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AP chem chapter 8,9,10 assignment

AP Chemistry

Chapter 8, 9, 10 Assignment

Part I

1984

Questions 8–9

(A) A network solid with covalent bonding

(B) A molecular solid with zero dipole moment

(C) A molecular solid with hydrogen bonding

(D) An ionic solid

(E) A metallic solid

8. Solid ethyl alcohol, C2H5OH

9. Silicon dioxide, SiO2

18. Hydrogen Halide Normal Boiling Point, °C

HF +19

HCl –85

HBr –67

HI –35

The liquefied hydrogen halides have the normal boiling points given above. The relatively high

boiling point of HF can be correctly explained by which of the following?

(A) HF gas is more ideal.

(B) HF is the strongest acid.

(C) HF molecules have a smaller dipole moment.

(D) HF is much less soluble in water.

(E) HF molecules tend to form hydrogen bonds.

40. The geometry of the SO3 molecule is best described as

(A) trigonal planar

(B) trigonal pyramidal

(C) square pyramidal

(D) bent

(E) tetrahedral

41. Which of the following molecules has the shortest bond length?

(A) N2

(B) O2

(C) Cl2

(D) Br2

(E) I2

51. Pi (π) bonding occurs in each of the following species EXCEPT

(A) CO2

(B) C2H4

(C) CN–

(D) C6H6

(E) CH4

54. Which of the following statements is always true about the phase diagram of any one–component

system?

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(A) The slope of the curve representing equilibrium between the vapor and liquid phases is positive.

(B) The slope of the curve representing equilibrium between the liquid and solid phases is negative.

(C) The slope of the curve representing equilibrium between the liquid and solid phases is positive.

(D) The temperature at the triple point is greater than the normal freezing point.

(E) The pressure at the triple point is greater than 1 atmosphere.

60. Which of the following has a zero dipole moment?

(A) HCN

(B) NH3

(C) SO2

(D) NO2

(E) PF5

66. Ca, V, Co, Zn, As

Gaseous atoms of which of the elements above are paramagnetic?

(A) Ca and As only

(B) Zn and As only

(C) Ca, V, and Co only

(D) V, Co, and As only

(E) V, Co, and Zn only

80. For which of the following molecules are resonance structures necessary to describe the bonding

satisfactorily?

(A) H2S

(B) SO2

(C) CO2

(D) OF2

(E) PF3

1989

Questions 11–14

(A) hydrogen bonding

(B) hybridization

(C) ionic bonding

(D) resonance

(E) van der Waals forces (London dispersion forces)

11. Is used to explain why iodine molecules are held together in the solid state

12. Is used to explain why the boiling point of HF is greater than the boiling point of HBr

13. Is used to explain the fact that the four bonds in methane are equivalent

14. Is used to explain the fact that the carbon-to-carbon bonds in benzene, C6H6, are identical

17. The Lewis dot structure of which of the following molecules shows only one unshared pair of

valence electrons?

(A) Cl2

(B) N2

(C) NH3

(D) CCl4

(E) H2O2

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21. Which of the following is true at the triple point of a pure substance?

(A) The vapor pressure of the solid phase always equal the vapor pressure of the liquid phase.

(B) The temperature is always 0.01 K lower that the normal melting point.

(C) The liquid and gas phases of the substance always have the same density and are therefore

indistinguishable.

(D) The solid phase always melts if the pressure increases at constant temperature.

(E) The liquid phase always vaporizes if the pressure increases at constant temperature.

31. The structural isomers C2H5OH and CH3OCH3 would be expected to have the same values for which

of the following? (Assume ideal behavior.)

(A) Gaseous densities at the same temperature and pressure

(B) Vapor pressures at the same temperature

(C) Boiling points

(D) Melting points

(E) Heats of vaporization

42. The SbCl5 molecule has trigonal bipyramid structure. Therefore, the hybridization of Sb orbitals

should be

(A) sp2

(B) sp3

(C) dsp2

(D) dsp3

(E) d2sp

3

43. Which of the following does NOT behave as an electrolyte when it is dissolved in water?

(A) CH3OH

(B) K2CO3

(C) NH4Br

(D) HI

(E) Sodium acetate, CH3COONa

47. CCl4, CO2, PCl3, PCl5, SF6

Which of the following does not describe any of the molecules above?

(A) Linear

(B) Octahedral

(C) Square planar

(D) Tetrahedral

(E) Trigonal pyramidal

Questions 49-51 refer to the phase diagram below.

