Title: STATES OF MATTER
1STATES OF MATTER
- Gases, Liquids and Solids
2The Nature of Gases
- Air What is it?
- All gases have similar physical behaviors.
- 1 mole of gas 22.4 L _at_ STP.
- Particles
- molecules (diatomic or polyatomic)
or - single atoms (noble or inert gases).
3The Physical Properties of Gases
- Have mass (volume-matter).
- Density mass / volume.
- Have compressibility.
- Fill their container completely.
- Exert pressure.
- Diffusion The rate at which gases mix.
4The Kinetic Molecular Theory
- Gases have mass.
- The distance separating gas particles is
relatively large. - Gases have constant random motion.
- Gases exert pressure because of collisions.
Collisions are perfectly elastic (no energy of
motion is lost). - The average KE of the gas is temperature
dependent. - Gas particles do not exert forces on another.
5Gas Measurement Involves 4 Variables
- Amount (n) Mole
- n mass / molar mass
- Volume (V)
- Gas fills the container volume of the container
Temperature (T) Kelvin vs. Celsius K OC
273 Pressure (P) The result of gas particle
collisions with the walls of the container.
6Atmospheric Pressure / Barometer
- Pressure exerted by the air in the atmosphere.
- Result of mass and gravity.
- Pressure force / unit area.
- Force Newton.
- Pressure Pascal (Pa).
- 1 atm pressure at sea level.
- 1 atm 101.3 kPa 14.7 lb/in2 760 mm Hg
(torr). - Atmospheric pressure decreases as altitude
increases!!!!
7Enclosed Gases
pressure
- Manometer a U tube used to measure
- pressure of an enclosed gas
- a higher mercury level in the gas arm indicates
that - the gas pressure is lower than atmospheric
pressure - a higher mercury level in the open arm
indicates - that the gas pressure is higher than atmospheric
pressure
liquid
8The Gas Laws
- Mathematical representations of the relationships
between the four variables - P, V, T, n
9Boyles Law
- The pressure and volume of a sample of gas at
constant temperature are inversely proportional
to each other. - At constant temperature, the volume of a fixed
amount of gas will decrease as the pressure
increases. - Spring of air or compressibility.
P1V1 P2V2 Sample Problems Pg. 433
10Charless Law
- The volume of gas is directly proportional to its
temperature. - At constant pressure, the volume of a fixed
amount of gas is directly proportional to its
absolute temperature. - Review absolute zero and the Kelvin scale.
V1T2 V2T1 Sample Problems Pg. 438
11Gay-Lussacs LawThe pressure temperature
relationship.
- At a constant volume , as the pressure increases
the temperature will increase. - A directly proportional, linear relationship.
P1 / T1 P2 / T2
12Combined Gas Law
- No Constants.
- V1 x P1 V2 x P2
- T1 T2
13Avogadros LawThe mole volume relationship.
- Equal volumes of gases at the same temperature
and pressure contain an equal number of
particles.
14Daltons Law of Partial Pressure
- The sum of the partial pressures of all the
components in a gas mixture is equal to the total
pressure of the gas mixture. - PT P1 P2 P3 .
15Ideal Gas Equation
- An ideal gas is described by the
kinetic-molecular theory postulates. - No such ideal gas exists.
- Real gases behave like ideal gases under many
ordinary conditions. - Exceptions are low temperatures and high
pressures. - PV nRT
- R universal gas constant
16Kinetic Molecular Theory
- Condensed states liquids and solids
- -higher densities then gases
- -variables are amount and temperature
- Liquids The attractive forces between the
particles are substantially stronger than a gas,
but are still not strong enough to hold the
particles in a fixed position. - Solids The attractive forces between particles
are stronger than in a gas or liquid and the
particles can not overcome these forces and move
away from each other.
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18Kinetic Molecular Theory
- According to the Kinetic Molecular Theory, the
state of a substance at room temperature depends
on the strength of the attractions between its
particles. - Temperature is a measure of the average KE of
the particles in a substance. - Different To different KE.
19Intermolecular Forces
- Ionic vs. Covalent bonds.
- Covalent bonds molecules.
- Intramolecular forces covalent bonds.
- Intermolecular forces The forces of attraction
between neighboring molecules - Substantially weaker than ionic or covalent
bonds. - Involved in the change of state.
- There are 3 types
- 1. Dispersion forces
- 2. Dipole-dipole forces
- 3. Hydrogen bonds.
20Dispersion Forces
- A force of attraction between induced dipoles.
- Induced dipole A dipole created by the presence
of neighboring dipole. - Perfectly symmetrical electron cloud vs.
temporary dipole. - Noble gas boiling points vs. dispersion forces.
21Dipole Dipole Forces
- Attractions between opposite charges of adjacent
permanent dipoles. - Polar bonding.
22Hydrogen Bonding
- Present when a covalent bond is formed between
hydrogen (weak electronegativity) and an element
with high electronegativity. - Fluorine, oxygen, and nitrogen.
23According to the kinetic molecular theory, the
state of a substance at room temperature depends
on the strength of the attractions between its
particles.
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25Properties of Liquids
- Liquid physical properties are determined by the
nature and strength of the intermolecular forces
present. - Viscosity
- A measure of resistance to motion that exists in
a liquid. - Strong intermolecular forces greater viscosity.
- Increases as the temperature decreases.
- Surface tension The resistance of a liquid to
an increase in its surface area. - An imbalance of forces at the surface of a
liquid. - The surface will behave as if a tight film was
stretched across it. - Increases with strong intermolecular forces.
26Water
- High boiling point.
- High specific heat.
- Density of solid is less than liquid.
- High surface tension.
- High heat of vaporization.
- The universal solvent.
- Heating Curves Phase Diagrams
27Changes of State Always involves a change in
energy.
MP BP Endothermic (energy absorbed) CP
FP Exothermic (energy released)
- BP GAS CP
- MP LIQUID FP
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- SOLID