AP Chemistry - Weebly
AP Chemistry
Unit 5 – The Gaseous State
This unit focuses on the gaseous state. It addresses Topic II (States of Matter) in the College Board's Advanced Placement Chemistry Topic Outline. In particular it focuses on Gasses.
Lesson 1 – Gases and Gas Laws
This lesson is an introduction to the basic equations that govern the behavior of gases. We'll start by reviewing the basic definitions of the three states of matter: solids, liquids, and gases. We'll then focus our attention on gases, and in particular on molecular descriptions of a gas. These molecular descriptions will serve as the basis for our discussion of three important empirical gas laws: Boyle's law, Charles's law, and Avogadro's law. All three empirical gas laws combine into a fundamental equation called the ideal gas law.
Objectives
Distinguish a gas from a solid and a liquid.
List the important physical characteristics of a gas.
Define the term pressure.
Explain gas pressure on the basis of the molecular nature of a gas.
Convert pressure from one unit to another.
State Boyle's law, Charles's law, and Avogadro's law in terms of the variables P, V, n, and T.
Explain Boyle's law, Charles's law, and Avogadro's law on the basis of the molecular nature of a gas.
Apply Boyle's law to calculate P or V under two sets of conditions.
Apply Charles's law to calculate V or T under two sets of conditions.
Apply Avogadro's law to calculate V or n under two sets of conditions.
Explain how the ideal gas law is derived from the empirical gas laws.
Define standard conditions of temperature and pressure.
Manipulate variables in the ideal gas law to solve for unknown terms.
Apply the ideal gas law in numerical calculations involving gas density and molar mass.
Lesson 2 – More about Gases
This lesson starts by looking at the properties of a mixture of gases and see how to apply Dalton's law of partial pressure to characterize the components of a gas mixture. We will look at the postulates of kinetic molecular theory and learn how to apply them to explain the empirical gas laws and the ideal gas law in molecular terms.
Objectives
Define the term "partial pressure."
Write an expression for the total pressure of a mixture of gases in terms of the partial pressures of its components.
Calculate the partial pressure of a gas in a mixture, given the total pressure and the mole fraction of the gas.
Explain the molecular basis for the correction in the pressure of a gas collected over water.
Summarize the key postulates of kinetic molecular theory.
Interpret the empirical gas laws and the ideal gas law on the basis of the postulates of kinetic molecular theory.
Define the terms "effusion" and "diffusion."
Apply Graham's laws of diffusion and effusion to explain the behavior of gases in terms of their molar mass.
Summarize the characteristics of an ideal gas.
Explain the physical conditions under which a gas does not behave in an ideal manner.
Evaluate the terms in the van der Waals equation for real gases in terms of the molecular basis of the correction to the ideal gas law.
Unit 8, Lesson 1: TUTORIAL: The Nature of Gases
Question 1
The breath that we exhale is a mixture of primarily which gases?
A. oxygen, nitrogen, and carbon dioxide
B. oxygen, argon, and carbon dioxide
C. carbon, oxygen, and nitrogen
D. carbon, hydrogen, and nitrogen
E. oxygen, nitrogen, and hydrogen
Question 2
What are the three states of matter?
A. Crystals, gelatin, and suspension
B. Solid, liquid, and gas
C. Solid and liquid only; gas is a vapor.
D. Solid and gas only; liquid is just a mixture of the two.
E. Liquid and gas; solid is a material.
Question 3
Properties of gases include all of the following except:
A. no fixed volume.
B. no fixed shape.
C. no compressibility.
D. low densities relative to solids and liquids.
E. significant thermal expansion.
Question 4
Which of the following statements is the best description of the molecular view of a gas sample?
A. Particles loosely attached to each other
B. Empty spaces between particles that collide at random with each other
C. Empty space between loosely attached groups of particles
D. Empty spaces within a rigid framework
E. Particles that are separated and rarely come in contact with one another
Question 5
What is the definition of pressure?
A. mass/volume
B. meters/sec
C. force/area
D. grams/cm3
E. distance/time
Question 6
What is the cause of a gas sample exerting pressure on its surroundings, such as a container?
