Reading Assignment 1:  KTU & Read Ch. 13.1 & 14.1- answer questions 13.1: 1-9 & 14.1: 1-6

I. Kinetic Theory & Nature of Gases  

Kinetic Theory -- Particles of any matter are in motion, possessing Kinetic Energy  

K. E. = 1/2 mv²    therefore  S K.E. = S (1/2 mv² )_n where n = # of particles  

A. Kinetic Theory of Gases-  Image

1.     A gas consists of large numbers of molecules in constant random motion.

Most gases are molecules with the exception of the noble gases whose representative particle is the atom.

2.     Gas molecules influence each other only by collision; they exert no other forces (attractive or repulsive) on each other.

3.     All collisions are perfectly elastic- all kinetic energy is conserved.

4. The volume actually occupied by the molecules of a gas is negligibly small; the vast majority of the volume of the gas is empty space through which the gas molecules are moving.

Introduction: The Four Gas Law Variables: Temperature, Pressure, Volume, and Moles

B. Kinetic Energy & Temperature  

1. Temperature is proportional to the average kinetic energy of the particles  

a. increasing temperature increases the velocity of the particles- increase in K.E.  

in-class Applet:  Changes in Kinetic Energy due to Temperature changes.

resource.  Maxwell-Boltzmann distribution of particles

Temperature is measured in units of Celsius, Fahrenheit or Kelvin

iamge: Comparison of Units of Temperature.

on-your-own Practice. Converting Temperatures

b. particles of different masses have different velocities at same temperatures. See Graham's Law.

2. Heat is the total kinetic energy of a mass- dependent upon mass (extensive physical property)

measured in calories or Joules  

1 cal  = 4.184 J  

3. Absolute Zero- temperature where S K.E. = 0.

Kelvin Scale (Lord Kelvin)- Establishes 0 at Absolute Zero.  = -273.15 ^oC  

Extension Question:  What happens to atoms at very low temperatures?  Bose-Einstein Condensation. 

This is part of the Physics 2000 series and will take some time, but it is real cool......

C. Pressure - the sum of all impact forces of the gas particles  

1. Vacuum- no particles → no pressure  

2. Atmospheric Pressure- pressure due to molecules in the atmosphere.

-air pressure decreases with elevation.  Why?  

a. Barometers- used to measure atmospheric pressure. Image

3. Pascal (Pa)- S.I. unit of pressure measure pressure.  

atmospheric pressure @ sea-level = 101.3 kPa

other measurements of pressure

1. mm Hg- used in mercury barometers  

1 atm = 760 mm Hg = 101.325 kPa = 101,325 Pascals = 1013 mb = 29.92 in Hg = 14.7 lbs/in²

handout: Conversion Table for common pressures.

3.     torr- (Evangelista Torricelli)  1 torr = 1 mm Hg  

STP- "Standard Temperature and Pressure" – 0 ^oC (273 K) and 1 atm (101.3 kPa)

Application: Scuba Diving Explained. Nitrogen necrosis, CO₂ & CO toxicities

D. Avogadros Hypothesis  

Equal volumes of gases at same temperature & pressure contain the same number of particles.  

Molar Volume = 22.4 L/ mole @ STP  

Assignment 1: Conversions

II.  The Behavior of Gases

PhET Applet: Gas Law simulation. Demonstrates the relationships between the gas variables.

Simulation: Gas Variables- How do the variables relate?

A. Particles vs. Pressure & Volume

1. Pressure- the sum of all gas collisions in a volume  

The pressure of a gas is proportional to the number of particles (Moles)

n / P = constant  (slope)  

"Diver's Law"- 

2. Avogadro's Law

The volume of a gas is directly proportional to the number of particles (Moles). Graph

In-Class Practice: Avogardro's Law

B. Volume vs. Pressure  

1. Increasing the volume decreases the number of particle collisions per time  

-pressure is inversely proportional to volume. Image

(P x V)  =  constant

2. Boyle’s Law. Named for Robert Boyle

The pressure of a confined gas varies inversely with the volume @ constant temperatures and moles  

Animated Tutorial:  Boyle's Law -- Boyle's Law

In-Class Practice: Boyle's Law. Solving for pressures and volumes of a confined gas at constant temperatures.

C. Temperature vs. Pressure  

Temperature is proportional to the average kinetic energy of the particles  

-increasing the temp. increases the kinetic energy of the particles (increase velocity)  

-increasing the kinetic energy increases the rate of collisions (pressure)  

(P / T)  =  constant

1. Gay-Lussac’s Law. Named for Joseph Louis Gay-Lussac

The pressure of a gas is directly proportional to the temperature @ constant volumes and moles

Animated Tutorial: Gay-Lussac Law. 

