Name____________________________________________period_____AP chemistry Unit 2 worksheet

1. List in order of increasing energy: 4f, 6s, 3d,1s,2p

2. Explain why the effective nuclear charge experienced by a 2s electron in boron is greater than that for the 2p electron.

3. Explain why the effective nuclear charge experienced by a 2s electron in aluminum is greater than that for the 2s electron
experienced by boron.

4. Which should experience the greater nuclear charge, a 2p electron in oxygen or a 2p electron in neon?

5. How many f orbitals have n=3?

6. Two electrons in an atom both occupy the 1s orbital. What quantity must be different for the two electrons?
7. How many unpaired electrons are there in an atom of tin in its ground state?

8. Of the following elements, which one is most likely to form an ion through the loss of two electrons?
a. strontium b. sulfur c. sodium d. chlorine e. aluminum

9. An atom has two electrons with principal quantum number (n) = 1, eight electrons with principal
quantum number (n) = 2 and seven electrons with principal quantum number (n) = 3. From this
data, supply the following values (if insufficient information is given, say so).
(a) The mass number. _________
(b) The atomic number. _________
(c) The electron configuration. ___________________________

10. What is the maximum number of electrons that can occupy each of the following
a. 3d b. 4s c. 2nd shell d. n=3
e. 2p f. 5f g. One 2p orbital h. n= 4
11. Write the orbital notation (can use noble gas) for each of the following
a. Sc b. Si
c. Sn d. Mn
12. Write the noble gas configuration for the following
b. Rb b. Se c. Zn
d. Pb e. Mn f. N

13. Write the full electronic configuration for argon

14. Identify the element from the electron configurations of atoms shown below. (3)
(a) [Ne] 3s2 3p2 _________ (b) [Ar] 4s2 3d7 _________ (c) [Xe] 6s2 _________

15. Which element could be represented by the complete PES spectrum to

the right?
A. Li b. B c. N d. Ne

16. Which of the following best explains the relative positioning and
intensity of the 2s peaks in the following spectra?

a)Be has a greater nuclear charge than Li and more electrons in the 2s orbital
b)Be electrons experience greater electron-electron repulsions than Li electrons
c)Li has a greater pull from the nucleus on the 2s electrons, so they are harder to remove
d)Li has greater electron shielding by the 1s orbital, so the 2s electrons are easier to
remove

17. Which will be closer to the nucleus, the n=3 electron shell in Ar or the n=3 shell in Kr?
18. Arrange the following atoms in order of increasing atomic radius: F, P, S, As and explain why.
19. Arrange the following atoms in order of increasing atomic radius: Al, Nb, Se, F, Mn and explain why.
20. An element having the configuration [Xe]6s1 belongs to the group:
a. alkaline earth metals b. alkali metals c. halogens d. noble gases e. none of these
21. Explain in terms of electron configurations, why hydrogen exhibits properties similar to both lithium and fluorine
22. Which of the following statements are true?
a. All are false
b. the first ionization energy of fluorine is greater than the first ionization energy of oxygen
c. as the atomic number increases within a group of the main group elements, the tendency is for first ionization energy to increase
d. it is easier to remove an electron from Na+ than from Na.
e. all particles with the electron configuration [Ar]4s2 have the same ionization energy.
23. Consider the element Scandium, atomic # 21.
(a) If the electronic configuration of the element were constructed "from scratch", into which orbital
(and into which shell) would the final electron be placed?
(b) When scandium forms an ion with a charge of +1, from which orbital (and from which shell) would the electron
be removed?
24. Based on their position on the periodic table, predict which atom of the following pairs will have the largest first ionization
energies. In each case explain with electron configuration and effective nuclear charge
a. O, Ne b. Mg, Sr c. K, Cr d. Br, Sb e. Ga, Ge

25. Identify two positive and two negative ions that are isoelectronic with argon. (4)
(a) Two Positive ions ________ ________ (b) Two Negative ions ________ ________

26. Compare the elements sodium and magnesium with respect to the following properties
a. Electron configuration b.Most common ionic charge
c. First ionization energy d. Atomic radius

27. Compare the elements fluorine and chlorine with respect to the following properties
a. Electron configuration b. Most common ionic charge c. Atomic radius

28. Write the noble gas configuration a. Fe3+ b. Ni2+

29. Arrange the atoms and ions in each of the following sets in order of increasing size
a. Br-, Na+, Mg2+ b. Ar, Cl-, S2-, K+

30. Using the periodic table, select the most electronegative atom in each of the following sets
a. B, Be, C, Si b. Zn, Ga, Ge, As c. Na, Mg, K, Ca
31. How many protons, neutrons, and electrons are in the following
a. 65Zn2+ b. 40Ar c. 14N3- d. 23Na+

32. Which ions are cations in the previous problem, which are anions?
33. How many valence electrons does each of the following atoms have?
a. C b. Ca c. H d. Pb e. Ar f. Cl
34. The ionization energies for an element are listed below
First second third fourth fifth
8 eV 15eV 80eV 109eV 141 eV
Based on the ionization energies, the element is most likely to be
a. Sodium b. magnesium c. aluminum d. silicon e. phosphorus

