AP Chemistry Chapter 5 Gases Ch

Total Page:16

File Type:pdf, Size:1020Kb

AP Chemistry Chapter 5 Gases Ch AP Chemistry Chapter 5 Gases Ch. 5 Gases Agenda 23 September 2017 Per 3 & 5 ○ 5.1 Pressure ○ 5.2 The Gas Laws ○ 5.3 The Ideal Gas Law ○ 5.4 Gas Stoichiometry ○ 5.5 Dalton’s Law of Partial Pressures ○ 5.6 The Kinetic Molecular Theory of Gases ○ 5.7 Effusion and Diffusion ○ 5.8 Real Gases ○ 5.9 Characteristics of Several Real Gases ○ 5.10 Chemistry in the Atmosphere 5.1 Units of Pressure 760 mm Hg = 760 Torr = 1 atm Pressure = Force = Newton = pascal (Pa) Area m2 5.2 Gas Laws of Boyle, Charles and Avogadro PV = k Slope = k V = k P y = k(1/P) + 0 5.2 Gas Laws of Boyle - use the actual data from Boyle’s Experiment Table 5.1 And Desmos to plot Volume vs. Pressure And then Pressure vs 1/V. And PV vs. V And PV vs. P Boyle’s Law 3 centuries later? Boyle’s law only holds precisely at very low pressures.Measurements at higher pressures reveal PV is not constant, but varies as pressure varies. Deviations slight at pressures close to 1 atm we assume gases obey Boyle’s law in our calcs. A gas that strictly obeys Boyle’s Law is called an Ideal gas. Extrapolate (extend the line beyond the experimental points) back to zero pressure to find “ideal” value for k, 22.41 L atm Look Extrapolatefamiliar? (extend the line beyond the experimental points) back to zero pressure to find “ideal” value for k, 22.41 L atm The volume of a gas at constant pressure increases linearly with the Charles’ Law temperature of the gas. All the gases extrapolate to zero volume at the same temperature. The slopes are different because the samples contain different numbers of moles of gas. V = bT 0K is called absolute zero Question. A sample of gas is heated from 50℃ to 100℃. What happens to the volume? Assume pressure remains constant. Question. A sample of gas is heated from 50℃ to 100℃. What happens to the volume? Assume pressure remains constant. V = bT V/T = constant V1 = V2 T1 T2 Question. A sample of gas is heated from 50℃ to 100℃. What happens to the volume? Assume pressure remains constant. V = bT ℃ V/T = constant T1 = 50 + 273 = 323K ℃ V1 = V2 T2 = 100 + 273 = 373K T1 T2 Question. A sample of gas is heated from 50℃ to 100℃. What happens to the volume? Assume pressure remains constant. V = bT ℃ V/T = constant T1 = 50 + 323 = K ℃ V1 = V2 T2 = 100 + 273 = 373K 323K 373K Question. A sample of gas is heated from 50℃ to 100℃. What happens to the volume? Assume pressure remains constant. V = bT V/T = constant 373KV1 = V2 1.15 V1 = V2 323K Avogadro’s Law V = an Where V is the volume of a gas, n is the number of moles of gas particles, and a is a proportionality constant. For a gas at constant temperature and pressure, the volume is directly proportional to the number of moles of gas. 5.3 The Ideal Gas Law PV = nRT Equation of state for a gas. R is the combined proportionality constant called the universal gas constant R is 0.08206 Latm/Kmol n = number of moles of gas Empirical equation, based on experimental data 5.3 The Ideal Gas Law PV = nRT Equation of state for a gas. A gas that obeys this equation is said to behave “ideally.” - Expresses behavior that real gases approach at low pressure and high temperatures. - Ideal gas is a hypothetical substance - Only minimal errors result from assuming ideal behavior for most gases for situations we encounter Question 6 moles of hydrogen gas in a container. 3 moles of oxygen gas in a separate container. The two gases are removed from their containers and placed into a third container where they react. If all three containers have the same volume, and all the gases are at the same temperature, which container has the lowest pressure? Sketch the situation: Use PV = nRT P = n (RT/V) P=6(Constant) P=3(Constant) P=6(Constant) Question If all three containers have the same volume, and all the gases are at the same temperature, which container has the lowest pressure? Container with oxygen in it before the reaction has the lowest pressure. 5.4 Gas Stoichiometry 1 mole of an ideal gas at 0℃ (273.2K) and 1atm. PV = nRT V = nRT/P = (1.000mol)(0.08206 L atm)(273.2K) Kmol V = 22.42 L is the molar volume Remember 0℃ (273.2K) and 1atm is called STP, standard temperature and pressure. 