condensation: the
phase/state change
from gas to liquid by
the removal of energy
in the form of heat.
Liquid Gas
Vaporization: the
phase/state change of a
liquid to a gas from adding energy in the form of
heat.
Occurs two different ways.
Evaporation: occurs only at the
surface of a liquid, when particles
escape the surface of a nonboiling
liquid and become a gas due to a
pressure change.
A necessary distinction
Liquid Gas
Particles have different KE. Higher KE particles
move faster and can overcome the IMF that keep
them in a liquid state.
Volatile liquids = liquids that readily evaporate due to
weak attractive forces between particles
Vapor pressure: the point in which pressure is in
equilibrium and thus molecules move between the
liquid and gas phases at the same rate.
Boiling: caused by a temperature
change, occurs throughout the
liquid as the liquid particles
change to bubbles of vapor.
A necessary distinction
Liquid Gas
Boiling point = temperature at which the vapor
pressure of the liquid is equal to atmospheric
pressure.
All energy absorbed gets used to evaporate the liquid
with a constant temperature and pressure.
Ex. At normal atmospheric pressure (= 1 atm = 760 torr = 101.3
kPa) water boils at 100℃.
Sublimation: the
phase/state change of
a solid directly to a
gas
Ex. Dry ice (solid CO2)
Solid Gas
Deposition: change of
a gas directly to a
solid
Ex. When frost forms on
a cold surface
Heating curve: a diagram that shows the phase
changes a substance goes through as energy, in
the form of heat, is added to it.
Heating Curve
Solid
Liquid
Gas
Melting
Boiling
Heat of fusion
Heat of vaporization
Phase diagram: a graph that shows pressure vs.
temperature; allows us to know what phase a
substance would be in at various pressures and
temperatures.
Phase Diagram
Shows how the
states of a system
change as
temperature and
pressure change
SOLI
D
LIQUI
D
GAS
= normal
freezing point
= normal
boiling point
= triple
point
= critical
point
Triple point: indicates the temperature and
pressure conditions necessary for a solid, liquid,
and gas of a substance to coexist at equilibrium
Critical point: indicates the critical temperature
and the critical pressure.
Critical temperature = the temperature above which a
substance cannot exist in a liquid state
Critical pressure = the lowest pressure that marks
where the substance can exist as a liquid at the critical
temperature
Overview
Pressure (P) = force per unit area on a
surface
Gas particles exert
pressure on every
surface they collide
with.
Increasing pressure
increases collisions
of particles.
Pressure is dependent
on volume,
temperature, and
number of particles.
Overview
Pressure (P) = force per unit area on a
surface
Can be measured in:
atm
mm Hg
Torr
Pa
kPa
psi
1 atm = 760 mm Hg = 760 torr = 1.01325 x 105
Pa = 101.325 kPa = 14.700 psi
Overview
Atmospheric
pressure: pressure
exerted by the
atmosphere (shell) that
surrounds Earth.
It is the sum of all of the
individual pressures of
the various gases that
make up the
atmosphere.
Standard atmospheric
pressure = 1 atm
Note: If a practice problem says
“STP”, it means that the conditions
are standard temperature and
pressure of 273 K and 1 atm.
Standard temperature = 273 K
Practice Time!
Convert 1140 mm Hg to atm.
Convert 202 kPa to psi.
Convert 19.0 psi to torr.
The reading on a tire pressure gauge says 35
psi. What is this in atm?
Pressure conversions
Dalton’s Law
The total pressure of a gas mixture is the
sum of the partial pressures of the
component gases.
Partial pressure: pressure of each gas in a
mixture
PT = P1 + P2 + P3…etc.
Dalton’s Law
Example: Oxygen gas with a partial pressure
of 301 mm Hg is mixed in a container with
chlorine gas that has a partial pressure of
0.649 atm. What is the total pressure inside
the container, in atm?
P1 = 301 mm Hg
P2 = 0.649 atm
PT = ?
1 atm = 760 mm Hg
PT = P1 + P2
301 mm Hg
760 mm Hg
1 atm
=
301 atm
760
=
0.396
atm
PT = 0.396 + 0.649
PT = 1.050 atm
Boyle’s Law
The volume of a gas is inversely
proportional to its pressure.
With a constant
mass and
temperature!
P1V1 = P2V2
Cartesian Diver
Boyle’s Law
Example: A sample of carbon dioxide gas has
a volume of 2.0 L with a pressure of 3.2 atm.
What volume would be needed to decrease
the pressure to 1.5 atm?
V1 = 2.0 L
P1 = 3.2 atm
P2 = 1.5 atm
V2 = ?
P1V1 = P2V2
P1V1 = P2V2
P2
P2
V2
=
P1V1
P2
V2
=
(3.2)(2.0)
1.5
V2
=
(6.4)
1.5
V2 = 4.3 L
Charles’s Law
The volume of a gas is directly proportional
to its temperature.
T1
T2
With a constant
mass and pressure
V1 = V2
Note: Temperature
MUST be in Kelvin!
K = ℃ + 273
Charles’s Law
Example: Nitrogen gas is cooled from 120℃
to 51℃. If its new volume is 65 mL, what was
its original volume?
T1 = 120℃
T2 = 51℃
V2 = 65 mL
V1 = ?
