1. States of Matter By: Ms. Buroker

2. Let’s Review …. Shall We?

3. Kinetic Theory 1.) All matter is made of atoms and molecules that act like tiny particles. 2.) These tiny particles are always in motion. The higher the temperature of the substance, the faster the particles move. 3.) At the same temperature, more- massive (heavier) particles move slower than less massive (lighter) particles. What does this suggest about solid particles vs. gas particles???

4. ) State of Matter: Solids Solids have … 1.) low KE - particles vibrate but can’t move around 2.) definite shape & volume 3.) crystalline - repeating geometric pattern 4.) amorphous - no pattern (e.g. glass, wax

5. State of Matter: Liquids Liquids have … 1.) higher KE - particles can move around but are still close together 2.) indefinite shape 3.) definite volume All liquids are considered fluids because their particles can move past one another!

6. State of Matter: Gases Gases have … 1.) high KE - particles can separate and move throughout container 2.) indefinite shape & volume All gases are considered fluids because their particles can move past one another!

7. State of Matter: Plasma Plasma has … 1.) very high KE - particles collide with enough energy to break into charged particles (+/-) 2.) gas-like, indefinite shape & volume 3.) stars, fluorescent light bulbs, TV tubes

8. Thermal Expansion Most matter expands when heated & contracts when cooled.  Temp causes  KE. Particles collide with more force & spread out.

9. Let’s Re-Visit The Kinetic Molecular Theory We make several assumptions when we talk about the kinetic molecular theory for gases: Particle Size: Gases are composed of mostly empty space, which means the volume of the particles is small compared to the volume of the empty space. Particle Motion: Gases are in constant, random motion. Gases have elastic collisions which means they loose NO kinetic energy when the bump into each other. Particle Energy: KE = 1/2mv 2 Temperature is a measure of the average kinetic energy of the particles in a sample of matter.

10. Gases Gases are Expandable and Compressible! Gases have Low Densities! Density is mass per unit volume. Gases can diffuse and effuse which describes the movement of one material through another! Think Perfume

11. Pressure Which shoes create the most pressure?

12. Key Units at Sea Level: 101.325 kPa (kilopascal) 1 atm 760 mm Hg 14.7 psi Pressure The pressure gases exert comes from them hitting the sides of the container they are in.

13. Barometer Atmospheric Pressure Manometer Contained Pressure Pressure

14. Dalton’s Law of Partial Pressures Dalton’s Law of Partial Pressures states that each gas present in a mixture exhibits its own individual pressure so that the sum of the total mixture of gases is equal to the sum of the individual pressures. P T = P 1 + P 2 + P 3 + …..

15. Let’s Try an Example Problem … A mixture of oxygen, carbon dioxide, and nitrogen has a total pressure of 0.97atm. What is the partial pressure of O 2, if the partial pressure of CO 2 is 0.70atm and the partial pressure of N 2 is 0.12atm?

16. Let Me Pose a Question ….. If all particles of matter at room temperature have the same kinetic energy, why then do we see some substances as solids, some as liquids, and some as gases?

17. The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together. Intramolecular Intermolecular Forces

18. It’s all about strength!!! The stronger the force … the stronger the bond … and the stronger the bond … the more closely together the molecules will be packed … the closer they are packed, the more likely you are to be a solid.

19. London Dispersion Forces Dispersion Forces are weak forces that result from temporary shifts in the density of electrons in electron clouds.

20. Dipole- Dipole Interactions Molecules that have permanent dipoles are attracted to each other. * The positive end of one is attracted to the negative end of the other and vice versa. * These forces are only important when the molecules are close to each other.

21. Hydrogen Bonding The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong. * We call these interactions hydrogen bonds.

22. Hydrogen Bonding Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

23. Name That Bond! Please Name the Intermolecular Force at work amongst the following molecules… H 2 NH 3 HCl HF Dispersion Forces Hydrogen Bonding Dipole- Dipole Hydrogen Bonding

24. Intermolecular Forces Affect Many Physical Properties The strength of the attractions between particles can greatly affect the properties of a substance or solution.

25. Viscosity Resistance of a liquid to flow is called viscosity. It is related to the ease with which molecules can move past each other. Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

26. Surface Tension Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

27. Energy’s Role Officially- energy is the ability to do work, but you can also think of it as the ability to change or move matter. Think about it … matter is made of atoms … and the state the matter is in depends on how fast the particles are moving … so energy and the state of matter are directly related. So, how does matter change its state?

28. Thermal Energy The total kinetic energy of the particles that make up an object. High Kinetic Energy = High Thermal Energy Note! Thermal energy also depends on the amount of substance …

29. Temperature Temperature is simply a measurement of a substance’s average kinetic energy!! We usually think of temperature as a measure of how hot or cold something is … can you come up with a different definition???

30. Changes of State When matter changes the state it’s in … this is an example of a physical change! Since the state of matter is directly related the amount of energy it has … then for the state to change- it stands to reason that the energy must change as well.

31. Changes of State Sublimation Melting Freezing Condensation Evaporation Some changes of state require energy to happen … while other require the removal of energy.

32. Phase Changes

33. Phase Change Diagram