Ch Kinetic Theory. 1.All matter is made of atoms and molecules that act like tiny particles.

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1 Ch Kinetic Theory 1.All matter is made of atoms and molecules that act like tiny particles.

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3 Kinetic Theory 2.These tiny particles are always in motion. The higher the temperature, the faster the particles move.

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5 Kinetic Theory 3.The more massive the particles, the slower the particles will diffuse and flow.

6 Intermolecular forces determine the phase of matter.

7 Gas Pressure The result of simultaneous collisions of billions of gas particles with an object. The more collisions, the greater the pressure

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10 Vacuum A controlled condition where no gas particles are present. So no gas pressure can exist.

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13 Atmospheric Pressure Results from collisions of air molecules with objects. Decreases as you climb a mountain because the air thins out at higher elevations Measured by a barometer

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16 Measuring Pressure STP (Standard Temp. Pressure) Standard Temperature at sea level is 0 0 C or 273 K Standard Pressure is kpa, 760 torr, 760 mm Hg, or 1 atm

17 Pressure Conversions How many kpa s are in 1.50 atm? 1 atm = kpa 1.50 atm x 101.3kPa = 152 kpa 1 atm How many kpa s are in 690 mm Hg? 690 mm Hg x kpa = 92 kpa 760 mm Hg

18 Energy and Temperature Kinetic Energy measures the average speed of particles. The higher the temperature the greater the particle speed. Temperature measures kinetic Energy SI base unit is Kelvin (K)

19 Converting from Celsius 0 degree Celsius is equal to 273 K K = 0 C C = K 273 Convert 191 K to Celsius 0 C = = C

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21 Absolute Zero Theoretical temperature at which all motion stops. Scientists have experimented a tenth of degree to 0 K, but have never gotten 0 K.

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23 Ch18.1 Ideal/Real Gases no definite shape. (both) no definite volume. (both) Particles move rapidly in constant random motion (both) State of disorder Low density (both) All collisions perfectly elastic, no attractive forces (Ideal)

24 Ideal/Real Gases Real gas particles will stick together and attract A real gas will behave like an ideal gas at low pressures or high temperatures Hydrogen and helium always behave ideally due to there small masses.

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28 Gas Relationships Relationships between pressure, volume, temperature and number of moles (amount) While examining relationships, two measurements will always be constant (unchanged)

29 Pressure vs Volume Real Gas

30 1. Pressure vs Volume Boyle s Law For a given mass of gas at a constant temperature, the volume of the gas varies inversely with pressure. Pressure increases, volume decreases As volume (space) decreases, the particles become closer and collide (pressure) more often. P 1 V 1 =P 2 V 2

31 Practice Problem If you had a gas that exerted 202 kpa of pressure and took up a space of ml. If you decide to expand the tank to 7.00 L, what would be the new pressure? (Assume constant temperature) P 1 V 1 =P 2 V 2 Check units 202 kpa x 3.00 liters = P 2 x 7.00 liters 606 = P 2 x 7.00 liters P 2 = 86.6 kpa

32 Ideal gas

33 2. Temperature vs Volume Charles Law For a given mass of gas at a constant pressure, the volume of the gas varies directly with its Kelvin temperature. Temperature increases, volume increases As temperature (speed of particles) increases, the particles move farther apart increasing volume (space) while maintaining a constant pressure. V 1 /T 1 =V 2 /T 2 or V 1 T 2 =V 2 T 1

34 Practice Problem If you took a balloon outside that was at C at 2.0 liters and heated up to C, what volume would the balloon occupy now? (Assume constant pressure) V 1 T 2 =V 2 T 1 Check units(remember KELVIN) 2.0 L x 302 K = V 2 x 293 K 604 = V 2 x 293 K V 2 = 2.1 L

35 CHARLES LAW: ΔT T and ΔV

36 CHARLES in charge was on TV

37 Ideal Gas

38 3. Temperature vs Pressure Gay-Lussac s Law For a given mass of gas at a constant volume, the pressure of a gas varies directly with its Kelvin temperature. Temperature increases, pressure increases As temperature (speed of particles) increases, the particles collide (pressure) more often in a set volume (space). P 1 /T 1 =P 2 /T or 2 P 1 T 2 =P 2 T 1

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40 Ostrich Egg in Microwave

41 Combined Gas Law Combines all three gas laws into one expression.

42 Practice Problem You have a 2.0 liter balloon that was at C and 1.5 atm. If you take this balloon and place it in a room at STP conditions, what volume would the balloon occupy? P 1 V 1 T 2 =P 2 V 2 T 1 (Remember KELVIN) 1.5 atm x 2.0 L x 273 K = 1atm x V 2 x 293 K 819 = V 2 x 293 K V 2 = 2.8 L

43 4. Moles (amount) vs Temp moles increases, temp. decreases Inverse relationship In a set volume (space), adding more moles of a gas (amount), will cause the particles to slow down (temp.) in order to maintain a constant pressure. Compress tanks become colder as you fill them

44 Ideal Gas

45 5. Moles (amount) vs pressure moles increases, pressure increases In a set volume (space), adding more moles of a gas (amount), will cause more collisions (pressure) between gas particles. Think of a super soaker or simply filling your tire

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48 6. Avogodro s s Law (ch19.1) Amount (moles) is directly proportional to the space occupied. The greater the moles of a gas (amount), the more volume (space) the particles will need in order to maintain constant pressure (particles collide) 1 mole of gas at STP= 22.4 liters of any gas

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50 Practice Problem How many liters of Hydrogen are in 6.2 grams of H 2 at STP? Molar of mass of is H 2 2 gram/mole 6.2 grams H 2 x 22.4 L H 2 / 2 gram of H 2 69 L of H 2

51 Practice Problem #2 What is the volume of hydrogen at STP can be produced when 6.54 grams of Zinc metal reacts with Hydrogen Chloride acid? Zn + 2HCl ZnCl 2 + H g Zn x 1 mol Zn x 1 mol HCl x 22.4 L H 2 65 g Zn 1 mol Zn 1 mol H 2 = 2.24 L H 2 will form

52 7. Ideal Gas Law (ch19.1) PV = nrt R = constant (L*atm)/(mol*K) 8.31 (L*kPa)/(mol*K) n = represents the number of moles. Can be used in determining densities of different gases.