35. Which of the following contains only atoms that are diamagnetic in their ground state?
a. Kr, Ca, and P b. Ne, Be, and Zn c. Ar, K, and Ba d. He, Sr, and C
36. Which of the following is the electron configuration of an excited atom that is likely to emit a quantum of energy?
(A) 1s2 2s22p63s23p1 (B) 1s2 2s22p6 3s23p5
(C) 1s2 2s22p6 3s2 (D) 1s2 2s22p6 3s13p1
37. The bond energy of fluorine in 159 kJ mol-1.
i. Determine the energy, in J, of a photon of light needed to break one F-F bond.
ii. Determine the frequency of this photon in s-1
iii. Determine the wavelength of this photon in nanometers

38. Barium imparts a characteristic green color to a flame. The wavelength of this light is 551 nm. Determine the energy
involved in kJ/mol
39. Explain the difference between metallic, ionic, and covalent bonding.
40. Why can metals conduct electricity?
41. Label each of the following as ionic, metallic, or covalent
a. NaOH b. N2O c. KCl d. HF e. O2 f. Al foil
42. Which of the following forms molecules?
a. K2CO3 b. F2 c. BaCl2 d. H2O e. Fe2O3

43. Predict the chemical formula of the ionic compound formed between the following pairs of elements
a. Al and F b. K and S c. Mg and N d. Ba and O
44. Arrange the following substances according to their expected lattice energies, listed them from lowest lattice energy to
highest: LiCl, KCl, KBr, CaO

45. Explain the following trends in lattice energy

a. MgO > MgS b. LiF > CsBr c. CaO > KF

46. How is bonding in Cl2 different than NaCl?

47. Draw the Lewis structure for O2. The bond in O2 is shorter than the O-O single bond. Explain this observation.

48. Predict whether the following compounds are molecular or ionic

a. B2H6 b. CH3OH c. LiNO3 d. Sc2O3 e. CsBr f. NOCl g. Ag2SO4

49. Which of the following bonds are polar?

a. P-O b. S-F c. Br-Br
50. Arrange the bonds in order of increasing polarity
a. C-F, O-F, Be-F b. N-Br, P-Br, O-Br
51. Label each compound as ionic, polar covalent, or nonpolar covalent
a. CO b. MgO c. Cl2 d. AlF3
52. Draw the Lewis structure for the following
a. CO b. N2 c. SF2 d. ClO2- e. PCl3 e. H2CO (both H bonded to C)

53. 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. Account for the fact that the carbon-oxygen bond length in CO32– is greater than the carbon-oxygen bond length in CO2.

54. How can the concept of resonance be used to explain that all six C-C bonds in benzene are equal in length?

55. 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) All the bond lengths in SO3 are identical and are shorter than a sulfur-oxygen single bond.

56. Draw the Lewis structure for each of the following molecules
a. CO32- b. BH3 c. I3- d. XeF4 e. AsF6-
57. In the Lewis structure for CH2Cl2, what is the number of unshared electron pairs
a. 2 b. 8 c. 10 d. 6 e. 4

58. Which one of the following molecules contains a triple bond?

a. PF3 b. NF3 c. C2H2 d. H2CO e. HOF

59. Which of the following has the greatest dipole moment?

a. H2 b. HCl c. HF d. CO

60. Which of the following will conduct electricity? (there can be more than one answer)
a. solid Mg b. solid NaCl c. aqueous MgCl2 d. liquid Sn e. solid CO2 f. liquid N2

61. Predict which of the following will have the highest boiling point
a. CO2 b. N2 c. NaCl

62. Which of the following ionic compounds has the smallest lattice energy?
a. Na2O b. LiF c. CaO d. CaCl2 e. MgS
63. Bronze (Cu and Sn); Steel (Fe and C). Which of the following correctly describes the malleability of both alloys compared
to their primary metal?
a. Bronze’s malleability would be comparable to that of copper, but steel’s malleability would be significantly lower
than that of iron.
b. Bronze’s malleability would be significantly higher than that of copper, but steel’s malleability would be
comparable to that of iron.
c. Both bronze and steel would have malleability values similar to those of their primary metals
d. Both bronze and steel would have malleability values greater than those of their primary metal.

Review
64. Which of the following species contain more electrons than neutrons?
a. 2H b. 11B c. 16O2- d. 19F-
65. How many protons, neutrons, and electrons are in an 5626Fe atom?
Protons Neutrons Electrons
b. 26 30 26
c. 26 56 26
d. 56 26 26
e. 56 82 56
66. Calculate the following to the correct number of significant figures.
a.1.27g/5.296cm3 b. 12.2mL + 0.38mL c. 7.355g - 2.785g d. 0.1 m x 3.21m

67. Let’s pretend you are holding two atoms of carbon that are isotopes. Describe what the two atoms have in common and
what they have different.

68. What is the mass, in grams, of 1.75 x 1020 molecules of caffeine, C8H10N4O2?

69. Determine the empirical formulas of the compounds with the following compositions by mass
10.4 percent C, 27.8 percent S, and 61.7 percent Cl

70. A 2.00g sample of limestone was dissolved in hydrochloric acid and all the calcium present in the sample was converted to
Ca2+(aq). Excess ammonium oxalate solution, (NH4)2C2O4(aq), was added to the solution to precipitate the calcium ions as
calcium oxalate, CaC2O4(s). The precipitate was filtered, dried and weighed to a constant mass of 2.43g. Determine the
percentage by mass of calcium in the limestone sample.

71. Naphthalene is a hydrocarbon that is used for moth balls. A sample was burned in pure oxygen to produce 1.100 g of
carbon dioxide and 0.1798 grams of water. If the molecular mass of the compound is 128 g/mole so what is the empirical
formula and molecular formula for this compound?