EXTRA PROBLEMS FOR PRACTICE - lcanderson.info

EXTRA PROBLEMS FOR PRACTICE INTERMOLECULAR FORCES AND SOLIDS & LIQUIDS - UNIT 07 _________________________________________________________

Questions 1 through 23 are multiple choice questions. Questions 24 through 33 are free response.

1. Which compound is most likely to form intermolecular hydrogen bonds?

2. Which best explains why bromine is soluble in mineral oil?

3. The strongest interaction between hexane and iodine is . . .

4. In general, the strongest interaction with water molecules in aqueous solution are for ions that have . . .

5. Water and ethanol are completely miscible largely due to which intermolecular forces?

6. The energy absorbed when dry ice sublimes is required to overcome which type of interaction?

7. A container is half-filled with a liquid and sealed at room temperature and atmospheric pressure. What happens inside the container?

8. Acetone, (CH3)2C=O, is a volatile, flammable liquid. The central carbon is sp2 hybridized. The strongest intermolecular forces present in acetone are . . .

Hydrogen bonding occurs between molecules that have H-O, H-N, or H-F bonds. The H on one molecule attracts a nonbonding pair of electrons on an N, O, or F of another molecule. The C=O bond in trigonal planar is very polar so there is a strong dipole, but they are not hydrogen bonds.

(a) C4H10 (b) NaH (c) C2H5OH (d) C2H5SH (e) CH4

(a) Both substances are liquid. (b) Both substances have similar densities.

(c) Both substances are made of nonpolar molecules. (d) Both substances dissolve in water.

(e) One substance is made of polar molecules and the other is made of nonpolar molecules.

(a) covalent bonds (b) London dispersion (c) ionic bonds

(d) hydrogen bonds (e) ion-dipole attractions

(a) large charge and large size. (b) large charge and small size. (c) small charge and large size.

(d) small charge and small size. (e) zero charge and small size.

(a) covalent bonds (b) London dispersion (c) ionic bonds

(d) hydrogen bonds (e) ion-dipole attractions

(a) covalent bonds (b) ion-dipole forces (c) dipole-dipole forces

(d) dispersion forces (e) hydrogen bonds

(a) Evaporation stops.

(b) Evaporation continues for a time then stops.

(c) The pressure inside the container remains constant.

(d) The pressure inside the container increases for a time and then remains constant.

(e) The liquid evaporates until it is all in the vapor phase.

(a) dipole-dipole forces. (b) London dispersion forces. (c) hydrogen bonds.

(d) covalent bonds. (e) ion-dipole forces.

Page of Revised 2018-2019 - from AP* Test Prep sources1 7


9. The molar masses of a series of similar polar molecules increases in this order: A < B < C < D < E. The boiling points, in degrees Celsius, of molecules A, B, C, D, and E are respectively, 20o, 50o, 150o, 100o, and 200o. Which molecule is likely to form hydrogen bonds?

Because of increased dispersion forces, the boiling points of a series of similar molecules will increase regularly with increasing molar mass. Molecule C is lighter than molecule D but has a significantly higher boiling point, so there must be some additional attractive forces, like a H-bond.

10. Which of the factors affect the vapor pressure of a liquid at equilibrium? I. Intermolecular forces of attractions within the liquid II. The volume and/or surface area of liquid present III. The temperature of the liquid

Questions 11 - 14 refer to the following information.

A solid is a poor conductor of electricity, is very hard, and has a high melting point. 11. The solid is probably . . .

12. The solid might be . . .

13. The solid could have the chemical formula . . .

14. The properties of this solid can be attributed to . . .

Examine the atomic-level structures of graphite and diamond to the right and answer Q 15-18. 15. The orbital hybridization of graphite and diamond,

respectively, are . . . (a) sp3 and sp3. (b) sp2 and sp2. (c) sp3 and sp2. (d) sp2 and sp3. (e) sp and sp2.

(a) A (b) B (c) C (d) D (e) E

(a) I only (b) II only (c) III only (d) I and II only (e) I and III only

(a) metallic. (b) an alloy. (c) ionic.

(d) molecular. (e) covalent network.

(a) quartz. (b) tin. (c) brass.

(d) table sugar. (e) rock salt.

(a) CaSO4. (b) Pb. (c) C20H42.

(d) SiC. (e) P2O5.

(a) an interlocking pattern of atoms. (b) intermolecular hydrogen bonding.

(c) the electron-sea model. (d) strong ion-ion interactions.

(e) interstitial atoms occupying the spaces between atoms.



16. Which contains delocalized electrons.

17. Which conducts electricity?

18. Graphite is characterized by . . .

19. List the substances, CCl4, CBr4, and CH4 in order of increasing boiling point.

20. How do viscosity and surface tension change (a) as temperature increases and (b) and intermolecular forces of attraction become stronger? (a) Viscosity increases as IMFs increase while surface tension decreases. Both viscosity and surface tension increase

with increasing temperature. (b) Viscosity decreases as IMFs increase while surface tension increases. Both viscosity and surface tension increase

with decreasing temperature. (c) Both viscosity and surface tension increase as IMFs increase and temperature decreases. (d) Both viscosity and surface tension decrease as IMFs increase and temperature increases.

