What type of intermolecular forces is active in the liquid state of each of the

You seal a container half-filled with water. Which best describes what occurs in the container? a. Water evaporates until the air becomes saturated with water vapor; at this point, no more water evaporates. b. Water evaporates until the air becomes overly saturated (supersaturated) with water, and most of this water recondenses; this cycle continues until a certain amount of water vapor is present, and then the cycle ceases. c. The water does not evaporate because the container is sealed. d. Water evaporates, and then water evaporates and recondenses simultaneously and continuously. e. The water evaporates until it is eventually all in vapor form. Justify your choice and for choices you did not pick, explain what is wrong with them.

Explain the following: You add 100 mL of water to a 500-mL round-bottomed flask and heat the water until it is boiling. You remove the heat and stopper the flask, and the boiling stops. You then run cool water over the neck of the flask, and the boiling begins again. It seems as though you are boiling water by cooling it.

Is it possible for the dispersion forces in a particular substance to be stronger than hydrogen-bonding forces in another substance? Explain your answer.

Does the nature of intermolecular forces change when a substance goes from a solid to a liquid, or from a liquid to a gas? What causes a substance to undergo a phase change?

How does vapor pressure change with changing temperature? Explain.

What occurs when the vapor pressure of a liquid is equal to atmospheric pressure? Explain.

What is the vapor pressure of water at 100 C? How do you know?

How do the following physical properties depend on the strength of intermolecular forces? Explain. a. melting point b. boiling point c. vapor pressure

Look at Figure 14.2. Why doesnt temperature increase continuously over time? That is, why does the temperature stay constant for periods of time?

Which are stronger, intermolecular or intramolecular forces for a given molecule? What observation(s) have you made that supports this position? Explain.

Why does water evaporate at all?

Sketch a microscopic picture of water and distinguish between intramolecular bonds and intermolecular forces. Which correspond to the bonds we draw in Lewis structures?

Which has the stronger intermolecular forces: N2 or H2O? Explain.

Which gas would behave more ideally at the same conditions of pressure and temperature: CO or N2? Why?

You have seen that the water molecule has a bent shape and therefore is a polar molecule. This accounts for many of waters interesting properties. What if the water molecule were linear? How would this affect the properties of water? How would life be different?

True or false? Methane (CH4) is more likely to form stronger hydrogen bonding than is water because each methane molecule has twice as many hydrogen atoms. Provide a concise explanation of hydrogen bonding to go with your answer

Why should it make sense that N2 exists as a gas? Given your answer, how is it possible to make liquid nitrogen? Explain why lowering the temperature works.

White phosphorus and sulfur both are called molecular solids even though each is made of only phosphorus and sulfur, respectively. How can they be considered molecular solids? If this is true, why isnt diamond (which is made up only of carbon) a molecular solid?

Why is it incorrect to use the term molecule of NaCl but correct to use the term molecule of H2O? Is the term molecule of diamond correct? Explain.

Which would you predict should be larger for a given substance: Hvap or Hfus? Explain why

In the diagram below, which lines represent the hydrogen bonds? a. The dotted lines between the hydrogen atoms of one water molecule and the oxygen atoms of a different water molecule. b. The solid lines between a hydrogen atom and oxygen atom in the same water molecule. c. Both the solid lines and dotted lines represent the hydrogen bonds. d. There are no hydrogen bonds represented in the diagram.

Use the heating/cooling curve below to answer the following questions. a. What is the freezing point of the liquid? b. What is the boiling point of the liquid? c. Which is greater: the head of fusion or the heat of vaporization? Explain.

Assume the two-dimensional structure of an ionic compound, MxAy, is Temperature (C) Heat added 0 20 40 60 80 100 120 140 160 H H O H H O H O H O H H O H H O H H H O H What is the empirical formula of this ionic compound?

What are London dispersion forces, and how do they arise?

What type of intermolecular forces is active in the liquid state of each of the following substances? a. Ne b. CO c. CH3OH d. Cl2

Discuss the types of intermolecular forces acting in the liquid state of each of the following substances. a. Xe b. NH3 c. F2 d. ICl

The boiling points of the noble gas elements are listed below. Comment on the trend in the boiling points. Why do the boiling points vary in this manner? He 272 C Kr 152.3 C Ne 245.9 C Xe 107.1 C Ar 185.7 C Rn 61.8 C

The heats of fusion of three substances are listed below. Explain the trend this list reflects. HI 2.87 kJ/mol HBr 2.41 kJ/mol HCl 1.99 kJ/mol

When dry ammonia gas (NH3) is bubbled into a 125-mL sample of water, the volume of the sample (initially, at least) decreases slightly. Suggest a reason for this.

