# 1.which pressure is different from the others?

Unit H Review—The Behavior of Gases

Summary

Gases have different physical and chemical properties. All gases have properties that can be explained by the Ideal Gas Model of the Kinetic Molecular Theory. This theory states: Gas molecules have no volume compared to the total space they occupy. There are no attractions between gas molecules. Gas molecules are constantly moving and colliding.  All gas molecules have the same average kinetic energy (AKE) at the same temperature.

Gas pressure is a result of moving gas molecules colliding with the walls of the container. Such pressure can be measured using pressure gauges and pressure sensors. Air pressure, specifically, is measured with a barometer.

The gas laws are a series of mathematical equations, in agreement with the Ideal Gas Model, that can be used to make predictions of pressure, volume, temperature, and moles of molecules.

Real gases do behave very closely (within a few percent) of what the Ideal Gas Model predicts at temperatures and pressures normally encountered in a typical high school laboratory.  At low temperatures and high pressures, however, real gases can be made to condense to liquids. This is because real gas molecules do occupy some space and do have some attraction for each other. The nature of these attractions forms a major portion of the content of the next unit.

 Gas Law Equation Gas Law Equation Boyle’s P1V1 = P2V2 (at constant T and n) Avogadro’s V1  =  V2 n1      n2 (at constant P and T) Charles’ V1  =  V2 T1     T2 (at constant P and n) Ideal PV = nRT (R = 8.314 kPaL mol-1 · K-1) Combined P1V1  =  P2V2 T1         T2 (at constant n) Dalton’s Ptotal = P1 + P2 + P3 + …

Review Questions:

Multiple-Choice:

1.      Which pressure is different from the others?

A.        790 mmHg

B.        105 kPa

C.        1.04 atm

D.        Choices A, B, and C are all the same pressure.

2.      As the temperature of a gas increases, the kinetic energy of its particles

A.        increases.

B.        decreases.

C.        remains the same.

3.      If the volume of a balloon expands, and the temperature remains constant, the pressure on the balloon

A.        increases.

B.        decreases.

C.        remains the same.

4.      If the temperature of the air inside an automobile tire increases, but the tire does not expand, the pressure of the air inside the tire

A.        increases.

B.        decreases.

C.        remains the same.

5.      A partially inflated weather balloon is released.  If the air temperature is constant, but the air pressure around the balloon drops as the balloon rises, the volume of the weather balloon

A.        increases.

B.        decreases.

C.        remains the same.

6.      A chemist has a certain volume of gas in a balloon.  If the volume of gas decreases overnight, but the pressure remains constant, the temperature of the gas

A.        increases.

B.        decreases.

C.        remains the same.

7.      All gases deviate from ideal gas behavior, particularly at high pressures and low

temperatures.

A.        True

B.        False

8.      Temperatures below absolute zero are not possible.

A.        True

B.        False

9.      For an ideal gas  is a constant.

A.        True

B.        False

10.  In a mixture of gases, each gas behaves independently of the other gases in the mixture.

A.        True

B.        False

11.  Which temperature scale provides a direct measure of the average kinetic energy of a substance?

A.        Celsius

B.        Kelvin

C.        Fahrenheit

D.        Reamur

12.  What is the name of a device used to measure gas pressure?

A.        Thermometer

B.        Vaporometer

C.        Calorimeter

D.        Barometer

13.  Which law describes the relationship between the volume and temperature of a gas?

A.        Boyle’s Law

B.        Dalton’s Law

C.        Charles’ Law

D.        Gay-Lussac’s Law

14.  According to Avogadro’s Principle, equal ____ of gas at the same temperature and pressure contain equal numbers of particles.

A.        masses                         B.        volumes

C.        samples                        D.        areas

15.  The individual molecules of gas at STP

A.        never collide.

B.        all have the same speed.

C.        travel at different speeds.

D.        have a constant speed.

16.  If kinetic energy is added to a sample of gas in a rigid container, which of the following takes place?

A.        The temperature increases.

B.        The pressure increases.

C.        The molecules of gas move faster.

D.        All the above occur.

17.  If two different gases have the same kinetic energy,

A.        the gas with more mass has a greater velocity.

B.        the gas with more mass has a lower velocity.

C.        the gases have the same velocity, regardless of mass.

D.        the gas with more mass has a higher temperature.

18.  Gases deviate from ideal behavior because gas particles

A.        are in continuous motion.

B.        move randomly in straight lines.

C.        have different amounts of kinetic energy.

D.        have some attraction for each other.

19.  When a closed tank of air is heated, the density of the air (neglecting any expansion of the tank)

A.        increases.

B.        decreases.

C.        remains the same.

20.  Which of the following is NOT one of the postulates of the Kinetic-Molecular Theory of gases?

A.        Gas molecules are in a state of constant, random motion.

B.         The diameter of a gas molecule is large when compared with the distance between gas molecules.

C.        There is no force of attraction between gas molecules.

D.        The temperature of a gas is a reflection of the average kinetic energy of the gas.

E.         Choices A, B, C, and D are all postulates of the Kinetic-Molecular Theory.

Problems:

21.  59.7 cm3 of helium gas at 95.0 kPa is expanded to 100.0 cm3 at constant temperature.  What is the new pressure on the gas?