5.4 Gas Stoichiometry What would the volume of 1 mole of an ideal gas be at a normal room temp. (22℃) and 1atm? PV = nRT 5.4 Gas Stoichiometry What would the volume of 1 mole of an ideal gas be at a normal room temp. (22℃) and 1atm? PV = nRT V = nRT/P = (1.000mol)(0.08206 L atm)(273.2 + 22)K Kmol V = 24.22 L Molar Mass of a Gas PV = nRT n = number of moles of gas = grams of gas Molar mass n = m/ molar mass P = m RT = density RT Molar mass V Molar Mass Molar Mass of a Gas P = m RT = density RT Molar mass V Molar Mass Or Molar mass = density RT P Question If 100.0g of Oxygen gas at STP is put into a balloon, what would the volume of the balloon be? Question If 100.0g of Oxygen gas at STP is put into a balloon, what would the volume of the balloon be? O2 Molar mass = 2(16.00) = 32.00 g/mol PV = mass RT Molar mass V = 100g(0.08206Latm mol-1 K-1) (273.2K) = 32.00 gmol-1 1 atm Question If 100.0g of Oxygen gas at STP is put into a balloon, what would the volume of the balloon be? O2 Molar mass = 2(16.00) = 32.00 g/mol PV = mass RT Molar mass V =100(0.08206 L)(273.2) = 70.05L 32.00 Question A certain mass of neon gas is contained in a rigid steel container. The same mass of Argon gas is added to this container. What will happen to the pressure in the container, assuming that the temperature stays constant? Question A certain mass of neon gas is contained in a rigid steel container. The same mass of Argon gas is added to this container. What will happen to the pressure in the container, assuming that the temperature stays constant? Say there is 1 mole of Neon, around 20 g. If we add 20g of Argon 20g x 1mol = 0.5 mol 39.95g New P = nRT/V = 1.5 (RT/V) 5.5 Dalton’s Law of Partial Pressures For a mixture of gases in a container, the total pressure exerted is the sum of the pressures that each of the gases would exert if it were alone. P total = P1 + P2 + P3 + … P = nRT/V the total number of moles of particles is what is important not identity of gas 5.5 Dalton’s Law of Partial Pressures P = nRT/V the total number of moles of particles is what is important not identity of gas Note this is telling us that for ideal gases, the volume of the individual gas molecules is not important in deciding the pressure exerted, and forces between particles do not influence the pressure either. Question 67 (p.221) Consider the flasks in the following diagram. What are the final partial pressures of H2 and N2 after the stopcock between the two flasks is opened? (Assume the final volume is 3.00L.) What is the total pressure (in torr)? 0.200 atm 475 torr 5.6 Kinetic Molecular Theory of Gases Postulates 1. The particles are so small compared with the distances between them that the volume of the individual particles can be assumed to be negligible (zero). 5.6 Kinetic Molecular Theory of Gases Postulates 2. The particles are in constant motion. The collisions of the particles with the walls of the container are the cause of the pressure exerted by the gas. 5.6 Kinetic Molecular Theory of Gases Postulates 3. The particles are assumed to exert no forces on each other; they are assumed neither to attract not to repel each other. 5.6 Kinetic Molecular Theory of Gases Postulates 4. The average kinetic energy of a collection of gas particles is assumed to be directly proportional to the Kelvin temperature of the gas. 5.6 Kinetic Molecular Theory of Gases How well do the predictions of the KMT (model) fit the experimental observations? Pressure and Volume P = (nRT) 1/V KMT says if volume decreased, particles will collide with container more often, P increases. 5.6 Kinetic Molecular Theory of Gases How well do the predictions of the KMT (model) fit the experimental observations? Pressure and Temperature P = (nR/V) T KMT says if temp. increased, particles will collide with container more often as they speed up, P increases. 5.6 Kinetic Molecular Theory of Gases How well do the predictions of the KMT (model) fit the experimental observations? Volume and Temperature V = (nR/P) T KMT says if temp. increased, particles will collide with container more often as they speed up, and so only way to keep P constant is to increase volume. 5.6 Kinetic Molecular Theory of Gases How well do the predictions of the KMT (model) fit the experimental observations? Volume and number of moles V = (RT/P) n KMT says if number of moles of gas increases, particles will collide with container more often, and so only way to keep P constant is to increase volume.