When 50 mL of liquid water at 25 C is added to 50 mL of ethanol (ethyl alcohol), also at 25 C, the combined volume of the mixture is considerably less than 100 mL. Give a possible explanation.

What is evaporation? What is condensation? Which of these processes is endothermic and which is exothermic?

If youve ever opened a bottle of rubbing alcohol or other solvent on a warm day, you may have heard a little whoosh as the vapor that had built up above the liquid escapes. Describe on a microscopic basis how a vapor pressure builds up in a closed container above a liquid. What processes in the container give rise to this phenomenon?

What do we mean by a dynamic equilibrium? Describe how the development of a vapor pressure above a liquid represents such an equilibrium.

Consider Figure 14.10. Imagine you are talking to a friend who has not taken any science courses, and explain how the figure demonstrates the concept of vapor pressure and enables it to be measured.

Which substance in each pair would be expected to have a lower boiling point? Explain your reasoning. a. CH3OH or CH3CH2CH2OH b. CH3CH3 or CH3CH2OH c. H2O or CH4

Which substance in each pair would be expected to show the largest vapor pressure at a given temperature? Explain your reasoning. a. H2O(l) or HF(l) b. CH3OCH3(l) or CH3CH2OH(l) c. CH3OH(l) or CH3SH(l)

Although water and ammonia differ in molar mass by only one unit, the boiling point of water is over 100 C higher than that of ammonia. What forces in liquid water that do not exist in liquid ammonia could account for this observation?

Two molecules that contain the same number of each kind of atom but that have different molecular structures are said to be isomers of each other. For example, both ethyl alcohol and dimethyl ether (shown below) have the formula C2H6O and are isomers. Based on considerations of intermolecular forces, which substance would you expect to be more volatile? Which would you expect to have the higher boiling point? Explain. dimethyl ether ethyl alcohol CH3OOOCH3 CH3OCH2OOH

What are crystalline solids? What kind of microscopic structure do such solids have? How is this microscopic structure reflected in the macroscopic appearance of such solids?

On the basis of the smaller units that make up the crystals, cite three types of crystalline solids. For each type of crystalline solid, give an example of a substance that forms that type of solid.

How do ionic solids differ in structure from molecular solids? What are the fundamental particles in each? Give two examples of each type of solid and indicate the individual particles that make up the solids in each of your examples.

A common prank on college campuses is to switch the salt and sugar on dining hall tables, which is usually easy because the substances look so much alike. Yet, despite the similarity in their appearance, these two substances differ greatly in their properties, since one is a molecular solid and the other is an ionic solid. How do the properties differ and why?

Ionic solids are generally considerably harder than most molecular solids. Explain.

Although crystals of table salt (sodium chloride) and table sugar (sucrose) look very similar to the naked eye, the melting point of sucrose (186 C) is several hundred degrees less than the melting point of sodium chloride (801 C). Explain.

The forces holding together a molecular solid are much (stronger/weaker) than the forces between particles in an ionic solid.

Explain the overall trend in melting points given below in terms of the forces among particles in the solids indicated. Hydrogen, H2 259 C Ethyl alcohol, C2H5OH 114 C Water, H2O 0 C Sucrose, C12H22O11 186 C Calcium chloride, CaCl2 772 C

What is a network solid? Give an example of a network solid and describe the bonding in such a solid. How does a network solid differ from a molecular solid?

Ionic solids do not conduct electricity in the solid state, but are strong conductors in the liquid state and when dissolved in water. Explain.

What is an alloy? Explain the differences in structure between substitutional and interstitial alloys. Give an example of each type.

The Chemistry in Focus segment Metal with a Memory discusses Nitinol, an alloy that remembers a shape originally impressed in it. Which elements compose Nitinol, and why is it classified as an alloy?