22.  Oxygen gas is collected over water at a temperature of 25.0°C and a pressure of 746.8 mmHg.  What is the pressure due to dry oxygen gas?

23.  What pressure will be exerted by 7.99 g of butane, C4H10, if the gas is in a container measuring 1.50 L and is at 27°C?

24.  5.18 L of nitrogen gas at 76.0°C are cooled to 6.0°C at constant pressure.  Calculate the new volume of the gas.

25.  What is the molar mass of a gas if 372.0 mL has a mass of 0.800 g at 99.8°C and 106.6 kPa?

26.  What volume of CO2(g) is produced when 1.05 kg of ethane, C2H6(g), is burned?  The temperature is 24.0°C and the pressure is 745.0 mmHg.

2 C2H6(g)  +  7 O2(g)  ¾®  4 CO2(g)  +  6 H2O(g)

Unit I Review—Liquids, Solids and Phase Changes (optional material in red)

Summary:

There are not only intermolecular attractions between molecules in the gas phase, but those forces of attraction are extremely important for liquids and solids.

The Kinetic-Molecular Theory is fundamental to the understanding of matter. The Kelvin temperature scale is an indicator of relative molecular motions, with the particles having no motion at absolute zero.

Intermolecular forces are called van der Waals forces: London dispersion forces and dipole-dipole forces. A special case of dipole-dipole forces is hydrogen bonding, which have higher than expected melting and boiling points, enhanced solubility, unique molecular shapes, etc.

Ionic bonding is the result of a 3-dimensional array of positive and negative ions attracting each other, resulting in very high melting and boiling points.

Metallic bonding occurs because the valence electrons in metal atoms are free to move to any of the empty orbitals in the crystal and, therefore, belong to the whole crystal rather than just between two atoms.

Substances that have network covalent bonding are generally very hard, strong, brittle, and have extremely high melting and boiling points.  Most do not conduct electricity.

In order to undergo a phase change, the attractions between whole particles (atoms, ions, or molecules) in the substance must overcome interparticle attractions (van der Waals, metallic bonds, ionic bonds or covalent bonds).

Vocabulary:

 absolute zero deposition dipole-dipole forces hydrogen bond hydrogen bond intermolecular forces interparticle attractions isoelectronic lattice energy London dispersion forces normal boiling point sublimation van der Waals forces vapor vapor pressure vaporization

Review Questions:

Multiple-Choice:

1.      Substances that are liquids at room temperature or below are

A.        ionic.

B.        nonpolar covalent molecular.

C.        metallic.

D.        covalent network.

2.      Crystals such as diamonds (very hard, high melting point, nonconductors) are classified as

A.        ionic crystals.

B.        covalent molecular crystals.

C.        covalent network crystals.

D.        metallic crystals.

3.      Metallic crystals characteristically have

A.        good electric conductivity.

B.        great hardness.

C.        low melting points.

D.        brittleness.

4.      Because of the hydrogen bonds in water, the hydrogen atom of one water molecule may be

A.        weakly attracted to the oxygen of a second water molecule.

B.        weakly attracted to the hydrogen of a second water molecule.

C.        strongly attracted to the second hydrogen of its own molecule.

D.        strongly attracted to the oxygen of a second water molecule.

5.      As ice is heated from a lower temperature towards its melting point, the hydrogen bonds

A.        get stronger.

B.        stretch.

C.        increase in number.

D.        cause the formation of hexagonal patterns.

6.      What is thought to cause dispersion forces?

A.        attraction between ions

B.        motion of electrons

C.        differences in electronegativity

D.        the formation of hexagonal patterns

7.      Why is hydrogen bonding only possible with hydrogen?

A.         because hydrogen is the only atom whose nucleus is not shielded by electrons when it is involved in a covalent bond

