3 Gas is one of the three main states of matter Gas particles may be atoms or molecules, depending on the type of substance (ie, element or compound)Gas particles have much more space between them than liquids or solids.Gases are said to be an expanded form of matter, solids and liquids are condensed forms of matter.3
4 General Properties of a Gas Gases do have mass (although it is sometimes difficult to measure).Gases have no definite volume,Gases have no definite shape.Gases are compressible, meaning they can be squeezed into smaller containers, or can expand to fill larger containers.Because gases compress, the density of gases can only be compared under specific conditions.4
5 Some Important GasesOxygen (O2): clear, breathable, supports combustion.Ozone (O3): poisonous, unstable form of oxygenNitrogen (N2): clear, low activity, most abundant gas in the Earth’s atmosphere.Hydrogen (H2): clear, lighter than air, flammable/explosiveCarbon dioxide (CO2): clear, but turns limewater cloudy. Does not support respiration but low toxicity. Heavier than air. Largely responsible for the greenhouse effect (global warming)Sulphur dioxide (SO2): smelly gas. When it combines with oxygen and water vapour it can form H2SO4, responsible for acid rain.5
6 Some Important GasesCarbon monoxide (CO): clear, colourless, but very toxic. It destroys the ability of blood to carry oxygen. About the same density as air.Ammonia (NH3): toxic, strong smell, refrigerant . Very soluble in water, forms a basic solution called ammonia-water (NH4OH) which is found in some cleaners.Freon® or CFC: Non-toxic refrigerant used in air-conditioners & freezers. Freon may catalyze ozone breakdown. The original Freon formula is now banned, but low chlorine versions are still in use.Methane (CH4): flammable gas, slightly lighter than air, produced by decomposition. Found in natural gas. Methane is also a “greenhouse” gas.Helium (He): inert, lighter than air. Used in balloons and in diver’s breathing mixtures.6
7 Acetylene (C2H2): AKA ethene, it is used as a fuel in welding, lanterns and other devices. Propane (C3H8): used as a fuel in barbecues, stoves, lanterns and other devices.Radon (Rn): A noble gas that is usually radioactive. It is heavier than air, and sometimes found in poorly ventilated basements.Neon (Ne) and Xenon (Xe): Noble gases found in fluorescent light tubes, and as insulators inside windows. They glow more brightly than other gases when electrons pass through them. Neon is slightly lighter than air, Xenon is quite a bit heavier.Compressed Air (78% N2, 21% O2): Not actually a pure gas, but a gas mixture that acts much like a pure gas. It is used by scuba divers (at shallow depths), and to run pneumatic tools, and for producing foam materials.7
8 Fun Gases (of no real importance) Nitrous Oxide (N2O)AKA: Laughing gas, Happy gas, Nitro, NOSOnce used as an anaesthetic in dentist offices, this sweet-smelling gas reduces pain sensitivity and causes euphoric sensations. It is an excellent oxidizer, reigniting a glowing splint much like oxygen would. It is used in racing where it is injected into the carburetor to temporarily increase an engine’s horsepower.Sulfur HexafluorideOne of the densest gases in common use. Fun with Sulfur hexafluoride8
9 Match the gas with the problem it causes Gas ProblemCarbon Dioxide Ozone layer depletionCFCs Global WarmingMethane Toxic poisoningCarbon monoxide Noxious smellSulfur dioxide Acid Rain9Next slide: Summary
10 Some Gases Classified by Relative Density Low Density gasesNeutral Density GasesHigh Density gases“lighter than air”<25 g/mol“similar to air” 29±4 g/mol“Denser than air” (>34 g/mol)Testable Property*: Balloon will float in airBalloon drops slowly through airBalloon drops quickly through airExamples:Hydrogen (H2) 2Helium (He) 4Methane (CH4) 16Ammonia (NH3) 17Neon (Ne) 20Hydrogen Fluoride (HF) 21“Cyanide“ (HCN) 27Acetylene (C2H4) 28Nitrogen (N2) 28Carbon monoxide 28Ethane (C2H6) 30Oxygen (O2) 32Fluorine (F2) 38Argon (Ar) 40Carbon dioxide (CO2) 44Propane (C3H8) 44Butane (C4H10) 58Sulphur Hexafluoride (SF6) 146*balloon test: Fill a large, lightweight balloon with the gas, then release it from a height of about 1.8 m in a room with still air. If the gas is lighter than air the balloon will float upwards. If it is close to air, the balloon will fall very slowly. If the gas is heavier than air, the balloon will fall quickly.
11 Some Gases Classified by Chemical Properties Combustible gases(combustion /explosion)Reactive- oxidizing Gases (support combustion)Non-Reactive gasesTestable property:Burning splint produces “pop”Glowing splint reignites, burning splint grows brighterBurning splint is extinguished, glowing splint is dimmedOther properties:Useful as fuelsCause metals and some other materials to corrode or oxidize. Can improve combustion.Can be used to preserve foods by slowing oxidationExamples:Hydrogen (H2)Methane (CH4)Propane (C3H8)Acetylene (C2H4)Oxygen (O2)Fluorine (F2)Chlorine (Cl2)Nitrous Oxide (NO2)Carbon dioxide (CO2)Nitrogen (N2)Argon (Ar)Helium (He)
12 Textbook Assignments Read Chapter 1: pp. 37 to 50 Do the exercises on pages 51 and 52Questions # 1 to 2212
13 Summary: Know the properties of gases Know the features of some important gases, esp:OxygenHydrogenCarbon dioxideKnow the environmental problems associated with some gases, eg.CFC’sSulfur dioxide13
14 Chapter 2 The Kinetic Theory “Moving, moving, moving, Keep those atoms moving...”14
15 Kinetic Theory Overview: The kinetic theory of gases (AKA. kinetic-molecular theory) tries to explain the behavior of gases, and to a lesser extent liquids and solids, based on the concept of moving particles or molecules.
