ap unit 2 worksheet answers

Name____________________________________________________________period_____AP chemistry
Unit 2 worksheet
Concepts to know:
Read Chapters 6, 7, 2.6, 8.3
The wave nature of light
Quantized energy and photons
Bohr’s model of the atom
The quantum model of the atom (this is the accepted model and PES is the proof)
Effective nuclear charge
Electron configuration
Sizes of atoms
Ionization energy
Electronegativity
Ions and ionic compounds
Practice problems
1. What are the SI units for
a. Wavelength of light
b. frequency of light
c. speed of light
-1
Meter
hertz (s )
m s-1 (m/s)
2. T/F (correct the statement if it is false)
a. Visible light is a form of electromagnetic radiation.
True
b. The frequency of radiation increases as the wavelength increases.
False, the frequency decreases as wavelength increases
c. Ultraviolet light has a shorter wavelength than visible light.
True
d. Electromagnetic radiation and sound waves travel at the same speed.
False, electromagnetic radiation travels at the speed of light
3. What is the wavelength of radiation that has a frequency for 5.11 x 1011 s-1? Would this be visible to the human
eye?
5.87 x 10-4 m; no
4. Exited mercury atoms emit light at 489 nm. What is the frequency of this radiation? Predict the color
associated with it.
6.12 x 1014 s-1; blue
5. Calculate the smallest increment of energy that can be emitted or absorbed at a wavelength of 812 nm.
2.45 x 10-19 J
6.
The most prominent line in the spectrum of neon is found at 865.438 nm. Other lines are found at 837.761 nm, 878.062
nm, 878.438 nm, and 1885.387 nm. Which of these lines represents the most energetic light?
at 865.438 nm
7. Is energy emitted or absorbed when the following electronic transitions occur in hydrogen?
a. From n=4 to n=2
Emitted
b. From an orbit of radius 2.12 A� to one of radius 8.48 A� .
Absorbed
8. In the Bohr model of the hydrogen atom, when the electron is in its ground state, it orbits the nucleus at a
specific radius of 0.53A� . In the quantum mechanical description of the hydrogen atom, the most probable
distance of the electron from the nucleus is 0.53 A� . Why are these statements different?
The Bohr (shell model) states with 100% certainty that the electron in hydrogen can be found at 0.53 A from the
nucleus. The quantum model is a statistical model that states the probability of finding the electron in certain
regions around the nucleus. While 0.53 A is the radius with the highest probability, that probability is less than
100%.
9. What are the similarities and differences between the hydrogen 1s and 2s orbitals?
Same spherical shape, but 2s has a larger radial extension than the 1s
10. List in order of increasing energy: 4f, 6s, 3d,1s,2p
1s, 2p, 6s, 4f
11. Explain why the effective nuclear charge experienced by a 2s electron in boron is greater than that for the 2p
electron.
The 2p electron in boron is shielded from the full charge of the nucleus by the 2s electrons, so the 2p electron
experiences a smaller effective nuclear charge
12. Explain why the effective nuclear charge experienced by a 2s electron in aluminum is greater than that for the 2s
electron experienced by boron. Aluminum has more protons than boron so the 2s electron will experience a
