1. Chapter 13 States of Matter
Solids, Liquids and Gasses

2. The Kinetic Molecular Theory Basic Assumptions
Particle Size
Gas particles have no volume (pin point particles)
The space between particles is extremely large compared to the volume of the particles. Due to this distance, there is no significant attractive or repulsive force acting on the particles.

3. The Kinetic Molecular Theory Basic Assumptions
Particle Motion
Gas particles are in constant random motion.
Collisions between particles are elastic (Energy can be transferred from one particle to another during a collision, but no energy is lost when particles collide)

4. The Kinetic Molecular Theory Basic Assumptions

5. Explaining the Behavior of Gases
Kinetic-molecular theory explains the behavior of gases.
Low density
Remember D = m/v
Compression and expansion
Diffusion and effusion
Diffusion – the movement of one material
through another.
Effusion – a gas escapes through a tiny
opening.
Glencoe. Chemistry: Matter & Change

6. Explaining the Behavior of Gases
Glencoe. Chemistry: Matter & Change
Diffusion and Effusion (cont.)
Graham’s law of effusion
Graham’s law also applies to diffusion

7. Explaining the Behavior of Gases
EXAMPLE:
Ammonia has a molar mass of 17.0 g/mol; hydrogen chloride has a molar mass of 36.5 g/mol. What is the ratio of their diffusion rates?

8. Gas Pressure
The force that a gas exerts per unit area as a result of the simultaneous collisions of many particles
No particles = no pressure =
The Mercury Barometer
a vacuum
Invented by Evangelista Torricelli
Two forces affect the height of the mercury column = Gravity and Atmospheric Pressure
Equivalent pressure units
760 mm Hg = kPa = 1atm = 760 torr = 14.7 psi

9. Explaining the Behavior of Gases.
Dalton’s law of partial pressures
The total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture.
The portion of the total pressure contributed by a single gas is called its partial pressure.
Partial pressure depends on the number of moles of gas, the size of the container, and the temperature of the mixture (not the identity).
Ptotal = P1 + P2 + P3 + … Pn

10. Glencoe. Chemistry: Matter & Change. 2005.
Gas Pressure
Glencoe. Chemistry: Matter & Change
EXAMPLE:
A mixture of oxygen, carbon dioxide, and nitrogen has a total pressure of 0.97 atm. What is the partial pressure of oxygen if the partial pressure of carbon dioxide is 0.70 atm and the partial pressure of nitrogen is 0.12 atm?

11. Forces of Attraction Intramolecular Attraction
Glencoe. Chemistry: Matter & Change

12. Forces of Attraction Intermolecular Attraction
Dispersion forces – weak forces that result from temporary shifts in the density of electrons in electron clouds.
Dipole-dipole forces – attractions between oppositely charges regions of polar molecules.

13. Forces of Attraction Intermolecular Attraction
Glencoe. Chemistry: Matter & Change
Hydrogen bonds – a dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a small, highly electronegative atom with at least one lone electron pair (fluorine, oxygen, or nitrogen atom).

14. 13.3 The Nature of Liquids Particle Spacing
Intermolecular attractions reduce the amount of space between particles in a liquid.
Particle Motion
Particles in a liquid have enough kinetic energy to flow
The tendency for particles move and their attraction for one another account for the physical properties of liquids

15. The Nature of Liquids Viscosity Viscosity & Temperature
Viscosity is a measure of the resistance of a liquid to flow.
It is determined by the type of intermolecular forces involved, the shape of the particles, and the temperature.
Viscosity & Temperature
When temperature increases, the average kinetic energy of the particles increases.
The added energy makes it easier for the molecules to overcome the intermolecular forces that keep the molecules from flowing.
Therefore, when temperature ↑, viscosity ↓.

16. The Nature of Liquids Surface Tension
The energy required to increase the surface area of a liquid by a given amount is called surface tension.
It is a measure of the inward pull by particles in the interior.
Intermolecular forces do not have an equal effect on all particles in a liquid.
Compounds that lower the surface tension of water are called surface active agents or surfactants.

17. The Nature of Liquids Capillary Action
Cohesion describes the force of attraction between identical molecules.
Adhesion describes the force of attraction between molecules that are different.
Movement of water being drawn upward is called capillary action, or capillarity.

18. Vaporization vs. Evaporation
Vaporization is the conversion of a liquid to a gas
Evaporation is vaporization that occurs at the surface of a liquid that is not boiling.
Evaporation depends on the intermolecular forces that hold the particles in a liquid together.
If the forces are weak, then the kinetic energy of the particles at the surface can overcome the intermolecular forces that hold them together.
Adding heat will increase the rate of evaporation of a liquid

19. Vapor Pressure
Vapor pressure is a measure of the forced exerted by a gas over a liquid.
Vapor pressure is created in a closed system as particles in a liquid evaporate and collide with the walls of the container.
There is a direct relationship between temperature and vapor pressure.
Manometer

20. Boiling
Boiling occurs at the temperature when the vapor pressure of a liquid equals the pressure exerted by the atmosphere.
The boiling point of a liquid is when the liquid changes from a liquid to a gas
Not all substances boil at the same temperature because of intermolecular attractions.

21. The Nature of Solids
Particles are arranged in an orderly fashion with fixed locations within a solid.
Heat increases the kinetic energy of particles in a solid which causes the organization of the solid to break-down = Melting.
The melting point of a solid is the temperature at which a solid changes into a liquid.

22. The Structure of Solids
Most solids are crystalline
The unit cell is the smallest group of particles in a crystal that retain the geometric shape of the crystal
There are 7 crystal systems

23. Allotropes and Amorphous Solids
Allotropes are solid substances that can exist in more than one form in the same physical state.
Allotropes of Carbon Amorphous Solid
Amorphous solids lack an ordered internal structure

24. Heating Curve and Change of State

25. Endothermic Phase Changes
Melting- solid absorbs energy until particles have enough speed to break free of IM forces holding them in place
Vaporization-liquid absorbs energy until particles have enough speed to break free of IM forces holding them close together
Sublimation – Solids are converted directly to gases without forming a liquid

26. Exothermic Phase Changes
Freezing – liquid particles release energy and particles become highly organized
Condensation-gases lose energy and particles come close enough together to experience intermolecular forces
Deposition – Process by which a gas turns into a solid without the formation of a liquid

27. Phase Changes

28. Phase Diagrams
Phase diagrams show the temperature and pressure conditions at which a substance exists as a solid, liquid, or gas
Triple oint

29. Phase Diagrams Variables that control the phase of a substance are…
Temperature
Pressure
A phase diagram is a graph of pressure vs. temperature that shows in which phase a substance exists 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 three phases of a substance can coexist.
The point that indicates the critical pressure and temperature above which a substance cannot exist as a liquid is called the critical point.