16 The Kinetic Theory of Gases (AKA: The Kinetic Molecular Theory) 2.1Page 54The Kinetic Theory of Gases (AKA: The Kinetic Molecular Theory)The Kinetic Theory of Gases tries to explain the similar behaviours of different gases based on the movement of the particles that compose them.“Kinetic” refers to motion. The idea is that gas particles* are in constant motion.* For simplicity, I usually call the gas particles “molecules”, although in truth, they could include atoms or ions.16
17 RThe Particle ModelNot in textThe Kinetic Theory is part of the Particle Model of matter, which includes the following concepts:All matter is composed of particles (ions, atoms or molecules) which are extremely small and have a varying space between them, depending on their state or phase.Particles of matter may attract or repel each other, and the force of attraction or repulsion depends on the distance that separates them.Particles of matter are always moving.++-17
18 Kinetic Molecular Theory And TemperatureThe absolute temperature of a gas (Kelvins) is directly proportional to the average kinetic energy of its molecules.In other words, when it is cold, molecules move slowly and have lower kinetic energy.When the temperature increases, molecules speed up and have more kinetic energy!18
19 Particle Motion and Phases of Matter 2.1.1Page 54Particle Motion and Phases of MatterRRecall that:In solids, the particles (molecules) are moving relatively slowly. They have low kinetic energyIn liquids, molecules move faster. They have higher kinetic energy.In gases, the particles move fastest, and have high kinetic energy.But, as we will find out later:Heavy particles moving slowly can have the same kinetic energy as light particles moving faster.19
20 Kinetic Theory Model of States LiquidParticles vibrate, move and “flow”, but cohesion (molecular attraction) keeps them close together.GasParticles move freely through container. The wide spacing means molecular attraction is negligible.SolidParticles vibrate but don’t “flow”. Strong molecular attractions keep them in place.20
21 Kinetic Motion of Particles 2.1.1Page 55Kinetic Motion of ParticlesParticles (ie. Molecules) can have 3 types of motion, giving them kinetic energyVibrational kinetic energy (vibrating)Rotational kinetic energy (tumbling)Translational kinetic energy (moving)21
22 Kinetic Theory and Solids & Liquids 2.1.1Page 56Kinetic Theory and Solids & LiquidsWhen it is cold, molecules move slowlyIn solids, they move so slowly that they are held in place and just vibrate (only vibrational energy)In liquids they move a bit faster, and can tumble and flow, but they don’t escape from the attraction of other molecules (more rotational energy, along with a little bit of vibration & translation)In gases they move so fast that they go everywhere in their container (more translational energy, with a little bit of rotation & vibration).22
23 Plasma, the “Fourth State” (extension material) When strongly heated, or exposed to high voltage or radiation, gas atoms may lose some of their electrons. As they capture new electrons, the atoms emit light—they glow. This glowing, gas-like substance is called “plasma”23
24 Kinetic Theory and the Ideal Gas 2.1.3Page 61Kinetic Theory and the Ideal GasAs scientists tried to understand how gas particles relate to the properties of gases, they saw mathematical relationships that very closely, but not perfectly, described the behaviour of many gases.They have developed theories and mathematical laws that describe a hypothetical gas, called “ideal gas.”24
25 2.1.3Page 61To make the physical laws (derived from kinetic equations from physics) work, they had to make five assumptions about how molecules work.Four of these are listed on page 61 of your textbookThe fifth one is not.25
26 Kinetic Theory Hypotheses about an Ideal Gas 2.1.3Page 61Kinetic Theory Hypotheses about an Ideal GasThe particles of an ideal gas are infinitely small, so the size is negligible compared to the volume of the container holding the gas.The particles of an ideal gas are in constant motion, and move in straight lines (until they collide with other particles)The particles of an ideal gas do not exert any attraction or repulsion on each other.The average kinetic energy of the particles is proportional to the absolute temperature.26
27 No Gas is IdealSome of the assumptions on the previous page are clearly not true.Molecules do have a size (albeit very tiny)Particles do exert forces on each other (slightly)As a result, there is no such thing as a perfectly “ideal gas”However, the assumptions are very good approximations of the real particle properties.Real gases behave in a manner very close to “ideal gas”, in fact so close that we can usually assume them to be ideal for the purposes of calculations.27
28 Other “Imaginary Features” of Ideal Gas 2.1.3Page 61Other “Imaginary Features” of Ideal GasAn ideal gas would obey the gas laws at all conditions of temperature and pressureAn ideal gas would never condense into a liquid, nor freeze into a solid.At absolute zero an ideal gas would occupy no space at all.28
29 Please Notice:Not all molecules move at exactly the same speed. The kinetic theory is based on averages of a great many molecules.Even if the molecules are identical and at a uniform temperature, a FEW will be faster than the average, and a FEW will be slower.If there are two different types of molecules, the heavier ones will be slower than the light ones – ON THE AVERAGE! – but there can still be variations. That means SOME heavy molecules may be moving as fast as the slowest of the light ones.Temperature is based on the average (mean) kinetic energy of sextillions of individual molecules.29
30 The mean & mode can help establish “average” molecules The range of kinetic energies can be represented as a sort of “bell curve.” Maxwell’s Velocity Distribution Curve.The mean & mode can help establish “average” moleculesMost molecules“Average”moleculesIncreasing # molecules“Slow”molecules“Fast”MoleculesmodemeanAveragekinetic energyIncreasing kinetic energy30
31 So, Given two different gases at the same temperature… What is the same about them? The AVERAGE kinetic energy is the same.Not the velocity of individual moleculesNot the mass of individual molecules.In fact, the lighter molecules will move fasterEk = mv2 kinetic energy of molecules2So, kinetic energy depends on both the speed (v) and on the mass (m) of the molecules.31
32 Distribution of Particles Around Average Kinetic Energies. Average kinetic energy of moleculesAverage kinetic energy of warmer moleculesAverage kinetic energy of colder moleculesNumber of moleculesSlowerthanaverage moleculesFasterthanaverage moleculesKinetic Energy of molecules(proportional to velocity of molecules)32
33 Kinetic Theory TriviaThe average speed of oxygen molecules at 20°C is 1656km/h.At that speed an oxygen molecule could travel from Montreal to Vancouver in three hours…If it travelled in a straight line.Each air molecule has about 1010 (ten billion) collisions per second10 billion collisions every second means they bounce around a lot!The number of oxygen molecules in a classroom is about:that’s more than there are stars in the universe!The average distance air molecules travel between collisions is about 60nm.m is about the width of a virus.33
34 Videos Kinetic Molecular Basketball Average Kinetic Energies Average Kinetic EnergiesThermo-chemistry lecture on kinetics34
35 AssignmentsRead pages 53 to 61Do Page 62 # 1-11
36 Chapter 2.2 Behaviors of Gases Compressibility Expansion Diffusion and EffusionGraham’s Law
37 2.2.1 Compressibility: 2.2.2 Expansion: Because the distances between particles in a gas is relatively large, gases can be squeezed into a smaller volume.Compressibility makes it possible to store large amounts of a gas compressed into small tanks2.2.2 Expansion:Gases will expand to fill any container they occupy, due to the random motion of the molecules.37
38 2.2.3 DiffusionDiffusion is the tendency for molecules to move from areas of high concentration to areas of lower concentration, until the concentration is uniform. They do this because of the random motion of the molecules.Effusion is the same process, but with the molecules passing through a small hole or barrier38Next slide:
39 Rate of Diffusion or Effusion It has long been known that lighter molecules tend to diffuse faster than heavy ones, since their average velocity is higher, but how much faster? heavy particle light particle39
40 Graham’s Law Internet demo of effusion Thomas Graham (c. 1840) studied effusion (a type of diffusion through a small hole) and proposed the following law:“The rate of diffusion of a gas is inversely proportional to the square root of its molar mass.”In other words, light gas particles will diffuse faster than heavy gas molecules, and there is a math formula to calculate how much faster.Where: v1= rate of gas 1v2= rate of gas 2M1= molar mass of gas1M2=molar mass of gas 240Next slide: Example
41 Graham’s Law Version #1, based on Effusion Rate The relationship between the rate of effusion or diffusion and the molar masses is:Note: See the inversion of the 1 and 2 in the 2nd ratio!Where: v1 is the rate of diffusion of gas 1, in any appropriate rate units*v2 is the rate of diffusion of gas 2, in the same units as gas 1M1 is the molar mass of gas 1M2 is the molar mass of gas 2*Rate units must be an amount over a time for effusion (eg: mL/s or L/min), or a distance over a time for diffusion (eg: cm/min or mm/s)
42 And in my spare time I invented dialysis, which has saved the lives of thousands of kidney patients Thomas Graham ( )Graham derived his law by treating gases as ideal, and applying the kinetic energy formula to them.Ek = ½ mv2All gases have the same kinetic energy at the same temperature,Therefore, mv2 for the first gas = mv2 for the second gas: m1v12 = m2v22.A bit of algebra then gave him his famous law.