greater effective nuclear charge
13. Which should experience the greater nuclear charge, a 2p electron in oxygen or a 2p electron in neon?
A 2p electron in neon experiences a greater effective nuclear charge
14. How many f orbitals have n=3? 0
15. Two electrons in an atom both occupy the 1s orbital. What quantity must be different for the two electrons?
They must have opposite spins
16. How many unpaired electrons are there in an atom of tin in its ground state? 2
17. 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
18. 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. __not enough info_______
(b) The atomic number. ___17______
(c) The electron configuration. _1s22s22p63s23p5______
19. What is the maximum number of electrons that can occupy each of the following
a. 3d
10
b. 4s 2
c. 2nd shell 8
e. 2p 6
f. 5f 14
g. One 2p orbital 2
d. n=3 18
h. n= 4 32
20. Write the orbital notation (can use noble gas) for each of the following
a. Sc
b. Si
↑↓
[Ar] 4s
c. Sn
↑↓
[Kr] 5s
↑ __ __ __ __
3d
↑↓ ↑↓ ↑↓ ↑↓ ↑↓
4d
↑↓
[Ne] 3s
d. Mn
↑ ↑ __
5p
[Ar]
21. Write the noble gas configuration for the following
a. Rb
b. Se
1
[Kr]5s
[Ar]4s23d104p4
d. Pb
e. Mn
2
10 14
2
[Xe]6s 5d 4f 6p
[Ar]4s23d5
↑_ __↑ ____
3p
↑↓
4s
↑_ ↑_ ↑_ ↑_ ↑_
3d
c. Zn
[Ar]4s23d10
f. N
[He]2s22p3
22. Write the full electronic configuration for argon 1s22s22p63s23p6
23. Identify the element from the electron configurations of atoms shown below. (3)
2
2
(a) [Ne] 3s 3p ___Si______
2
7
(b) [Ar] 4s 3d ___Co______
2
(c) [Xe] 6s _____Ba____
24. H, He, Li, B, Ne, Na
a. For each atom, tell how many peaks we will see in PES 1,1,2,3,3,4
b. Tell what peaks go with what orbitals. (See in class)
c. Draw what a PES would look like for each atom. (See in class)
25. Which will be closer to the nucleus, the n=3 electron shell in Ar or the n=3 shell in Kr?
Kr
26. Arrange the following atoms in order of increasing atomic radius: F, P, S, As and explain why.
F, S, P, As
(atoms get bigger when a shell is added or less protons in the same number of shells)
27. Arrange the following atoms in order of increasing atomic radius: Al, Nb, Se, F, Mn and explain why.
F, Al, Se,Mn, Nb, (atoms get bigger when a shell is added or less protons in the same number of shells)
28. 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
29. What is the trend in the first ionization energy as one proceeds down group 1? Explain how this relates to size
of the atoms. It decreases, the atoms have more shells as you go down a group so they have a larger radius
and the nucleus does not hold on to them as well.
30. Arrange the following pure solid elements in order of increasing electrical conductivity: P, Ag, and Sb
P, Sb, Ag
31. Explain in terms of electron configurations, why hydrogen exhibits properties similar to both lithium and
fluorine It has one valence electron like lithium, but only needs one valence electron to have a full shell like
fluorine
32. 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.
33. 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? 3d
(b) When scandium forms an ion with a charge of +1, from which orbital (and from which shell)
would the electron be removed? 4s
34. 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
Ne
Mg
Cr
Br
Ge
Smaller atoms have a harder time losing their electrons because the valence electrons feel a greater
effective nuclear charge so it takes more energy to remove the electron.
35. For each pair, which element will have greater metallic character
a. Li or Be
b. Li or Na
c. Sn or P
36. Predict whether each of the following oxides is ionic or molecular: SO2, MgO, Li2O, N2O, XeO3
Molecular, ionic, ionic, molecular, molecular
37. Identify two positive and two negative ions that are isoelectronic with argon. (4)
+
2+____
(a) Two Positive ions _K _______ ____Ca
-_____
2-______
__S
(b) Two Negative ions ___Cl
38. Compare the elements sodium and magnesium with respect to the following properties
a. Electron configuration 1s22s22p63s1; 1s22s22p63s2
b. Most common ionic charge+1; +2
c. First ionization energy magnesium has a higher first ionization energy
d. Atomic radius magnesium is smaller than sodium
39. Compare the elements fluorine and chlorine with respect to the following properties
a. Electron configuration
b. Most common ionic charge
c. Atomic radius
2 2
5
2 2
6 2
5
1s 2s 2p ; 1s 2s 2p 3s 3p
-1;-1
Chlorine is larger than fluorine
40. Why are monatomic cations smaller than their corresponding neutral atom? Lost one electron so the
remaining electrons have less repulsion and a greater effective nuclear charge
41. Write the noble gas configuration
a. Fe3+
b. Ni2+
[Ar]3d5
[Ar]3d8
42. Which neutral atom is isoelectric with each of the following
a. Clb. Se2c. Mg2+
Ar
Kr
Ne
43. 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+
44.