A. Weight of the particles on the container
B. Sum of the volumes of all the gas particles in the sample
C. Chemical reactions with the container walls
D. Collisions of particles of gas
E. Energy obtained when the gas vaporized from a liquid
Question 7
The device that is used to measure atmospheric pressure is a:
A. manometer.
B. calorimeter.
C. pressure sensor.
D. digital balance.
E. barometer.
Question 8
What is considered to be normal air pressure at sea level?
A. It varies too much to consider any one value as normal.
B. One atmosphere of pressure is the value.
C. Whatever is measured by the weather service on the first day of each month.
D. It has different measurements on each continent.
E. Pressure can only be measured in a laboratory.
Question 9
The device that is used to measure the pressure of a particular sample of gas is a:
A. manometer.
B. calorimeter.
C. pressure sensor.
D. digital balance.
E. barometer.
Question 10
All of the following are units of pressure except:
A. newton.
B. torr.
C. atmosphere.
D. pascal.
E. millimeters of mercury.
Unit 8 Lesson 1 – Online Exercise: Gas Pressure
Question 1
All gases are the same in that they take the ___ and ___ of their container.
A. shape; volume
B. mass; color
C. density; odor
D. phase; state
E. chemical properties; physical properties
Question 2
Gases can be differentiated from solids and liquids because gases have:
A. no color.
B. a spherical shape.
C. average densities on the order of 10–3 g/mL.
D. reduced chemical reactivity.
E. life-giving properties.
Question 3
Which of the following is not an example of a common substance that is a gas at room temperature?
A. Nitrogen
B. Iodine
C. Oxygen
D. Water vapor
E. Carbon dioxide
Question 4
How is it that your breath can create a bubble in bubble gum, even though the exhaled breath has such low density?
A. Its density is less than the gum.
B. The gum is compressible.
C. The molecules of the exhaled breath collide frequently on the inside walls of the bubble, pushing it out.
D. The molecules of the expelled air chemically react with the molecules of the gum, increasing its mass.
E. The gum has a density very close to that of the gases in the expelled air.
Question 5
Uranium hexafluoride, UF6, is one of the densest gases known. What is its density at STP, given that the standard molar volume is 22.4 L/mole? Use a periodic table as needed.
A. 15.7 g/L
B. 352 g/mol
C. 22.4 L/mol (since one mole of any gas has the same volume).
D. 7880 g/L
E. 15.7 g/mL
Question 6
Which of the following is a characteristic of room temperature gases that makes them significantly different from the other states of matter?
A. Tendency to form ions
B. Same volume at all temperatures
C. Thermal expansion
D. Electrical conductivity
E. More molecules per mL
Question 7
As a sample of gas is heated, its volume increases. Why does this occur?
A. The gas softens when heated.
B. The gas decomposes into more molecules.
C. The entire universe is expanding.
D. Increased speed of motion makes for increased force of collisions with the walls of the container.
E. The lid on the container loosens when warmed.
Question 8
A closed-end manometer has a U-tube with a sample of hydrogen in the closed end. The difference in the height of the columns of mercury is 18.0 cm. What is the pressure of the hydrogen?
A. 18.0 atm
B. 1.82 x 106 pascals
C. 1.80 mmHg
D. 0.237 pascals
E. 180 torr
Question 9
The air in a laboratory was measured with a barometer to be 7.00 x 102 mmHg. What is the laboratory air pressure in atm?
A. 0.700
B. 0.921
C. 0.00691
D. 6.91
E. 700
Question 10
At an international science convention, a sample of gas was labeled 10.13 kPa (kilopascal). How should this pressure be written in atmospheres (atm)?
