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Chapter 6
States of Matter:
Gases, Liquids, and Solids
Denniston Topping Caret
4th Edition
Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
6.1 The Gaseous State
Ideal Gas Concept
• Ideal gas - a model of the way that particles of a gas behave at the microscopic level.
• We can measure the following of a gas:– temperature,– volume,– pressure and – quantity (mass)
We can systematically change one of the properties and see the effect on each of the others.
6.1
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teMeasurement of Gases
• The most important gas laws involve the relationship between– number of moles (n) of gas– volume (V)– temperature (T)– pressure (P)
• Pressure - force per unit area.
• Gas pressure is a result of force exerted by the collision of particles with the walls of the container.
6.1
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te• Barometer - measures atmospheric
pressure.
– Invented by Evangelista Torricelli
• A commonly used unit of pressure is the atmosphere (atm).
• 1 atm is equal to:
760 mmHg
760 torr
76 cmHg
6.1
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teBoyle’s Law
• Boyle’s Law - volume of a gas is inversely proportional to pressure if the temperature and number of moles is held constant.
PV = k1
or
PiVi = PfVf
1
6.1
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6.1
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1. A 5.0 L sample of a gas at 25oC and 3.0 atm is compressed at constant temperature to a volume of 1.0 L. What is the new pressure?
2. A 3.5 L sample of a gas at 1.0 atm is expanded at constant temperature until the pressure is 0.10 atm. What is the volume of the gas?
2
16.
1 T
he
Gas
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tate Charles’ Law
• Charles’ Law - volume of a gas varies directly with the absolute temperature (K) if pressure and number of moles of gas are constant.
2T
Vk or
f
f
i
i
T
V
T
V
6.1
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6.1
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1. A 2.5 L sample of gas at 25oC is heated
to 50oC at constant pressure. Will the
volume double?
2. What would be the volume in question 1?
3. What temperature would be required to
double the volume in question 1?
2
6.1
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f
ff
i
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T
VP
T
VP
• This law is used when a sample of gas undergoes change involving volume, pressure, and temperature simultaneously.
1
2
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Calculate the temperature when a 0.50 L sample of gas at 1.0 atm and 25oC is compressed to 0.05 L of gas at 5.0 atm.
6.1
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te Avogadro’s Law
• Avogadro’s Law - equal volumes of an ideal gas contain the same number of moles if measured under the same conditions of temperature and pressure.
3
Vk
n or
f
f
i
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nn
VV
1
6.1
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Assuming no change in temperature and pressure, how many moles of gas would be needed to double the volume occupied by 0.50 moles of gas?
2
6.1
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teMolar Volume of a Gas
• Molar Volume - the volume occupied by 1 mol of any gas.
• STP - Standard Temperature and Pressure T = 273 K (or 0oC) P = 1 atm
• At STP the molar volume of a gas is 22.4 L– We will learn to calculate the volume
later.
6.1
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te Gas Densities
• We know: density = mass/volume
• Let’s calculate the density of H2.
• What is the mass of 1 mol of H2?• 2.0 g
• What is the volume of 1 mol of H2?• 22.4 L
• Density = 2.0 g/22.4 L = 0.089 g/L
6.1
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te The Ideal Gas Law1
• Combining Boyle’s Law, Charles’ Law and Avogadro’s Law gives the Ideal Gas Law.
PV=nRT
R (ideal gas constant) = 0.08206 L.Atm/mol.K
6.1
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teLet’s calculate the molar volume at STP
using the ideal gas law:
PV = nRT
• What would be the pressure? 1 atm
• What would be the temperature? 273 K
• What would be the number of moles? 1 mol
atm 1
K 273)Kmol
atmL6mol(0.08201
P
RTV
n22.4 L
2
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1. What is the volume of gas occupied by 5.0 g CH4 at 25oC and 1 atm?
2. What is the mass of N2 required to occupy 3.0 L at 100oC and 700 mmHg?
6.1
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• Dalton’s Law - a mixture of gases exerts a pressure that is the sum of the pressures that each gas would exert if it were present alone under the same conditions.
Pt=p1+p2+p3+...
• For example, the total pressure of our atmosphere is equal to the sum of the pressures of N2 and O2.
22 ONair ppP
6.1
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1. Gases are made up of small atoms or molecules that are in constant and random motion.
2. The distance of separation is very large compared to the size of the atoms or molecules.
– the gas is mostly empty space.
Provides an explanation of the behavior of gases that we have studied in this chapter. Summary follows:
6.1
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– No attractive or repulsive forces exist between them.