45.
46.
47.
48.
49.
50.
Mg, Na, Br
K, Ar, Cl, S
What is a covalent bond?
A bond that forms when atoms share electrons
What is an ionic bond?
A bond that forms when atoms give and take electrons
What is a metallic bond?
A bond that forms when nuclei are attracted to a “sea of electrons”
Why can metals conduct electricity?
There are electrons moving around
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
C
As
Mg
Label each of the following as ionic, metallic, or covalent
a. NaOH
b. N2O
c. KCl
d. HF
e. O2
f. Al foil
ionic
covalent
ionic
covalent
covalent
metallic
Which of the following forms molecules?
a. K2CO3
b. F2
c. BaCl2
d. H2O
e. Fe2O3
51. How many protons, neutrons, and electrons are in the following
a. 65Zn2+
b. 40Ar
c. 14N3d. 23Na+
30,35,28
18,22,18
7,7,10
11,12,10
52. Which ions are cations in the previous problem, which are anions? A,d
53. How is bonding in Cl2 different than NaCl? In Cl2, electrons are being shared, and in NaCl sodium gives an
electron to chlorine
54. Which of the following species contain more electrons than neutrons?
a. 2H
b. 11B
c. 16O2d. 19F55. How many valence electrons does each of the following atoms have?
a. C
b. Ca
c. H
d. Pb
e. Ar
4
2
1
4
8
56. What is the most common charge when each become ions
a. Mg
b. O
c. Al
d. P
e. K
+2
-2
+3
-3
+1
57. The ionization energies for an element are listed below
First
second
third
fourth
8 eV
15eV
80eV
109eV
f. Cl
7
f. Br
-1
fifth
141 eV
Based on the ionization energies, the element is most likely to be
a. Sodium
b. magnesium
c. aluminum
d. silicon
e. phosphorus
Practice MQ and FRQ
58. 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
59. How many protons, neutrons, and electrons are in an
Protons
Neutrons
Electrons
a. 26
30
26
b. 26
56
26
c. 56
26
26
d. 56
82
56
56
26Fe
atom?
60. Which of the following is the electron configuration of an excited atom that is likely to emit a quantum of energy?
2
2
6 2
1
2
2
6
2
5
(B) 1s 2s 2p 3s 3p
(A) 1s 2s 2p 3s 3p
2
2
6
2
2
2
6
1
1
(C) 1s 2s 2p 3s
(D) 1s 2s 2p 3s 3p
-1
61. The bond energy of fluorine in 159 kJ mol .
-19
i. Determine the energy, in J, of a photon of light needed to break an F-F bond. 2.64 x 10 J
-1
14 -1
3.98 x 10 s
ii. Determine the frequency of this photon in s
iii. Determine the wavelength of this photon in nanometers 750 nm
b.
Barium imparts a characteristic green color to a flame. The wavelength of this light is 551 nm. Determine the
energy involved in kJ/mol
220 kJ/mol
Review
62. How many significant figures are in each of the following?
a.0.030100 kJ
b. 6.022 x 1023 atoms
c. 100
d. 1001
5
4
1
4
63. 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
0.240 g/mL
12.6 mL
4.570 g
0.3 m2
3
64. The density of pure silver is 10.5 g/cm . If 5.25 g of pure silver pellets are added to a graduated cylinder
containing 11.2 mL of water, to what level will the water in the cylinder rise?
11.7 mL
65. 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.
They have the same number of protons, but one would have more mass than the other one.
66. What is the mass, in grams, of 1.75 x 1020 molecules of caffeine, C8H10N4O2?
0.0563 g
67. Determine the empirical formula of the compound with the following compositions by mass
10.4 percent C, 27.8 percent S, and 61.7 percent Cl
CSCl2