Name: ____________________________
Block: _____ Date: _________________
Unit 2 Test Review
Complete the following questions on a separate piece of paper. THIS IS AN ASSIGNMENT AND
YOU GET CREDIT FOR COMPLETING IT!
Polarity
1. Define pure covalent, polar covalent, and ionic bonds in terms of polarity. Give two examples of
each type of bond (6 total).
• Pure Covalent – bond between two nonmetals that have very similar electronegativity
values (basically identical elements) Examples: O-O and N-N and there are many more
• Polar Covalent – bond between two nonmetals that have different electronegativity values
(two different elements) Examples: C-O and H-Cl and there are many more
• Ionic – bond between a metal and nonmetal. They have very different electronegativity
values. Examples: Na-Cl and Ca-O and there are many more
2. Use the periodic table to predict the polarity of the following bonds. Put in order from least polar
bond to most polar bond.
a. OBr, OF, OCl, OI
OI, OBr, OCl, OF
b. NS, NSe, NTe, NCl NTe, NSe, NS, NCl
c. CN, CO, CF, CC
CC, CN, CO, CF
3. Identify the following bonds as either polar covalent, covalent, or ionic AND EXPLAIN WHY
a. LiBr
Ionic
Why? Metal and nonmetal
b. Na2O
Ionic
Why? Metal and nonmetal
c. Br2
Pure Covalent Why? Two or more of the same nonmetals
d. CI3
Polar Covalent Why? Two or more different nonmetals
e. O3
Pure Covalent Why? Two or more of the same nonmetals
Polar Covalent Why? Two or more different nonmetals
f. H2O
g. F2
Pure Covalent Why? Two or more of the same nonmetals
h. TeO
Polar Covalent Why? Two or more different nonmetals
i. Magnesium and phosphorus Ionic Why? Metal and nonmetal
j. Sulfur and Iodine
Polar Covalent
Why? Two or more different nonmetals
Molecular Formulas
Complete the following for each of the following molecules:
4. CO2
a. Valence electrons
C – 4x1=4
O – 6x2=12 16 total
b. Lewis Dot/Structural Formula
c. Number of electron groups (around the CENTRAL atom) - 2
d. Number of lone pairs (around the CENTRAL atom) - 0
e. Molecular geometry - Linear
f. Polarity - Nonpolar
5. PI3
a. Valence electrons
P – 5x1=5
I – 7x3=21
26 total
b. Lewis Dot/Structural Formula
c. Number of electron groups (around the CENTRAL atom) - 4
d. Number of lone pairs (around the CENTRAL atom) - 1
e. Molecular geometry – Trigonal pyramidal
f. Polarity - Polar
6. Br2
a. Valence electrons
Br – 7x2=14
b. Lewis Dot/Structural Formula
c. Number of electron groups (around the CENTRAL atom) – No central atom
d. Number of lone pairs (around the CENTRAL atom) – No central atom
e. Molecular geometry - Linear
f. Polarity - Nonpolar
7. HCP
a. Valence electrons
H – 1x1=1
C – 4x1=4
P – 5x1=5
10 total
b. Lewis Dot/Structural Formula
c. Number of electron groups (around the CENTRAL atom) - 2
d. Number of lone pairs (around the CENTRAL atom) - 0
e. Molecular geometry - Linear
f. Polarity - Polar
8. P2
a. Valence electrons
P – 5x2=10
b. Lewis Dot/Structural Formula
c. Number of electron groups (around the CENTRAL atom) – No central atom
d. Number of lone pairs (around the CENTRAL atom) – No central atom
e. Molecular geometry - Linear
f. Polarity - Nonpolar
9. SF6 (this is an exception to the octet rule and sulfur will have 12 electrons around it)
a. Valence electrons
S – 6x1=6
F – 7x6=42 48 total
b. Lewis Dot/Structural Formula
P
c.
d.
e.
f.
Number of electron groups (around the CENTRAL atom) - 6
Number of lone pairs (around the CENTRAL atom) - 0
Molecular geometry - Octahedral
Polarity - Nonpolar
Electron Configurations
Determine if the following configurations are correct. If the configuration is correct name the element; if
not state what is wrong.