21. An important difference between crystalline solids and amorphous solids is that crystalline solids have . . .

22. In an interstitial alloy, the solute radii are ??? the solvent radii.

23. If a certain metal forms in a face centered cubic unit cell with an edge length of 3.52Å, what would you calculate as the radius of one atom?

24. Use concepts of chemical bonding and/or intermolecular forces to account for each of the following observations: a. The boiling points of water, ammonia, and methane are 100oC, -33oC, and -164oC, respectively.

H2O and NH3 have a higher BP than CH4 because H2O and NH3 both form H-bonds while CH4 does NOT form H-bonds. H-bonds are relatively strong IMF which hold the molecules together in liquids requiring more energy to boil them. H2O has a higher BP than NH3 because H2O has 2 lone pairs of electrons per molecule (versus 1 lone pair per molecule of NH3). Additionally, oxygen is more electronegative than nitrogen.

b. At 25oC and 1.0 atm, chlorine is a gas, bromine is a liquid, and iodine is a solid. Cl2, Br2, and I2 are all nonpolar molecules and have only London dispersion forces. London dispersion forces become stronger as mass increases, because an increase in mass means there are more electrons, and the molecules have a greater polarizability. I2 is the most massive molecule and has the highest boiling point.

(a) Neither graphite nor diamond. (b) Graphite only. (c) Diamond only.

(d) Both graphite and diamond. (e) There is not enough info.

(a) Neither graphite nor diamond. (b) Graphite only. (c) Diamond only.

(d) Both graphite and diamond. (e) There is not enough info.

(a) adjacent layers that can slide past each other. (b) London dispersions that hold the layers together.

(c) strong covalent bonds between the C atoms. (d) a high melting, relatively soft material.

(e) all of these.

(a) CH4 < CBr4 < CCl4 (b) CCl4 < CH4 < CBr4 (c) CH4 < CCl4 < CBr4 (d) CBr4 < CCl4 < CH4

(a) flat surfaces. (b) variables colors. (c) repeating lattice patterns. (d) low melting points.

(a) greater than (b) equal to (c) less than (d) sequestered by

(a) 4.60x10−8 cm (b) 4.98x10−8 cm (c) 8.80x10−9 cm (d) 1.24x10−8 cm



c. Calcium oxide (2615oC) melts at a much higher temperature than does potassium chloride (770o). CaO - larger charge and smaller size, so there is a greater Coulombic force. KCl - smaller charge and larger size, so there is a smaller Coulombic force.

d. Propane is a gas and ethanol is a liquid, even though they have similar molar masses. Propane - CH3CH2CH3 - nonpolar - only London dispersion forces Ethanol - CH3CH2OH - polar molecule - dipole-dipole forces and H-bonds - stronger forces means that ethanol would have a higher boiling point so propanol would be a gas while ethanol is a liquid.

25. Answer the following questions about water using principles of solids, liquids, and gases and intermolecular forces. a. Why does water boil at a lower temperature in Denver, Colorado, than in New York City?

BP is temperature where atmospheric pressure equals vapor pressure. In Denver, there is an increased elevation so there is a decrease in atmospheric pressure so the liquid will not have to have as great a vapor pressure and therefore does not have to reach as high a temperature. So its boiling point is lower.

b. For substances of similar molar mass, why does water have unusually high values for boiling point, heat of vaporization, and surface tension?

H2O forms very strong H-bonds between its molecules which holds it molecules together. The fact that the molecules are tightly held so it will require a large quantity of energy to vaporize water and break apart its molecules as it boils. This also means that its molecules are strongly attracted to each other meaning there is a high surface tension.

c. What structural features of ice cause it to float on water?

H-bonds between molecules cause it to form a relatively open hexagonal structure that is less dense than the more loosely held liquid.

d. Why does calcium chloride dissolve exothermically in water.

Dissolving ionic solid = breaking ion-ion bonds = endothermic process Dissolving ionic solid = forming ion-dipole forces = exothermic process Since CaCl2 dissolves exothermically, the ion-dipole forces formed between solute and solvent are stronger than the ion-ion bonds of CaCl2.

26. Use Coulomb’s law to qualitatively explain two reasons why hydrogen bonds, which are especially strong dipole-dipole forces of attraction, form only between hydrogen on one molecule and either nitrogen oxygen, or fluorine on another molecule. Coulomb’s law states that the force of attraction between two particles is directly proportional to their charges and inversely related to the distance between them. The small sizes of N, O, and F allow them to get very close to H, another small atom, minimizing the distance between them which increases the the force. Also, nitrogen, oxygen, and fluorine are highly



electronegative, whereas hydrogen is electropositive making the charges large and the force large.