4. Gas particles collide with each other and with the walls of the container without losing energy. – The energy is transferred from one atom or
molecule to another.
5. The average kinetic energy of the atoms or molecules is proportional to absolute temperature. – K.E. = 1/2mv2 so as temperature goes up, the
speed of the particles goes up.
6.1
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of Gases explain the following statements?
• Gases are easily compressible.
• Gases will expand to fill any available volume.
• Gases have low density.
4
6.1
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te• Remember: pressure is a force per unit
area resulting from collision of gas particles with the walls of the container.
• If pressure remains constant why does volume increase with temperature?
6.1
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Ideal Gases Vs. Real Gases
• In reality there is no such thing as an ideal gas. Instead this is a useful model to explain gas behavior.
• Non-polar gases behave more ideally than polar gases because attractive forces are present in polar gases.
6.2 The Liquid State
• Liquids are practically incompressible.– Enables brake fluid to work in your car
• Viscosity - a measure of a liquids resistance to flow.– Flow occurs because the molecules can easily
slide past each other.– Glycerol - example of a very viscous liquid.– Viscosity decreases with increased temperature.
5
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• Surface tension - a measure of the attractive forces exerted among molecules at the surface of a liquid.
• Surface molecules are surrounded and attracted by fewer liquid molecules than those below.
• Net attractive forces on surface molecules pull them downward.
– Results in “beading”
• Surfactant - substance added which decreases the surface tension
– example: soap
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Vapor Pressure of a Liquid6
• What happens when you put water in a sealed container?
• Both liquid water and water vapor will exist in the container.
• How does this happen below the boiling point?
• Kinetic Theory - Liquid molecules are in continuous motion, with their average kinetic energy directly proportional to the Kelvin temperature.
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energy + H2O(l) H2O(g)
The green line represents the minimum energy required to break the intermolecular attractions.
Even at the cold temp, some molecules can be converted.
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• Once there are molecules in the vapor phase, they can be converted back to the liquid phase
H2O(g) H2O(l) + energy
• evaporation - the process of conversion of liquid to gas, at a temperature too low to boil
• condensation - conversion of the gas to the liquid state.
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• When the rate of evaporation equals the rate of condensation, the system is at equilibrium.
• Vapor pressure of a liquid - the pressure exerted by the vapor at equilibrium
H2O(g) H2O(l)
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• Boiling point - the temperature at which the vapor pressure of the liquid becomes equal to the atmospheric pressure.
• Normal boiling point - temperature at which the vapor pressure of the liquid is equal to 1 atm.
• What happens when you go to a mountain where the atmospheric pressure is lower than 1 atm?
• The boiling point lowers.
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• Boiling point is dependant on the intermolecular forces
– Polar molecules have higher b.p. than nonpolar molecules.
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Van der Waals Forces 7
• Van der Waals Forces are types of
intermolecular forces.• Consists of:
1 Dipole-dipole interactions (section 4.5)2 London forces
• London forces:– exist between all molecules,– is the only attractive force between
nonpolar atoms or molecules,– and dipole-dipole attractions occur between
polar molecules.
• Electrons are in constant motion.• Electrons can be, in an instant, arranged
in such a way that they have a dipole. (Instantaneous dipole)
• The temporary dipole interacts with other temporary dipoles to cause attraction.
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6.2
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Hydrogen Bonding8
• Hydrogen bonding:
– not considered a Van der Waals Force
– is a special type of dipole-dipole attraction
– is a very strong intermolecular attraction causing higher than expected b.p. and m.p.
• Requirement for hydrogen bonding:
– molecules have hydrogen directly bonded to O, N, or F
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Examples of hydrogen bonding:
H2O
NH3
HF
6.3 The Solid State
• Particles highly organized and well defined fashion
• Fixed shape and volume
• Properties of Solids:– incompressible– m.p. depends on strength of attractive force
between particles– Crystalline solid - regular repeating structure– Amorphous solid - no organized structure.
9
6.3
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1. Ionic Solids• held together by electrostatic forces
• high m.p. and b.p.
• hard and brittle
• if dissolves in water, electrolytes
• NaCl
2. Covalent Solid• Held together entirely by covalent bonds• high m.p. and b.p.• extremely hard• Diamond
6.3
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te3. Molecular solids
• molecules are held together with intermolecular forces
• often soft• low m.p.• often volatile• ice
4. Metallic solids• metal atoms held together with metal bonds• metal bonds
– overlap of orbitals of metal atoms
– overlap causes regions of high electron density where electrons are extremely mobile - conducts electricity
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The EndChapter 6
