AP Chemistry: Chemical Bonding I
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AP Chemistry: Chemical Bonding I
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AP Chemistry: Chemical Bonding I
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•Warm-ups and problems will be collected before you take the test.
•Read Chapter 9: Chemical Bonding I: Basic Concepts
•Read Chapter 20 section 3: Band Theory of Conductivity
•Know how to name and draw structural formulas for organic functional groups in Regents Table R.
Answer the following problems in the space provided. For problems involving an equation, carry out the
following steps: 1. Write the equation. 2. Substitute numbers and units. 3. Show the final answer with units.
There is no credit without showing work.
Valence Electrons
1. How many electrons does a silver atom have? How many of a silver atom’s electrons are involved in bonding?
2. Is bond formation endothermic or exothermic? Explain.
Lewis Dot Symbols
3. Write Lewis dot symbols for the following atoms and ions:
(a) I
(b) I(c) P (d) P3(e) Mg
(f) Mg2+
(g) Pb
(h) Pb2+
Ionic Bond
4. Define “ionic bond.” Why does ionic bonding usually involve a metal and a nonmetal atom?
5. List two common physical properties of most ionic compounds. Explain each in terms the nature of an ionic
bond.
6. Beryllium forms a compound with chlorine that has the empirical formula BeCl2. How would you
experimentally determine whether it is an ionic or molecular compound? (The compound is not soluble in
water.)
7. Write the Lewis dot symbols of the reactants and products in the following reactions. (First predict products
and balance the equations.)
(a) Sr + Se 
(b) Mg + P 
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Lattice Energy
8. Define lattice energy. Why does lattice energy correlate with melting point?
9. Write the formula for Coulomb’s law. Explain how lattice energy can be explained using Coulomb’s law.
10. Specify which compound in the following pairs of ionic compounds has higher lattice energy. In each case,
explain your choice:
(a) KCl or MgO
(b) LiF or LiBr
(c) BaCl2 or LiCl
Metallic Bond
11. Define metallic bond. Draw a two dimensional model of metallic bonding for aluminum showing at least 12
aluminum atoms.
12. Explain how the definition of metallic bonding explain:
(a) high electrical conductivity of metals.
(b) metals are malleable.
(c) metals are ductile.
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13. Predict which metal in each pair has higher melting point. Explain your prediction in terms of Coulomb’s law.
(a) Be or Mg
(b) Na or Mg
14. Describe how the band theory of metals explains why metals conduct electricity.
15. Describe the difference between an interstitial alloy and a substitutional alloy. Give several physical properties
of each that would distinguish between these two types of alloys.
Covalent Bonding
16. For bond formation in F2, sketch a plot showing potential energy vs. interatomic distance. Label which side of
the PE axis is attractive force and which is repulsive force, and label bond distance and bond strength.
17. For the physical properties listed in the table, state whether the value is relatively high (H) or low (L) for a
typical ionic and molecular compound. Explain the difference in terms of the types of forces holding particles
together.
Physical
Ionic
Molecular
Explain
Property
Compound
Compound
melting point
heat of fusion
conductivity
(solid)
conductivity
(aqueous)
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Organic Chemistry
Alkanes
18. Draw the structural formulas for the alkyl groups derived from propane and pentane.
19. Name the alkanes that have the following condensed structural formulas:
a. CH3CH2CH3
b. CH3(CH2)6CH3
20. Draw the structural formula for 2-methylbutane.
21. Why are alkane molecules nonpolar? What type of intermolecular force(s) holds alkanes together in the liquid
state?
