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Liquids and Solids Chapter 13.

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Presentation on theme: "Liquids and Solids Chapter 13."— Presentation transcript:

1 Liquids and Solids Chapter 13

2 I. Comparison of Solids, Liquids and Gases
Shape Molecular Forces Density Compressibility Fluidity Rate of Diffusion

3 I. Comparison of Solids, Liquids and Gases
Shape Intermolecular Forces Density Compressibility Fluidity Rate of Diffusion definite geometric shapes shape of container fill the container very strong moderate weak most dense (except H2O) relatively high least dense relatively incompressible very compressible incompressible none yes, it flows yes, they flow yes, they diffuse very well yes, but very slowly yes, but slower than gases

4 II. Changes of State gas solid liquid

5 II. Changes of State most energy gas least energy solid liquid
deposition evaporation moderate energy least energy sublimation condensation melting solid liquid freezing Question: Which phase changes ADD energy, and which phase changes RELEASE energy?

6 III. Properties of Liquids
A. Surface Tension - pulling together of a liquids SURFACE molecules, creating tighter arrangement B. Vapor Pressure - the gaseous pressure resulting from the vaporization of SURFACE molecules of liquid C. Boiling Point - the temperature at which a liquids vapor pressure equals the atmospheric pressure D. Viscosity – resistance to flow Note: 1) Vaporization occurs at the SURFACE of liquids 2) Boiling occurs WITHIN, as well as the surface of a liquid 3) Boiling Point temperatures can be raised or lowered by a corresponding change to the atmospheric pressure

7 Surface Tension pulling together of a liquids SURFACE molecules, creating tighter arrangement

8 Dipole-Dipole Interaction

9 Hydrogen Bonding

10 IV. Water and It’s Properties
A. Intermolecular hydrogen bonding is responsible for the high boiling point of water: O H note: the H-bondings are additional bonds -- weaker than covalent bonds -- that must be broken before boiling can occur O H O H O H

11 B. When water freezes, the water molecules arrange themselves into hexagonal patterns:
notes: 1) the wider spacing between molecules causes ice to be LESS DENSE than liquid water 2) water is most dense at 4 oC… warmer and colder than this causes expansion 3) the volume of the hexagonal spaces in ice is equal to the volume of ice that is above the water line when floating O H O H O H O H O H

12 C. Energy requirements for phase changes of water:
1) water requires 1 calorie (4.184 joules) to heat 1 gram of water 1 oC. 2) ice requires 80 calories per gram (6.02 kJ/mol)to melt without any temperature change 3) water requires 540 calories per gram (40.8 kJ/mol) to boil without any temperature change Ex: How many calories are required to melt 50 grams of ice, heat it to 100 oC, then boil it?

13 V. Phase Diagrams A. Conveniently shows the relationship between temperature, pressure and the physical state of a substance. B. Water has the only phase diagram with a negative slope to the solid/liquid boundary 1) as pressure increases on ice, it liquifies 2) higher pressures cause smaller volumes … 3) … for water, liquid has smaller volume than solid C. Triple Point – when all 3 phases coexist D. Critical Point – 1) temp at which gas can no longer become liquid no matter what the pressure is 2) pressure required to produce liquefaction at the critical temperature

14 Phase Diagram for Water

15 Phase Diagram for CO2

16 VI. Types of Solids A. Crystalline Solids - 1) ionic crystals - metal + nonmetal ex: NaCl; Fe2O3 2) Covalent Network - covalently bonded atoms or molecules forming a large crystal ex: C (diamond); SiO2 (sand) 3) Metallic - metal atoms sharing valence e- ex: Fe; Cu; 4) Covalent Molecular - intermolecular forces between covalent solids ex: ice; frozen gases (dry ice), S8, P4, I2 B. Amorphous Solids - non regularly shaped solids ex: plastics; glass

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