Recommended publications
  • THE SOLUBILITY of GASES in LIQUIDS Introductory Information C
    THE SOLUBILITY OF GASES IN LIQUIDS Introductory Information C. L. Young, R. Battino, and H. L. Clever INTRODUCTION The Solubility Data Project aims to make a comprehensive search of the literature for data on the solubility of gases, liquids and solids in liquids. Data of suitable accuracy are compiled into data sheets set out in a uniform format. The data for each system are evaluated and where data of sufficient accuracy are available values are recommended and in some cases a smoothing equation is given to represent the variation of solubility with pressure and/or temperature. A text giving an evaluation and recommended values and the compiled data sheets are published on consecutive pages. The following paper by E. Wilhelm gives a rigorous thermodynamic treatment on the solubility of gases in liquids. DEFINITION OF GAS SOLUBILITY The distinction between vapor-liquid equilibria and the solubility of gases in liquids is arbitrary. It is generally accepted that the equilibrium set up at 300K between a typical gas such as argon and a liquid such as water is gas-liquid solubility whereas the equilibrium set up between hexane and cyclohexane at 350K is an example of vapor-liquid equilibrium. However, the distinction between gas-liquid solubility and vapor-liquid equilibrium is often not so clear. The equilibria set up between methane and propane above the critical temperature of methane and below the criti­ cal temperature of propane may be classed as vapor-liquid equilibrium or as gas-liquid solubility depending on the particular range of pressure considered and the particular worker concerned.
    [Show full text]
  • THE SOLUBILITY of GASES in LIQUIDS INTRODUCTION the Solubility Data Project Aims to Make a Comprehensive Search of the Lit- Erat
    THE SOLUBILITY OF GASES IN LIQUIDS R. Battino, H. L. Clever and C. L. Young INTRODUCTION The Solubility Data Project aims to make a comprehensive search of the lit­ erature for data on the solubility of gases, liquids and solids in liquids. Data of suitable accuracy are compiled into data sheets set out in a uni­ form format. The data for each system are evaluated and where data of suf­ ficient accuracy are available values recommended and in some cases a smoothing equation suggested to represent the variation of solubility with pressure and/or temperature. A text giving an evaluation and recommended values and the compiled data sheets are pUblished on consecutive pages. DEFINITION OF GAS SOLUBILITY The distinction between vapor-liquid equilibria and the solUbility of gases in liquids is arbitrary. It is generally accepted that the equilibrium set up at 300K between a typical gas such as argon and a liquid such as water is gas liquid solubility whereas the equilibrium set up between hexane and cyclohexane at 350K is an example of vapor-liquid equilibrium. However, the distinction between gas-liquid solUbility and vapor-liquid equilibrium is often not so clear. The equilibria set up between methane and propane above the critical temperature of methane and below the critical temperature of propane may be classed as vapor-liquid equilibrium or as gas-liquid solu­ bility depending on the particular range of pressure considered and the par­ ticular worker concerned. The difficulty partly stems from our inability to rigorously distinguish between a gas, a vapor, and a liquid, which has been discussed in numerous textbooks.
