Liquids and Solids Chapter 13. I. Comparison of Solids, Liquids and Gases SolidsLiquidsGases Shape Molecular Forces Density Compressibility Fluidity Rate.

Liquids and Solids

Chapter 13

I. Comparison of Solids, Liquids and Gases

Solids Liquids Gases

Shape

MolecularForces

Density

Compressibility

Fluidity

Rate ofDiffusion

II. Changes of State

solid liquid

gas

sublimation

melting

freezing

condensation

evaporationdeposition

Question: Which phase changes ADD energy, and which phase changes RELEASE energy?

most energy

least energy

moderate energy

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

Surface Tension pulling together of a liquids SURFACE molecules, creating

tighter arrangement

Dipole-Dipole Interaction

Hydrogen Bonding

IV. Water and It’s Properties A. Intermolecular hydrogen bonding is responsible for the high boiling point of water:

O

H H O

H H

O

H H

O

H H

note: the H-bondings are additional bonds -- weaker than covalent bonds -- that must be broken before boiling can occur

B. When water freezes, the water molecules arrange themselves into hexagonal patterns:

O

H

H

O

H

H

O

H

H

O

HH

O

H

H O

H H

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

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?

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

Phase Diagram for Water

Phase Diagram for CO2

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