43 Graham’s Law Version #2, Based on Effusion Time Sometimes it’s easier to measure the time it takes for a gas to effuse completely, rather than the rate. Graham’s law can be changed for this, but the relationship between time and molar mass is direct as the square root:Note: In this variant law, the relationship is not inverted!Where: t1 is the time it takes for the first gas to effuse completely.t2 is the time it takes for an equal volume of the 2nd gas to effuseM1 is the molar mass of the first gasM2 is the molar mass of the second gas.
44 Example of Graham’s Law: How much faster does He diffuse than N2? MN2=2x14.0=28 g/molNitrogen (N2) has a molar mass of 28.0 g/molHelium (He) has a molar mass of 4.0 g/molThe difference between their diffusion rates is:Notice the reversal of order!So helium diffuses 2.6 times faster than nitrogenMHe=1x4.0=4 g/mol44Next slide: 2.3 Pressure of Gases
45 AssignmentsRead pages 63 to 67Do Questions 1 to 10 on page 68
46 Chapter 2.3 Pressure of Gases What is Pressure Atmospheric Pressure Outer Space (immeasurable)Spaceship 1 (2006)X15 (1963)100 km < kPaChapter 2.3Edge of SpacePressure of GasesWhat is PressureAtmospheric PressureMeasuring Pressure40 km 1 kPaHighest Jet 4 kPa20 km 6 kPa10 km 25 kPaMt Everest 31 kPa5 km 55 kPa46Mr. Smith0 km kPa
47 Pressure Pressure is the force exerted by a gas on a surface. The surface that we measure the pressure on is usually the inside of the gas’s container.Pressure and the Kinetic TheoryGas pressure is caused by billions of particles moving randomly, and striking the sides of the container.Pressure Formula:Pressure = force divided by area47
48 Atmospheric PressureThis is the force of a 100 km high column of air pushing down on us.Standard atmospheric pressure is1.00 atm (atmosphere), or101.3 kPa (kilopascals), or760 Torr (mmHg), or14.7 psi (pounds per square inch)Pressure varies with:Altitude. (lower at high altitude)Weather conditions. (lower on cloudy days)48
50 Measuring Pressure Barometer: measures atmospheric pressure. Two types:Mercury BarometerAneroid BarometerManometer: measures pressure in a container (AKA. Pressure guage)Dial Type: Similar to an aneroid barometerU-Tube: Similar to a mercury barometerPiston type: used in “tire guage”50
51 the Mercury BarometerA tube at least 800 mm long is filled with mercury (the densest liquid) and inverted over a dish that contains mercury.The mercury column will fall until the air pressure can support the mercury.On a sunny day at sea level, the air pressure will support a column of mercury 760 mm high.The column will rise and fall slightly as the weather changes.Mercury barometers are very accurate, but have lost popularity due to the toxicity of mercury.51
52 The Aneroid BarometerIn an aneroid barometer, a chamber containing a partial vacuum will expand and contract in response to changes in air pressureA system of levers and springs converts this into the movement of a dial.
53 Manometers (Pressure Gauges) Manometers work much like barometers, but instead of measuring atmospheric pressure, they measure the pressure difference between the inside and outside of a container.Like barometers they come in mercury and aneroid types. There is also a cheaper “piston” type used in tire gauges, but not in science.Tire gauge(piston manometer)U-tube manometer Pressure gauge(mercury manometer) (aneroid)You Tube manometer
54 Reading U-tube manometers When reading a mercury U-tube manometer, you measure the difference in the heights of the two columns of mercury.If the tube is “closed” then the height (h) is the gas pressure in mmHg. P(mmHg)=h(mmHg)If the tube is “open” and h is positive (the pressure you are measuring is greater than the atmosphere) then you must add atmospheric pressure in mmHg. Pgas(mmHg) = Patm(mmHg)+h(mm)Atm. pressureMust be in mmHg, not cm or kPa!After you finish, you can convert your answer to kPa, or atm. Or whatever.
55 Manometer Examples on a day when the air pressure is 763mmHg (101 Manometer Examples on a day when the air pressure is 763mmHg (101.7 kPa)Closed tube: Pgas(mm Hg)=h (mm Hg)Pgas = h = 4 cm = 40 mm HgPgas =h= 4 cm4 cmOpen: Pgas(mmHg)=P atm(mmHg) +h (mmHg)Pgas = mm Hg =823 mm HgPgas =6 cm Higher6Open: Pgas(mmHg)=P atm(mmHg) -h (mmHg)Pgas = mm Hg =703 mm HgPgas =96 cm Lower
