States of Matter

I. Review: Phases of Matter

A. Solid

B. Liquid

C. Gas

II. Forces of Attraction

A. Intramolecular Forces

B. Intermolecular Forces

Short range forces between molecules

1. Hydrogen Bonding

-  A particularly ______force

-  Occurs in compounds in which a hydrogen atom is attached to ______, ______, or ______atom.

-  The hydrogen atom will act like a ______

-  The hydrogen atom will then be attracted to ______on a nearby molecule.

+ H O-

H+

+ H O-

H+

+ H O-

H+

-  Hydrogen bonding ______melting and boiling points because ______energy is required to break the forces between molecules.

2. Dipole-Dipole Forces

-  A ______force

-  Occurs in ______molecules (dipoles) because they have ______charge distribution (a positive end and a negative end)

-  ______end of one molecule is attracted to a ______end of a nearby molecule.

H+ - F --- H+ - F- --- H+ - F-

-  Dipole – dipole forces will also ______melting and boiling points.

-  A dipole can also temporarily attract electrons from another molecule causing an ______dipole.

3. London Dispersion Forces

-  A ______intermolecular force (a type of induced dipole).

-  Occurs in ______molecules and ______gases.

-  Electrons are constantly ______, therefore electron density ______.

-  This effect ______with increasing number of ______.

F2 Gas

Br2 Liquid

I2 Solid

III. Liquids and Solids

A. Liquids

1. Density and Compression

2. Fluidity

3. Viscosity

B. Solids

1. Density of Solids

2. Crystalline Solids

3. Network Covalent Solids

4. Metallic Solids

5. Amorphous Solid

IV. Gases and the Kinetic Molecular Theory

A. Kinetic Molecular Theory

B. Assumptions:

1. Particle Size

2. Particle Motion

a.  Collisions

1.  elastic

2.  inelastic

b. Constant, rapid, random motion

In all directions

c. No forces of attraction or repulsion between gas particles

Therefore, do not stick together

3. Particle Energy

Depends on temperature

Increase Temp, Increase speed & KE of gas particles

Decrease Temp, Decrease speed & KE of gas particles

C.  Temperature

1.  Definition: A measure of the average KE of the particles in an object

2. Three scales:

a. Celsius – Anders Celsius, Swedish astronomer – Based on when water freezes and boils. (Δ 100o)

b. Kelvin – Lord Kelvin (William Thomson)

S.I. unit – Absolute zero – all motion stops (Δ 100o)

c. Fahrenheit – Gabriel Fahrenheit

Alcohol thermometer in 1709

Hg in 1714 (Δ 180o)

3. Conversions

K = oC + 273

oF = 1.8(oC)+32

oC = (oF-32)/1.8

4. Practice:

a. 28 o C = ______K

b. 200 K = ______o C

c. – 15 o C = ______o F

d.  10 K = ______o F

D.  Physical Properties (of gases)

1. Density - low density à typically a small amout of mass per unit of volume

2.  Fluidity

-Gases are more fluid than liquids

3.  Compression

Reduce the volume by squeezing the molecules closer together

4.  Expansion

-Increase volume, molecules move further apart

5.  Diffusion

Movement of molecules. Mixingof gases – (high à low concentrations)

6.  Effusion

Movement of molecules through a small opening (ex: puncture a tire or balloon)

E. Graham’s Law of Diffusion/Effusion

Velocity A √ Mass B

= at same T and P

Velocity B √ Mass A

Essentially, _Smaller__ molecules move _Faster_

than _larger_ ones.

Sample: Which gas will diffuse the fastest: ammonia, NH3, or hydrogen chloride, HCl?

F. Gas Pressure

1. Definition of Pressure:

Gas molecules hitting container walls

2. Mathematical Formula:

P = Force/area = N/m2 = Pascal (Pa)

3. Air Pressure Conversions-Conversion Factors

760 mmHg (torr) = 1 atm = 101.3 kPa = 30.0 in Hg = 1 torr = .133 kPa

Sample: Convert 0.70 atm to mmHg.

Practice:

Convert 2.0 atm to torr (mmHg).

Convert 725 torr to atm.

3. Atmospheric (air) Pressure

The mixture of gases surrounding the Earth exerting pressure (78%N2, 21%O2, Ar, CO2…)

4. Measuring Atmospheric Pressure

a. Hg Barometer = glass tube sealed at one end with the other end immersed in a container of Hg

5. Measuring Pressure of Gases

a. Closed Manometer

Instrument used to measure pressure of a gas in a closed container

b. Open Manometer

Open to air so that atmospheric pressure must be taken into account

5. Standard Pressure (STP = standard temperature and pressure (1 atm and 0oC)

1 Atm = 760 torr (mmHg)

6. Air Pressure Varies with Altitude

Higher altitudes = lower pressure

7. Dalton’s Law of Partial Pressures

Ptotal = P1 + P2 + P3 …….

The sum of all the individual pressures = the total pressure

Practice:

1. Determine the total pressure, in mm Hg, for a mixture that contains four gases with partial pressures of 5.0 atm, 4.56 atm, 3.02 atm, and 1.2 atm.

2. What is the partial pressure of hydrogen gas in a mixture of hydrogen and helium if the total pressure is 600 mmHg and the partial pressure of helium is 439 mmHg?

V. Phase Changes

A. Phase changes that require energy (endothermic)

1.  Melting (fusion)

Solid to liquid

2. Vaporization (evaporation)

Liquid to gas

3. Sublimation

Solid to gas

B. Phase changes that release energy (exothermic)

1. Condensation

gas to liquid

2. Deposition

gas to solid

3. Freezing

liquid to solid

Heating Curve

Diagonal lines (1 state of matter) = inc in temp, inc in KE

Horizontal lines (2 states of matter) = Phase change

No change in temp, therefore, no change in KE