22. Why do the melting points of the alkanes increase with increasing molar mass?
Alkenes and Alkynes
23. Differentiate between unsaturated and saturated hydrocarbons.
24. Draw the carbon skeleton for each of the following alkenes.
a. 2-pentene
b. 2-methyl-2-pentene
c. 3-methyl-3-hexene
Isomers
25. Draw condensed structural formulas for all structural isomers of C5H10. Name each compound.
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Hydrocarbon Rings
26. Draw structural formulas for cyclohexane and benzene. Why is benzene less reactive than cyclohexane?
Functional Groups
27. What is a functional group?
28. Identify the functional group in each structure. Name each compound.
a. CH3-OH
b. CH3-CH2-NH2
c. CH3-CH2-COOH
d. CH3-CH2-CH2-Br
e. CH3-CH2-O-CH2-CH3
29. Name the following halocarbons.
a. CH3CH2Cl
b. CH3CHCH=CH2
|
Cl
30. Name the following aldehydes and ketones:
a. CH3CH2CHO
b. CH3CH2CH2COCH2CH3 (how do you know this compound is a ketone and not an ether?)
c. CH3CH2CH2CH2CHO
31. Name the following ethers:
a. CH3OCH2CH3
32. Draw the full structural formula for propanoic acid.
b. CH3CH2CH2OCH3
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33. Explain why 1-butanol has a higher boiling point than diethyl ether. Which compound would you expect to be
more soluble in water? Explain.
Infrared Spectroscopy
34. The energy of IR radiation matches which type of molecular motion?
35. Explain how IR spectroscopy can be used to help determine the structural formula of an unknown compound.
36. The C-H bond in a typical alkane absorbs electromagnetic radiation with a frequency of 9.4E13 Hz when this
bond undergoes a stretching vibration. Calculate the energy of this vibration in J/bond and kJ/mol.
37. Name the physical process taking place within an atom/molecule that is measured by each of the following
types of spectroscopy: UV/visible spectroscopy, X-ray photoelectron spectroscopy, and IR spectroscopy.
Electronegativity and Bond Type
38. For bonds between each of these atoms, state whether the bond is expected to be nonpolar covalent (NC), polar
covalent (PC), or ionic (I). The arrange these bonds in order of increasing ionic character:
(a) C to H
(b) F to H
(c) Br to H
(d) Na to I
(e) Al to S
(f) Li to Cl
Lewis Structure of Covalent Compounds
39. Write Lewis structures for the following molecules:
(a) methane
(b) methanamine
(c) dimethyl ether
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40. The following structures are incorrect. Explain what is wrong with each one and give a correct Lewis structure
for the molecule. (Relative positions of atoms are shown correctly.)
(a) H—C=N
(b) H=C=C=H
(c)H—O=F
41. Write the Lewis structure for acetic acid.
42. Write Lewis structures for the following ions. Show formal charges.
(a) O22(b) C22(c) NO+
(d) NH4+
43. Write Lewis structures for the following four iso-electronic species. Show formal charges.
(a) CO
(b) NO+
(c) CN(d) N2
Resonance
44. Draw two resonance structures for the nitrite ion. Show formal charges.
45. Draw two resonance structures for diazomethane, CH2N2. Show formal charges. The skeletal structure of the
molecule is:
H
C N N
H
46. Draw three resonance structures for the molecule N2O in which the atoms are arranged in the order NNO.
Indicate formal charges.
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Exceptions to the Octet Rule
47. In the vapor phase, beryllium chloride consists of discrete BeCl2 molecules. Is the octet rule satisfied for Be in
this compound? If not, can you form an octet around Be by drawing another resonance structure? How
plausible is this structure?
48. Draw the Lewis structure for NO2.
49. Draw three resonance structures for the chlorate ion. Show formal charges.
50. Write a Lewis structure for SbCl5. Does this molecule obey the octet rule?
51. Write Lewis structures for the reaction:
Bond Dissociation Energy
52. For the reaction:
O(g) + O2(g)  O3(g)
A1C13 + Cl-  AlCl4-
H° = -107.2 kJ
Calculate the average bond energy in O3.
53. For the reaction
2C2H6(g) + 7O2(g)  4CO2(g) + 6H2O(g)
a. Predict the enthalpy of reaction from the average bond energies in Table 9.4.
b. Calculate the enthalpy of reaction from the standard enthalpies of formation (see Appendix 3) of the
reactant and product molecules, and compare the result with your answer for part (a).