Liquids and Solids

1. Liquids and Solids H2O (g) H2O (s) H2O ()

2. Three States of Matter The state (or phase) of matter is determined by the arrangement and motion of particles. The motion of particles is governed by the kinetic energy (KE) of the particles (Remember that KE = 1/2mv2) Solids Liquids Gases Increase KE Increase KE

3. Changes of State sublimation sublimation boiling melting vaporization condensation freezing deposition

4. Three States of Matter, cont.

5. Attractive Forces In Molecules Intermolecular Forces: Attractive forces between molecules Types of inter-molecular forces • dipole-dipole (1% as strong as covalent bonds) A special type of dipole-dipole force is the hydrogen bond. These bonds form between molecules that contain a hydrogen atoms bonded to a very electronegative element like N, O or F. Hydrogen bonds are very strong compared to an ordinary dipole-dipole bond. E.g HF, NH3, H2O all form hydrogen bonds Hydrogen bonding 10% as strong as covalent bonds

6. Water molecules are polar molecules. The  - oxygen forms intermolecular bonds with the  + hydrogen of another water molecules. This is an example of a special type of intermolecular bond called a hydrogen bond. Inter-molecular forces

7. 2. London dispersion forces (instantaneous and induced dipoles) • NON-POLAR MOLECULES

8. Non-polar molecule This instantaneous dipole will effect any nearby molecules Movement of electrons causes an instantaneous dipole This induces a dipole in a nearby molecule

9. Properties of Liquids As we consider the properties of liquids (and solids) that KE and intermolecular attractions are governing the behavior of the substance

10. Properties of Liquids Vaporization: Change from liquid to gas via boiling process Evaporation: Change from liquid to gas at the surface of a liquid, not caused by boiling --This happens because the molecules at the top a of the liquid don’t have as strong of an attraction to the other molecules. (If they have high enough KE, they can escape)

11. Open Containers: Evaporation causes liquid molecules to leave as gases and escape (amount of liquid decreases) Closed Containers: Evaporation causes liquid molecules to vaporize, but they get caught in the container, creating : VAPOR PRESSURE Evaporation: A Closer Look

12. Vapor Pressure This is known as Dynamic Equilibrium because the rates of evaporation and condensation are EQUAL . In a sealed container, molecules will start to evaporate and the liquid’s volume will decrease. Evaporation and Condensation DO NOT stop happening once the flask has reached equilibrium But, after the air above the liquid becomes “saturated”, some of these molecules will then condense. After a short time, the volume of the liquid will not change. The rate of evaporation = the rate of condensation

13. Properties of Liquids, cont. Boiling: When all the molecules of a liquid have enough kinetic energy to vaporize, the liquid is said to be boiling. Boiling Point (bp): The temperature at which the vapor pressure of a liquid is just equal to the external pressure on the liquid. Normal Boiling Point: The temperature at which a substance boils at atmospheric pressure (101.3kPa)

14. Boiling point and Vapor pressure When water is heated, the kinetic energy of the molecules increases and eventually bubbles of vapor form within it. The vapor pressure in the bubble is the same as the vapor pressure of the water at that temperature. When the temperature of the water reaches a point that the vapor pressure of the bubbleequals atmospheric pressure, the bubbles get larger, rise to the surface, and escape as steam. The water begins to boil.

15. Boiling point and Vapor pressure REMEMBER: vapor pressure of the bubbleequals atmospheric pressure The water begins to boil atmospheric pressure 450mm Hg 58o C At lower atmospheric pressures, the kinetic energy does not have to be as high to make the vapor pressure in the bubble equal to atmospheric pressure. 450 mm Hg By reducing the atmospheric pressure, The water begins to boil at a lower temperature.

16. Other Properties of Liquids 1. Why, when you pour a liquid onto a surface does it form droplets? 2. Why do some liquids exhibit capillaryaction? Hg H2O 3. Why are some liquids more viscous than others?

17. Viscosity:is the resistance to motion of a liquid. Maple syrup is more viscous than water. But water is much more viscous than gasoline or alcohol. The stronger the attraction between molecules of a liquid, the greater itsresistance to flow and so the more viscous it is. Consider the following substances a) molasses b) water c) ethyl alcohol • Which is the least viscous? • Which substance has the strongest intermolecular attractions? Ethyl alcohol Molasses

18. Surface tension This water strider uses surface tension to his advantage The inward force or pull which tends to minimize the surface area of any liquid is surface tension.

19. Surface tension occurs because the molecules on the surface of the liquid cannot bond to the outside molecules. As a result, they look for something else to bond to (in order to increase stability). They get “pulled” in towards each other until their surface area is minimized, thus minimizing the contact with the outside. Surface tension in water is caused by hydrogen bonding between polar molecules. The more polar a liquid the stronger its surface tension. Hg pure H2O H2O with detergent **Surfactants are compounds that reduce the surface tension of a liquid (soaps and detergents are examples)** The smallest surface area a liquid can form is a sphere.

20. Capillary action is the spontaneous rising of a liquid in a narrow tube. Two forces are responsible for this action: Cohesive forces:the intermolecular forces between molecules of the liquid Adhesive forces: the attractive forces between the liquid molecules and their container If the container is made of a substance that has polar bonds then a polar liquid will be attracted to the container. This is why water forms a concave meniscus while mercury forms convex meniscus Hg H2O

21. Properties of Solids Solids generally have an orderly arrangement of atoms Melting: When the kinetic energy of all the atoms in a solid is increased to a point where the atoms are able to freely flow around one another, the solid is said to have melted Melting Point (mp): The temperature at which a solid turns into a liquid.

22. Crystalline -Most solids are crystalline -Contain particles arranged in an orderly, repeating, 3-D pattern called a crystal lattice Non-Crystalline - Amorphous solids have no set crystal structure Examples: 1. Glass 2. Asphalt 3. Rubber 4. Plastic 5. Candles (Wax) Types of Solids

23. Allotropes Some crystalline solids (pure substances) occur in a variety of different forms, known as Allotropes. Each allotrope has a different crystalline pattern that connects the atoms of the solid Carbon, Sulfur, Phosphorus, Oxygen, Boron and Antimony all have allotropes The most common examples of allotropes are found in elemental Carbon: Diamond Graphite Buckminsterfullerene (Bucky Balls)

24. ICE Solids are almost always more dense than their liquid forms, however, there is one exception: ICE Ice molecules are locked in fixed positions, held by intermolecular-bonds. Ice is less dense than liquid water because the molecules are further apart than in liquid water.

25. Tracking Changes in State sublimation sublimation boiling melting vaporization condensation freezing deposition

26. Tracking Changes in State Melting/Vaporization Curves: tracks temperature changes as a function of time and shows all state changes boiling

27. Tracking Changes in State Phase Diagrams: Shows the various conditions at which each state of a substance can occur boiling