1. The Boring States of Matter CH11

2. Kinetic Energy E K The energy an object has because of its motion. Temperature is a measurement of average kinetic energy. Kinetic Theory- Tiny particles in all forms of matter are in constant motion.

3. Liquids vs. Gases Liquids, the molecules are moving and touching. They interact They take up less space than gasses They overall have less E k than a gas of the same substance Gases, the molecules bounce off each other, but do not stay in contact There is very little interaction between molecules They take up a lot of space

4. Which of these 2 parts are compressible? Can you “squish” the gas? Can you “squish” the liquid? gas liquid

5. Evaporation: conversion of a liquid to a gas At the surface, there are a few molecules that have enough E k to escape to gas. By increasing the temperature, more molecules will have the necessary E k and evaporation will occur faster than before. The vapor pressure increases with more heat/ Ek

6. Average kinetic energy Temperature is a measurement of average kinetic energy. If a beaker of water reads 20°C, do all molecules in the beaker have kinetic energy = 20 ⁰ C? There is a broad range of kinetic energies. Most of the molecules are “around 20C”, but there are some significantly less energetic, and some significantly more energetic.

7. Solids Molecules in solids only vibrate in place, they do not slide past each other They interact in their FIXED position More dense than gas, and most liquids Salt crystal, atoms are in a set position

8. Heating a solid The vibrations increase If they vibrate enough, some of the bonds holding the solid together will break. This is called the melting point. S  L melting L  S freezing L  G vaporization (or evaporation) G  L condensation

9. Evaporation vs. Boiling point Atm is pressing down on the surface

10. Solid to a Gas?…Sublimation Solids also have a vapor pressure. When vapor pressure is high enough, the solid will go to gas, without stopping at liquid. This is called sublimation. Gas to a solid – DEPOSITION Dry ice, is a classic example of sublimation

11. Phase Transition names SLSL LSLS L  G GLGL S  G G  S d Melting Freezing Vaporization Condensation Sublimation deposition Give the change of state for each term

12. Ionic Solids Strong forces between oppositely charged ions. HIGH melting points HIGH boiling points Non conductors as solids, conductors while molten. Often water soluble  Depends upon attractive forces for each other and other molecules.

13. Molecular Combination of 2 or more non metals Molecular substances have 3 important types of weak intermolecular forces 1. Dispersion Forces 2. Dipole Forces 3. Hydrogen bonds Inter-molecular (between molecule) forces are weak. Therefore easy to separate: Have low melting and boiling points  Think Oxygen and Water  They’re gases and liquids at room temperature Non-Conductors of electricity

14. Dispersion (London) Forces Most common type of intermolecular force. Caused by temporary induced- dipoles formed in adjacent molecules. All molecules have dispersion forces, the strength depends on 2 factors:  *The # of electrons in the molecule As molar mass increases, dispersion forces become stronger, the boiling pt of non-polar molecules increases. Think of the electron cloud being agitated F 2 -188 Cl 2 -34 Br 2 59 I 2 184

15. Dipole Forces Electrically attractive forces between + and – end of adjacent polar molecules. Boiling points of  N 2 -196C  O 2 -183C NO -151C The Nitrogen Monoxide is slightly polar and therefore has weak dipole forces. This explains the relatively higher melting points.

16. Hydrogen Bonding Unusually strong DIPOLE forces. This is due to the very small Hydrogen atom’s Electronegativity difference with: F fluorine O oxygen N nitrogen The strongest of the ‘weak forces’ Water H 2 O bp = 100 C, H 2 S -61C

17. E k = ½ x mass x velocity 2

18. Kinetic Theory and Gases A gas is composed of particles  Small spheres…no real volume for each particle The gas particles move in constant random motion  Very fast ( 1000 mph), straight line until collision occurs All collisions are perfectly elastic  No Ek is lost in the collision

19. Gas Pressure Moving bodies exert force when they collide with other bodies. The result of the billions of particles colliding is gas pressure

20. Atmospheric pressure (Atm) Gravity holds air molecules in the atmosphere Atm is the result air colliding with objects The average pressure under all that air is 1 atmosphere. Pressure= Force Area

21. Air Pressure Barometer created by Evangelista Torricelli in 1646 Inverted a tube filled with mercury into a dish until the force of the Hg inside the tube balanced the force of the atmosphere on the surface of the liquid outside the tube

22. ATM and pressure The average height of the column at sea level was 760millimeters of Mercury or 760mmHg We later gave this average pressure the name TORR, for Torricelli

23. 3 Ways to express measures of pressure Millimeters of Mercury (or Torr)  mmHg, Torr Atmospheric Pressure Units  Atm Pascal (or KiloPascals) we don’t use this unit  Pa, or kPa 760 mmHg = 1.0 atm = 760 Torr

24. Convert the following pressure measurments: 720 Torr to atm.1.45 atm. to mmHg

25. Putting it all together How can you know if a substance will melt, or sublimate? The transitions depend upon both pressure and temperature. At a given Temperature T:  At a low pressure, something might sublimate, at a higher pressure it would likely melt

26. Phase diagram solid liquid gas Pressure temperature Triple point, all 3 phases exist simultaneously

27. Phase Diagram for Water Triple point.006 atm.001C

28. 4 th state of matter: Plasma Plasma Occurs at super hot temperatures Gas atoms are stripped of their electrons Mix of loose electrons, and + gas ions is called PLASMA Hot plasmas make up the stars, and can be 10 million degrees Not very relevant to HS Chemistry.

29. Maxwell-Boltzman Distrubution At high T More molecules are in a higher energy state