A. 0.100
B. 0.133
C. 7698.8
D. 101.3
E. 100.0
Unit 8 Lesson 1:Tutorial – Empirical Gas Laws
|Question 1 | |
| | |
| What is the relationship between the pressure on a gas sample and its| |
|volume? | |
|[pic] | |
|A. | |
|Pressure and volume are directly proportional. | |
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|[pic] | |
|B. | |
|Pressure and volume are only indirectly proportional. | |
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|[pic] | |
|C. | |
|Pressure can affect volume, but volume does not affect pressure. | |
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|[pic] | |
|D. | |
|Pressure and volume are both related to temperature but not to each other.| |
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|[pic] | |
|E. | |
|Pressure and volume are inversely proportional. | |
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|Question 2 | |
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| How is the relationship between pressure and temperature stated | |
|mathematically as Boyle's law? ("k" is a constant.) | |
|[pic] | |
|A. | |
|PV = k | |
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|[pic] | |
|B. | |
|P/V = k | |
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|[pic] | |
|C. | |
|V/P = k | |
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|[pic] | |
|D. | |
|Pk = V | |
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|[pic] | |
|E. | |
|Vk = P | |
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|Question 3 | |
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| Which of the following graphs correctly illustrates Boyle's law? | |
|[pic] | |
|A. | |
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|[pic] | |
|B. | |
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|[pic] | |
|C. | |
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|D. | |
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|E. | |
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|Question 4 | |
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| What does the graph of Boyle's law look like if we graph 1/P versus | |
|V? | |
|[pic] | |
|A. | |
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|B. | |
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|C. | |
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|D. | |
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|E. | |
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|Question 5 | |
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| What is the relationship between the absolute temperature of a gas | |
|sample and its volume? | |
|[pic] | |
|A. | |
|Temperature and volume are inversely proportional. | |
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|[pic] | |
|B. | |
|Temperature and volume are not related. | |
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|[pic] | |
|C. | |
|Temperature and volume are directly proportional. | |
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|[pic] | |
|D. | |
|When temperature increases, the volume of the gas decreases. | |
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|[pic] | |
|E. | |
|Temperature and volume are only related when the Celsius scale is used. | |
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|Question 6 | |
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| How is the relationship between temperature and volume represented | |
|mathematically as Charles' law? | |
|[pic] | |
|A. | |
|TV = k | |
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|[pic] | |
|B. | |
|T = V | |
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|[pic] | |
|C. | |
|1/T = V | |
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|[pic] | |
|D. | |
|V = kT | |
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|E. | |
|VT = k | |
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|Question 7 | |
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| If we plot the volume of a gas versus temperature in degrees Celcius | |
|and extrapolate the curve, we see that the x intercept (zero volume) is | |
|not reached until -273°C. What does this temperature represent? | |
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|[pic] | |
|A. | |
|This is the temperature on Venus. | |
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|[pic] | |
|B. | |
|The temperature does not make sense because a substance cannot have zero | |
|volume. | |
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|[pic] | |
|C. | |
|Absolute zero, or zero kelvins. | |
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|[pic] | |
|D. | |
|It is the temperature at which the solid phase appears. | |
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|E. | |
|It is not a temperature; it is a measurement of volume only. | |
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|Question 8 | |
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| If we graph volume versus absolute temperature (in kelvins), we get a| |
|straight line that passes through the origin. What does the slope of this | |
|line represent? | |
|[pic] | |
|A. | |
|The slope is equal to the proportionality constant, k. | |
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|[pic] | |
|B. | |
|It is the rate of increase in temperature over a standard time frame. | |
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|[pic] | |
|C. | |
|It is equal to the standard temperature. | |
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|[pic] | |
|D. | |
|The slope is not representative of anything in particular. | |
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|[pic] | |
|E. | |
|The slope represents the available empty space around the gas molecules. | |
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|Question 9 | |
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| What is the relationship between the amount of gas (number of moles | |
|of gas particles) and the volume of the sample? | |
|[pic] | |
|A. | |
|The number of moles of gas particles is inversely proportional to the | |
|volume of the sample. | |
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|[pic] | |
|B. | |
|The number of gas particles varies directly with the volume. | |
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|[pic] | |
|C. | |
|The number of moles of gas is normally independent of volume. | |
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|[pic] | |
|D. | |
|The number of gas particles is related to volume, only if temperature and | |
|pressure also vary. | |
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|E. | |
|The moles of a gas cannot be determined because the particles are too | |
|small. | |
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|Question 10 | |
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| What is the molar volume of a sample of a gas when collected at | |
|standard conditions (0°C and 1.0 atm pressure)? | |
|[pic] | |
|A. | |
|The molar volume can be determined only when the kelvin temperature is | |
|known. | |
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|[pic] | |
|B. | |
|The molar volume is equal to 1.0 atmosphere. | |
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|[pic] | |
|C. | |
|The molar volume is 273 Kelvin. | |
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|[pic] | |
|D. | |
|The molar volume of a gas cannot be determined. | |
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|[pic] | |
|E. | |
|The molar volume is 22.414 L. | |
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Unit 8, Lesson 1 – TUTORIAL: The Ideal Gas Law
Question 1
What are the parameters of a sample of gas that are all interrelated and affect each other when any one of them changes?
