States of Matter
The Solid State
•Particles are tightly packed, very close together (strong cohesive forces)
•Low kinetic energy (energy of motion)
•Fixed shape and volume
•Crystalline or amorphous structure
The Liquid State
•Particles are close to each other (making them mostly
incompressible)
•Attractive forces keep molecules close, but not so close to
restrict movement
The Gas State
•Gas particles move randomly and rapidly.
•Size of gas particles is small compared to the space between the
particles.
•Gas particles exert no attractive forces on each other.
• Kinetic energy of gas particles increases with increasing
temperature.
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a
The symbol “~” means approximately.
Gases and Pressure
•When gas particles collide with the walls of a container,
they exert a pressure.
•Pressure (P) is the force (F) exerted per unit area (A).
Pressure
=
Force
=
Area
F
A
760. mm Hg
1 atmosphere (atm) =
760. torr
14.7 psi
101,325 Pa
Gas Laws
Mathematical relationships describing the behavior of gases
with regard to mixing, diffusion, changes in pressure,
changes in temperature
Boyle’s Law: Describes the relation between pressure and
volume of a gas, under a constant temperature
PiVi = PfVf
where i = initial condition and f = final condition
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Boyle’s Law: Inverse relation between Pressure and Volume
Example:
Freon-12, CCl2F2, is used in refrigeration systems. What is the
new volume (L) of a 8 L sample of Freon gas initially at 50 mm
Hg after its pressure is changed to 200 mm Hg at constant T?
1.
Set up a data table
Conditions 1
P1 = 50 mm Hg
V1 = 8 L
P2
V2
Conditions 2
= 200 mm Hg
= ?
2. Solve Boyle’s Law for V2:
P1V1 = P2V2
V2
= V1P1
P2
V2 = 8 L x 50 mm Hg
=
200 mm Hg
2L
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Learning Check
A sample of helium gas has a volume of 6.4 L at
a pressure of 0.70 atm.
What is the new volume when the pressure is
increased to 1.40 atm (T constant)?
Solution
P1V1 = P2V2
Solve for V2:
V2 = V1P1
P2
V2
= 6.4 L x 0.70 atm = 3.2 L
1.40 atm
Volume decreases when there is an increase in the pressure
(Temperature is constant).
Learning Check
A sample of oxygen gas has a volume of 12.0 L at 600.
mm Hg. What is the new pressure when the volume
changes to 36.0 L? (T and n constant.)
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Solution
Conditions 1
Conditions 2
P1 = 600. mm Hg
V1 = 12.0 L
P2
=
P2 = ?
V2 = 36.0 L
P1 V1
V2
600. mm Hg x 12.0 L = 200. mm Hg
36.0 L
Charles’ Law: Describes relation between temperature and
volume of a gas, under constant pressure
Vi/Ti = Vf/Tf
Charles’s Law: Direct relationship between Volume and Temperature
Charles’ Law
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Example:
A balloon has a volume of 785 mL at 21°C. If the
temperature drops to 0°C, what is the new volume of the
balloon (P constant)?
1. Set up data table:
Conditions 1
Conditions 2
V1 = 785 mL
T1 = 21°C = 294 K
V2 = ?
T2 = 0°C = 273 K
Be sure that you always use the Kelvin (K)
temperature in gas calculations!
2. Solve Charles’ law for V2
V1 = V2
T1
T2
V2 = V1 T2
T1
V2 = 785 mL x 273 K = 729 mL
294 K
Learning Check
A sample of oxygen gas has a volume of 420 mL at a
temperature of 18°C. At what temperature (in °C) will
the volume of the oxygen be 640 mL (P and n constant)?
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Solution
T2 =
T2
T1V2
V1
= 291 K x
640 mL
420 mL
= 443 K – 273 K
= 443 K
= 170°C
Combined Gas Law: Describes relation between pressure,
temperature and volume of a gas
PiVi/Ti = PfVf/Tf
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Example:
A sample of helium gas has a volume of 0.180 L, a pressure of 0.800
atm and a temperature of 29°C. At what temperature (°C) will the
helium have a volume of 90.0 mL and a pressure of 3.20 atm (n
constant)?
1. Set up Data Table
Conditions 1
Conditions 2
P1 = 0.800 atm
P2 = 3.20 atm
V1 = 0.180 L (180 mL)
V2 = 90.0 mL
T1 = 29°C + 273 = 302 K T2 = ?
2. Solve for T2
P1 V1
T1
=
P2 V2
T2
T2 = T1 P2V2
P1V1
T2 = 302 K x 3.20 atm x 90.0 mL = 604 K
0.800 atm 180.0 mL
T2 = 604 K – 273 = 331 °C
Learning Check
A gas has a volume of 675 mL at 35°C and 0.850 atm
pressure. What is the volume(mL) of the gas at –95°C
and a pressure of 802 mm Hg (n constant)?
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Solution
Data Table
T1 = 308 K
V1 = 675 mL
P1 = 646 mm Hg
T2 = -95°C + 273 = 178K
V2 = ?