27. Using some common laboratory equipment and materials, devise a plan to deduce the type of bonding in a sample of unknown solid. Specify the tests you would perform and discuss what the results mean. Test the sample for conductivity. If it conducts electricity, it is either a metal of a covalent network solid. If it has high conductivity, it is probably a metal. Add water to a portion and if it dissolves, it is probably an ionic or a molecular solid. Test the solution for conductivity. Ionic solids in water conduct electricity whereas molecular solids do not. Determine the melting point. Molecular solids usually melt well below 200oC. Add acid to a portion of the sample, some metals dissolve in acid with the evolution of a gas.

28. (a) Draw an atomic representation of a binary ionic compound. Use the interactions of the particle to justify the structure and stability of the solid. Each cation (usually the smaller species) attracts multiple anions (usually larger species) to form strong ionic bonds creating a very rigid, stable, crystalline structure. (b) Use your model to explain these macroscopic properties of ionic solids:

i. melting point Ionic solids have very high melting points because the ionic bonds are very strong requiring large amounts of heat/energy to overcome the forces of attraction that hold the crystal in place.

ii. conductivity Each ion is locked in place with respect to the other ions by the strong ionic bonds, so the ions cannot move to conduct an electric current.

iii. solubility ionic compounds will dissolve only if their ions have a relatively low charge and a corresponding low force of attraction.

(c) Make a two-dimensional drawing of the solid showing at the atomic level why it is brittle. When a stress causes slippage along a plane, the ions realign in a cation-cation and anion-anion arrangement. The results in repulsive forces that can cause the crystal to break abruptly.

29. Draw a molecular depiction of the structure of ice. Label and name the major forces of attraction between the molecules and tell why these forces operate.

See figure 11.12 on page 472 The dotted lines are the H-bonds that hold the structure together. A H-bond between two water molecules is an especially strong dipole-dipole interaction between an electropositive hydrogen atom from one molecule and an electronegative oxygen atom on another molecule.



30. (a) Draw an atomic representation of a metal and describe the relative positions of atoms and electrons. There is an ordered array of metal atom cores surrounded by a sea of “mobile” delocalized electrons.

(b) Use your depiction to explain on an atomic level the following macroscopic properties of metals: i. electrical conductivity

A metal conducts electricity well because its delocalized electrons are “mobile”, not tightly held to any one metal atom. If a voltage is applied to a metal, the negative electrons will flow toward the positive end.

ii. malleability A metal deforms easily because it mobile, delocalized, loosely held electrons easily redistribute themselves as the positions of the atoms change.

iii. high thermal conductivity A metal has the property of high thermal conductivity because heat applied to one end will be transferred throughout the metal by the kinetic energy of the delocalized loosely held electrons.

31. Draw a model of an interstitial alloy showing the relative positions of the atoms and the electrons. Explain the composition of a steel alloy and why it is harder and stronger than pure iron. Steel is an alloy in which small atoms of carbon fit into the spaces between the larger iron atoms. The nonmetal carbon atoms form covalent bonds with the iron atoms making the crystal structure larger and stronger.

A Blast from the Past . . . 32. The two stable isotopes of chlorine have masses of 34.969amu and 36.966amu.

a. What are the mass numbers of the two isotopes of chlorine?

Cl-35 and Cl-37

b. Calculate the percent abundance of the lighter isotope. Hint: You know that Cl = 35.4527 amu.

35.4527amu = 34.969 x + 36.966 (1 − x) x = 0.7578 Cl-35 = 75.78% Cl-37 = 24.22%

c. How many types of molecules with different masses exist in a sample of chlorine gas is the sample exists entirely as diatomic molecules? Explain your answer. Cl-35 & Cl-35 Cl-35 & Cl-37 Cl-37 & Cl-37 There would be three different masses.

d. Calculate the mass of a single chlorine molecule (in grams) having the largest molecular mass. 37Cl−37Cl so 2(36.966amu) = 73.932 amu 73.932 g/mol ÷ 6.02x1023 molecules/mol = 1.23x10−22 g/molecule



e. What is the mass (in amu and also in grams) of the most abundant molecule? Explain. 35Cl−35Cl so 2(34.969amu) = 69.938 amu 69.938 g/mol ÷ 6.02x1023 molecules/mol = 1.16x10−22 g/molecule

33. Like chlorine, iodine is a halogen and forms similar compounds. Write the names and formulas of the four oxyanions and the four oxyacids of iodine.

IO4− = periodate HIO4 = periodic acid IO3− = iodate HIO3 = iodic acid IO2− = iodite HIO2 = iodous acid IO− = hypoiodite HIO = hypoiodous acid


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EXTRA PROBLEMS FOR PRACTICE - lcanderson.info