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49. The normal boiling point of the substance represented by the phase diagram above is

(A) –15°C

(B) –10°C

(C) 140°C

(D) greater than 140°C

(E) not determinable from the diagram

50. The phase diagram above provides sufficient information for determining the

(A) entropy change on vaporization

(B) conditions necessary for sublimation

(C) deviations from ideal gas behavior of the gas phase

(D) latent heat of vaporization

(E) latent heat of fusion

51. For the substance represented in the diagram, which of the phases is most dense and which is least

dense at –15°C.

Most Dense Least Dense

(A) Solid Gas

(B) Solid Liquid

(C) Liquid Solid

(D) Liquid Gas

(E) The diagram gives no information about densities.

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59. Which of the following compounds is ionic and contains both sigma and pi covalent bonds?

(A) Fe(OH)3

(B) HClO

(C) H2S

(D) NO2

(E) NaCN

1994

26. Which of the following actions would be likely to change the boiling point of a sample of a pure

liquid in an open container?

I. Placing it in a smaller container

II. Increasing the number of moles of the liquid in the container

III. Moving the container and liquid to a higher altitude

(A) I only

(B) II only

(C) III only

(D) II and III only

(E) I, II, and III

32. CH3CH2OH boils at 78°C and CH3OCH3 boils at –24°C, although both compounds have the same

composition. This difference in boiling points may be attributed to a difference in

(A) molecular mass

(B) density

(C) specific heat

(D) hydrogen bonding

(E) heat of combustion

58. On a mountaintop, it is observed that water boils at 90°C not at 100°C as at sea level. This

phenomenon occurs because on the mountaintop the

(A) equilibrium water vapor pressure is higher due to the lower atmospheric pressure

(B) equilibrium water vapor pressure is lower due to the lower atmospheric pressure

(C) equilibrium water vapor pressure equals the atmospheric pressure at a lower temperature

(D) water molecules have a higher average kinetic energy due to the lower atmospheric pressure

(E) water contains a greater concentration of dissolved gases

1999

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25. The cooling curve for a pure substance as it changes from a liquid to a solid is shown above. The

solid and the liquid coexist at

(A) point Q only

(B) point R only

(C) all points on the curve between Q and S

(D) all points on the curve between R and T

(E) no point on the curve

2002

Questions 3-5 refer to the following molecules.

(A) CO2

(B) H2O

(C) CH4

(D) C2H4

(E) PH3

3. The molecule with only one double bond

4. The molecule with the largest dipole moment

5. The molecule that has pyramidal geometry

Questions 15-16 relate to the graph below. The graph shows the temperature of a pure substance as it is

heated at a constant rate in an open vessel at 1.0 atm pressure. The substance changes from a solid to the

liquid to the gas phase.

15. The substance is at its normal freezing point at time

(A) t1

(B) t2

(C) t3

(D) t4

(E) t5

16. Which of the following best describes what happens to the substance between t4 and t5 ?

(A) The molecules are leaving the liquid phase.

(B) The solid and liquid phase coexist in equilibrium.

(C) The vapor pressure of the substance is decreasing.

(D) The average intermolecular distance is decreasing.

(E) The temperature of the substance is increasing.

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28. Of the following compounds, which is the most ionic?

(A) SiCl4

(B) BrCl

(C) PCl3

(D) Cl2O

(E) CaCl2

29. The best explanation for the fact that diamond is extremely hard is that diamond crystals

(A) are made up of atoms that are intrinsically hard because of their electronic structures

(B) consist of positive and negative ions that are strongly attracted to each other

(C) are giant molecules in which each atom forms strong covalent bonds with all of its neighboring

atoms

(D) are formed under extreme conditions of temperature and pressure

(E) contain orbitals or bands of delocalized electrons that belong not to single atoms but to each

crystal as a whole

48. Sodium chloride is LEAST soluble in which of the following liquids?

(A) H2O

(B) CCl4

(C) HF

(D) CH3OH

(E) CH3COOH

53. According to the VSEPR model, the progressive decrease in the bond angles in the series of

molecules CH4, NH3, and H2O is best accounted for by the

(A) increasing strength of the bonds

(B) decreasing size of the central atom

(C) increasing electronegativity of the central atom

(D) increasing number of unshared pairs of electrons

(E) decreasing repulsion between hydrogen atoms

56. The boiling points of the elements helium, neon, argon, krypton, and xenon increase in that order.

Which of the following statements accounts for this increase?

(A) The London (dispersion) forces increase.