B.        because hydrogen is the only atom that is the same size as an oxygen atom

C.        because hydrogen has the highest electronegativity of any element in the periodic table

8.      What is the basis of a metallic bond?

A.        the attraction of metal ions for mobile electrons

B.        the attraction between neutral metal atoms

C.        the neutralization of protons by electrons

D.        the attraction of oppositely charged ions

9.      What occurs during the dissolving of an ionic crystal?

A.        Ions separate from molecules.

B.        Molecules surround ions.

C.        Molecules bind covalently to molecules.

D.        Ionic compounds are formed.

10.  Why are two nonpolar substances able to dissolve each other?

A.        There is no repulsive force between them.

B.        They combine to produce a polar substance.

C.        There is no attractive force between them.

D.        Nonpolar molecules cannot dissolve in each other.

11.  Chlorine is a gas, bromine is a liquid, and iodine a solid because of differences in the strength of their

A.        hydrogen bonds.

B.        dispersion forces.

C.        dipole interactions.

D.        polar bonds.

12.  The high surface  tension of water is due to the

A.        small size of water molecules.

B.        high kinetic energy of water molecules.

C.        hydrogen bonding between water molecules.

D.        covalent bonds in water molecules.

13.  It has been said, “There are forces of attraction between molecules in all chemical systems.”  Which of the following experimental observations supports this statement?

A.        Gases can be condensed to form liquids.

B.        Solids are very difficult to compress.

C.        Liquids have an indefinite shape.

D.        Not all solids are ionic crystals.

14.  Ionic solids

A.        are soft and have low melting points.

B.        melt to form liquids that conduct electricity.

C.        are malleable.

D.        conduct electricity.

15.  Solid sodium metal and molten sodium chloride conduct electricity because both contain

A.        mobile electrons.

B.        mobile ions.

C.        active metals, which are good conductors.

D.        mobile charged particles.

For questions #16 to 26, write the letter of the bond or attractive force, chosen from the list below, that is most closely associated with that compound or phrase.

 A.                ionic bond B.                 network covalent bonds C.                 dipole-dipole attractions D.                hydrogen bonds E.                 metallic bond F.                  covalent bonds G.                dispersion forces

16.  Are weak enough to permit solid iodine to sublime readily upon heating.

17.  Al(s)

18.  Bond noble gas atoms in the liquid phase.

19.  Responsible for the extremely high melting point of diamond (above 3500º C).

20.  Link the atoms in a molecule of a diatomic gaseous element.

21.  Allows methanol to be infinitely soluble in ethanol.

22.  SiO2(s)

23.  Positive ions immersed in a “sea of mobile (delocalized) electrons.”

24.  This substance will not conduct as a solid, but will conduct as a liquid or in solution.

25.  OCl2

26.  (diagram)

27.  Which of the following processes requires the least energy?

A.        Breaking the bond between Na+ and Cl in NaCl.

B.        Breaking the bond between H and Cl in HCl.

C.        Separating two CO molecules.

D.        Separating two H2O molecules.

28.  Which of the following transformations is sublimation?

A.   Solid ® Gas

B.   Gas ® Solid

C.   Liquid ® Solid

D.   Solid ® Liquid

29.  Using intermolecular force theory, explain why a substance will change from a gas to a liquid if the temperature is lowered sufficiently.

30.  In terms of electron mobility and electronegativity, explain why Na(s) is a very good conductor of electric current while NaCl(s) is a nonconductor.

Exercises: Phase Changes (Use after Section I.35)

1.      Why is H2O(g) usually called water vapor?

2.      What is the difference between evaporation and boiling?

3.      How would it be possible to boil water at room temperature?

4.      When water is heated to boil, tiny bubbles form on the bottom surface.  What gas is inside the bubbles a) initially and b) at full boil?

5.      a) If we have no frost or built up ice in a freezer and introduce water vapour into the freezer (by opening it on a humid day) what will happen?
b) Conversely, I once put a unwrapped snowball in my mother’s freezer in January and by July it was almost shriveled in size to a small pea. What happened?.

6.      People will often refer to the boiling point of water as 100°C.  What is actually meant by this?

7.      Some non-chemists have argued that the ocean must be less than 10 000 years old based on the following logic:
a) water carrying salts flows to the ocean from rivers running over mineral rich rocks
b) this water and mineral salt solution flows into the sea
c) only water evaporates from the oceans (the salt is left behind)
d) thus the salt concentration of the ocean should increase yearly and the measured salinity of the ocean is much less than if the ocean has been here for millions and billions of years.
Using LeChatelier’s principle and the reversibility of dissolving and salt crystallization explain how to respond to this charge.

8.      Consider the following reversible reaction:

N2O4 (g) à 2 NO2 (g)

If we add in more NO2 we’ll force the reaction further to the _____. Explain using LeChatelier’s Principle. What would happen if added more N2O4?