1. Chapter 13: States of Matter
Honors Chemistry

2. 1. The Kinetic-Molecular Theory
Section 1 - Gases
1. The Kinetic-Molecular Theory
The term gas comes from the Greek word chaos, which means without order.
Atomic composition affects chemical and physical properties of solids and liquids.
Substances that are gases at room temperature display similar physical properties despite different compositions
Kinetic-molecular theory explains the different properties of solids, liquids, and gases.
Kinetic comes from the Greek word meaning “to move”.
The kinetic-molecular theory describes the behavior of matter in terms of particles in motion.
The model makes several assumptions about size, motion, and energy of gas particles

3. The Kinetic-Molecular Theory (cont.)
Particle size
Gases consist of small particles separated by empty space.
Gas particles are too far apart to experience significant attractive or repulsive forces.
Particle motion
Gas particles are in constant random motion.
Particles move in a straight line until they collide with other particles or walls
Collisions between gas particles are elastic
An elastic collision is one in which no kinetic energy is lost, but may be transferred.

4. The Kinetic-Molecular Theory (cont.)
Particle Energy
Kinetic energy of a particle depends on two factors, mass and velocity.
Velocity reflects both speed and direction of motion
All particles of a single gas have the same mass but not same velocity
Kinetic energy and temperature are related
Temperature is a measure of the average kinetic energy of the particles in a sample of matter.
At a given temperature, all gases have the same average kinetic energy

5. 2. Explaining the Behavior of Gases
The constant motion of gas particles allows a gas to expand and fill its container
Low Density (mass/volume)
Great amounts of space exist between gas particles.
Gases with greater density have fewer particles than those with lower density in a given volume
Compression and Expansion
Compression reduces the empty spaces between particles.
Particles fill the available space during expansion

6. Explaining the Behavior of Gases (cont.)
Diffusion and Effusion
Gases easily flow past each other because there are no significant forces of attraction.
The random motion of gas particles causes the gases to mix until they are evenly distributed.
Diffusion is the movement of one material through another.
Rate of diffusion depends on mass of particle
Lighter particles diffuse faster than heavier ones
In order for all particles to have same kinetic energy, lighter particles must have greater velocity than heavier particles

7. Explaining the Behavior of Gases (cont.)
Diffusion and Effusion (cont.)
Effusion is a gas escaping through a tiny opening.
Graham’s law of effusion states that the rate of effusion for a gas is inversely proportional to the square root of its molar mass.
Graham’s law also applies to diffusion of gases at the same temperature.

8. 3. Gas Pressure Pressure is defined as force per unit area.
Gas particles exert pressure when they collide with the walls of their container.
The particles in the earth’s atmosphere exert pressure in all directions called air pressure or atmospheric pressure.
There is less air pressure at high altitudes because there are fewer particles pressing down than at sea level.

9. Gas Pressure (cont.) Measuring Air Pressure
Torricelli invented the barometer.
Barometers are instruments used to measure atmospheric air pressure.
The height of mercury in a barometer is always about 760 mm.
The exact height is determined by two forces, gravity and air pressure.
Manometers measure gas pressure in a closed container.

10. Gas Pressure (cont.) Units of Pressure
The SI unit of pressure is the pascal (Pa).
The pascal is derived from the SI unit of force, the newton (N), which is derived from three SI base units: the kilogram, the meter, and the second.
One pascal is equal to a force of one Newton per square meter or 1 N/m2.
At sea level, the average air pressure is mm Hg when the temperature is 0 ⁰C.
One atmosphere is equal to 760 mm Hg (torr), or kilopascals (kPa).

11. Gas Pressure (cont.) Dalton’s Law of Partial Pressures
Dalton’s law of partial pressures states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases of the mixture.
The partial pressure depends on the number of moles, size of the container, and temperature, and is independent of the type of gas.
At a given temperature and pressure, the partial pressure of one mole of any gas is the same.
Ptotal = P1 + P2 + P3 +...Pn
Partial pressure can be used to calculate the amount of gas produced in a chemical reaction.