(B) The hydrogen bonding increases.

(C) The dipole-dipole forces increase.

(D) The chemical reactivity increases.

(E) The number of nearest neighbors increases.

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67. Which of the following describes the changes in forces of attraction that occur as H2O changes phase

from a liquid to a vapor?

(A) H–O bonds break as H–H and O–O bonds form.

(B) Hydrogen bonds between H2O molecules are broken.

(C) Covalent bonds between H2O molecules are broken.

(D) Ionic bonds between H+ ions and OH

– ions are broken.

(E) Covalent bonds between H+ ions and H2O molecules become more effective.

70. Of the following pure substances, which has the highest melting point?

(A) S8

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(B) I2

(C) SiO2

(D) SO2

(E) C6H6

2008

7. Which of the following pure substances has molecules with a pyramidal shape?

(A) NH3(g)

(B) BH3(g)

(C) H2(g)

(D) H2S(g)

(E) HBr(g)

20. In solid methane, the forces between neighboring CH4 molecules are best characterized as

(A) ionic bonds

(B) covalent bonds

(C) hydrogen bonds

(D) ion-dipole forces

(E) London (dispersion) forces

22. Which of the following is a nonpolar molecule that contains polar bonds?

(A) F2

(B) CHF3

(C) CO2

(D) HCl

(E) NH3

28. Which of the following molecules contains only single bonds?

(A) CH3COOH

(B) CH3CH2COOCH3

(C) C2H6

(D) C6H6

(E) HCN

37. Which of the following elements combines with oxygen to form a covalent network solid?

(A) Si

(B) S

(C) C

(D) Mg

(E) Cs

40. On the basis of strength of intermolecular forces, which of the following elements would be expected

to have the highest melting point?

(A) Br2

(B) Cl2

(C) F2

(D) Kr

(E) N2

43. A pure liquid in an open vessel boils at the temperature at which the

(A) molar entropy of the liquid becomes equal to that of the gas

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(B) vapor pressure of the liquid becomes equal to the equilibrium pressure at the triple point

(C) vapor pressure of the liquid becomes equal to the atmospheric pressure on the surface of the

liquid

(D) molar heat capacity of the liquid becomes equal to that of the gas

(E) average kinetic energy of the liquid molecules becomes equal to that of the gas molecules

56. The London (dispersion) forces are weakest for which of the following gases under the same

conditions of temperature and pressure?

(A) H2

(B) O2

(C) Xe

(D) F2

(E) N2

71. Of the following single bonds, which is the LEAST polar?

(A) N–H

(B) H–F

(C) O–F

(D) I–F

(E) O–H

73. The figure above shows two closed containers. Each contains the same volume of acetone in

equilibrium with its vapor at the same temperature. The vapor pressure of the acetone is

(A) higher in container 1 because the surface area of the liquid is greater

(B) higher in container 1 because the volume of vapor is greater

(C) lower in container 1 because the level of the liquid is lower

(D) the same in both containers because the volume of the liquid is the same

(E) the same in both containers because the temperature is the same

Part II – Free Response

1973

Discuss briefly the relationship between the dipole moment of a molecule and the polar character of the

bonds within it. With this as the basis, account for the difference between the dipole moments of CH2F2

and CF4.

1974

The boiling points of the following compounds increase in the order in which they are listed below:

CH4 < H2S < NH3

Discuss the theoretical considerations involved and use them to account for this order.

1975

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Suppose that a molecule has the formula AB3. Sketch and name two different shapes that this molecule

may have. For each of the two shapes, give an example of a known molecule that has that shape. For one

of the molecules you have named, interpret the shape in the context of a modern bonding theory.

1976

NF3 and PF5 are stable molecules. Write the electron-dot formulas for these molecules. On the basis of

structural and bonding considerations, account for the fact that NF3 and PF5 are stable molecules but NF5

does not exist.

1977

The state of aggregation of solids can be described as belonging to the following four types:

(1) ionic (3) covalent network

(2) metallic (4) molecular

For each of these types of solids, indicate the kinds of particles that occupy the lattice points and identify

forces among these particles. How could each type of solid be identified in the laboratory?

1979 (1)

Draw Lewis structures for CO2, H2, SO3 and SO32-

and predict the shape of each species.

1979 (2)

Butane, chloroethane, acetone, and 1-propanol all have approximately the same molecular weights. Data

on their boiling points and solubilities in water are listed in the table below.

