Molecular Kinetic Theory S. Staron 2-11

KINETIC THEORY OF MATTER ALL PARTICLES (ATOMS OR MOLECULES)OF MATTER ARE IN CONSTANT MOTION! Kinetic – comes from Greek word meaning “to move” Kinetic Energy – energy object has due to its motion http://www.chem.purdue.edu/gchelp/atoms/states.html http://www.media.pearson.com.au/schools/cw/au_sch_whalley_sf1_1/int/matter.html

Separation Between Particles Gases separation between particles is very large compared to their size there are no attractive or repulsive forces between the molecules Liquids particles are still far apart but now they are close enough that attractive forces confine the material to the shape of its container Solids particles are so close that the forces of attraction confine the material to a specific shape.

USING THE KINETIC THEORY TO EXPLAIN BEHAVIOR OF MATTER Gases Particles in constant, random motion – allows gas to fill container Motion of particles unaffected by other particles unless collision Forces of attraction between particles can be ignored Liquid Particles can flow to new locations Force of attraction between particles keeps them close Solid Particles vibrate around fixed locations due to stronger forces of attraction between particles http://phet.colorado.edu/simulations/sims.php?sim=States_of_Matter

Particles are in Constant Motion Each particle is in constant motion Gases: the movement of the particles is assumed to be random and free Liquids: the movement is somewhat constrained by the volume of the liquid Solids: the motion of the particles is severely constrained to a small area, in order for the solid to maintain its shape. The velocity of each particle determines its kinetic energy.

Collisions Transfer Energy The numerous particles often collide with each other. If a gas or liquid is confined in a container, the particles collide with the particles that make up the walls of a container.

Temperature Temperature of an object or collection of matter is the average kinetic energy of the particles. Faster moving particles - a higher temperature. A thermometer is used to measure the temperature Puts it into temperature degrees instead of kinetic energy units.

Compare average KE of particles in 3 states of matter? For a Given Substance: Solid: low KE  slower speed  vibrate around fixed locations Why? Particles “stuck” due to forces of attraction between particles Liquid: greater KE than solid  greater speed  not “stuck” in place so can flow to new locations Forces of attraction? Still there but not as strong Gas: greater KE than liquid  even greater speed  particles can spread out far from each other Forces of attraction? Very weak (can be ignored)

How is one state changed into another? Energy is ABSORBED by the system from the surroundings  Endothermic Process Energy is RELEASED by the system to the surroundings  Exothermic Process This change in the state of matter is called a PHASE CHANGE

States of Matter State Movement speed Position Plasma Extremely fast – so much energy that electrons come off All over “Bumper cars” Gas Fast (500 – 1000 m/s) Liquid Slower than in a gas (~200 m/s) Slide around each other (Hallway full of people) Solid A. Crystalline solid (crystal) B. Amorphous solid (glass)   No change of position Vibrate in place Stay in one place A. Orderly arrangement B. Not very ordered BEC Bose-Einstein Condensate VERY cold! So close they act as one big atom/molecule

States of Matter http://phet.colorado.edu/simulations/sims.php?sim=Gas_Properties

Solids The kinetic theory says solids are closely packed atoms or molecules (groups of atoms). Solids don’t have enough space between the particles to move around, but do have energy to vibrate in place.

Picture of a solid:

Liquids The kinetic theory says liquids are closely packed atoms or molecules (groups of atoms), but do have enough energy to keep some amount of space between their particles. Liquids do have enough space between them for particles to slide around each other, but not enough energy to overcome gravity.

Picture of a Liquid

Gas The kinetic theory says gases are widely spaced atoms or molecules. Gases have a lot of space between the particles to move around.

Picture of a Gas

Equivalent Units Many of the properties of gases can be measured in different ways. Conversion from one unit of pressure to another is very important. To achieve this there has to be a conversion factor to move from one unit to another. Here is a list of equivalent amounts of pressure: 1 atm 760 mmHg 76 cmHg 101.3 kPa 760 torr 29.92 inches Hg

General Properties of Gases Gas particles can be monatomic (Ne), diatomic (N2), or polyatomic (CH4) – but they all have these characteristics: 1. Gases have mass. 2. Gases diffuse to fill their containers uniformly and completely. 3. Gases diffuse and mix rapidly. 4. Gases exert pressure. 5. A gas’s pressure is dependent on its temperature.

Gases are not dense, and fill up their containers.

The Kinetic Theory The Kinetic Theory states that the tiny particles in all forms of matter are in constant motion. This theory is used to explain the behaviors common among gases There are 3 basic assumptions made by the Kinetic Theory as it applies to gases.

Large Separation Between Particles Gas Separation between particles is very large compared to their size There are no attractive or repulsive forces between the molecules Liquid Particles are still far apart, but now are close enough that attractive forces confine the material to the shape of its container Solid Particles are so close that the forces of attraction confine the material to a specific shape.

Kinetic Theory Assumption 1: An “Ideal” Gas A gas is composed of small hard particles. The particles have an insignificant volume and are relatively far apart from one another. There is empty space between particles. There is no attractive or repulsive forces acting between particles.

Kinetic Theory Assumption 2: An “Ideal” Gas The particles in a gas move in constant random motion. Particles move in straight paths and are completely independent of each of other A particle’s path is only changed by colliding with another particle or the sides of its container.

Kinetic Theory Assumption 3: An “Ideal” Gas All collisions a gas particle undergoes are perfectly elastic. No energy is lost from one particle to another, and the total kinetic energy remains constant.

Deviations from the “Ideal” Gas 1) Real molecules have volume. The ideal gas consumes the entire amount of available volume. It does not account for the volume of the molecules themselves. 2) There are intermolecular forces. An ideal gas assumes there are no attractions between molecules. In reality, attractions slow down the molecules and reduce the amount of collisions. (Otherwise a gas could not condense to become a liquid.)

Gas Compresses Since one of the properties of a gas is compressibility, a gas at a certain volume can be compressed by applying more pressure. The mass of the gas will remain unchanged. Since the mass remains the same and the volume decreases, the density of the gas is greater.

If you squeeze a gas, its volume can be reduced considerably.

Example If the mass of the gas is .50 grams and the volume of the gas is one liter then the density of the gas is .50 grams/liter. However, if the gas is compressed to only take up one quarter of a liter then the density will change to 2 g/l.

If I opened up a bag of popcorn in front of the class you would soon be able to smell it in the back. The popcorn smell is a high- energy group of molecules in a gaseous state. This property of gases spreading out until they have filled the room is called diffusion (spreading from an area of high concentration to one of low concentration).

It is because of all of the empty space between gas molecules that another gas molecule can pass between two of them until each gas is spread out over the entire container.

If the gases are in constant random motion the fact that they are moving and colliding with everything around them means they will mix with other gases uniformly. This doesn’t happen at the same speeds for all gases though. Some gases diffuse more rapidly then other gases – this is based on their size and their energy.