# The Kinetic Molecular Theory

## Presentation on theme: "The Kinetic Molecular Theory"— Presentation transcript:

The Kinetic Molecular Theory
Basic Assumptions Particle Size Gas particles have no volume (pin point particles) The space between particles is extremely large compared to the volume of the particles. Due to this distance, there is no significant attractive or repulsive force acting on the particles.

The Kinetic Molecular Theory
Basic Assumption Particle Motion Gas particles are in constant random motion. Collisions between particles are elastic (Energy can be transferred from one particle to another during a collision, but no energy is lost when particles collide)

Basic Assumptions Particle Energy
The mass and velocity of a particle determine the kinetic energy of a particle Temperature is a measure of the average kinetic energy of particles in a sample.

The Kinetic Molecular Theory
Mass/Velocity Relationship Questions Condition #1: Two particles (one heavy and one light) traveling at the same velocity. Which exhibits the greatest kinetic energy? Condition #2: Two particles of the same size traveling at different velocities (fast and slow).

Explaining the Behavior of Gases
Properties Low Density (pinpoint mass/volume of empty space) Random Motion Behaviors Compression Gases can be compressed due to the large space that exists between particles Gases expand to fill their containers due to constant random motion

Explaining the Behavior of Gases
Properties No attractive or repulsive forces acting on particles Particles exhibit constant random motion Behaviors Particles can flow easily past each other in a process called diffusion. The rate of diffusion is dependent on the mass of the particles. Question: Based on this equation which particles diffuse faster, heavy or light particles?

Explaining the Behavior of Gases
Property Particles exhibit constant random motion Behavior Effusion (similar to diffusion, where particles escape through a tiny opening) Graham’s Law of Effusion Rate of effusion

Gas Pressure An important property of any gas is its pressure
A linear momentum of the gas molecules Pressure is defined as force per unit area Gas particles exert pressure when they collide with the walls of their container The less mass, the less pressure it exerts

Gas Pressure (Continued)
Temperature, volume and number of moles affect the pressure that a gas exerts The Earth is surrounded by an atmosphere that extends into space for hundreds of kilometers. The particles in air move in every direction which exerts pressure in all of these directions; this is known as atmospheric pressure Air Pressure varies at different points on earth

Measuring Gas Pressure
Evangelista Torricelli ( ) was the first to demonstrate that air exerted He designed experiments in order to test his theories about liquid and density He invented a barometer: an instrument used to measure atmospheric pressure

Units of Pressure The SI unit of pressure is called the pascal (Pa)
The pressures measured by barometers and monometers can be reported in: Millimeters of mercury (mm Hg) Also the Torr, which is named after Torricelli Air Pressure is often reported in a unit called an atmosphere (atm) 1 atm= 760 mm Hg 760 torr 101.3 kPa (kilopascals) 14.7 psi (pounds per square inch)

Dalton’s Law of Partial Pressure
Dalton’s Law explains that the totally pressure is equal to the sum of all the pressures of the parts This is only true for ideal gases, but the error is small for real gases PressureTotal = Pressure1 + Pressure2 + Pressuren The portion of the total pressure contributed by a single gas is called its partial pressure Partial pressure of a gas depends on Number of moles of gas Size of the container temperature

Force holds all matter together. The force in a molecule that is between atoms is a chemical or intramolecular force. The force between molecules is called a physical or intermolecular force.

Dipole-dipole Forces are attractive forces between the positive end of one polar molecule and the negative end of another polar molecule. Dipole-dipole forces have strengths that range from 5 kJ to 20 kJ per mole. They are much weaker than ionic or covalent bonds and have a significant effect only when the molecules involved are close together.

Figure: two arrangements of polar iodine monochloride (ICl) molecules that give rise to dipole-dipole attractions.

Hydrogen Bonds special type of dipole-dipole attraction Definition: a dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a small, highly electronegative atom with at least one lone electron pair. only form to fluorine, oxygen, or nitrogen atoms. REASON: they are electronegative enough to cause a large partial positive charge on the hydrogen atom, but small enough that the lone pairs of electrons can come close to hydrogen atoms.

Hydrogen Bonds in a water molecule:
Water has two hydrogen atoms and one oxygen atom. Two molecules of water can form a hydrogen bond between them. The oxygen of one water molecule has two lone pairs of electrons, each of which can form a hydrogen bond with hydrogen on two other water molecules.

Density & Compression Density- Density is a measurement of the mass per unit volume Compression- This unit of measure expresses the degree of firmness or softness Liquids can be compressed but the change in volume for liquids is much smaller because liquid particles are already tightly packed Liquids are much denser than gases

Fluidity Fluidity- The ability to flow
A liquid can diffuse through another liquid Liquids diffuse more slowly than gases even at the same temp Liquids are less fluid than gases

Viscosity Viscosity- The measure of the resistance of a liquid to flow
Viscosity of a liquid is determined by the type of intermolecular forces, the shape of the particles, and temperature The stronger the forces, the higher the viscosity Viscosity decreases with temperature The added energy from a higher temperature makes it easier for the molecules to overcome the intermolecular forces that keep the molecules from flowing Higher viscosity with higher molecular mass