56 AssignmentsRead pages 69 to 73.Do Page 74, Questions 1 to 4.
57 Chapter 2.4 The Simple Gas Laws Other Simple Laws that are a Gas: Boyle’s Law Relates volume & pressureCharles’ Law Relates volume & temperatureGay-Lussac’s Law Relates pressure & temperatureAvogadro’s Law Relates to the number of molesOther Simple Laws that are a Gas:Cole’s Law Relates thinly sliced cabbage to vinegarMurphy’s Law Anything that can go wrong will.Clarke’s Laws Relates possible and impossible57
58 Clarke’s Laws of the impossible* Clarke’s 1st Law: If an elderly and respected science teacher (like me) tells you that something is possible, he is probably right. If he tells you something is impossible, he’s almost certainly wrong.Clarke’s 2nd Law: The only way to find the limits to what is possible is to go beyond them.Clarkes 3rd Law: Any sufficiently advanced technology is indistinguishable from magic.*these are slightly paraphrased, I quote them from memory. They were developed by science fiction writer Arthur C. Clarke
59 Lesson 2.4.1 Boyle’s Law Robert Boyle (1662) For Pressure and Volume“For a given mass of gas at a constant temperature, the volume varies inversely with pressure.”59Next slide: Air in Syringe
60 Robert BoyleBorn: 25 January 1627 Lismore, County Waterford, Ireland Died 31 December 1691 (aged 64) London, EnglandFields: Physics, chemistry; Known for Boyle's Law. Considered to be the founder of modern chemistryInfluences: Robert Carew, Galileo Galilei, Otto von Guericke, Francis BaconInfluenced: Dalton, Lavoisier, Charles, Gay-Lussack, Avogadro.Notable awards: Fellow of the Royal Society60
61 PressureGas pressure is the force placed on the sides of a container by the gas it holdsPressure is caused by the collision of trillions of gas particles against the sides of the containerPressure can be measured many waysStandard PressureAtmospheres (atm) 1 atmKilopascals (kPa)or(N/m2) kPa = N/m2Millibars (mB) mBTorr (torr) or mm mercury 760 torr = 760 mmHgCentimetres of mercury 76 cmHgInches of mercury (inHg) 29.9 inHg (USA only)Pounds per sq. in (psi) psi (USA only)61
62 Example of Boyle’s Law: Air trapped in a syringe If some air is left in a syringe, and the needle removed and sealed, you can measure the amount of force needed to compress the gas to a smaller volume.62Next slide: Inside syringe
63 Inside the syringe… Read- don't copy The harder you press, the smaller the volume of air becomes. Increasing the pressure makes the volume smaller!The original pressure was low, the volume was large. The new pressure is higher, so the volume is small.Click Here for an internet demo using psi (pounds per square inch) instead of kilopascals (1kPa=0.145psi)lowhigh63Next slide: PV
64 This means that:As the volume of a contained gas decreases, the pressure increasesAs the volume of a contained gas increases, the pressure decreasesThis assumes that:no more gas enters or leaves the container, andthat the temperature remains constant.The mathematical formula for this is given on the next slide64Next slide: Example
65 Boyle’s Law Relating Pressure and Volume of a Contained Gas By changing the shape of a gas container, such as a piston cylinder, you can compress or expand the gas. This will change the pressure as follows:Where: P1 is the pressure* of the gas before the container changes shape.P2 is the pressure after, in the same units as P1.V1 is the volume of the gas before the container changes, in L or mLV2 is the volume of the gas after, in the same units as V1*appropriate pressure units include: kPa, mmHg, atm. Usable, but inappropriate units include psi, inHg.
66 Example 1You have 30 mL of air in a syringe at 100 kPa. If you squeeze the syringe so that the air occupies only 10 mL, what will the pressure inside the syringe be?P1 × V1 = P2 × V2, so..100 kPa × 30 mL = ? kPa × 10 mL3000 mL·kPa ÷ 10 mL = 300 kPaThe pressure inside the syringe will be 300 kPa66Next slide: Graph of Boyle’s Law
67 Graph of Boyle’s Law The Pressure-Volume Relationship Boyle’s Law produces an inverse relationship graph.P(kpa) x V(L)100 x 8 = 800200 x 4 = 800Volume (L) 300 x 2.66 = 800400 x 2 = 800500 x 1.6 = 800600 x 1.33 = 800700 x 1.14 = 800800 x 1 = 80067Pressure (kPa) Next slide: Real Life Data
68 Example 2: Real Life Data In an experiment Mr. Taylor and Tracy put weights onto a syringe of air.At the beginning, Mr. Taylor calculated the equivalent of 4 kgf of atmospheric pressure were exerted on the syringe.0+4= 4kg : 29 mL (116)2+4= 6kg : 20 mL (120)4+4=8kg : 15 mL (120)6+4=10kg: 12 mL (120)8+4=12kg: 10.5 mL (126)68Next slide: Boyle’s Law Experiment or skip to: Lesson 2.3 Charles’ Law:
69 Summary: Boyle’s law P1V1=P2V2 The volume of a gas is inversely proportional to its pressureFormula: P1V1=P2V2Graph: Boyle’s law is usually represented by an inverse relationship graph (a curve)P1V1=P2V2Volume (L) 69Pressure (kPa)
71 Assignments on Boyle’s Law Read pages 75 to 79Do questions 1 to 10 on page 97
72 Boyle’s Law Lab Activity Read, Don’t WriteWe will use the weight of a column of mercury to compress and expand air (a gas) sealed in a glass tube.Read the handout for details of the procedure. (Note: You may shorten the procedure section in your report by including and referring to this handout as part of a complete sentence.)You should still write all other report sections (purpose, materials, diagram, observations etc.) in full, as normal.
73 Diagram of Boyle’s Law Apparatus #1. Horizontal#2 Open end up#3 Open end down
74 Collecting DataYou will need to find the length of the mercury column with the tube held horizontal:You also need this atmospheric information:(a) Position of “right”side of mercury___ mm(b) Position of “left” of mercury column(c) Height of mercury column (a) – (b)(c) mm(d) today’s temperature*___ °C(e) today’s barometric pressure (blackboard)(e) mmHg*used to calibrate the barometer, not used in calculations
75 Collecting Data (continued) Data set 1 - Horizontal Tube:(f) Position of “left” side of columnmm(g) Position of closure(h) “volume” of gas (f) – (g)(h) MmData set 2 - Open End Up:(i) Position of bottom of column(j) Position of closure should be same as (g)(k) “volume” of gas (i) – (j)(k) mm
76 Collecting Data (continued) Data set 3 - Open End Down:l) Position of Top of columnmmm) Position of closure should be same as (g)n) “volume” of gas (l) – (m)This concludes the collection of data, now we must process it and calculate the PV (pressure x volume) values at each of the three conditions.
77 Calculations (e) (c) (h) (e)+(c) (k) (e)- (c) (n) Barometric pressure Item (e)Column HeightItem (c)“Pressure”P“Volume”VPVPxVHorizontal(e)(c)(h)Open End Up(e)+(c)(k)Open End Down(e)- (c)(n)Since we are using analogues for pressure & volume, the units don’t matter.
78 Conclusion and Discussion According to Boyle’s law, the PV values should all be identical. In the real world they will not be identical, but they should be very close.Analyze your results. While doing this you should find the percentage similarity between your largest and smallest result (smallest over largest x 100%). This can help you conclude if your results have supported Boyle’s Law or not.Discuss sources of error, and explain if they were significant in your results.Discuss the meaning of Boyle’s law as it relates to this activity.