-The temperatures must be in units of Kelvin

In-Class Practice:  Gay-Lussac Equation

D. Volume vs. Temperature

1. Charles’ Law . Named for Jacques Alexandre César Charles

The volume of a gas is directly proportional to the temperature @ constant pressures and moles. Graph

Animated Tutorial: Charles' Law -- Charles' Law

-The temperatures must be in units of Kelvin.

In-Class Practice: Charles' Law

E. Combined Gas Law  

In-Class Practice: Combined Gas Law

Assignment 2: Gas Laws Worksheet 2  -- answers

F. Dalton’s Law of Partial Pressures

In a gaseous mixture, the total pressure of the mixture is equal to the sum of the individual pressures of each gas (partial pressure)

 -The percent composition of a gas is equivalent to the percent of the total pressure  

-The % composition is constant within the mixture @ any pressure, temperature or volume changes  

Image. Collecting a gas over water

G. Ideal Gas Law

-where R is the ideal gas law constant =  8.21 x 10^-2 L atm / mol K  (8.31 L kPa / mol K)

In-Class Practice Ideal Gas Law

Assignment 3: Gas Law Wkst 3; - answers

H. Real vs. Ideal Gases  

Ideal Gas- any gas that follows the gas laws at all pressures and temperatures  

3 Presumptions of Ideal Gases  

1.     The S of the particle volume is zero. (negligible compared to volume of the gas)

2.     The gas particles have no intermolecular attractions  

3.     The gases do not change state- all particles have the same kinetic energy

I. Ideal Gases vs. Real Gases- (Deviation from Ideal Gas Law)  

Real gases behave as ideal gases only over narrow pressure and temperature ranges  

1.     @ high pressure- decreasing volume causes particles to associate (intermolecular attractions)- creates a contraction of volume below predicted  

n  =  P  V /  R  T            ratio of n decreases more than expected (see graph)

2.     @ very high pressures- decreasing volume causes particles to become very close- the particle volume now becomes a factor

n  =  P  V /  R  T            ratio of n doesn't decrease as expected (see graph)

volume does not contract as much as predicted  

J. Diffusion and Graham’s Law  

1. Diffusion- the net movement of particles from regions of high concentration to regions of low concentration.  

2. Graham’s Law- the rate of effusion of a gas is inversely proportional to the square root of the molar mass  

Animated Tutorial: Effusion Rates

Assignment 4: Graham's Law Problems; #s 135-139 -- answers

Own-your-own Review:  Gases and Gas Laws with Practice Problems

Own-your-own practice: Review of Gas Laws.  These are good review type questions. 

Assignment 5: Ch 12 Standardized Test Prep- pg 485

Assignment 6:  Stoichiometry worksheet involving gases. 

Reading Assignment 3:  KTU & read 8.4- answer questions 29-35

III. Intermolecular Attractions

The state of a substance is determined by magnitudes of the Kinetic Energy of the particles and the magnitude of the intermolecular attractions, the force that holds the individual pieces of matter together.

These attractions are based on the chemical make-up of the particles (i.e. polarity & geometry)

A. Van der Waals Interactions (Forces)- The attractions between neutral particles. 

1. Dipole-Dipole Forces- The attractive force that occurs between neutral polar molecules.

The strength of the dipole-dipole force is proportional polarity.  (For molecules of similar size)

2. Hydrogen Bonding- Dipole-dipole interactions that involve a hydrogen atom in a polar bond (F, O, & N)

This strongest of the Van der Waals Interactions, but still considerably weaker than covalent & ionic bonds

3. London Dispersion Forces- The attraction between nonpolar particles or atoms

This is due to instantaneous dipoles (temporary) which arises from random movements of electrons

Induced Dipole- The temporary dipole in one particle and induce a dipole on a second.

B. Ion-Dipole Forces- The attraction between an ion and the charged end of a polar molecule

Common in solutions containing soluble ionic salts.

Study Sheet:  Steps to Determine Type of IMA. This is a good supplement to your notes. 

Reading Assignment 4: KTU & read 13.2-13.4- answer questions 10-14 & 25-29

IV. The Nature of Liquids  

1^st Condensed state of matter (solid the other)- distance between particles is reduced due to intermolecular forces.

