1 / 39

KINETIC THEORY

KINETIC THEORY. Unit 7 Chemistry Langley. *Corresponds to Chapter 13 (pgs. 384-409) in Prentice Hall Chemistry textbook. KINETIC THEORY. Kinetic Theory states that the tiny particles in all forms of matter are in constant motion. Kinetic refers to motion

emele
Download Presentation

KINETIC THEORY

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. KINETIC THEORY Unit 7 Chemistry Langley *Corresponds to Chapter 13 (pgs. 384-409) in Prentice Hall Chemistry textbook

  2. KINETIC THEORY • Kinetic Theory states that the tiny particles in all forms of matter are in constant motion. • Kinetic refers to motion • Helps you understand the behavior of solid, liquid, and gas atoms/molecules as well as the physical properties • Provides a model behavior based off three principals

  3. KINETIC THEORY • 3 Principles of Kinetic Theory • All matter is made of tiny particles (atoms) • These particles are in constant motion • When particles collide with each other or the container, the collisions are perfectly elastic (no energy is lost)

  4. STATES OF MATTER • 5 States of Matter • Solid • Liquid • Gas • Plasma • Bose-Einstein Condensates http://www.plasmas.org/E-4phases2.jpg

  5. SOLIDS • Particles are tightly packed and close together • Particles do move but not very much • Definite shape and definite volume (because particles are packed closely and do not move) • Most solids are crystals • Crystals are made of unit cells (repeating patterns) • The shape of a crystal reflects the arrangement of the particles within the solid

  6. SOLIDS • Unit cells put together make a crystal lattice (skeleton for the crystal) • Crystals are classified into seven crystal systems: cubic, tetragonal, orthorhombic, monoclinic, triclinic, hexagonal, rhombohedral • Unit cell  crystal lattice  solid

  7. SOLIDS • Amorphous Solid: • A solid with no defined shape (not a crystal) • A solid that lacks an ordered internal structure • Examples: Clay, PlayDoh, Rubber, Glass, Plastic, Asphalt • Allotropes: • Solids that appear in more than one form • 2 or more different molecular forms of the same element in the same physical state (have different properties) • Example: Carbon • Powder = Graphite • Pencil “lead” = graphite • Hard solid = diamond

  8. SOLIDS www.ohsu.edu/research/sbh/resultsimages/crystalvsglass.gif

  9. SOLIDS Allotropes of Carbon: a) diamond, b) graphite, c) lonsdaleite, d)buckminsterfullerene (buckyball), e) C540, f) C70, g) amorphous carbon, and h) single-walled (buckytube) www.wikipedia.org

  10. LIQUIDS • Particles are spread apart • Particles move slowly through a container • No definite shape but do have a definite volume • Flow from one container to another • Viscosity – resistance of a liquid to flowing • Honey – high viscosity • Water – low viscosity chemed.chem.purdue.edu/.../graphics

  11. GASES • Particles are very far apart • Particles move very fast • No definite shape and No definite volume http://www.phy.cuhk.edu.hk/contextual/heat/tep/trans/kinetic_theory.gif

  12. PLASMA • Particles are extremely far apart • Particles move extremely fast • Only exists above 3000 degrees Celsius • Basically, plasma is a hot gas • When particles collide, they break apart into protons, neutrons, and electrons • Occurs naturally on the sun and stars

  13. BOSE-EINSTEIN CONDENSATE • Particles extremely close together • Particles barely move • Only found at extremely cold temperatures • Basically Bose-Einstein is a cold solid • Lowest energy of the 5 states/phases of matter

  14. GASES AND PRESSURE • Gas pressure is the force exerted by a gas per unit surface area of an object • Force and number of collisions • When there are no particles present, no collisions = no pressure = vacuum • Atmospheric Pressure is caused by a mixuture of gases (i.e. the air) • Results from gravity holding air molecules downward in/on the Earth’s atmosphere; atmospheric pressure decreases with altitude, increases with depth • Barometers are devices used to measure atmospheric pressure (contains mercury) • Standard Pressure is average normal pressure at sea level • As you go ABOVE sea level, pressure is less • As you go BELOW sea level, pressure is greater

  15. GASES AND PRESSURE • Standard Pressure Values • At sea level the pressure can be recorded as: • 14.7 psi (pounds per square inch) • 29.9 inHg (inches of Mercury) • 760 mmHg (millimeters of Mercury) • 760 torr • 1 atm (atmosphere) • 101.325 kPa (kilopascals) • All of these values are EQUAL to each other: • 29.9 inHg = 101.325 kPa • 760 torr = 760 mmHg • 1 atm = 14.7 psi • and so on………. • Say hello to Factor Label Method!!!!!!!!!!!!