12. End of Section 13.1

13. 4. Intermolecular Forces
Section 2 - Forces of Attraction
4. Intermolecular Forces
Attractive forces between molecules cause some materials to be solids, some to be liquids, and some to be gases at the same temperature.
Intermolecular forces (van der Waals forces) can hold together identical particles
The three intermolecular forces, dispersion, dipole-dipole, and hydrogen bonds, are all weaker than intramolecular forces

14. Intermolecular Forces (cont.)
Molecules are nonpolar because electrons are evenly distributed between the equally electronegative atoms
Under the right conditions these molecules can become compressed.
In order for these molecules to be compressed there must be some force of attraction between them
Dispersion forces (London forces) are weak forces that result from temporary shifts in density of electrons in electron clouds.

15. Intermolecular Forces (cont.)
Due to the temporary nature of the dipoles, dispersion forces are the weakest intermolecular force.
Dispersion forces play a significant role only when there are no stronger forces of attraction between identical nonpolar molecules
These forces can have a noticeable effect as the number of electrons involved increase.
This explains why fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid at room temperature.

16. Intermolecular Forces (cont.)
Polar molecules contain permanent dipoles.
Dipole-dipole forces are attractions between oppositely charged regions of polar molecules.
Dipole-dipole forces are stronger than dispersion forces as long as the molecules being compared have approximately the same mass

17. Intermolecular Forces (cont.)
Hydrogen bonds are special dipole-dipole attractions that occur between molecules that contain a hydrogen atom bonded to a small, highly electronegative atom with at least one lone pair of electrons, such as fluorine, oxygen, or nitrogen.
These atoms are electronegative enough to cause a large partial positive charge on the hydrogen atom, yet small enough that their lone pairs can come close to hydrogen atoms

18. End of Section 13.2

19. Section 3 - Liquids & Solids
Forces of attraction keep molecules closely packed in a fixed volume, but not in a fixed position.
Density & Compression:
At the same temperature, both liquid and gas particles have the same average kinetic energy
Liquids are much denser than gases because of the stronger intermolecular forces holding the particles together.
Large amounts of pressure must be applied to compress liquids to very small amounts because they are already tightly packed.

20. Liquids (cont.) Fluidity: Viscosity:
Fluidity is the ability to flow and diffuse; liquids and gases are fluids.
Liquids diffuse more slowly than a gas at the same temperature because intermolecular attractions interfere with the flow.
Viscosity:
Viscosity is a measure of the resistance of a liquid to flow and is determined by the type of intermolecular forces, size and shape of particles, and temperature.
The stronger the attractive forces, the higher the viscosity
The larger the molecule, the slower it moves, the closer they are to each other

21. Liquids (cont.) Viscosity & Temperature:
The stronger the intermolecular attractive forces, the higher the viscosity.
Larger molecules create greater viscosity.
Long chains of molecules result in a higher viscosity.
Increasing the temperature decreases viscosity because the added energy allows the molecules to overcome intermolecular forces and flow more freely.

22. Liquids (cont.) Surface Tension:
Intermolecular forces do not have an equal effect on all particles in a liquid.
Particles in the middle of a liquid can be attracted to particles around them
For surface particles, there are no attractions from above to balance the attractions from below, thus there is a net attractive force pulling down on particles at the surface.
Surface tension is the energy required to increase the surface area of a liquid by a given amount.
The surface tends to have the smallest possible area.
For surface area to increase, particles from the interior must move to the surface.

23. Liquids (cont.) Capillary Action:
Surfactants are compounds that lower the surface tension of water.
Capillary Action:
Cohesion is the force of attraction between identical molecules.
Adhesion is the force of attraction between molecules that are different.
The surface of the water in a graduated cylinder is concave because the water molecules are more strongly attracted to the silicon dioxide in glass (adhesion) than to other water molecules (cohesion)
Capillary action is the upward movement of liquid into a narrow cylinder, or capillary tube.

24. 6. Solids Density of Solids:
The particles in a solid are in constant motion, in order for a substance to be a solid, it must contain particles with strong attractive intermolecular forces.
These forces limit the motion of the particles to vibration in a fixed position.
Density of Solids:
Most solids are more dense than most liquids.
There is about a 10% difference in density between the solid and liquid states of most substances
Ice is not more dense than water.

