Kinetic Theory

1. Kinetic Theory Solids, Liquids and Gases

2. Kinetic Theory • Most compounds exist in only one state at a time – • Water can be • Ice (solid) • Water (liquid) • Steam (gas)

3. Kinetic Theory • However, the state of existence is easily changed through • Changes in temperature • Cool water down and it freezes into ice • Heat water up and it boils into steam

4. Kinetic Theory • The word kinetic means • Motion • Kinetic energy is • The energy held by an object because of its motion

5. Kinetic Theory • Kinetic theory is • The concept that the particles in all matter are in constant motion.

6. Kinetic Theory • For gases, the following assumptions have been applied: • The particles of a gas are small, hard spheres with insignificant volume and large distances between one another. These particles are not attracted to one another; nor are they repulsed from one another.

7. Kinetic Theory • For gases, the following assumptions have been applied: • Each particle in a sample of a gas is in constant, straight-line motion. They will only change direction when they run into something, like another particle of gas or the container holding them. This straight-line motion is called a “random walk.”

8. Kinetic Theory • For gases, the following assumptions have been applied: • Energy is not lost when the particles hit each other. This means all of the collisions between the particles are called perfectly elastic. Overall, this mean the gas does not lose the energy it has because it is always running into itself… it doesn’t wear itself out.

9. Gas Pressure • Everything that is moving will • Exert a force when it collides with another object • Imagine getting hit by a car, a bicycle and a rubber ball that are all moving at the same speed • Which one hits you the hardest? • Which one hits you the softest? • Why? car rubber ball More mass = harder hit

10. Gas Pressure • Well, gas particles exert force too. • The force applied by one gas particle is extremely tiny • A container of gas will have lots of particles, so billions of gas particles will run into the container simultaneously. • All of these forces hitting at the same time add up to a measurable amount of force.

11. Gas Pressure • Gas pressure is • The sum of the forces applied by all the gas particles divided by the area over which the forces are applied. • This is a measurement of how “hard” the gas is hitting the container. • How could we increase the gas pressure? • More particles, make them move faster, smaller area, etc.

12. Gas Pressure • What if there were no gas particles present to run into anything, like out in space? What would the pressure of that sort of system be? • The pressure would be zero. • This is called a vacuum.

13. Gas Pressure • In this class, we will express pressure using the units of the SI system, Pascals (Pa); and an older unit called the atmosphere. • We will sometimes see other units also; just convert them to one that is easy to use. • 101325 Pa = 1 atm = 14.7 psi = 760 mmHg

14. Gas Pressure • As a review, what does STP stand for? • Standard temperature and pressure • What is STP • 1 atm and 0oC • What is that temperature in Kelvin? • 273.15 K

15. Kinetic Energy and Kelvin Temperature • When you heat something up, a little of the energy is stored as potential energy while the rest speeds the particles up. • That is, it increases the kinetic energy of the particles in the matter. • The particles will NOT all have the same kinetic energy • Some (about half) are moving faster than average • Some (the other half) are moving slower than average

16. Kinetic Energy and Kelvin Temperature • Theoretically, if you cooled something down enough all molecular motion would stop. This would happen at • Absolute zero (0 Kelvin) • We’ve never been able to cool anything down to absolute zero in a lab

17. Kinetic Energy and Kelvin Temperature • In the Kelvin scale, temperature and kinetic energy are directly related to each other. What does this mean?

18. The Nature of Liquids • Objectives: • Describe the nature of a liquid in terms of the attractive forces between the particles • Differentiate between evaporation and boiling of a liquid, using kinetic theory

19. The Nature of Liquids • What makes a liquid different from a solid?

20. A Model for Liquids • Does the kinetic theory apply to liquids? • Yes – the particles in a liquid are in motion. • The particles move around more slowly than gas particles. • The particles are attracted to each other. • These attractions are called intermolecular forces. • These forces are what cause liquids to have a defined volume. • The forces are too weak to lock the particles into place, which is why liquids flow.