Compound

Formula

Boiling

Pt.(ºC) Solubility

in water

Butane CH3CH2CH2CH3 0 insoluble

Chloroethane CH3CH2Cl 12 insoluble

Acetone

CH3C CH3

56 completely

miscible

1-Propanol

CH3CH2CH2OH

97

completely

miscible

On the basis of dipole moments (molecular polarities) and/or hydrogen bonding, explain in a qualitative

way the differences in the

(a) boiling points of butane and chloroethane.

(b) water solubilities of chloroethane and acetone.

(c) water solubilities of butane and 1-propanol.

(d) boiling points of acetone and 1-propanol.

1982

(a) Draw the Lewis electron-dot structures for CO32–

, CO2, and CO, including resonance structures

where appropriate.

(b) Which of the three species has the shortest C–O bond length? Explain the reason for your answer.

(c) Predict the molecular shapes for the three species. Explain how you arrived at your predictions.

1984

Give a scientific explanation for the following observations. Use equations or diagrams if they are

relevant.

C

||

O

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(a) It takes longer to cook an egg until it is hard-boiled in Denver (altitude 1 mile above sea level) than it

does in New York City (near sea level).

(b) Perspiring is a mechanism for cooling the body.

1985

Substance Melting Point, ºC

H2 -259

C3H8 -190

HF -92

CsI 621

LiF 870

SiC >2,000

(a) Discuss how the trend in the melting points of the substances tabulated above can be explained in

terms of the types of attractive forces and/or bonds in these substances.

(b) For any pairs of substances that have the same kind(s) of attractive forces and/or bonds, discuss the

factors that cause variations in the strengths of the forces and/or bonds.

1988 (1)

The normal boiling and freezing points of argon are 87.3 K and 84.0 K, respectively. The triple point is at

82.7 K and 0.68 atmosphere.

(a) Use the data above to draw a phase diagram for argon. Label the axes and label the regions in which

the solid, liquid and gas phases are stable. On the phase diagram, show the position of the normal

boiling point.

(b) Describe any changes that can be observed in a sample of solid argon when the temperature is

increased from 40 K to 160 K at a constant pressure of 0.50 atmosphere.

(c) Describe any changes that can be observed in a sample of liquid argon when the pressure is reduced

from 10 atmospheres to 1 atmosphere at a constant temperature of 100 K, which is well below the

critical temperature.

(d) Does the liquid phase of argon have a density greater than, equal to, or less than the density of the

solid phase? Explain your answer, using information given in the introduction to this question.

1988 (2)

Using principles of chemical bonding and/or intermolecular forces, explain each of the following.

(a) Xenon has a higher boiling point than neon has.

(b) Solid copper is an excellent conductor of electricity, but solid copper chloride is not.

(c) SiO2 melts at a very high temperature, while CO2 is a gas at room temperature, even though Si and

C are in the same chemical family.

(d) Molecules of NF3 are polar, but those of BF3 are not.

1989 (1)

CF4 XeF4 ClF3

(a) Draw a Lewis electron-dot structure for each of the molecules above and identify the shape of each.

(b) Use the valence shell electron-pair repulsion (VSEPR) model to explain the geometry of each of

these molecules.

1989 (2)

The melting points of the alkali metals decrease from Li to Cs. In contrast, the melting points of the

halogens increase from F2 to I2.

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(a) Using bonding principles, account for the decrease in the melting points of the alkali metals.

(b) Using bonding principles, account for the increase in the melting points of the halogens.

(c) What is the expected trend in the melting points of the compounds LiF, NaCl, KBr, and CsI? Explain

this trend using bonding principles.

1990

Use simple structure and bonding models to account for each of the following.

(a) The bond length between the two carbon atoms is shorter in C2H4 than in C2H6.

(b) The H–N–H bond angle is 107.5º, in NH3.

(c) The bond lengths in SO3 are all identical and are shorter than a sulfur-oxygen single bond.

(d) The I3– ion is linear.

1991

Experimental data provide the basis for interpreting differences in properties of substances.

TABLE 1

Compound

Melting

Point (ºC) Electrical Conductivity

of Molten State (ohm-1

)

BeCl2 405 0.086 MgCl2 714 > 20 SiCl4 -70 0 MgF2 1261 > 20

TABLE 2 Substance Bond Length

(angstroms) F2 1.42 Br2 2.28 N2 1.09

Account for the differences in properties given in Tables 1 and 2 above in terms of the differences in

structure and bonding in each of the following pairs.