For Exercises 5160 choose one of the following terms to match the definition or description given. a. alloy b. specific heat c. crystalline solid d. dipoledipole attraction e. equilibrium vapor pressure f. intermolecular g. intramolecular h. ionic solids i. London dispersion forces j. molar heat of fusion k. molar heat of vaporization l. molecular solids m. normal boiling point n. semiconductor boiling point at pressure of 1 atm

For Exercises 5160 choose one of the following terms to match the definition or description given. a. alloy b. specific heat c. crystalline solid d. dipoledipole attraction e. equilibrium vapor pressure f. intermolecular g. intramolecular h. ionic solids i. London dispersion forces j. molar heat of fusion k. molar heat of vaporization l. molecular solids m. normal boiling point n. semiconductor energy required to melt 1 mole of a substance

For Exercises 5160 choose one of the following terms to match the definition or description given. a. alloy b. specific heat c. crystalline solid d. dipoledipole attraction e. equilibrium vapor pressure f. intermolecular g. intramolecular h. ionic solids i. London dispersion forces j. molar heat of fusion k. molar heat of vaporization l. molecular solids m. normal boiling point n. semiconductor forces between atoms in a molecule

For Exercises 5160 choose one of the following terms to match the definition or description given. a. alloy b. specific heat c. crystalline solid d. dipoledipole attraction e. equilibrium vapor pressure f. intermolecular g. intramolecular h. ionic solids i. London dispersion forces j. molar heat of fusion k. molar heat of vaporization l. molecular solids m. normal boiling point n. semiconductor forces between molecules in a solid

For Exercises 5160 choose one of the following terms to match the definition or description given. a. alloy b. specific heat c. crystalline solid d. dipoledipole attraction e. equilibrium vapor pressure f. intermolecular g. intramolecular h. ionic solids i. London dispersion forces j. molar heat of fusion k. molar heat of vaporization l. molecular solids m. normal boiling point n. semiconductor instantaneous dipole forces for nonpolar molecules

For Exercises 5160 choose one of the following terms to match the definition or description given. a. alloy b. specific heat c. crystalline solid d. dipoledipole attraction e. equilibrium vapor pressure f. intermolecular g. intramolecular h. ionic solids i. London dispersion forces j. molar heat of fusion k. molar heat of vaporization l. molecular solids m. normal boiling point n. semiconductor lining up of opposite charges on adjacent polar molecules

For Exercises 5160 choose one of the following terms to match the definition or description given. a. alloy b. specific heat c. crystalline solid d. dipoledipole attraction e. equilibrium vapor pressure f. intermolecular g. intramolecular h. ionic solids i. London dispersion forces j. molar heat of fusion k. molar heat of vaporization l. molecular solids m. normal boiling point n. semiconductor maximum pressure of vapor that builds up in a closedcontainer

For Exercises 5160 choose one of the following terms to match the definition or description given. a. alloy b. specific heat c. crystalline solid d. dipoledipole attraction e. equilibrium vapor pressure f. intermolecular g. intramolecular h. ionic solids i. London dispersion forces j. molar heat of fusion k. molar heat of vaporization l. molecular solids m. normal boiling point n. semiconductor mixture of elements having metallic properties overall

For Exercises 5160 choose one of the following terms to match the definition or description given. a. alloy b. specific heat c. crystalline solid d. dipoledipole attraction e. equilibrium vapor pressure f. intermolecular g. intramolecular h. ionic solids i. London dispersion forces j. molar heat of fusion k. molar heat of vaporization l. molecular solids m. normal boiling point n. semiconductor repeating arrangement of component species in a solid

For Exercises 5160 choose one of the following terms to match the definition or description given. a. alloy b. specific heat c. crystalline solid d. dipoledipole attraction e. equilibrium vapor pressure f. intermolecular g. intramolecular h. ionic solids i. London dispersion forces j. molar heat of fusion k. molar heat of vaporization l. molecular solids m. normal boiling point n. semiconductor solids that melt at relatively low temperatures

Given the densities and conditions of ice, liquid water, and steam listed in Table 14.1, calculate the volume of 1.0 g of water under each of these circumstances.

In carbon compounds a given group of atoms can often be arranged in more than one way. This means that more than one structure may be possible for the same atoms. For example, both the molecules diethyl ether and 1-butanol have the same number of each type of atom, but they have different structures and are said to be isomers of one another. diethyl ether CH3OCH2OOOCH2OCH3 1-butanol CH3OCH2OCH2OCH2OOH Which substance would you expect to have the larger vapor pressure? Why?