A. mass, density, color, volume, and pressure
B. pressure, volume, number of moles, and temperature
C. shape, hardness, transparency, and odor
D. temperature, pressure, number of moles, and molar mass
E. molar mass, volume, pressure, and temperature
Question 2
What are the three empirical gas laws?
A. The ABC's of gases: Ambercrombie's law, Boyle's law, and Charles law
B. ideal gas law, law of conservation of mass, and law of conservation of gas volume
C. law of definite proportion, law of direct proportion, and law of inverse proportion
D. Boyle's law, Charles' law, and Avogadro's law
E. law of pressure, law of temperature, and law of number of moles
Question 3
When the product of two variables is a constant, they are said to be ____ proportional and when the product of variable and a constant is equal to another variable, the variables are said to be ___ proportional.
A. inversely; directly
B. directly; inversely
C. directly; indirectly
D. indirectly; directly
E. multiplicatively; equally
Question 4
What is the equation for the ideal gas law?
A. V = kT
B. PV = k
C. k = PVnRT
D. PRT = nV
E. PV = nRT
Question 5
What are the values for STP, standard temperature and pressure?
A. 25 K, 760 mmHg
B. 25°C, 1 mmHg
C. 0°C, 1 atm
D. 273 K, 760 atm
E. It varies from one gas to the next, depending on how ideal it is.
Question 6
What are the physical conditions that allow a gas to behave in an ideal way?
A. high temperature, low pressure
B. high density, high volume
C. high temperature, high pressure
D. low temperature, high pressure
E. low thermal expansion, high conformity
Question 7
What are the properties of an "ideal" gas?
A. no odor, no color
B. no volume for individual particles, no attraction between particles
C. no mass for individual particles, no thermal expansion
D. high boiling point, high melting point
E. no pressure, no temperature
Question 8
Which of the following have units that cannot vary when using the ideal gas law?
A. pressure and volume
B. temperature and volume
C. pressure and number of moles
D. mass, density, and temperature
E. temperature and number of moles
Question 9
When experimentally determining the molar mass of a sample of gas using the ideal gas law, which measurement is not needed?
A. mass of container, volume of gas collected
B. room temperature and pressure
C. mass of container with the gas, mass of container without the gas
D. total volume of the container, mass of the water displaced
E. volume of water displaced, temperature of the room
Question 10
How can the ideal gas law be integrated into a stoichiometry problem?
A. It cannot be integrated, it must always be used separately.
B. They both only work when measuring gases.
C. Stoichiometry deals with mole ratios in equations. The ideal gas law can calculate the moles of a gas (n).
D. Stoichiometry, like the ideal gas law, applies only to reactions run under standard laboratory conditions (STP).
E. The ideal gas law converts grams to moles, which can be used in stoichiometry.
Unit 8 Lesson 1 – Online Exercise: Using the Ideal Gas Law
Question 1
A 100.0 mL sample of oxygen was collected at 27.0°C and 836 torr. What would be the volume at 127°C and standard pressure?
A. 68.2 mL
B. 82.5 mL
C. 411 mL
D. 147 mL
E. 121 mL
Question 2
The ideal gas law wasn't originally determined through experiment, but was derived by combining the three empirical gas laws:
PV=k1, V= Tk2, V= nk3
What would be the value of the universal gas constant, R, in terms of the constants in the above equations?
A. k1k2k3
B. (k1k2)/k3
C. (k1)/(k2k3)
D. k1 + k2 + k3
E. (k1)/(k2 + k3)
Question 3
The volume of 1.000 x 103 millimole of gas at 0.000°C and 101.3 kPa is 22,400 cm3. What is the value for R using these units?