  16. GASES AND PRESSURE • STP • Standard Temperature and Pressure • Standard Pressure values are the values listed on the previous slides • Standard Temperature is 0°C or 273 K • If temperature is given to you in Farenheit, must convert first! • °F = (9/5)°C + 32 • °C = (5(°F-32)) / 9 Remember order of operation rules • K = 273 + °C • °C = K – 273

  17. GASES AND PRESSURE • Pressure Conversions • Example 1: 421 torr = ? Atm • Step 1: Write what you know • Step 2: Draw the fence and place the given in the top left • Step 3: Arrange what you know from step 1 such that the nondesired units canceling out so that you are only left with the units you want (i.e. atm) • Step 4: Solve • Step 5: Report final answer taking into account the appropriate significant figures

  18. GASES AND PRESSURE • Pressure Conversions • Example 2: 32.0 psi = ? torr

  19. TEMPERATURE • Temperature is the measure of the average kinetic energy of the particles. • 3 Units for Temperature: • Celsius • Farenheit • Kelvin • Has an absolute zero • Absolute lowest possible temperature • All particles would completely stop moving • Temperature Conversions: • Example 1: Convert 35°C to °F • Example 2: Convert 300 Kelvin to °C

  20. MEASURING PRESSURE • Manometers: • Measure pressure • 2 kinds: open and closed • Open Manometers: • Compare gas pressure to air pressure • Example: tire gauge • Closed Manometer: • Directly measure the pressure (no comparison) • Example: barometer

  21. KINETIC ENERGY AND TEMPERATURE • Energy of motion • Energy of a moving object • Matter is made of particles in motion • Particles have kinetic energy • KE = (mv2)/2 OR KE = (ma)/2 • Kinetic Energy is measured in Joules • 1 J = 1kg•m2/s2 • The mass must be in kg • The velocity must be in m/s OR acceleration must be in m2/s2

  22. KINETIC ENERGY AND TEMPERATURE • Calculate the KE of a car with a mass of 1500 kg and a speed of 50 m/s

  23. KINETIC ENERGY AND TEMPERATURE • Calculate the KE of a car with a mass of 6780 grams and a speed of 36 km/h

  24. KINETIC ENERGY AND TEMPERATURE • Temperature-measure of the average kinetic energy of the particles • Kelvin Scale: • Has an absolute zero (0K) • Absolute lowest possible temperature • In theory, all particles would completely stop moving • Speed of Gases: • If two gases have the same temperature (particles moving at the same speed) how can you tell which gas has a greater speed? • The only difference is mass! • To find mass, use the periodic table

  25. KINETIC ENERGY AND TEMPERATURE • Speed of Gases • Example 1: If CH4 and NH3 are both at 284 K, which gas has a greater speed? • Step One: Add up the mass of each gas using the periodic table. • Step Two: The lighter gas moves faster (think about a race between a 100-pound man and a 700-pound man, the lighter man would move faster) • Example 2: Which gas has a faster speed between Br2 and CO2 if both are at 32°F?

  26. TERMINOLOGY for PHASE CHANGES • Melting-commonly used to indicate changing from solid to liquid • Normal melting point-The temperature at which the vapor pressure of the solid and the vapor pressure of the liquid are equal • Freezing-Changing from a liquid to a solid • Melting and freezing occur at the same temperature • Liquifaction-Turning a gas to a liquid • Only happens in low temperature and high pressure situations

  27. TERMINOLOGY for PHASE CHANGES • Difference in Gas and Vapor • Gas-state of matter that exists at normal room temperature • Vaport-produced by particles escaping from a state of matter that is normally liquid or solid at room temperature • Boiling-used to indicate changing from a liquid to a gas/vapor • Normal boiling point - temperature at which the vapor pressure of the liquid is equal to standard atmospheric pressure, which is 101.325 kPa • Boiling point is a function of pressure. • At lower pressures, the boiling point is lower