V1 = V2
T1
T2
K = ℃ + 273
K = ℃ + 273
K = 120℃ + 273
K = 393 K
K = 51℃ + 273
K = 324 K
V1 = V2
T1
T2
(T1)
(T1)
V1 = V2
T2
(T1)
V1 = 65
324
(393)
V1 = 79 mL
Gay-Lussac’s Law
The pressure of a gas is directly
proportional to its temperature.
T1
T2
With a constant
mass and volume
P1 = P2
Note: Temperature
MUST be in Kelvin!
K = ℃ + 273
*Can demo
Gay-Lussac’s Law
Example: If a gas is heated from 231 K to 317
K with volume being held constant, what would
be the final pressure if the initial pressure
was 619 mm Hg?
T1 = 231 K
T2 = 317 K
P1 = 619 mm Hg
P2 = ?
P1 = P2
T1
T2
P1 = P2
T1
T2
(T2)
(T2)
P2 = P1
T1
(T2)
P2 = 619
231
(317)
P2 = 849 mm Hg
Combined Gas Law
Represents the relationship between
pressure, volume, and temperature of a
fixed amount of gas.
Amount of gas must stay constant for this to
hold true!
P1V1 = P2V2
T1
T2
Combined Gas Law
Example: A sample of an unknown gas has a volume of
2.50 L, a pressure of 0.861 kPa, and a temperature of
299 K. If the volume is doubled to 5.00 L and the
pressure reduced to 0.551 kPa, what would you expect
the temperature to be?
V1 = 2.50 L
P1 = 0.861 kPa
T1 = 299 K
V2 = 5.00 L
P2 = 0.551 kPa
T2 = ?
P1V1 = P2V2
T1
T2
P1V1 = P2V2
T1
T2
(T2)
(T2)
P1V1 = P2V2
T1
(T2)
(T1)
(T1)
P1V1 = P2V2
(T2)
(T1)
P1V1
P1V1
P2V2
T2 =
T1
P1V1
P2V2
T2 =
T1
P1V1
(0.551)(5.00)(299)
T2 =
(0.861)(2.50)
824
T2 =
2.15
T2 = 383 K
Gas Laws and Volume
Gay-Lussac’s law of combining volumes of gases:
The volumes of gaseous reactants and products can
be expressed as small whole number ratios
With a constant temperature and pressure
Avogadro’s Law: Equal volumes of gases contain
equal numbers of molecules
Meaning, gas volume is directly proportional to the
amount of gas
With a constant temperature and pressure.
V = kn
V = volume
k = constant
n = amount of gas, in moles
Why this matters: We can use coefficients
in chemical equations to tell us the relative
number of moles AND molecules AND
volumes!
Ex. 2H2(g) + O2(g) 🡪 2H2O(g)
2 : 1 : 2
Molar Volume
Molar volume: the volume that 1 mole of a
gas occupies at STP
STP (standard temperature and pressure) = a
temperature of 0.00 ℃ (273 K) and a pressure
of 1.00 atm.
**At STP the volume of 1 mol of any gas is
22.4 L**
We can use this in stoichiometry calculations as a
conversion factor!
22.4 L of O2 at STP = 1 mol of O2
22.4 L of O2
1 mol of O2
1 mol of O2
22.4 L of O2
or
Molar Volume
Example: What is the volume of 6.50 moles
of H2 gas at STP?
6.50 mol H2
1 mol H2 = 22.4 L H2
V = ?
6.50 mol H2
1 mol H2
22.4 L
=
145.6 L
1
=
146 L of H2
Molar Volume
Example: What is the mass of 28.0 L of
helium gas at STP?
28.0 L He
1 mol He = 22.4 L He
g = ?
28.0 L He
22.4 L He
1 mol He
Molar mass of He =
4.00 g/mol
1 mol He
4.00 g He
=
112 g He
22.4
=
5.00 g He
Practice Time!
Find the volume of 0.513 moles of Cl2 gas at
STP.
Find the volume of 12.0 g of neon gas at
STP.
How many molecules are in 2.73 L of H2 gas
at STP?
Molar volume calculations
Ideal Gas Law
Shows the mathematical relationship
between pressure, volume, temperature,
and number of moles of a gas
PV = nRT
P = pressure, in atm
V = volume, in L
n = amount of gas, in mol
R = ideal gas constant; 0.0821 L·atm
mol · K
T = temperature, in Kelvin
Ideal Gas Law
Example: How many moles of gas would be
contained in 2.75 L of gas at 30.0 °C and
2.00 atm?
V = 2.75 L
T = 30.0 °C
P = 2.00 atm
R = 0.0821 L·atm
mol·K
= 303 K
n = ?
PV = nRT
PV = nRT
RT
RT
n = PV
RT
n = (2.00)(2.75)
(0.0821)(303)
n = 5.50
24.9
n = 0.221 mol
Ideal Gas Law
Example: What mass of CO2 would fill an
80.0 L tank at STP?
V = 80.0 L
T = 273 K
P = 1.00 atm
R = 0.0821 L·atm
mol·K
n = ?
PV = nRT
PV = nRT
RT
RT
n = PV
RT
n = (1.00)(80.0)
(0.0821)(273)
n = 80.0
22.4
n = 3.57 mol CO2
g = ?
Molar mass of CO2 =
1(12.01) + 2(16.00) =
44.01 g/mol
1 mol CO2
44.01 g CO2
=
157 g CO2
1
=
157 g CO2