    [Show full text]
  • CEE 370 Environmental Engineering Principles Henry's
    CEE 370 Lecture #7 9/18/2019 Updated: 18 September 2019 Print version CEE 370 Environmental Engineering Principles Lecture #7 Environmental Chemistry V: Thermodynamics, Henry’s Law, Acids-bases II Reading: Mihelcic & Zimmerman, Chapter 3 Davis & Masten, Chapter 2 Mihelcic, Chapt 3 David Reckhow CEE 370 L#7 1 Henry’s Law Henry's Law states that the amount of a gas that dissolves into a liquid is proportional to the partial pressure that gas exerts on the surface of the liquid. In equation form, that is: C AH = K p A where, CA = concentration of A, [mol/L] or [mg/L] KH = equilibrium constant (often called Henry's Law constant), [mol/L-atm] or [mg/L-atm] pA = partial pressure of A, [atm] David Reckhow CEE 370 L#7 2 Lecture #7 Dave Reckhow 1 CEE 370 Lecture #7 9/18/2019 Henry’s Law Constants Reaction Name Kh, mol/L-atm pKh = -log Kh -2 CO2(g) _ CO2(aq) Carbon 3.41 x 10 1.47 dioxide NH3(g) _ NH3(aq) Ammonia 57.6 -1.76 -1 H2S(g) _ H2S(aq) Hydrogen 1.02 x 10 0.99 sulfide -3 CH4(g) _ CH4(aq) Methane 1.50 x 10 2.82 -3 O2(g) _ O2(aq) Oxygen 1.26 x 10 2.90 David Reckhow CEE 370 L#7 3 Example: Solubility of O2 in Water Background Although the atmosphere we breathe is comprised of approximately 20.9 percent oxygen, oxygen is only slightly soluble in water. In addition, the solubility decreases as the temperature increases.
    [Show full text]
  • Ideal Gasses Is Known As the Ideal Gas Law
    ESCI 341 – Atmospheric Thermodynamics Lesson 4 –Ideal Gases References: An Introduction to Atmospheric Thermodynamics, Tsonis Introduction to Theoretical Meteorology, Hess Physical Chemistry (4th edition), Levine Thermodynamics and an Introduction to Thermostatistics, Callen IDEAL GASES An ideal gas is a gas with the following properties: There are no intermolecular forces, except during collisions. All collisions are elastic. The individual gas molecules have no volume (they behave like point masses). The equation of state for ideal gasses is known as the ideal gas law. The ideal gas law was discovered empirically, but can also be derived theoretically. The form we are most familiar with, pV nRT . Ideal Gas Law (1) R has a value of 8.3145 J-mol1-K1, and n is the number of moles (not molecules). A true ideal gas would be monatomic, meaning each molecule is comprised of a single atom. Real gasses in the atmosphere, such as O2 and N2, are diatomic, and some gasses such as CO2 and O3 are triatomic. Real atmospheric gasses have rotational and vibrational kinetic energy, in addition to translational kinetic energy. Even though the gasses that make up the atmosphere aren’t monatomic, they still closely obey the ideal gas law at the pressures and temperatures encountered in the atmosphere, so we can still use the ideal gas law. FORM OF IDEAL GAS LAW MOST USED BY METEOROLOGISTS In meteorology we use a modified form of the ideal gas law. We first divide (1) by volume to get n p RT . V we then multiply the RHS top and bottom by the molecular weight of the gas, M, to get Mn R p T .