Refer to Chapter 14.13.  Intermolecular Forces. (Van der Waals Interactions)  

Produces constant volume and non-constant shape  

A. Vaporization  

Conversion from liquid state to gaseous state- energy has to be added to overcome intermolecular forces  

1. Evaporation is a cooling process- releases the higher energy molecules by absorbing energy from the other molecules.

a. Vapor pressure – the  pressure produced by evaporating gas inside a confined volume  

Vapor pressure is directly proportional to temperature  

Table:  Vapor Pressure of Water

b. Condensation- Gaseous molecules that lose energy and convert to liquid state  

c. Dynamic equilibrium- the equilibrium associated with evaporating and condensing  molecules  

2. The Boiling Point

The temperature when the vapor pressure is equal to the external pressure

a. The boiling point is directly proportional pressure

Under normal pressures a liquid will not experience a higher temperature than the  boiling point of the liquid.

b. Normal boiling point- the boiling point of a liquid at a pressure of 1 atm

c. Factors influencing boiling points

1. Molecular Mass- More massive particles have less velocity (½ mv²) at common temperature.  This requires more heat in order to overcome IMAs.

See Boiling Points: straight-chained vs. branched-chained alkanes

2. Intermolecular Attractions (as measured by dipole moments)- Particles that are attracted to each other exhibit an increase in "apparent" mass; they act much larger.  The velocity of these attracted particles decrease at common temperatures thus requiring more heat in order to vaporize.

V. The nature of solids  

2^nd condensed state of matter  

Tutorial: Solid Structures.

A. Solid- constant volume & shape- particles and held in place.

1. kinetic energy is measured by the amount of vibration  

2. melting point- the temperature at which a solid turns into a solid  

B. Crystalline structures- the repeating patterns of atoms, ions or molecules in a solid  

See. Types of crystalline structures

1. unit cell- the smallest group of particles which represents the crystals if repeated in 3 dimensions  

Coordination Number- the number of other particles that associate to a given particle.

a. Types of Crystalline structures

1. Close-Packed- Coordination Number of 12.

i. Hexagonal Close-Packed- Alternating layers are superimposable- ABABAB

ii. Face-Centered Close-Packed- Alternating layers are not superimposable- ABCBABC

2. Body Centered- Coordination Number of 8

3. Simple Cubic- Coordination Number of 6

Tutorial: Understanding Crystal Structures.

2. Amorphous solids- solids that lack regular crystalline structures.  

3. Ionic crystals- The crystalline structure is determined by the relative sizes and ratios of the cations and anions. 

Larger cations can touch a more anions where small cations can touch only a few. It is common where the larger anions usually form a closed packed crystal with the smaller cations filling the holes or "interstices" between the anions.

a. Radius Rule: The ratio of cation radius (r⁺) to anion radius (r^-) determines the general geometry of the crystal. 

Tutorial: Examples of Crystalline Structures.

VI. Phase Changes:  changes from one state of matter to another.

a phase change requires an change in the kinetic energy of the particles. 

A. Melting/Freezing- the transformation between solid and liquid  

1. melting point- the temperature at which melting/freezing takes place  

B. Boiling/Condensation- the transformation between liquid and gas

1. boiling point- the temperature at which boiling/condensation takes place  

(vapor pressure = atmospheric pressure)  

C. Sublimation- the transformation between solid and gas which doesn’t involve the liquid state

D. Energy and Phase Change  

-A phase change requires a change in the kinetic energy of the particles, but the kinetic energy is required to decrease/increase the intermolecular forces.

-This is defined as Latent Heat

Table:  Latent heats of common materials

1.     Latent Heat of Fusion- the heat required to change 1 gram of a solid to liquid  

 L_f for water is 80 ^cal/g

2.     Latent Heat of Vaporization- the heat required to change 1 gram of a liquid to gas  

L_v for water is 540 ^cal/g

3.     Specific Heat- the heat required to raise 1 gram of a substance by 1 Celsius degree  

Specific heat for water is 1 ^cal/gC^o

Tutorial:  Calorimetry. Virtual Chemistry

Experiment: Introduction to Calorimetry.  

Assignment 7: Phase Change/Temperature Change worksheet

E. Phase Diagrams- 

-Shows the relationship between the phase changes associated with temperature and pressure conditions. The following is a phase diagram for water:

See. Typical Phase Diagram

1. Equilibrium- The condition where the conversion of one state to another occurs at the same time and at the same rate.  Identified as the lines on the phase diagram.

2. Triple point- The specific temperature and pressure where all 3 phases occur in equilibrium.  

For water- 0^oC and 4.6 torr

Practice:  Heat changes

Sample Quiz:  Heat changes

Assignment 8: Unit VI Review Worksheet