79 Answers to Boyle’s Law Sheet 1.00 L of a gas at standard temperature and pressure (101 kPa) is compressed to 473 mL. What is the new pressure of the gas?1 markformulaP1 • V1 =P2 • V21 markKnownP1= 101 kPaV1= 1.00x103 mLP2= unknownV2= 473 mL101kPa • 1000 mL = P2 kPa • 473 mLP2 = 101•1000 kPa•mL = kPa473 mL1 mark1 markAnswer: the pressure will be about 214 kilopascals
80 In a thermonuclear device the pressure of 0. 050 L of gas reaches 4 In a thermonuclear device the pressure of L of gas reaches 4.0x108kPa. When the bomb casing explodes, the gas is released into the atmosphere where it reaches a pressure of 1.00x102kPa. What is the volume of the gas after the explosion?formulaP1 • V1 =P2 • V21 markKnownP1= 4.0x108kPaV1= LP2= 1x102kPaV2=unknown1 mark4.0x108kPa • 0.050L = 1x102kPa • V2LV2 = 4x108•0.05 kPa•L = 2.00x105 L1x102kPa1 mark1 markAnswer: there will be 2.00x105Litres (or L) of gas
81 The volume would be 3.33x10-5 Litres synthetic diamonds can be manufactured at pressures of 6.00x104 atm. If we took 2.00L of gas at 1.00 atm and compressed it to 6.00x104 atm, what would the volume be?KnownP1= 1.00 atmV1= 2.00 LP2= 6.0x104 atmV2= unknown1 markFormulaP1V1=P2V21 mark1.00•2.00 = 6.0•104 • V2V2 = 2.00 ÷ 6.0x104V2 = 3.33 x10-5 LorP1=1.01x102kPa,P2=6.06x106kPa.1 markThe volume would be 3.33x10-5 Litres1 mark
82 Divers get the bends if they come up too fast because gas in their blood expands, forming bubbles in their blood. If a diver has L of gas in his blood at a depth of 50m where the pressure is 5.00x103 kPa, then rises to the surface where the pressure is 1.00x102kPa, what will the volume of gas in his blood be? Do you think this will harm the diver?1 markKnownP1=5.00x103 kPaV1= LP2= 1.00x102 kPaV2= UnknownFormulaP1V1=P2V21 mark5.0x103kPa • L = 1x102kPa • V2LV2 = 5x103•0.05 kPa•L = 2.50 L1x102kPa1 markThe sudden appearance of 2½ litres of gas in the diver’s bloodstream could be quite deadly.1 mark
83 Lesson Charles’ LawThe Relationship between Temperature and Volume.“Volume varies directly with Temperature”83Next slide: Jacques Charles
84 Jacques Charles (1787)“The volume of a fixed mass of gas is directly proportional to its temperature (in kelvins) if the pressure on the gas is kept constant”This assumes that the container can expand, so that the pressure of the gas will not rise.Born: November 12, 1746 ( ) Beaugency, OrléanaisDied: April 7, ( ) (aged 76), ParisNationality: FranceFields: physics, mathematics, hot air ballooningInstitutions: Conservatoire des Arts et MétiersNext slide: The Mathematical formula for this law
85 Charles’ Law Relating Volume and Temperature of a Gas If you place a gas in an expandable container, such as a piston or balloon, as you heat the gas its volume will increase, as you cool it the volume will decrease.Where: T1 is Temperature of the gas before it is heated, in kelvins.T2 is Temperature of the gas after it is heated, in kelvinsV1 is the volume of the gas before it was heated, in L or mLV2 is the volume of the gas after it was heated, in the same units.
86 Charles Law EvidenceCharles used cylinders and pistons to study and graph the expansion of gases in response to heat.See the next two slides for diagrams of his apparatus and graphs.Lord Kelvin (William Thompson) used one of Charles’ graphs to discover the value of absolute zero.86Next slide: Diagram of Cylinder & Piston
87 Charles Law Example Piston Cylinder Trapped Gas Click Here for a simulated internet experiment87Next slide: Graph of Charles’ Law
88 Graph of Charles Law Expansion of an “Ideal” Gas Charles discovered the direct relationship6LLord Kelvin traced it back to absolute zero.5LExpansion of an “Ideal” Gas4L3L2LExpansion of most real gasescondensationfreezeLiquid state1L273°CSolid state-250°C-200°C-150°C-100°C-50°C0°C50°C100°C150°C200°C250°C°CNext slide: Example
89 ExampleIf 2 Litres of gas at 27°C are heated in a cylinder, and the piston is allowed to rise so that pressure is kept constant, how much space will the gas take up at 327°C?Convert temperatures to kelvins: 27°C =300k, 327°C = 600kUse Charles’ Law: (see below)Answer: 4 LitresNext slide: Lesson 2.4 Gay Lussac’s Law
90 Standard Temperature & Pressure (STP) Since the volume of a gas can change with pressure and temperature, gases must be compared at a specific temperature and pressure. The long-standing standard for comparing gases is called Standard Temperature and Pressure (STP)Standard Temperature =0°C = 273 KStandard Pressure =101.3 kPa
91 (SATP) Ambient Temperature Ambient Temperature = 25°C = 298 K Some chemists prefer to compare gases at 25°C rather than 0°C. At zero it is freezing, a temperature difficult to maintain inside the lab. This alternate set of conditions is known as Standard Ambient Temperature and Pressure (SATP). Although not widely used, you should be aware of it, and always watch carefully in case a question uses AMBIENT temperature instead of STANDARD temperature.Ambient Temperature = 25°C = 298 KStandard Pressure = kPa(SATP)
92 Comparison Standard and Ambient Conditions Standard Temperature & Pressure (STP)Ambient Temperature & Pressure (SATP)Pressure101.3 kPaTemperature °C0 °C25 °CTemperature KKKMolar Volume22.4 L/mol24.5 L/mol
93 Summary: Charles’ lawThe volume of a gas is directly proportional to its temperatureFormula:Graph: Charles’ law is usually represented by a direct relationship graph (straight line)Video1Volume (L) Absolute zero0°C=273KTemp
94 Charles’ Law Worksheet 1. The temperature inside my fridge is about 4˚C, If I place a balloon in my fridge that initially has a temperature of 22˚C and a volume of 0.50 litres, what will be the volume of the balloon when it is fully cooled? (for simplicity, we will assume the pressure in the balloon remains the same)Data:T1=22˚CT2=4˚CV1=0.50 LTo find:V2= unknownTemperatures must be converted to kelvin=295K=277KSo:V2=V1 x T2 ÷ T1V2=0.5L x 277K295KV2=0.469 Lmultiplydivide94The balloon will have a volume of 0.47 litres
95 Answer: The balloon’s volume will be 0.71 litres A man heats a balloon in the oven. If the balloon has an initial volume of 0.40 L and a temperature of 20.0°C, what will the volume of the balloon be if he heats it to 250°C.DataV1= 0.40LT1= 20°CT2= 250°CV2= ?Convert temperatures to kelvin20+273= 293K, =523k=293 KUse Charles’ Law=523 K0.7139L0.40L x 523 K ÷ 293 K = LAnswer: The balloon’s volume will be 0.71 litres95
96 3. On hot days you may have noticed that potato chip bags seem to inflate. If I have a 250 mL bag at a temperature of 19.0°C and I leave it in my car at a temperature of 60.0°C, what will the new volume of the bag be?Convert temperatures to kelvin19+273= 292K, =333KData:V1=250 mLT1= 19.0°CT2=60.0°CV2= ?Use Charles’ Law=292 K=333 KmL250mL x 333 K ÷ 292 K = mLAnswer: The bag will have a volume of 285mL
97 4. The volume of air in my lungs will be 2.35 litres Although only the answers are shown here, in order to get full marks you need to show all steps of the solution!4. The volume of air in my lungs will be 2.35 litresBe sure to show your known informationChange the temperature to Kelvins and show them.Show the formula you used and your calculationsState the answer clearly.5.6. The temperature is K, which corresponds to 6.70 C. A jacket or sweater would be appropriate clothing for this weather.