(a) MgCl2 and SiCl4 (c) F2 and Br2

(b) MgCl2 and MgF2 (d) F2 and N2

1992 (1)

Explain each of the following in terms of atomic and molecular structures and/or intermolecular forces.

(a) SbCl3 has measurable dipole moment, whereas SbCl5 does not.

(b) The normal boiling point of CCl4 is 77ºC, whereas that of CBr4 is 190ºC.

(c) NaI(s) is very soluble in water, whereas I2(s) has a solubility of only 0.03 gram per 100 grams of

water.

1992 (2)

NO2 NO2– NO2

+

Nitrogen is the central atom in each of the species given above.

(a) Draw the Lewis electron-dot structure for each of the three species.

(b) List the species in order of increasing bond angle. Justify your answer.

(c) Select one of the species and give the hybridization of the nitrogen atom in it.

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(d) Identify the only one of the species that dimerizes and explain what causes it to do so.

1995 (1)

The conductivity of several substances was tested using the apparatus represented by the diagram below.

The results of the tests are summarized in the following data table.

AgNO3 Sucrose Na H2SO4 (98%)

Melting

Point

(ºC)

212º 185º 99º Liquid at Room Temp.

Liquid

(fused) ++ – ++ +

Water

Solution ++ – ++

(1) ++

(2)

Solid – – ++ Not Tested

Key: ++ Good conductor + Poor conductor

– Nonconductor

(1) Dissolves, accompanied by evolution of flammable gas (2) Conduction increases as the acid is added slowly and carefully to water

Using models of chemical bonding and atomic or molecular structure, account for the differences in

conductivity between the two samples in each of the following pairs.

(a) Sucrose solution and silver nitrate solution.

(b) Solid silver nitrate and solid sodium metal.

(c) Liquid (fused) sucrose and liquid (fused) silver nitrate.

(d) Liquid (concentrated) sulfuric acid and sulfuric acid solution.

1995 (2)

Explain the following in terms of the electronic structure and bonding of the compounds considered.

(a) Liquid oxygen is attracted to a strong magnet, whereas liquid nitrogen is not.

(b) The SO2 molecule has a dipole moment, whereas the CO2 molecule has no dipole moment. Include

the Lewis (electron-dot) structures in your explanation.

(c) Halides of cobalt(II) are colored, whereas halides of zinc(II) are colorless.

(d) A crystal of high purity silicon is a poor conductor of electricity; however, the conductivity increases

when a small amount of arsenic is incorporated (doped) into the crystal.

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1995 (3)

The phase diagram for a pure substance is shown above. Use this diagram and your knowledge about

changes of phase to answer the following questions.

(a) What does point V represent? What characteristics are specific to the system only at point V?

(b) What does each point on the curve between V and W represent?

(c) Describe the changes that the system undergoes as the temperature slowly increases from X to Y to Z

at 1.0 atmosphere.

(d) In a solid-liquid mixture of this substance, will the solid float or sink? Explain.

1996

Explain each of the following observations in terms of the electronic structure and/or bonding of the

compounds involved.

(a) At ordinary conditions, HF (normal boiling point = 20ºC) is a liquid, whereas HCl (normal boiling

point = –114ºC) is a gas.

(b) Molecules of AsF3 are polar, whereas molecules of AsF5 are nonpolar.

(c) The N–O bonds in the NO2– ion are equal in length, whereas they are unequal in HNO2.

(d) For sulfur, the fluorides SF2, SF4, and SF6 are known to exist, whereas for oxygen only OF2 is known

to exist.

2000

Answer the following question about the element selenium, Se (atomic number 34).

Selenium reacts with fluorine to form SeF4. Draw the complete Lewis electron-dot structure for SeF4 and

sketch the molecular structure. Indicate whether the molecule is polar or nonpolar, and justify your

answer.

2002

Use the principles of atomic structure and/or chemical bonding to explain each of the following. In each

part, your answer must include references to both substances.

(a) The carbon-to-carbon bond energy in C2H4 is greater than it is in C2H6.

(b) The boiling point of Cl2 is lower than the boiling point of Br2.

2003

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For the following, use appropriate chemical principles to explain the observation. Include chemical

equations as appropriate.

Water droplets form on the outside of a beaker containing an ice bath.

2005 (1)

Answer the following questions that relate to chemical bonding

(a) Draw the complete Lewis structure (electron-dot diagram) for the molecules CF4, PF5, and SF4.

(b) On the basis of the Lewis structures drawn in part (a), answer the following questions about the

particular molecule indicated.