A. 1.23 x 103
B. 8.31
C. 0.120
D. 0.0821
E. Its value is undefined when dividing by zero.
Question 4
What is the density in g/L of of fluorine gas at 7.00 x 102 torr and 27.0°C? (molar mass of F2 = 38.0 g/mol)?
A. 1.42
B. 0.0352
C. 1.20 x 104
D. 1.08 x 103
E. 26.7
Question 5
An unknown gas was released from a container that weighed 10.00 g when empty. 224 mL of gas was collected and let stand in a room that was at 1.00 atm and 0.00°C. If the container weighed 10.28 g before the gas was released, what is the molar mass of this gas?
A. 1.00 x 102 g/mol
B. 62.7 g/mol
C. 341 g/mol
D. 32 g/mol
E. 28 g/mol
Question 6
If a chemistry student collected 834 mL of gas at 25°C and 8.00 x 102 torr, how many moles did she collect?
A. 27.3
B. 35.9
C. 0.0359
D. 0.428
E. 26.1
Question 7
25.0 g of dry ice (solid CO2, molar mass = 44.0 g/mol) was dropped into a balloon and put in a freezer at –5.00 °C. In the morning, it was found that it hadn't been cold enough to prevent the dry ice from subliming into gas. It filled the balloon to 15.0 L in size. What was the pressure in the balloon?
A. 633 torr
B. 17.0 atm
C. -0.0155 atm
D. 833 torr
E. 36.7 kPa
Question 8
In a controlled greenhouse at 30.0°C and 1.0 atm, a crop of seedling tomato plants was found to consume 250 L of carbon dioxide gas during photosynthesis. Assuming 100% yield, what mass of glucose (molar mass = 180 g/mol) was produced for the growing plants?
6 CO2 (g) + 6 H2O (l) C6H12O6 (s) + 6 O2 (g)
A. 1.0 x 101 g
B. 1.7 g
C. 5.0 x 101 g
D. 3.0 x 102 g
E. 1.0 x 102 g
Question 9
100.0 g of copper (II) carbonate (molar mass = 123.56 g/mol) was heated until it decomposed completely. The gas was collected and cooled to room temperature and normal pressure, 25°C and 1.00 atm. What volume of carbon dioxide was produced? The reaction is:
CuCO3 (s) CuO (s) + CO2 (g)
A. 2,450 L
B. 1.66 L
C. 19.8 L
D. 0.0261 L
E. 30.2 L
Question 10
12 liters of ammonia were produced by the Haber process. What volume of hydrogen was used during the reaction? The reaction took place at STP (0°C, 1.0 atm).
N2 (g) + 3 H2 (g) 2 NH3 (g)
A. 12 L
B. 3.0 L
C. 18 L
D. 0.80 L
E. 0.54 L
Unit 8 Lesson 2 – Tutorial: Gas Mixtures
Question 1
The ideal gas law can be applied to:
A. only a sample of monoatomic gases such as helium, neon, and argon.
B. only natural gases such as carbon dioxide, oxygen, and butane.
C. only pure samples of gases, consisting of one type of gas.
D. only gases, but they can be pure samples or mixtures.
E. only substances that can exist in the gas phase, including volatile liquids.
Question 2
The partial pressure of a gas is the pressure of:
A. the liquid phase, which has less pressure than the gas.
B. only part of a sample of a pure gas that is collected.
C. the gas itself, since pressure is only part of what makes up the ideal gas law.
D. one gas in a mixture, as if it was measured by itself.
E. the gas at lower temperature, which is predictably at a lower pressure than high temperature.
Question 3
How can the pressure of a mixture of gases be determined?
A. It is equal to the sum of the individual partial pressures.
B. The total pressure is equal to the average pressure of all the gases.
C. The square root of the inverse of the molecular weights is equal to the total pressure.
D. It can be determined only through direct measurement.
E. The pressure of the mixture is approximately equal to the pressure of the individual gas having the highest mole fraction.
Question 4
Dalton's law can be expressed in equation form. Select the choice that most closely matches this relationship.