98 Charles’ Law Assignments Read pages 80 to 84Do questions 11 to 21 on pages 97 and 98
99 Gay-Lussac’s Law Lesson 2.4.3 For Temperature-Pressure changes. “Pressure varies directly with Temperature”99Next slide:’
100 Joseph Gay-Lussac (1802)“The pressure of a gas is directly proportional to the temperature (in kelvins) if the volume is kept constant.”Born 6 December 1778Saint-Léonard-de-NoblatDied 9 May Saint-Léonard-de-NoblatNationality: FrenchFields: ChemistryKnown for Gay-Lussac's law100100Next slide:’
101 Gay-Lussac’s Law Relating Pressure and Temperature of a Gas Where: P1 is the pressure* of the gas before the temperature change.P2 is the pressure after the temperature change, in the same units.T1 is the temperature of the gas before it changes, in kelvins.T2 is the temperature of the gas after it changes, in kelvins.*appropriate pressure units include: kPa, mmHg, atm.
102 Gay-Lussac’s LawAs the gas in a sealed container that cannot expand is heated, the pressure increases.For calculations, you must use Kelvin temperatures:K=°C+273pressure102102
103 Remove irrelevant fact ExampleA sealed can contains 310 mL of air at room temperature (20°C) and an internal pressure of 100 kPa. If the can is heated to 606 °C what will the internal pressure be?Remove irrelevant factData:P1= 100kPaV1=310 mLT1=20˚CP2=unknownT2=606˚C˚Celsius must be converted to kelvins20˚C = 293 K ˚C = 879 K=293K=879KFormula:multiplydividex = ÷ 293x = 300Answer: the pressure will be 300 kPa103103Next slide: T vs P graph
104 Temperature & Pressure Graph The graph of temperature in Kelvin vs. pressure in kilopascals is a straight line. Like the temperature vs. volume graph, it can be used to find the value of absolute zero.104104
106 Summary: Gay-Lussac’s law The pressure of a gas is directly proportional to its temperatureFormula:Graph: Gay-Lussac’s law is usually represented by an direct relationship graph (straight line)Pressure Absolute zero0°C=273KTemp
107 Assignment on Gay-Lussac’s Law Read pages 85 to 87Answer questions #22 to 30 on page 98
108 Avogadro’s Law Lesson 2.4.4 For amount of gas. “The volume of a gas is directly related to the number of moles of gas”108108Next slide: Lorenzo Romano Amedeo Carlo Avogadro di Quaregna
109 Lorenzo Romano Amedeo Carlo Avogadro di Quaregna “Equal volumes of gas at the same temperature and pressure contain the same number of moles of particles.”Amedeo AvogadroBorn: August 9, 1776Turin, ItalyDied: July 9, 1856Field: PhysicsUniversity of TurinKnown for Avogadro’s hypothesis, Avogadro’s number.
110 You already know most of the facts that relate to Avogadro’s Law: That a mole contains a certain number of particles (6.02 x 1023)That a mole of gas at standard temperature and pressure will occupy 22.4 Litres (24.5 at SATP)The only new thing here, is how changing the amount of gas present will affect pressure or volume.Increasing the amount of gas present will increase the volume of a gas (if it can expand),Increasing the amount of gas present will increase the pressure of a gas (if it is unable to expand).110
111 It’s mostly common sense… If you pump more gas into a balloon, and allow it to expand freely, the volume of the balloon will increase.If you pump more gas into a container that can’t expand, then the pressure inside the container will increase.111
112 Avogadro’s Laws Relating Moles of Gas to Volume or Pressure Where: V1 = volume before, in appropriate volume units.V2 = volume after, in the same volume unitsP1=pressure before, in appropriate pressure units.P2=pressure after, in the same pressure units.n1 = #moles beforen2 = #moles after112
113 Assignments on Avogadro’s Law Read pages 92 to 96Do Questions 31 to 36 on page 98113
114 The General Gas Law and the Ideal Gas Law Lesson 2.5The General Gas Law and the Ideal Gas Law114Next slide:
115 The Combined or General Gas Law The general (or combined) gas law replaces the four simple gas laws. It puts together:Boyle’s LawCharles’ LawGay-Lussac’s LawAvogadro’s LawAdvantages of the Combined Gas Law:It is easier to remember one law than four.It can handle changing more than one variable at a time (eg. Changing both temperature and pressure)= General Gas Law115
116 The General Gas Law Relating all the Simple Laws Together MemorizeWhere: P1 P2 are the pressure of the gas before and after changes.V1, V2 are the volume of the gas before and after changes.T1 T2 are the temperatures, in kelvinsn 1, n2 is the number of moles of the gas.
117 The neat thing about the General gas law is that it can replace the three original gas laws. Just cross out or cover the parts that don’t change, and you have the other laws:If the temperature is constant, then you have Boyle’s law.If, instead, pressure remains constant, you have Charles’ LawAnd finally, if the volume stays constant, then you have Gay-Lussac’s LawMost of the time, the number of moles stays the same, so you can remove moles from the equation.117
118 FYI: Deriving a Formula, The Ideal Gas LawThe Ideal Gas Law is derived from the General Gas Law in several mathematical steps.First, start with the general gas law, including P, V, T, and the amount of gas in moles (n) .FYI: Deriving a Formula,no need to copy all of itNext slide:
119 Remember Standard Temperature & Pressure (STP) Standard Temperature is 0°C or more to the point, 273K = 25°C = 298K)Standard Pressure is kPa (one atmospheric pressure at sea level)At STP one mole of an ideal gas occupies exactly 22.4 Litres = 24.5 L)OK, You should already know this part.If you don't, record it now!