Surface Tension Surface Tension- The measure of the inward pull by particles in the interior if the liquid For the surface area to increase particles from the interior must move to the surface The stronger the attractions between particles, the greater the surface tension Surfactants- Compounds that lower the surface tension of water, also referred to as surface agents Water has a high surface tension because its molecules can form multiple hydrogen bonds

Capillary Action Capillary Action- The movement of a liquid along the surface of a solid caused by the attraction of molecules of the liquid to the molecules of the solid (adhesion) Molecules of water are naturally attracted to each other and form temporary hydrogen bonds with each other (cohesion); their attraction for each other on the surface of a liquid, for example, gives rise to surface tension Results in the elevation or depression of liquids in capillaries (small narrow tubes)

Density of Solids Except water
There is about a 10% difference in density between the solid and liquid states of most substances In general, the particles in solids are more closely packed than those in liquids Most solids are more dense than most liquids Except water the water molecules in ice are less closely packed together than in liquid water

Crystalline Solids A solid whose atoms, ions, or molecules are arranged in an orderly, geometric, three-dimensional structure Crystals -- The individual pieces of a crystalline solid Unit Cell -- the smallest arrangement of connected points that can be repeated in 3 dimensions to form the lattice The shape of a crystalline solid is determined by the type of unit cell from which its lattice is built Crystal shapes differ because the surfaces of unit cells do not always meet at right angles and the edges of the faces vary in length

Molecular Solids The molecules are held together by dispersion forces,
ex: table sugar The molecules are held together by dispersion forces, dipole - dipole forces or hydrogen bonds Most are not solid at room temperature Molecular compounds are solids at room temperature because of their large molar mass Because they have no ions, molecular solids are poor conductors of heat and electricity Fairly soft and have low to moderately high melting points

Covalent Network of Solids
ex: quartz and diamond Atoms that can form multiple covalent bonds are able to form covalent network solids Very high melting points Usually have poor conductivity

Ionic Solids ex: table salt Each ion is surrounded by ions of opposite charge The type of ions and the ratio of ions determine the structure of the lattice and the shape of the crystal There is a network of attractions that extends throughout an ionic crystal This gives these compounds their high melting points and hardness Ionic crystals are strong but brittle

Metallic Solids Strength varies by the physical property
Consists of positive metal ions surrounded mobile electrons Strength varies by the physical property Mobile electrons make metals malleable and ductile They also make metals good conductors of electricity and heat When force is applied to a metal, electrons shift and keep the metal ions bonded in their new positions All metallic elements can form a metallic solid

Amorphous Solids A solid in which there is no long range order of the positions of the atoms ex: glass and rubber Has no shape Any liquid can be made into an amorphous solid if given a sufficiently high cooling rate Cooling reduces molecular mobility An amorphous solid often forms when a molten material cools too quickly to allow enough crystals to form Not all particles are arranged in specific order

Endothermic Processes
Endothermic- A chemical reaction in which a greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds form in the product molecules

Melting The amount of energy required to melt one mole of a solid depends on the strength of the forces keeping the particles together in the solid The melting point of a crystalline solid is the temperature at which the forces holding its crystal lattice together are broken and it becomes a liquid

Vaporization vs. Evaporation
Vaporization is the process by which a liquid changes to a gas or a vapor Evaporation is when vaporization occurs only at the surface of a liquid

Boiling The temperature at which the vapor pressure of a liquid equals the external or atmospheric pressure is called the boiling point

Process of Boiling Molecules throughout the liquid have energy to vaporize Bubbles of vapor collect below the surface of the liquid and rise to the surface

Sublimation Sublimation is the process by which a solid changes directly to a gas without first becoming a liquid

Exothermic Changes Exothermic- A chemical reaction in which more energy is released than is required to break bonds in the initial reaction ature=related

Condensation The energy-releasing process by which a gas or vapor becomes a liquid Condensation is the reverse of vaporization Different causes of the condensation of water vapor in which all involve a transfer of energy

Condensation The vapor molecules can come in contact with a cold surface such as the outside of a glass containing ice water Heat from the vapor molecules is transferred to the glass as the water vapor condenses The water vapor that condenses on blades of grass or the car forms liquid droplets called dew

Condensation When a layer of air near the ground cools, water vapor in the air condenses and forms fog Clouds form when layers of air high above the surface of Earth cool

Deposition The energy-releasing process by which a substance changes from a gas or vapor to a solid without first becoming a liquid Deposition is the reverse of sublimation When water vapor comes in contact with a cold window in winter, it forms a solid deposit on the window called frost

Freezing Freezing is the reverse of melting
The freezing point is the temperature at which a liquid is converted into a crystalline solid

Phase Diagrams Two variables control the phase of a substance: temperature and pressure Phase Diagram- a graph of pressure versus temperature that shows which phase substance exists under different conditions of temperature and pressure

Phase Diagrams Triple Point- the point on a phase diagram representing the temperature and pressure at which three phases of a substance (solid, liquid, and gas) can coexist Critical Point- indicates the critical pressure and critical temperature above which water can not exist as a liquid

Additional Example of a Phase Diagram
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