A. Ptot = (nHe + nO2)RT/V
B. Xtotal = PA + PB + PC + .... Pz
C. PA = nART/V
D. Ptotal = PA + PB + PC + ....Pz
E. Ptotal = Xtotal
Question 5
Dalton's law of partial pressures makes sense according to the concept of ideal gases because:
A. individual particles take up no space.
B. individual particles of any ideal gas do not interact whether they are the same or not.
C. real gases behave in an ideal manner at room temperature.
D. ideal gases can only be mixed by adding individual pressures.
E. ideal gases can be separated into partial pressures more easily than real gases.
Question 6
The number of moles of one substance in a mixture, divided by the total number of moles in the mixture, is known as:
A. partial pressure.
B. gas density.
C. mole fraction.
D. mole proportion.
E. molarity.
Question 7
The partial pressure of a component gas in a mixture of gases can be determined by using what method?
A. Only by separating the gas from the mixture
B. By multiplying the mole fraction of that component by the total pressure of the mixture
C. By measuring only part of the pressure
D. By using a special partial pressure manometer
E. By dividing the total pressure by the number of gases in the mixture
Question 8
In a mixture containing oxygen at a partial pressure of 400.0 torr and nitrogen at a partial pressure of 800.0 torr, what is the mole fraction of the oxygen?
A. 1200.0
B. 2.000
C. 0.6667
D. 0.5000
E. 0.3333
Question 9
When gas is collected over water, the vapor pressure of the evaporated water is mixed in with the gas. Dalton's law is used to determine the pressure of the "dry" gas alone. However, the pressure of the water vapor is rarely measured, yet is easily determined. How is this accomplished?
A. The vapor pressure of water is always constant.
B. The vapor pressure of water is determined by the volume of the container.
C. The vapor pressure of water is related to temperature and can be found in a table.
D. The vapor pressures of all substances are the same.
E. The water vapor is cooled until it condenses.
Question 10
A sample of hydrogen was collected over water at 25.0°C. It was found to have a volume of 250.0 mL and a pressure of 1.03 atm. The vapor pressure of water was found to be 23.8 torr. What is the pressure of the hydrogen at this temperature and volume?
A. 1.06 atm
B. 1.03 atm
C. 783 torr
D. 759 torr
E. 22.8 torr
Unit 8 Lesson 2 - Online Exercise: Dalton: Applying the Law of Partial Pressure
Question 1
A 342.8 mL sample of the gas above a solution of water and dissolved carbon dioxide was collected at 25°C. The sample contained 0.030 moles. What is the partial pressure of the carbon dioxide? (The vapor pressure of water at 25°C = 23.8 mmHg.)
A. 1603 mmHg
B. 0.0313 atm
C. 1626 mmHg
D. 2.14 atm
E. 160.3 mmHg
Question 2
Several gases were collected together and their individual pressures were determined. It was found that the partial pressure of the oxygen in the mixture was 0.40 atm. There was also 0.30 atm of nitrogen gas and a variety of other gases in the mixture. The total pressure of the mixture was 1.5 atm. What is the mole fraction of the oxygen?
A. 0.27
B. 0.18
C. 0.80
D. 0.40
E. 0.47
Question 3
Dry air is typically 75.5% nitrogen by mass, 23.1% oxygen by mass, and 1.3% argon by mass. The remainder of the air is a mixture of gases present at trace amounts. Assuming that air contains only nitrogen, oxygen, and argon, what is the mole fraction of nitrogen in dry air? NOTE: Express the answer as a percentage.
A. 75.5
B. 3.45
C. 78.3
D. 2.70
E. 1.28
Question 4
Dry air is typically 75.5% nitrogen by mass, 23.1% oxygen by mass, and 1.3% argon by mass. This means that the mole fraction of nitrogen in dry air is .783. What is the partial pressure of N2 in dry air, when the atmospheric pressure is exactly 1.00 atm?
A. 0.22 atm
B. 28.0 g/mol
C. 596 torr
D. 78.2 atm
E. 0.78 torr
Question 5
A mixture of three gases has an equal number of moles of xenon and hydrogen, each of which is equal to twice the number of moles of fluorine. The total pressure of the mixture is 1000.0 torr. What is the partial pressure of the fluorine gas?