120 The Ideal Gas Law: Calculating the Ideal Gas Constant. We are going to calculate a new constant by substituting in values for P2, V2, T2 and n2At STP we know all the conditions of the gas.Substitute and solve to give us a constantNext slide: R-- The Ideal Gas Constant
121 The Ideal Gas Constant is the proportionality constant that makes the ideal gas law work The Ideal Gas Constant has the symbol RR=8.31 L· kPa / K·molThe Ideal Gas constant is 8.31 litre-kilopascals per kelvin-mole.MemorizeNext slide: Ideal Gas Formula
122 FYI: Deriving a Formula, no need to copy all of itSo, ifThen, by a bit of algebra: P1V1=n1RT1Since we are only using one set of subscripts here, we might as well remove them: PV=nRT
123 The Ideal Gas Law Relating Conditions to the Ideal Gas Constant Where: P=Pressure, in kPaV=Volume, in Litresn= number of moles, in molR= Ideal Gas constant, 8.31 LkPa/KmolT = Temperature, in kelvins
124 The Ideal gas law is best to use when you don’t need a “before and after” situation. Just one set of data (one volume, one pressure, one temperature, one amount of gas)If you know three of the data, you can find the missing one.
125 Sample Problem8.0 g of oxygen gas is at a pressure of 2.0x102 kPa (ie: 200 Kpa w. 2 sig fig) and a temperature of 15°C. How many litres of oxygen are there? Formula: PV = nRTVariables: P=200 kPaV=? (our unknown)= xn= 8.0g ÷ 32 g/mol =0.25 molR=8.31 L·kPa/K·mol (ideal gas constant)T= 15°C = 288K200 x = (0.25)(8.31)(288) , thereforex= (0.25)(8.31)(288) ÷ 200=2.99 LThere are 3.0 L of oxygen (rounded to 2 S.D.)
126 Sample problem8.0 g of oxygen gas is at a pressure of 2.0x102 kPa (ie: 200KPa) and a temperature of 15°C. How many litres of oxygen are there? (assume 2 significant digits)Temperature has been converted to kelvinsData:P=200 kPaV=unknown = Xn= not givenR=8.31 L·kPa/K·molT= 15°C = 288K---m (O2) = 8gM (O2) = 32.0 g/molCalculate the value of n using the mole formula:0.25 mol200 x = (0.25)(8.31)(288) , thereforex= (0.25)(8.31)(288) ÷ 200=2.99 LThere are 3.0 L of oxygen (rounded to 2 S.D.)
127 Sample Problem8.0 g of oxygen gas is at a pressure of 2.0x102 kPa (ie: 200KPa) and a temperature of 15°C. How many litres of oxygen are there? (give answer to 2 significant digits)Data:P = 200 kPaR = 8.31 L·kPa/K·molT = = 288Km(O2)= 8.0 gM(O2)= 32.0 g/moln =To find:VFormula:Work:Next slide: Ideal vs. Real
128 Ideal vs. Real GasesThe gas laws were worked out by assuming that gases are ideal, that is, that they obey the gas laws at all temperatures and pressures. In reality gases will condense or solidify at low temperatures and/or high pressures, at which point they stop behaving like gases. Also, attraction forces between molecules may cause a gas’ behavior to vary slightly from ideal.A gas is ideal if its particles are extremely small (true for most gases), the distance between particles is relatively large (true for most gases near room temperature) and there are no forces of attraction between the particles (not always true)At the temperatures where a substance is a gas, it follows the gas laws closely, but not always perfectly.For our calculations, unless we are told otherwise, we will assume that a gas is behaving ideally. The results will be accurate enough for our purposes!Next slide: Summary
129 Testing if a gas is ideal If you know all the important properties of a gas (its volume, pressure, temperature in kelvin, and the number of moles) substitute them into the ideal gas law, but don’t put in the value of R. Instead, calculate to see if the value of R is close to 8.31, if so, the gas is ideal, or very nearly so. If the calculated value of R is quite different from 8.31 then the gas is far from ideal.
130 ExampleA sample of gas contains 1 mole of particles and occupies 25L., its pressure 100 kPa is and its temperature is 27°C. Is the gas ideal?Convert to kelvins: 27°C+273=300KPV=nRT (ideal gas law formula)100kPa25L=1molR300K, so…R=100kPa25L÷(300K1mol)R=8.33 kPaL /Kmol expected value: 8.31 kPaL /KmolSo the gas is not ideal, but it is fairly close to an ideal gas,It varies from ideal by only 0.24%
131 Gas Laws OverviewWhen using gas laws, remember that temperatures are given in Kelvins (K)Based on absolute zero: –273°CThe three original gas laws can be combined, and also merged with Avogadro’s mole concept to give us the Combined Gas Law.Rearranging the Combined Gas Law and doing a bit of algebra produces the Ideal Gas Law.Substituting in the STP conditions we can find the Ideal Gas Constant.“Ideal gases” are gases that obey the gas laws at all temperatures and pressures. In reality, no gas is perfectly ideal, but most are very close.
132 Gas Laws: Summary R=8.31 Lkpa/Kmol Simple gas laws Boyle’s Law: Charles’ Law:Gay-Lussac’s Law:Combined gas law:Ideal gas law:The ideal gas constant:R=8.31 Lkpa/Kmol
134 Assignments on the Simple Gas Laws Finish Exercises p. 99 #37 to 52
135 Extra Assignments Old text References: Textbook Chapter 10: pp. 221 to 240Student Study Guide pp. 2-4 to 2-11Old Textbook: page 241 # 25 to 30Do these in your assignments folder.Extra practice:Study guide: pp 2.12 to 2.17 # 1 to 22There is an answer key in the back for theseDo these on your own as review
136 Exercise AnswersThe pressure will double, since there is twice as much gas occupying the same space. (I answered this using logic and Avagadro’s hypothesis rather than math. It stands to reason that twice as much gas in the same space will increase the pressure.)The pressure will be four times as high, since the volume is one quarter what it was before: P1V1 = P2V2 so… P1V1 = 4P1 x ¼V1 (again, although you can do it with math, logic works better)The pressure will be one third as great as it was before, since there is three times the volume: P1V1 = P2V2, so = 1/3 P1 x 3V1
137 The gas cannot expand, so it exerts force on its container The gas cannot expand, so it exerts force on its container. As the temperature increases, the gas particles move faster, hitting the container sides more frequently and with more force. This causes greater pressure. You can also explain this using Gay-Lussac’s law; P1/T1 = P2/T2Make sure you use the KELVIN temperatures. The formula is P1/T1 = P2/T2 or 300 kpa/300K = xkPa/100K, so the pressure will be 100 kPaAn ideal gas obeys the gas laws at all temperatures and pressures (no real gas is perfectly ideal. More ideal properties will be discussed in the next section).