    [Show full text]
  • Masteringphysics: Assignmen
    MasteringPhysics: Assignment Print View http://session.masteringphysics.com/myct/assignmentPrint... Manage this Assignment: Chapter 18 Due: 12:00am on Saturday, July 3, 2010 Note: You will receive no credit for late submissions. To learn more, read your instructor's Grading Policy A Law for Scuba Divers Description: Find the increase in the concentration of air in a scuba diver's lungs. Find the number of moles of air exhaled. Also, consider the isothermal expansion of air as a freediver diver surfaces. Identify the proper pV graph. Associated medical conditions discussed. SCUBA is an acronym for self-contained underwater breathing apparatus. Scuba diving has become an increasingly popular sport, but it requires training and certification owing to its many dangers. In this problem you will explore the biophysics that underlies the two main conditions that may result from diving in an incorrect or unsafe manner. While underwater, a scuba diver must breathe compressed air to compensate for the increased underwater pressure. There are a couple of reasons for this: 1. If the air were not at the same pressure as the water, the pipe carrying the air might close off or collapse under the external water pressure. 2. Compressed air is also more concentrated than the air we normally breathe, so the diver can afford to breathe more shallowly and less often. A mechanical device called a regulator dispenses air at the proper (higher than atmospheric) pressure so that the diver can inhale. Part A Suppose Gabor, a scuba diver, is at a depth of . Assume that: 1. The air pressure in his air tract is the same as the net water pressure at this depth.
    [Show full text]
  • Chapter 14: Gases
    418-451_Ch14-866418 5/9/06 6:19 PM Page 418 CHAPTER 14 Gases Chemistry 2.a, 3.c, 3.d, 3.e, 4.a, 4.c, 4.d, 4.e, 4.f, 4.g, 4.h I&E 1.b, 1.c, 1.d What You’ll Learn ▲ You will use gas laws to cal- culate how pressure, tem- perature, volume, and number of moles of a gas will change when one or more of these variables is altered. ▲ You will compare properties of real and ideal gases. ▲ You will apply the gas laws and Avogadro’s principle to chemical equations. Why It’s Important From barbecuing on a gas grill to taking a ride in a hot-air balloon, many activities involve gases. It is important to be able to predict what effect changes in pressure, temperature, volume, or amount, will have on the properties and behavior of a gas. Visit the Glencoe Chemistry Web site at chemistrymc.com to find links about gases. Firefighters breathe air that has been compressed into tanks that they can wear on their backs. 418 Chapter 14 418-451_Ch14-866418 5/9/06 6:19 PM Page 419 DISCOVERY LAB More Than Just Hot Air Chemistry 4.a, 4.c I&E 1.d ow does a temperature change affect the air in Ha balloon? Safety Precautions Always wear goggles to protect eyes from broken balloons. Procedure 1. Inflate a round balloon and tie it closed. 2. Fill the bucket about half full of cold water and add ice. 3. Use a string to measure the circumference of the balloon.
    [Show full text]
  • Gases Boyles Law Amontons Law Charles Law ∴ Combined Gas
    Gas Physics FICK’S LAW Adolf Fick, 1858 Properties of “Ideal” Gases Fick’s law of diffusion of a gas across a fluid membrane: • Gases are composed of molecules whose size is negligible compared to the average distance between them. Rate of diffusion = KA(P2–P1)/D – Gas has a low density because its molecules are spread apart over a large volume. • Molecules move in random lines in all directions and at various speeds. Wherein: – The forces of attraction or repulsion between two molecules in a gas are very weak or negligible, except when they collide. K = a temperature-dependent diffusion constant. – When molecules collide with one another, no kinetic energy is lost. A = the surface area available for diffusion. • The average kinetic energy of