A. 100.00 torr
B. 200.00 torr
C. 250.00 torr
D. 333.33 torr
E. 400.00 torr
Question 6
In a mixture of gases, the partial pressures of helium, neon, and argon are 150.0 mmHg, 250.0 mmHg and 400.0 mmHg respectively. What is the total pressure of the mixture?
A. 1.050 atm
B. 0.9500 atm
C. 0.5000 atm
D. 0.3125 atm
E. 0.1875 atm
Question 7
There are two gases in a mixture. They are in a 500.0 mL flask at 32°C. It was found that the first gas has a partial pressure of 250.0 torr. The other gas was found to contain 0.0500 moles. What is the mole fraction of the first gas?
A. 0.00657
B. 0.116
C. 2.50
D. 0.152
E. 0.329
Question 8
There is a mixture of 2 gases in which the mole fraction of gas A is equal to 0.20. The pressure of the mixture is 800.0 mmHg. What is the partial pressure of gas B?
A. 0.21 atm
B. 0.80 torr
C. 160 torr
D. 640 torr
E. 800.80 torr
Question 9
When some water was decomposed, 1.0 mole of hydrogen and 0.50 mole of oxygen were collected. The total pressure of the mixture of oxygen and hydrogen was 1.0 atm and the final temperature was 20°C. What is the volume of this sample? 2 H2O 2 H2 (g) + O2 (g)
A. 1.5 moles
B. 1.66 L
C. 12 L
D. 24 L
E. 36 L
Question 10
A 250.0 mL sample of oxygen gas was collected over water at a temperature of 25°C and 750.0 mmHg. How many moles of gas were there in this sample? The vapor pressure of water at 25°C = 23.8 mmHg.
A. 0.0101
B. 0.00976
C. 7.42
D. 0.116
E. 9.76
Unit 8 Lesson 2 – Tutorial: Kinetic Molecular Theory
Question 1
Kinetic energy is the type of energy that is associated with:
A. motion.
B. kinds of gases.
C. electricity.
D. position.
E. electromagnetic radiation.
Question 2
Why are bromine and iodine useful as vapors for study in the laboratory?
A. Color
B. Odor
C. Molecular mass
D. Rapid motion
E. Easy chemical symbols to remember
Question 3
Kinetic Molecular Theory (KMT) deals specifically with:
A. molecular shape.
B. molecular bonding.
C. molecular motion.
D. behavior of ions.
E. similarities of the states of matter.
Question 4
If a collision is considered to be elastic, then after the collision:
A. there is no motion.
B. there is no change in motion.
C. there is no change in total kinetic energy.
D. there is much stretching of the bonds.
E. each molecule moves at exactly the same speed.
Question 5
The postulates of the Kinetic Molecular Theory (KMT) include all of the following except:
A. Gas particles are in constant, random, linear motion.
B. Particles have elastic collisions with each other and with the walls of their container.
C. Gas particles have negligible volume as individual particles.
D. The density of a gas is the mass divided by volume of the sample.
E. The average kinetic energy of the particles is proportional to the absolute temperature of the sample.
Question 6
The Kinetic Molecular Theory (KMT) can be used to explain all of the following except:
A. Dalton's law.
B. Avogadro's law.
C. Boyle's law.
D. Charles' law.
E. Beer's law.
Question 7
Which of these statements explains how the Kinetic Molecular Theory (KMT) can be used to explain Dalton's law of partial pressures?
A. The total pressure depends on the total number of collisions inside the container, and is unrelated to the chemical nature of the gas particles themselves.
B. The increase in average kinetic energy means that the particles will collide with more force.
C. The increased number of particles results in an increase in the volume.
D. The increase in frequency and strength of molecular collisions results in an increased volume in order to keep the temperature the same.
E. The particles collide with each other as well as with the walls of the container, so the measured pressure is only a partial measurement of the total pressure present in the container.