21. Evaporation • What happens to a saucer of water left sitting out for a long time? • The water turns from a liquid into a gas – it evaporates

22. Evaporation • Vaporization: • The conversion of a liquid to a gas or vapor. • Evaporation: • Vaporization that occurs at the surface of a liquid that is not boiling • Faster than average liquid particles at the surface of the liquid can “escape” into the gas state • Cools a liquid down (the hot particles leave)

23. Evaporation • Vapor pressure: • The pressure produced when vapor particles above a liquid in a closed container collide with the container walls • A dynamic equilibrium exists between the vapor and the liquid.

24. Evaporation • What does dynamic mean? • In motion • What is the opposite of dynamic? • Static (still) • What does equilibrium mean? • A state of balance or equality between opposing forces

25. Evaporation • How can an equilibrium be dynamic? • Everything going one direction has an opposite that goes the other way • In a dynamic equilibrium, the two opposing sides are equal – essentially canceling each other out • For evaporation, a dynamic equilibrium means the rate of particles going from liquid to gas is equal to the rate of particles going from gas to liquid

26. Boiling Point • How do you make a liquid evaporate faster? • Heat it up – hotter particles move faster and are more likely to escape into the gas state

27. Boiling Point • Boiling point: • The temperature at which the vapor pressure of a liquid is just equal to the external pressure on the liquid

28. Boiling Point • Boiling: • Vaporization throughout a liquid • Bubbles of vapor form all over the sample • Normal boiling point: • Boiling point at a pressure of 1 atm (air pressure at sea level) • While boiling the temperature will remain constant • The addition of more heat will vaporize more liquid rather than raise the temperature

29. The Nature of Solids • Objectives: • Describe how the degree of organization of particles distinguishes solids from liquids and gases • Distinguish between a crystal lattice and a unit cell

30. Review • Name three molecular compounds • Carbon dioxide, water, sulfur trioxide • Name three ionic compounds • Sodium chloride, magnesium hydroxide, calcium carbonate • How are molecular and ionic compounds different? • Molecular compounds exist as molecules. Ionic compounds are huge collections of ions.

31. A Model for Solids • Do the particles in solids move? • Yes, but with much less freedom than the particles in liquids. • They vibrate around a fixed point. • These particles are locked into position, so they can’t flow – solids have both a definite volume and a definite shape.

32. A Model for Solids • Melting point: • The temperature at which a solid turns into a liquid. • Also the temperature at which a liquid turns into a solid • Melting point = freezing point

33. A Model for Solids • Melting is a result of adding heat to an object. • More heat = higher temperature = increased kinetic energy of the particles = faster movement of the particles • At the melting point, the particles are moving so fast that they are able to break free of the intermolecular forces holding them together. • They begin to slide past one another as in a liquid

34. A Model for Solids • Why are the melting points of ionic compounds higher than molecular compounds? • The ionic compounds have stronger forces holding them together. • Stronger forces = higher melting point • Can all compounds be melted? • No – some, like wood and various plastics, will decompose when heated. They never melt.

35. Crystal Structure and Unit Cells • Many solids are crystalline in nature. • The atoms, ions or molecules of the solid are arranged in an orderly, repeating, 3D pattern called the crystal lattice. • Crystals are classified into seven systems. • You don’t have to learn their names.

36. Crystal Structure and Unit Cells • The smallest group of particles within a crystal that retains the shape of the crystal is called • The unit cell

37. Crystal Structure and Unit Cells • Three types of unit cells • Simple cubic • Body-centered cubic • Face-centered cubic

38. Crystal Structure and Unit Cells • An allotrope is • A different form of an element. • For example, carbon has 3 allotropes: • Graphite • Coal • Diamond

39. Crystal Structure and Unit Cells • Not all solids are crystalline – • Some lack the organized structure found in crystalline solids • These unorganized solids are called amorphous solids • Amorphous means “without shape”

40. Crystal Structure and Unit Cells • There are two key difference between crystalline and amorphous solids: • Crystalline solids have a sharply defined melting point. Amorphous solids don’t – they will have temperature range over which they soften and then melt. • When struck, crystals will fragment into pieces with the same shape as the original. Amorphous solids will break into pieces with sharp edges (like glass) that are shaped randomly.