a molecule is proportional to the absolute temperature. (P2–P1) = The difference in concentraon (paral • Gases are easily expandable and compressible (unlike solids and liquids). pressure) of the gas across the membrane. • A gas will completely fill whatever container it is in. D = the distance over which diffusion must take place. – i.e., container volume = gas volume • Gases have a measurement of pressure (force exerted per unit area of surface). – Units: 1 atmosphere [atm] (≈ 1 bar) = 760 mmHg (= 760 torr) = 14.7 psi = 0.1 MPa Boyles Law Amontons Law • Gas pressure (P) is inversely proportional to gas volume (V) • Gas pressure (P) is directly proportional to absolute gas • P ∝ 1/V temperature (T in °K) ∴ ↑V→↓P ↑P→↓V ↓V →↑P ↓P →↑V • 0°C = 273°K • P ∝ T • ∴ P V =P V (if no gas is added/lost and temperature
    [Show full text]
  • The Behavior and Applications of Gases
    329670_ch10.qxd 07/20/2006 11:32 AM Page 394 The Behavior and Applications Contents and Selected Applications 10.1 The Nature of Gases of Gases 10.2 Production of Hydrogen and the Meaning of Pressure 10.3 Mixtures of Gases—Dalton’s Law and Food Packaging Chemical Encounters: Dalton’s Law and Food Packaging 10.4 The Gas Laws—Relating the Behavior of Gases to Key Properties Chemical Encounters: Balloons and Ozone Analysis 10.5 The Ideal Gas Equation 10.6 Applications of the Ideal Gas Equation Earth’s atmosphere is a very thin layer of gases (only 560 km thick) that surrounds Chemical Encounters: Automobile Air Bags Chemical Encounters: Acetylene the planet. Although the atmosphere makes up only a small part of our planet, 10.7 Kinetic-Molecular Theory it is vital to our survival. 10.8 Effusion and Diffusion 10.9 Industrialization: A Wonderful, Yet Cautionary, Tale Chemical Encounters: Ozone Chemical Encounters: The Greenhouse Effect 394 329670_ch10.qxd 6/19/06 1:45 PM Page 395 Our planet is surrounded by a relatively thin layer known as an atmosphere. It supplies all living organisms on Earth with breathable air, serves as the vehicle to deliver rain to our crops, and protects us from harmful ultraviolet rays from the Sun. Although it does contain solid particles, such as dust and smoke, and liquid droplets, such as sea TABLE 10.1 Composition of Dry Air spray and clouds, much of the atmosphere we live in is Percent Parts per Million actually a mixture of the gases shown in Table 10.1.
    [Show full text]
  • The Ideal Gas Law
    The Ideal Gas Law Of the three phases of matter, solid liquid and Putting all these dependencies together, we can gas, the simplest to understand is gas. This is write an equation for the pressure of the gas in because the particles that make up a gas, atoms the box. and molecules, simply rattle around like ping pong balls in a vigorously shaken box. = 푛푅푇 푃 We see that pressure is 푉proportional to the temperature of the gas, and to , the number 푃 of moles of gas particles. It’s also inversely 푇 푛 proportional to , the volume of the box. is the constant of proportionality. We call it the 푉 푅 universal gas constant. The equation is usually written this way. = This is called the Ideal푃푉 Gas푛푅푇 Law. The remarkable thing about this equation is that it holds for all gases, whether it’s air, hydrogen The balls bounce against the inside of the box, gas, carbon dioxide, or whatever. It even holds imparting tiny impulses so that together they for the gas that makes up the Sun and all other exert a force on each wall. The amount of force stars, and the gas between stars that comprises per unit area is what we call pressure. most normal matter in the universe. The pressure that’s exerted (whether it’s ping The Number of Gas Particles in a Gallon of Air pong balls or gas particles) is proportional the Let’s use the Ideal Gas Law to make a number of particles in the box. calculation: the number of gas particles in a It’s also dependent on the speed that each gallon of air at the surface of the Earth.