Which of the substances in each of the following sets would be expected to have the highest boiling point? Explain why. a. Ga, KBr, O2 b. Hg, NaCl, He c. H2, O2, H2O

Which of the substances in each of the following sets would be expected to have the lowest melting point? Explain why. a. H2, N2, O2 b. Xe, NaCl, C (diamond) c. Cl2, Br2, I2

When a person has a severe fever, one therapy to reduce the fever is an alcohol rub. Explain how the evaporation of alcohol from the persons skin removes heat energy from the body.

What is steel?

Some properties of potassium metal are summarized in the following table: Normal melting point 63.5 C Normal boiling point 765.7 C Molar heat of fusion 2.334 kJ/mol Molar heat of vaporization 79.87 kJ/mol Specific heat of the solid 0.75 J/g C a. Calculate the quantity of heat required to heat 5.00 g of potassium from 25.3 C to 45.2 C. b. Calculate the quantity of heat required to melt 1.35 moles of potassium at its normal melting point. c. Calculate the quantity of heat required to vaporize 2.25 g of potassium at its normal boiling point.

What are some important uses of water, both in nature and in industry? What is the liquid range for water?

Describe, on both a microscopic and a macroscopic basis, what happens to a sample of water as it is cooled from room temperature to 50 C below its normal freezing point.

Cake mixes and other packaged foods that require cooking often contain special directions for use at high elevations. Typically these directions indicate that the food should be cooked longer above 5000 ft. Explain why it takes longer to cook something at higher elevations.

Why is there no change in intramolecular forces when a solid is melted? Are intramolecular forces stronger or weaker than intermolecular forces?

What do we call the energies required, respectively, to melt and to vaporize 1 mole of a substance? Which of these energies is always larger for a given substance? Why?

The molar heat of vaporization of carbon disulfide, CS2, is 28.4 kJ/mol at its normal boiling point of 46 C. How much energy (heat) is required to vaporize 1.0 g of CS2 at 46 C? How much heat is evolved when 50. g of CS2 is condensed from the vapor to the liquid form at 46 C?

Which is stronger, a dipoledipole attraction between two molecules or a covalent bond between two atoms within the same molecule? Explain.

For a liquid to boil, the intermolecular forces in the liquid must be overcome. Based on the types of intermolecular forces present, arrange the expected boiling points of the liquid states of the following substances in order from lowest to highest: NaCl(l), He(l), CO(l), H2O(l).

What are London dispersion forces and how do they arise in a nonpolar molecule? Are London forces typically stronger or weaker than dipoledipole attractions between polar molecules? Are London forces stronger or weaker than covalent bonds? Explain.

Discuss the types of intermolecular forces acting in the liquid state of each of the following substances. a. N2 b. NH3 c. He d. CO2 (linear, nonpolar)

Explain how the evaporation of water acts as a coolant for the earth.

What do we mean when we say a liquid is volatile? Do volatile liquids have large or small vapor pressures? What types of intermolecular forces occur in highly volatile liquids?

Although methane, CH4, and ammonia, NH3, differ in molar mass by only one unit, the boiling point of ammonia is over 100 C higher than that of methane (a nonpolar molecule). Explain.

Which type of solid is likely to have the highest melting pointan ionic solid, a molecular solid, or an atomic solid? Explain.

What types of intermolecular forces exist in a crystal of ice? How do these forces differ from the types of intermolecular forces that exist in a crystal of solid oxygen?

Discuss the electron sea model for metals. How does this model account for the fact that metals are very good conductors of electricity?

Water is unusual in that its solid form (ice) is less dense than its liquid form. Discuss some implications of this fact.

Describe in detail the microscopic processes that take place when a liquid boils. What kind of forces must be overcome? Are any chemical bonds broken during these processes?

Water at 100 C (its normal boiling point) could certainly give you a bad burn if it were spilled on the skin, but steam at 100 C could give you a much worse burn. Explain.

What is a dipoledipole attraction? Give three examples of liquid substances in which you would expect dipoledipole attractions to be large.

What is meant by hydrogen bonding? Give three examples of substances that would be expected to exhibit hydrogen bonding in the liquid state.

Although the noble gas elements are monatomic and could not give rise to dipoledipole forces or hydrogen bonding, these elements still can be liquefied and solidified. Explain.

Describe, on a microscopic basis, the processes of evaporation and condensation. Which process requires an input of energy? Why?