P2 = 802 mm Hg
Solve for V2
V2 =
V1 P1 T2
P2T1
V2 = 675 mL x 646 mm Hg x 178K
802 mm Hg x 308K
= 314 mL
Gay–Lussac’s Law: Describes the relation between pressure
and temperature of a gas, at a constant volume
•Pressure and temperature are directly related
Pressure
Temperature
P
=
P1
T1
constant
=
=
k
T
P2
T2
Note: Temperature must be expressed in kelvins.
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Avogadro’s Law: Equal volumes of gases measured at the same
temperature and pressure contain equal number of molecules
Vi/ni = Vf/nf
where n = number of moles
Avogadro’s Law Example:
If 0.75 mole of helium gas occupies a volume of 1.5 L, what volume
will 1.2 moles of helium occupy at the same temperature and
pressure?
Conditions 2
V2 = ?
n2 = 1.2 moles He
Conditions 1
V1 = 1.5 L
n1 = 0.75 mole He
V1/n1 =V2/n2
V2 = V1n2
n1
V2 = 1.5 L
x 1.2 moles He =
0.75 mole He
2.4 L
Ideal Gas Law: Describes relation between pressure,
volume, temperature and the number of molecules in an ideal
gas sample
PV = nRT
where R = universal gas constant (0.0821 L atm/K mol)
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Ideal Gas Law Example:
A cylinder contains 5.0 L of O2 at 20.0°C and 0.85 atm. How many
grams of oxygen are in the cylinder?
P = 0.85 atm, V = 5.0 L, T = 293 K, n (or g =?)
PV = nRT Æ n = PV
RT
= (0.85 atm)(5.0 L)(mole K) = 0.18 mole O2
(0.0821atm L)(293 K)
= 0. 18 mole O2 x 32.0 g O2 = 5.8 g O2
1 mole O2
Partial Pressure: Pressure an individual gas in a mixture would
exert were it alone in the same container
Dalton’s Law: Total pressure exerted by a mixture of gases equals
the sum of the partial pressures
P(total) = P(gas 1) + P(gas 2) + P(gas 3)
etc.
Dalton’s Law of Partial Pressures
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Summary of Gas Laws
Boyle: PiVi = PfVf
Charles: Vi/Ti = Vf/Tf
Avogradro: Vi/ni = Vf/nf
Gay-Lussac: Pi/Ti = Pf/Tf
Dalton: P(total) = P(gas 1) + P(gas 2) + P(gas 3)
*Combined: PiVi/Ti = PfVf/Tf
*Ideal:
PV = nRT
* Memorize for exam
Intermolecular Forces
Intermolecular forces: attractive forces that exist
between molecules.
In order of increasing strength, these are:
•London dispersion forces
•dipole–dipole interactions
•hydrogen bonding
London Dispersion Forces
London dispersion forces: weak interactions due to the
momentary changes in electron density in a molecule.
•Change in electron density creates a temporary dipole.
•The weak interaction between these temporary dipoles constitutes
London dispersion forces.
•All covalent compounds exhibit London dispersion forces.
•The larger the molecule, the larger the attractive force, and the
stronger the intermolecular forces.
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Dipole-dipole interaction: attraction between positive end of one
polar molecule and negative end of a different polar molecule
Hydrogen bonding: Specific type of dipole-dipole force,
between the partial positive charge on H and partial negative
charge on an electronegative element such as O, N, F
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Intermolecular Forces:
Boiling Point and Melting Point
Boiling point: temperature at which a liquid is converted to a gas
Melting point: temperature at which a solid is converted to a liquid
The stronger the intermolecular forces on a substance, the higher
its boiling point and melting point are.
Examples of Intermolecular Forces and
Boiling, Melting Points:
•Both molecules have London dispersion forces and nonpolar bonds.
•In this case, the larger molecule will have stronger attractive forces.
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Vapor Pressure
Evaporation: the conversion of liquids into the gas phase.
•Evaporation is endothermic—it absorbs heat from
the surroundings.
Condensation: the conversion of gases into the liquid phase.
•Condensation is exothermic—it gives off heat to
the surroundings.
Viscosity and Surface Tension
Viscosity: a measure of a fluid’s resistance to flow freely
•Compounds with strong intermolecular forces
tend to be more viscous than compounds with
weaker forces.
•Substances composed of large molecules tend
to be more viscous, too, because large molecules
do not slide past each other as freely.
Surface tension: a measure of the resistance of a liquid to spread out.
•Interior molecules in a
liquid are surrounded by
intermolecular forces on
all sides.
•Surface molecules only
experience intermolecular
forces from the sides and
from below.
The stronger the intermolecular forces, the higher the surface tension.
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The Solid State: Types of Solids
Crystalline solid: has a regular arrangement of particles—
atoms, molecules, or ions—with a repeating structure.
There are four different types of crystalline solids—
ionic, molecular, network, and metallic.
Crystalline Solids
Ionic solid: composed of oppositely charged ions
Molecular solid: composed of individual molecules arranged regularly
Network solid: composed of a vast number of atoms covalently
bonded together (SiO2).
Metallic solid: a lattice of metal cations surrounded by a cloud of e−
that move freely (Cu).
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Amorphous Solids
Amorphous solid: has no regular arrangement of its closely
packed particles.
•They can be formed when liquids cool too quickly
for regular crystal formation.
•Very large covalent molecules tend to form
amorphous solids, because they can become
folded and intertwined.
Examples: rubber, glass, and plastic.
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