Question 8
The equation KE = (1/2)mv2 for kinetic energy:
A. works only for large objects.
B. isn't compatible with quantum mechanics.
C. can be applied to molecules of a gas sample.
D. is useful only over a narrow range of temperatures.
E. can be used in place of Avogadro's law.
Question 9
Ideal gas particles:
A. have no odor and no color.
B. have no mass or volume.
C. have no density and do not form bonds.
D. have no mutual attraction and no individual volume.
E. do not fill their containers or take their shapes.
Question 10
The conditions under which real gases deviate most from ideal gases are:
A. high temperature and high humidity.
B. low temperature and high pressure.
C. low pressure and low molecular mass.
D. high temperature and low pressure.
E. high number of moles and high volume.
Unit 8 Lesson 2 – Online Exercise: The Behavior of Gases
Question 1
The van der Waals equation adjusts the ideal gas law by raising the pressure and lowering the volume. This is needed because:
A. the volume of a gas shouldn't include the volume of the gas molecules themselves.
B. the pressure of a gas needs to be taken as an average.
C. the volume of a gas has to be adjusted for standard temperature.
D. the pressure of a gas needs to be calibrated to the average at sea level.
E. the volume of a gas should't include the thickness of the walls of the container.
Question 2
The van der Waals equation for real gases:
is often used for gases that deviate significantly from ideal, such as chlorine. The values for van der Waals constants for chlorine are: a = 6.490 (atm · L2/mol2) and b = 0.05620 L/mol. If exactly one mole of chlorine fills a 20.00 L container at 27.0oC, what is the exact pressure of the gas sample?
A. 1.251 atm
B. 1.230 atm
C. 1.218 atm
D. 1.211 atm
E. 1.203 atm
Question 3
The equation: v = shows us that:
A. the heavier a gas is, the slower it moves.
B. the hotter a gas is, the slower it moves.
C. the heavier a gas is, the faster it moves.
D. the cooler a gas is, the faster it moves.
E. the velocity of a gas can only be calculated with complex mathematics.
Question 4
We can use the equation
to determine:
A. rates of diffusion only.
B. qualitative comparison of velocities only.
C. time of effusion only.
D. molar mass of an unknown gas only.
E. all of the above.
Question 5
If two samples of helium and argon gas are at the same temperature, what is the ratio of velocities of the atoms of helium and argon? (VHe/VAr)
A. 40
B. 10
C. 3.2
D. 0.50
E. 0.25
Question 6
It takes 20.0 seconds for a sample of helium to effuse from a container. About how long will it take for a similar sample of nitrogen to effuse from the same container at the same temperature?
A. 140 sec
B. 20 sec
C. 53 sec
D. 2.6 sec
E. More information is needed.
Question 7
A sample of hydrogen requires 1 minute, 20 seconds to to effuse from a container. A similar sample of an unknown gas takes 5 minutes, 20 seconds to effuse. What is the identity of the unknown gas?
A. He
B. Ne
C. Ar
D. O2
E. Cl2
Question 8
Gaseous ammonia (NH3) and hydrogen chloride gas (HCl) are introduced at opposite ends of a tube at the same time. After a short wait, a cloud of white smoke of ammonium chloride is noticed in the tube, marking the point where the two gases met as they diffused. The tube is 100.0 cm long. About how far in from the ammonia end did the gases meet?
A. 31.8 cm
B. 40.6 cm
C. 46.4 cm
D. 59.4 cm
E. 68.2 cm
Question 9
Which of these statements explains how the kinetic molecular theory can be used to explain Charles's law?
A. The total pressure depends on the total number of collisions inside the container, and is unrelated to the chemical nature of the gas particles themselves.
B. The increase in average kinetic energy means that the particles will be colliding with more force. Volume must increase if pressure is to be kept constant.
C. The increased number of particles results in an increase in the number of collisions.
D. The increase in frequency and strength of molecular collisions results in an increased volume in order to keep temperature the same.
E. The increase in volume creates more space between particles and causes them to collide with less frequency.
Question 10
Which of these statements explains how the kinetic molecular theory can be used to explain Boyle's law?
A. The total pressure depends on the total number of collisions inside the container, and is unrelated to the chemical nature of the gas particles themselves.
B. The increase in average kinetic energy means that the particles will be colliding with more force.
C. The increased number of particles results in an increase in the number of collisions.
D. The increase in frequency of molecular collisions resulting from a decreased volume will show an increase in overall pressure.
E. The increase in the number of collisions per second and an increase in average force per collision will cause the volume to expand.
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