    [Show full text]
  • Ideal Gas Law
    Ideal Gas Law September 2, 2014 Thermodynamics deals with internal transformations of the energy of a system and exchanges of energy between that system and its environment. A thermodynamic system refers to a speciied collection of matter. Such a system is said to be closed in no mass is exchanged with its surround- ings and open otherwise. The air parcel that will serve as our system is in principle closed. In practice, mass can be exchanged with its surroundings through entrainment and mixing across the system’s country, which is the control surface. A thermodynamic state of a system is defined by the various properties characterizing it. Two types of thermodynamic properties characterize the state of a system. A property that does not depend on the mass of the system is said to be intensive (a specific property), and extensive otherwise. Intensive properties are denoted with lower case, and extensive with upper case. An intensive property can be derived from an extensive property by dividing by mass z = Z=m. In general, describing the thermodynamic state of a system requires us to specify all of its properties. However, that requirement can be relaxes for gases an other substances at normal tem- peratures and pressures. An ideal gas is a pure substance whose thermodynamic state is uniquely determined by any two intensive properties, which are then referred to as state variables. That means the z3 = g(z1; z2). The ideal gas law is the equation of state for an ideal gas. 1. Ideal Gas Law It is convenient to express the amount of a gas as the number of moles n.
    [Show full text]
  • Ideal Gas Law Problems
    Ideal Gas Law Name ___________ 1) Given the following sets of values, calculate the unknown quantity. a) P = 1.01 atm V = ? n = 0.00831 mol T = 25°C b) P = ? V= 0.602 L n = 0.00801 mol T = 311 K 2) At what temperature would 2.10 moles of N2 gas have a pressure of 1.25 atm and in a 25.0 L tank? 3) When filling a weather balloon with gas you have to consider that the gas will expand greatly as it rises and the pressure decreases. Let’s say you put about 10.0 moles of He gas into a balloon that can inflate to hold 5000.0L. Currently, the balloon is not full because of the high pressure on the ground. What is the pressure when the balloon rises to a point where the temperature is -10.0°C and the balloon has completely filled with the gas. 4) What volume is occupied by 5.03 g of O2 at 28°C and a pressure of 0.998atm? 5) Calculate the pressure in a 212 Liter tank containing 23.3 kg of argon gas at 25°C? Answers: 1a) 0.20 L 1b) 0.340 atm 2) 181 K 3) 0.043 atm 4) 3.9 L 5) 67.3 atm Using the Ideal Gas Equation in Changing or Constant Environmental Conditions 1) If you were to take a volleyball scuba diving with you what would be its new volume if it started at the surface with a volume of 2.00L, under a pressure of 752.0 mmHg and a temperature of 20.0°C? On your dive you take it to a place where the pressure is 2943 mmHg, and the temperature is 0.245°C.
    [Show full text]
  • Boyle's Law: Registering the Pressure Changes of Gas Held in a Closed
    TM DataHub Boyle’s Law: Registering the Pressure Changes of Gas Held in a Closed Container at a Constant Temperature, in Relation to Variations in Volume (Teacher’s Guide) © 2012 WARD’S Science. v.11/12 For technical assistance, All Rights Reserved call WARD’S at 1-800-962-2660 OVERVIEW Students will investigate the effect of volume changes on the pressure inside a syringe with a fixed amount of air at a constant temperature. They will measure the air pressure and then proceed to build a graph plotting their results in order to analyze them. MATERIALS NEEDED Ward’s DataHub USB connector cable* Luer-Lock syringe 60 mL Plastic tube * – The USB connector cable is not needed if you are using a Bluetooth enabled device. NUMBER OF USES This demonstration can be performed repeatedly. © 2012 WARD’S Science. v.11/12 1 For technical assistance, All Rights Reserved Teacher’s Guide – Boyle’s Law call WARD’S at 1-800-962-2660 FRAMEWORK FOR K-12 SCIENCE EDUCATION © 2012 * The Dimension I practices listed below are called out as bold words throughout the activity. Asking questions (for science) and defining Use mathematics and computational thinking problems (for engineering) Constructing explanations (for science) and designing Developing and using models solutions (for engineering) Practices Planning and carrying out investigations Engaging in argument from evidence Dimension 1 Analyzing and interpreting data Obtaining, evaluating, and communicating information Science and Engineering Science and Engineering Patterns Energy and matter: Flows,
    [Show full text]