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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

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KINETIC THEORY

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  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

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