  28. TERMINOLOGY for PHASE CHANGES • 2 types of boiling: boiling and evaporation • Evaporation takes place only at the surface of a liquid or solid while boiling takes place throughout the body of a liquid • Particles have high kinetic energy • Particles escape and become vapor • Condensation-used to indicate changing from a vapor to a liquid

  29. TERMINOLOGY for PHASE CHANGES • Sublimation - when a substance changes directly from a solid to a vapor • The best known example is "dry ice", solid CO2 • Deposition-when a substance changes directly from a vapor to a solid (opposite of sublimation) • Example-formation of frost • Dynamic equilibrium - when a vapor is in equilibrium with its liquid as one molecule leaves the liquid to become a vapor, another molecule leaves the vapor to become a liquid. An equal number of molecules will be found moving in both directions • Equilibrium - When there is no net change in a system

  30. TERMINOLOGY for PHASE CHANGES • Points to Know: • Melting Point-Temperature when solid turns to a liquid • Freezing Point-Temperature when liquid turns to a solid • Boling Point-Temperature when a liquid turns to a vapor • Doesn’t boil unitl vapor pressure coming off liquid is equal to the air pressure around it • Since air pressure changes with height, water does not always boil at 100°C • Condensing Point-Tempeature when vapor turns to liquid

  31. ENTROPY • A measure of the disorder of a system • Systems tend to go from a state of order (low entropy) to a state of maximum disorder (high entropy) • Entropy of a gas is greater than that of a liquid; entropy of a liquid is greater than that of a solid • Solids=low entropy; plasma=high entropy • Entropy tends to increase when temperature increases • As substances change from one state to another, entropy may increase or decrease

  32. Le CHATELIER’S PRINCIPLE • Anytime stress is placed on a system, the sytem will readjust to accommodate that stress • If a chemical system at equilibrium experiences a change in concentration, temperature, volume, or total pressure, then the equilibrium shifts to partially counteract the imposed change • Can be used to predict the effect of a change in conditions on a chemical equilibrium • Is used by chemists in order to manipulate the outcomes of reversible reactions, often to increase the yield of reactions

  33. Le CHATELIER’S PRINCIPLE • When liquids are heated (stress) they produce vapor particles (adjust) • When liquids are cooled (stress) the particles inside tighten to form a solid (adjust)

  34. Le CHATELIER’S PRINCIPLE • Le Chatelier’s Principle explaining boiling and condensation using covered beaker partially filled with water • At a given temperature the covered beaker constitutes a system in which the liquid water is in equilibrium with the water vapor that forms above the surface of the liquid. • While some molecules of liquid are absorbing heat and evaporating to become vapor, an equal number of vapor molecules are giving up heat and condensing to become liquid. • If stress is put on the system by raising the temperature, then according to Le Châtelier's principle the rate of evaporation will exceed the rate of condensation until a new equilibrium is established

  35. PHASE DIAGRAMS • A diagram showing the conditions at which substance exists as a solid, liquid, or vapor • Shows the temperature and pressure required for the 3 states of matter to exist • Conditions of pressure and temperature at which two phases exist in equilibrium are indicated on a phase diagram by a line separating the phases • Draw the phase diagram for water

  36. PHASE DIAGRAM-WATER

  37. PHASE DIAGRAM-WATER • Explanation of Phase Diagram: • X axis-Temperature (°C) • Y axis- Pressure (kPa) • Line AB – line of sublimation • Line BD – boiling point line • Line BC – melting point line • Point B – triple point (all 3 states of matter exist at the same time) • Tm – melting point at standard pressure • Tb – boiling point at standard pressure

  38. HEAT in CHANGES of STATE • Energy Diagrams (also referred to as Heating Curves) • Graphically describes the enthalpy (the heat content of a system at sonstant pressure) changes that take place during phase changes • X axis is Energy (Heat supplied) • Y axis is Temperature

  39. HEAT in CHANGES of STATE • Constructing Energy Diagrams • Step 1: Determine/Identify the melting and boiling points for the specified substance • Step 2: Draw x and y axis (energy vs temp) • Step 3: Calculations • First diagonal line: Q = mcDT • First horizontal line: Q = mHf • Second diagonal line: Q = mcDT • Second horizontal line: Q = mHv • Third horizontal line: Q = mcDT • Add up all values!!! • Draw the energy diagram for 10 grams of water as it goes from –25°C to 140°C

More Related