138 20) PCO2 = 3.33 kPa, since all the partial pressures will add up to the total pressure ( =33.3)21) Use Boyle’s law: P1V1=P2V2, therefore 91.2kpa4.0L=20.3kpaxL so therefore x=91.2x4÷20.3 the new volume is 17.9 L22) Use Boyle’s law: P1V1=P2V2 ,so x=100kPa6L÷25.3kPa. The new volume will be 23.7L
139 23) Use Charles’Law: V1/T1=V2/T2, convert the temperature from °CK, so -50°C223K and 100 °C373K so… 5L/223K = x/373K so… x=5373÷223.The new volume will be about 8.36 L24) Use Gay-Lussack’s law: P1/T1=P2/T2, don’t forget to change 27°C300K. So… 200kPa/300K=223kPa/x.The new temperature will be 61.5°C(converted from 334.5K)
140 ANSWERS The combined gas laws: (this answer is straight from the lesson)26) Convert the temperatures to kelvin, set up equation, leaving out n1 and n2 (moles don’t change), cross multiply:Answer:The new pressure is kPaMultiply these togetherThen divide by these
141 27) Data given: need to find: m=12g(O2) M(O2) P=52.7kPa V=x L R=8.31LkPa/Kmol n in molT= 25°C T in kelvinFind the number of moles of O2: n=m/MM(O2)=32g/mol so: 12g ÷ 32g/mol = 0.375mol.Convert CK, 25°C+273=298Kformula: PV=nRTso: kPaxL=0.375mol8.31Lk•Pa/Kmol298Kso: x = (0.375 mol 8.31L•kPa 298 K) • __1_K mol kpaAnswer: The volume will be about 17.6 L32g/mol0.375mol298K
142 #28-30, answers (with brief explanation) (see me at lunch if you need more explanation) 28) Litres at STP56 L b) 6.72 L c) 7.84 L(remember: each mole of29) Answer: The pressure will be 1714 kPa(use the formula PV=nRT)30) Answer: The volume will be 16.8 L
143 Dalton’s Law of partial pressures Lesson 2.8Dalton’s Law of partial pressures
144 John DaltonBesides being the founder of modern atomic theory, John Dalton experimented on gases. He was the first to reasonably estimate the composition of the atmosphere at 21% oxygen, 79% NitrogenBorn6 September Eaglesfield, Cumberland, EnglandDied27 July 1844 Manchester, EnglandNotable studentsJames Prescott JouleKnown forAtomic Theory, Law of Multiple Proportions, Dalton's Law of Partial Pressures, DaltonismInfluencesJohn Gough
145 Partial Pressure Many gases are mixtures, eg. Air is 78% nitrogen, 21% Oxygen, 1% other gasesEach gas in a mixture contributes a partial pressure towards the total gas pressure.The total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the individual gases in the mixture.101.3 kPa (Pair) = 79.1 kPa (N2) kPa (O2) kPa(Other)Next slide:
146 Kinetic Theory Connection Hypothesis 3 of the kinetic theory states that gas particles do not attract or repel each other.Dalton established that each type of gas in a mixture behaved independently of the other gases.The pressure of each gas contributes towards the total pressure of the mixture.
147 Dalton’s Law The Law of Partial Pressures of Gases Where: PT is the total pressure of mixed gasesP1 is the pressure of the 1st gasP2 is the pressure of the 2nd gasetc...
148 Variant of Dalton’s Law (used for finding partial pressure of a gas in a mixture) Where: PA=Pressure of gas AnA = moles of gas AnT= total moles of all gasesPT= Total Pressure of all gases
149 Uses of Dalton’s Law Story Don’t copy In the 1960s NASA used the law of partial pressures to reduce the launch weight of their spacecraft. Instead of using air at 101 kPa, they used pure oxygen at 20kPa.Breathing low-pressure pure oxygen gave the astronauts just as much “partial pressure” of oxygen as in normal air.Lower pressure spacecraft reduced the chances of explosive decompression, and it also meant their spacecraft didn’t have to be as strong or heavy as those of the Russians (who used normal air).. This is one of the main reasons the Americans beat the Russians to the moon.
150 Carelessness with pure oxygen, however, lead to the first major tragedy of the American space program…At 20 kPa, pure oxygen is very safe to handle, but at 101 kPa pure oxygen makes everything around it extremely flammable, and capable of burning five times faster than normal.On January 27, 1967, during a pre-launch training exercise, the spacecraft Apollo-1 caught fire. The fire spread instantly, and the crew died before they could open the hatch.
151 Crew of Apollo 1Gus Grissom, Ed White, Roger Chaffee
152 Exercises : Page 113 in new textbook, # 1 to 8 Extra practice (if you haven’t already started):Study guide: pp 2.12 to 2.17 # 1 to 22There is an answer key in the back for theseDo these on your own as review
153 Summary:Dalton’s Law: The total pressure of a gas mixture is the sum of the partial pressures of each gas.PT = P1 + P2 + …Graham’s Law: light molecules diffuse faster than heavy onesAvogadro’s hypothesisA mole of gas occupies 22.4L at STP and contains 6.02x1023 particles
154 Summary of Kinetic Theory Hypotheses (re. Behaviour of gas molecules):1. Gases are made of molecules moving randomly2. Gas molecules are tiny with lots of space between.3. They have elastic collisions (no lost energy).4. Molecules don’t attract or repel each other (much)Results:The kinetic energy of molecules is related to their temperature (hot molecules have more kinetic energy because they move faster)Kinetic theory is based on averages of many molecules (graphed on the Maxwell distribution “bell” curve)Pressure is caused by the collision of molecules with the sides of their containers.Hotter gases and compressed gases have more collisions, therefore greater pressure.
155 Gases are made of particles Particles move randomly! PressureEnergy of a particle:KE = ½ mV 2Pressure is the result of particles colliding with the container walls.P = F /A
156 Assigned Activities References: Practice problems: Read Textbook ppPractice problems:Textbook: p199 #1-3Student study guide: pp to 2-20(practice problems are for self-correction)Assignments (to be collected in your folder):Page 241: all questions from 25 to 34Handout #1: “combined gas law” #52-58Handout #2: “gases & gas laws” 5 questions (on the back.)
157 Answers (sheet 1) 52: The volume of gas will be 36.5 L 53: The temperature will be 908K or 635C54: The volume will be 250 mL or 0.25L55: The pressure will be 251 kPa56: The pressure will stay the same57: The pressure will be 42.2 kPa58: The volume will be 10.2 L
158 Answers (sheet 2) 1: The volume is about 32.5 L 2: The mass is about 1.53 x 10-7 g3: The pressure is about kPa4: The pressure will increase by 168 kPa (tricky: most students say 268kPa, but that’s what it ends at, NOT how much it changes!)5: The total pressure is about 172kPa
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