25. Solids (cont.) Crystalline Solids:
Crystalline solids are solids with atoms, ions, or molecules arranged in an orderly, geometric, three-dimensional structure.
A unit cell is the smallest arrangement of atoms in a crystal lattice that has the same symmetry as the whole crystal.

26. Solids (cont.)
Crystalline solids can be classified into five categories based on types of particles they contain:
Atomic solids
Molecular solids
Molecules held together by van der Waals forces
Not solids at room temperature
Poor conductors of heat and electricity
Covalent network solids
Formed by atoms that can form multiple covalent bonds

27. Solids (cont.) Ionic solids Metallic solids
High melting points and hardness
Strong but brittle
Metallic solids
Consist of positive metal ions surrounded by a sea of mobile electrons
Variable bond strength, malleable, ductile, excellent conductivity

28. Solids (cont.) Amorphous solids:
Amorphous solids are solids in which the particles are not arranged in a regular, repeating pattern, therefore no crystals.
Amorphous solids form when molten material cools too quickly to allow enough time for crystals to form.

29. End of Section 13.3

30. 7. Phase Changes That Require Energy
Section 13.4: Phase Changes
7. Phase Changes That Require Energy
Melting:
Occurs when heat flows into a solid object.
Heat is the transfer of energy from an object at a higher temperature to an object at a lower temperature.
When ice is heated, the energy is not used to raise the temperature of the ice, it is used to disrupt bonds
The ice eventually absorbs enough energy to break the hydrogen bonds that hold the water molecules together.
When the bonds break, the particles move apart and ice melts into water.

31. Phase Changes That Require Energy (cont.)
The melting point of a crystalline solid is the temperature at which the forces holding the crystal lattice together are broken and it becomes a liquid.
It is difficult to specify an exact melting point for an amorphous solid because they tend to act like liquids when they are still in the solid state

32. Phase Changes That Require Energy (cont.)
Vaporization:
A vapor is the gas phase of a substance that is ordinarily a liquid at room temperature
Vaporization is the process by which a liquid changes to a gas or vapor.
Evaporation is vaporization only at the surface of a liquid.
In a closed container, the pressure exerted by a vapor over a liquid is called vapor pressure.
Particles with enough energy escape from the liquid and enter the gas phase.

33. Phase Changes That Require Energy (cont.)
The boiling point is the temperature at which the vapor pressure of a liquid equals the atmospheric pressure.
At the boiling point, molecules throughout the liquid have enough energy to vaporize

34. Phase Changes That Require Energy (cont.)
Sublimation:
Sublimation is the process by which a solid changes into a gas without becoming a liquid.
If ice cubes are left in a freezer for a long time, they shrink because the ice sublimes
This property of ice is used in a process called freeze drying

35. 8. Phase Changes That Release Energy
Condensation:
As energy flows from water vapor, the velocity decreases which allows the molecules to form hydrogen bonds with other water molecules.
The formation of hydrogen bonds signals the change from the vapor phase to the liquid phase.
The process by which a gas or vapor becomes a liquid is called condensation.
The reverse of vaporization
When liquid forms energy is released.

36. Phase Changes That Release Energy (cont.)
Deposition:
Deposition is the process by which a gas or vapor changes directly to a solid, and is the reverse of sublimation.
As heat flows from water to the surroundings, the particles lose energy.
Freezing:
The freezing point is the temperature at which a liquid is converted into a crystalline solid.

37. 9. Phase Diagrams
Temperature and Pressure combine to control the phase of a substance.
A phase diagram is a graph of pressure versus temperature that shows in which phase a substance will exist under different conditions of temperature and pressure.
The triple point is the point on a phase diagram that represents the temperature and pressure at which all three phases of a substance can coexist.
The phase diagram for different substances are different from water.

38. End of Section 13.4

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60. CIM Table 12.4 Unit Cells Table 12.5 Types of Crystalline Solids
Figure Phase Diagrams

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