53 Practice Problem A propane tank that holds g of C 3 H 8. How much larger a container would be needed to hold the same amount of propane if the gas is at 25 0 C and a pressure of 2280 mm Hg?

54 Solution PV=nRT. First solve for n g of C 3 H 8. x 1 mole = 44 grams C 3 H moles 2280 mmhg x 1 atm/760 mm Hg = 3.0 atm V*3.0 atm= moles x x298K V = 560 L C 3 H 8

55 Practice Problem 2.0 grams of N 2 is kept under a pressure of 0.95 atm, and a temperature of C. What is the density of the gas under these conditions?

56 D =2.0 g / 1.9 L = 1.1g/L N Solution PV=nRT. First solve for n 2.0 g of N 2 x 1 mole = 28 grams N moles V*0.95 atm= moles x x303K = 1.9 L D=m/v

57 8. Graham s s Law of Diffusion (ch18.2) Diffusion is the random scattering of gas molecules. The longer they diffuse the more evenly distributed they will become in the container. The heavier the gas the slower the rate of diffusion.

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60 LAST LAW! I PROMISE (ch18)

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62 Ch Changing States

63 Changes of State Endothermic Process solid liquid gas Exothermic Process

64 Melting and Freezing Melting Point is the temperature at which a solid becomes a liquid Melting and freezing take place at the same threshold temperature. According to Kinetic Theory, almost all solids and liquids expand and become disordered when the temperature is raised.

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68 Evaporation Conversion of a liquid to a gas or vapor below its boiling point. It occurs only at the surface. Remember the difference between a vapor and gas. Vapor is normally a liquid or solid at room temperature

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70 Vapor Pressure The pressure exerted by a vapor in equilibrium with its liquid state. Vapor pressure measures how easily a liquid changes into vapor Liquids with high vapor pressures turn into vapors very easily. (Volatile liquids) Ex. Gasoline, perfume

71 Dynamic Equilibrium Once equilibrium is reached, the vapor particles will begin to condense back to a liquid at the same rate they change into a vapor.

72 Vapor Equilibrium reached

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74 Boiling Point The temperature at which the vapor pressure of the liquid equals the atmospheric pressure The entire liquid is changing state, not just the surface. Liquids with low boiling points are considered volatile

75 Difference between Evaporation and Boiling

76 Super Heated Water

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78 Distillation A method of separating substance with different boiling points. Used in desalinating sea water.

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80 Sublimation Process where solid goes directly to a gas (vapor), because the vapor pressure is so high, liquid phase does not exist. Ex. Iodine, Dry Ice

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82 Condensation The changing of a vapor to a liquid

83 Liquefaction Changing a gas into a liquid. A gas can be changed into a liquid by two methods: must be placed under tremendous pressure (compressing) Placed in really cold temperature conditions

84 Intermolecular Forces The forces holding molecules to each other. What phase is strongest? Solids What phase is the weakest? Gases (vapor)

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86 STRONG INTERMOLECULAR FORCES Don t change phase easily High melting points Low vapor pressure Nonvolatile Substances High boiling points High viscosity High surface tension

87 WEAK INTERMOLECULAR FORCES Do change phase easily Low melting points High vapor pressure Volatile Substances Low boiling points Low viscosity Low surface tension

88 Heating Curve Used to show how much enthalpy energy (Heat transfer) is needed to change phase. Enthalpy (heat) of Fusion- energy required to change from solid to liquid Enthalpy (heat) of Vaporization- energy required to change from liquid to vapor

89 Heating Curve

90 Enthalpy of Vaporization Enthalpy of Fusion

91 Phase Diagram Shows how states of matter of a substance are affected by changes in temperature and pressure. Triple Point- point where all three states of matter meet. Critical Point- Point where only the vapor can exist.

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93 Ch Liquids Definite volume no definite shape (takes shape of container) Difficult to compress disorderly arrangement on particles Flowing motion of particles

94 Liquid Properties Viscosity- the resistance of a fluid to flow Thick fluids have high viscosity Ex. Syrup

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97 Liquid Properties Surface Tension- Ability of liquid molecules to hold on to each other. Apparent skin affect Ex. Over filling a liquid in a glass with out the liquid spilling

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100 Liquid Properties Capillary Rise- the tendency of a liquid to rise in a small diameter tube due to the surface tension of the liquid. Used to measure surface tension of a liquid

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102 Hydrogen Bonding Causes water to be very polar. This bonding also causes water to decrease in density and expand as it freezes (increases space between molecules)

103 Ch. 16 Solids Definite shape Definite volume Difficult to compress Orderly arrangement of particles Smallest amount of movement of particles.

104 Metal Solids Solid Structures Crystal structure (repeating patterns) Allotropes (different forms of same element, ex. Carbon) Amorphous (no crystal structure) Glasses, rubber, plastics