10. 1s22s22p63s23p64s24d104p65s1 Should be 3d10
11. 1s22s22p63s3 Can’t have more than 2 electrons in the s orbital
12. [Ar] 5s24d105p5 Should be [Kr] if you’re starting with 5s2
13. [Xe] 6s24f10 Correct
14. How many electrons can the following orbitals hold?
a. “s” orbital – 2
c. “d” orbital – 10
b. “p” orbital – 6
d. “f” orbital – 14
Complete the following for the each of the following atoms:
15. Carbon
a. Complete Electron Configuration – 1s22s22p2
b. Spin (Orbital) Diagram – 1s↑↓ 2s↑↓ 2p ↑ ↑ __
c. Noble Gas Configuration [He]2s22p2
d. Ion Configuration – Carbon will not form an ion
e. Valence Electron Configuration – 2s22p2
f. Lewis Dot Structure 16. Calcium
a. Complete Electron Configuration – 1s22s22p63s23p64s2
b. Spin (Orbital) Diagram – 1s↑↓ 2s↑↓ 2p↑↓ ↑↓ ↑↓ 3s↑↓ 3p↑↓ ↑↓ ↑↓ 4s↑↓
c. Noble Gas Configuration – [Ar]4s2
d. Ion Configuration – Ca2+: 1s22s22p63s23p6
e. Valence Electron Configuration – 4s2
f. Lewis Dot Structure 17. Silver
a. Complete Electron Configuration – 1s22s22p63s23p64s23d104p65s24d9
b. Spin (Orbital) Diagram - 1s↑↓ 2s↑↓ 2p↑↓ ↑↓ ↑↓ 3s↑↓ 3p↑↓ ↑↓ ↑↓ 4s↑↓ 3d↑↓ ↑↓ ↑↓ ↑↓ ↑↓
4p↑↓ ↑↓ ↑↓ 5s↑↓ 4d↑↓ ↑↓ ↑↓ ↑↓ ↑
c. Noble Gas Configuration – [Kr] 5s24d9
d. Ion Configuration – Ag1+: 1s22s22p63s23p64s23d104p65s14d9
e. Valence Electron Configuration – 5s2
f. Lewis Dot Structure 18. Silicon
a. Complete Electron Configuration – 1s22s22p63s23p2
b. Spin (Orbital) Diagram - 1s↑↓ 2s↑↓ 2p↑↓ ↑↓ ↑↓ 3s↑↓ 3p↑ ↑ __
c. Noble Gas Configuration – [Ne]3s23p2
d. Ion Configuration – Silicon will not form an ion
e. Valence Electron Configuration – 3s23p2
f. Lewis Dot Structure 19. Xenon
a. Complete Electron Configuration – 1s22s22p63s23p64s23d104p65s24d105p6
b. Spin (Orbital) Diagram – 1s↑↓ 2s↑↓ 2p↑↓ ↑↓ ↑↓ 3s↑↓ 3p↑↓ ↑↓ ↑↓ 4s↑↓ 3d↑↓ ↑↓ ↑↓ ↑↓ ↑↓
4p↑↓ ↑↓ ↑↓ 5s↑↓ 4d↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 5p↑↓ ↑↓ ↑↓
c. Noble Gas Configuration – [Kr]5s24d105p6
d. Ion Configuration – Xenon will not form an ion
e. Valence Electron Configuration – 5s25p6
f. Lewis Dot Structure 20. Tungsten
a. Complete Electron Configuration – 1s22s22p63s23p64s23d104p65s24d105p66s24f145d4
b. Spin (Orbital) Diagram – 1s↑↓ 2s↑↓ 2p↑↓ ↑↓ ↑↓ 3s↑↓ 3p↑↓ ↑↓ ↑↓ 4s↑↓ 3d↑↓ ↑↓ ↑↓ ↑↓ ↑↓
4p↑↓ ↑↓ ↑↓ 5s↑↓ 4d↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 5p↑↓ ↑↓ ↑↓ 6s↑↓ 4f↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 5d↑ ↑ ↑ ↑ __
c. Noble Gas Configuration – [Xe]6s24f145d4
d. Ion Configuration – W2+: 1s22s22p63s23p64s23d104p65s24d105p64f145d4
e. Valence Electron Configuration – 6s2
f. Lewis Dot Structure -
Periodic Trends
21. Explain the relationship between electrons, protons, and energy levels that cause the ionization
energy trend to occur the way it does (increases up and to the right).
Going across the periodic table the ionization energy will increase because the nuclear charge
(number of protons) increases, pulling the valance electrons tighter to the nucleus and therefore
harder to lose an electron. Going down the periodic table the ionization decreases because the
electrons are further from the nucleus and there is more shielding which makes it so less energy is
required to remove an electron.
22. Which of the following atoms have the highest ionization energy: K, Ca, As, S, or Cl? Cl
23. Explain the relationship between electrons, protons, and energy levels that causes the atomic
radius trend to occur the way it does (increases down and decreases right).
Going across the periodic table, atomic radius tends to get smaller because the nuclear charge
(more protons) increases, which has a greater attraction to the valence electrons which pulls the
valence electrons closer to the nucleus. Going down the periodic table, atomic radius gets bigger
within a group because new energy levels are added which move the valence electrons further
away.
24. Choose the two that are in correct order for atomic radius.
a. P, Si, Na , Cl, Al
d. As, P, N, Sb
b. Cl, P, Si, Al, Na - increasing
e. Sb, As, P, N - decreasing
c. Na, Al, P, Si, Cl
f. N, P, Sb, As
25. Explain HOW the radius of an atom changes when it becomes a cation or an anion and why this
occurs.
Cation – the radius will get smaller because cations will lose their electrons and drop down to
the next energy level
Anion – the radius will get slightly larger because anions gain electrons which will repel each
other and push them a little further away
26. What element is the most electronegative? Fluorine
27. Typically, are metals or nonmetals more electronegative? Nonmetals are more electronegative
28. Choose the correct order from least to greatest electronegativity
a. F < O < Be < Li < B
c. F < Be < B < O < Li
b. Be < O < Li < F < B
d. Li < Be < B < O < F
Electromagnetic Spectrum
29. Define the following:
a. Wavelength – Distance between two consecutive crests
b. Photon – Particle of electromagnetic radiation, emitted when electrons return to the
ground state
c. Frequency – Number of waves (cycles) per second that pass a certain point in space
d. Ground state – Lowest possible energy state
e. Valence electron – the electrons in the outer most occupied energy level of an atom
30. Put the electromagnetic light spectrum in order from greatest energy to least energy.
Violet, Blue, Green, Yellow, Orange, Red
31. Put the visible light spectrum from longest to shortest wavelength.
Red, Orange, Yellow, Green, Blue, Violet
32. Explain the relationship between wavelength, color, frequency, and energy are related. Use what
you saw in the flame test lab to explain your answer. For example, why potassium ions have a
yellow color and copper is green when held in a flame.
When the wavelength gets shorter, the frequency increases which means the energy will increase.
Red has a longer wavelength than violet, so elements that emit a red color will have less energy
than elements that emit a violet color.