(i) What is the F–C–F bond angle in CF4?

(ii) What is the hybridization of the valence orbitals of P in PF5?

(iii) What is the geometric shape formed by the atoms in SF4?

(c) Two Lewis structures can be drawn for the OPF3 molecule, as shown below.

Structure 1 Structure 2

(i) How many sigma bonds and how many pi bonds are in structure 1?

(ii) Which one of the two structures best represents a molecule of OPF3? Justify your answer in

terms of formal charge.

2005 (2)

Use principles of atomic structure, bonding and/or intermolecular forces to respond to each of the

following. Your responses must include specific information about all substances referred to in each

question.

(a) At a pressure of 1 atm, the boiling point of NH3(l) is 240 K, whereas the boiling point of NF3(l) is

144 K.

(i) Identify the intermolecular forces(s) in each substance.

(ii) Account for the difference in the boiling points of the substances.

(b) The melting point of KCl(s) is 776˚C, whereas the melting point of NaCl(s) is 801˚C.

(i) Identify the type of bonding in each substance.

(ii) Account for the difference in the melting points of the substances.

2008 (1)

Using principles of atomic and molecular structure, answer the following questions about compounds of

xenon, fluorine, and oxygen.

(a) Xenon can react with oxygen and fluorine to form compounds such as XeO3 and XeF4. Draw the

complete Lewis electron-dot diagram for each of these molecules.

(b) On the basis of the Lewis electron-dot diagrams you drew for part (a), predict the following:

(i) The geometric shape of XeO3 molecule

(ii) The hybridization of the valence orbitals of xenon in XeF4

(c) Predict whether the XeO3 molecule is polar or nonpolar. Justify your prediction.

2008 (2)

Answer the following questions by using principles of molecular structure and intermolecular forces.

P

O

FF

F

:

::

:....

..

..

..

: :

P

O

FF

F

:

::

:....

....

..

..

::

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(a) Structures of the pyridine molecule and benzene molecule are shown below. Pyridine is soluble

in water, whereas benzene is not soluble in water. Account for the difference in solubility. You

must discuss both of the substances in your answer.

(b) Structures of the dimethyl ether molecule and the ethanol molecule are shown below. The normal

boiling point of dimethyl ether is 250 K, whereas the normal boiling point of ethanol is 351 K.

Account for the difference in boiling points. You must discuss both of the substances in your

answer.

(c) SO2 melts at 201 K, whereas SiO2 melts at 1,883 K. Account for the difference in melting points.

You must discuss both of the substances in your answer.

(d) The normal boiling point of Cl2(l) (238 K) is higher than the normal boiling point of HCl(l) (188

K). Account for the differences in normal boiling points based on the types of intermolecular

forces in the substances. You must discuss both of the substances in your answer.

2009

Answer the following questions related to sulfur and one of its compounds.

(a) Which of the two species, S and S2–

, would be attracted into a magnetic field? Explain.

(b) In the H2S molecule, the H–S–H bond angle is close to 90°. On the basis of this information,

which atomic orbitals of the S atom are involved in bonding with the H atoms?

(c) Two types of intermolecular forces present in liquid H2S are London (dispersion) forces and

dipole-dipole forces.

(i) Compare the strength of the London (dispersion) in liquid H2S to the strength of the London

(dispersion) forces in liquid H2O. Explain.

(ii) Compare the strength of the dipole-dipole forces in liquid H2S to the strength of the dipole-

dipole forces in liquid H2O. Explain.

2010

Use the information in the table below to respond to the statements and questions that follow. Your

answers should be in terms of principles of molecular structure and intermolecular forces.

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(a) Draw the complete Lewis electron-dot diagram for ethyne.

(b) Which of the four molecules contains the shortest carbon-to-carbon bond? Explain.

(c) A Lewis electron-dot diagram of a molecule of ethanoic acid is given below. The carbon atoms

in the molecule are labeled x and y, respectively.

Identify the geometry of the arrangement of atoms bonded to each of the following.

(i) Carbon x

(ii) Carbon y

(d) Energy is required to boil ethanol. Consider the statement “As ethanol boils, energy goes into

breaking C–C bonds, C–H bonds, C–O bonds, and O–H bonds.” Is the statement true or false?

Justify your answer.

(e) Identify a compound from the table above that is nonpolar. Justify your answer.

(f) Ethanol is completely soluble in water, whereas ethanethiol has limited solubility in water.

Account for the difference in solubilities between the two compounds in terms of intermolecular

forces.

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