Unit 6 – Modern Atomic Theory
Knowledge Objectives:
Students will know…
1. The electromagnetic spectrum and the visible spectrum within it.
2. The relationship between energy, color, and wavelength of light.
3. The wave-particle nature of light.
4. What happens when electrons move between ground states and excited states.
5. Energy levels of all atoms are quantized and how this relates to the various line-emmision spectra of each element.
6. The wave mechanical model of the atom.
7. How to predict the location of an element on the periodic table based on ionization energies.
8. The various types of electron orbitals.
9. The quantum numbers.
10. Pauli’s Exclusion Principle, Aufbau’s Principle, and Hund’s Rule and how they apply to locating electrons around an atom.
Skill Objectives:
Students will be skilled at…
1. Explaining how the periodic table of the elements is organized and labeling several common families of elements.
2. Differentiating between metals, nonmetals, and metaloids.
3. Discussing the following periodic trends: atomic radius, ionic radius, ionization energy, electron affinity, electronegativity, and
metalic character.
4. Writing full electron configuration for various elements.
5. Writing orbital notation electron configuration for various elements.
6. Writing noble gas electron configuration for various elements.
7. Using the electron configurations to identify the core and valence electrons for an element.
Important Vocabulary:
Alkali Metal
Alkaline Earth Metal
Atomic Radius
Aufbau Principle
Bohr Model
Core Electrons
Diatomic Element
Electromagnetic Radiation
Electromagnetic Spectrum
Electron Configuration
Electronegativity
Excited State
Frequency
Ground State
Group
Halogen
Hund’s Rule
Ionic Radius
Ionization Energy
Line-emmision Spectrum
Main Group Elements
Metal
Metalic Character
Metaloid
Noble Gas
Noble Gas Configuration
Noble Metal
Nonmetal
Orbital
Orbital Notation
Pauli Exclusion Principle
Period
Periodic Table of the Elements
Photon
Principle Energy Level
Quantized
Quantum Number
Representative Elements
Sublevel
Transition Element
Valence Electron
Wave Mechanical Model
Wavelength
64. List several common types of electromagnetic radiation in order from longest wavelength to shortest wavelength.
65. As the frequency of a wave increases, the wavelength _____________________.
66. Waves with a higher frequency have __________________ energy.
67. Explain what is meant by the wave-particle nature of light.
68. What is the difference between an excited and a ground state? When an atom returns to its ground state, what
happens to the excess energy of the atom? Draw a diagram to demonstrate the change.
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69. Are the colors of flame tests due to taking in energy or releasing energy? Explain.
70. Do atoms in excited states emit radiation randomly, at any wavelength? Explain.
71. What does it mean when we say energy levels are quantized?
72. What are the essential points of Bohr’s theory of the structure of the hydrogen atom?
73. Write the location on of each of the following families or classifications of elements on a periodic table: metals,
nonmetals, metalloids, alkali metals, alkaline earth metals, halogens, noble gases, transition metals, & diatomic
elements.
74. Complete the following table:
Complete the following table:
Property
Trend Period
Atomic Radius
Ionic Radius
Electronegativity
Ionization Energy
Trend Group
75. Arrange the following atoms from largest to smallest atomic radius, and from highest to lowest ionization
energy.
a. Cs, K, Li
b. Ba, Sr, Ca
c. I, Br, Cl
d. Mg, Si, S
76. Explain the difference among the terms energy level, sublevel, and orbital.
77. How many electrons can be placed in
a. a given s subshell? _________
b. a given p subshell? _________
c. a specific p orbital? _________
d. a given d subshell? _________
e. a given d orbital? _________
f. a given f subshell? _________
78. What are the differences between the 2s orbital and the 1s orbital for an atom? How are they similar?
79. How does the energy of a principal energy level depend on the value of n? Does a higher value of n mean a
higher or lower energy?
80. The number of sublevels in a principal energy level (increases/decreases) as n increases.
81. Define Pauli’s Exclusion Principle and explain how it is used in creating electron configurations.
82. Define Aufbau’s Principle and explain how it is used in creating electron configurations.
83. Define Hund’s Rule and explain how it is used in creating electron configurations.
84. Write complete electron configuration diagrams for the following:
19
81. Define Pauli’s Exclusion Principle and explain how it is used in creating electron configurations.
82. Define Aufbau’s Principle and explain how it is used in creating electron configurations.
83. Define Hund’s Rule and explain how it is used in creating electron configurations.
84. Write complete electron configuration diagrams for the following:
a. As
b. C
85. Write orbital configuration diagrams for the following:
a. Mg
b. P
86. Write noble gas configuration diagrams for the following:
a. Sr
b. Si
c. Fe
d. Lu
87. Why are the valence electrons more important to the atom’s chemical properties than the core electrons? How is
the number of valence electrons in an atom related to the atom’s position on the periodic table?
Knowledge Objectives:
Students will know…
1. How to draw Lewis structures for elements and compounds.
2. The difference between bonding pairs of electrons and lone pairs of electrons.
3. How to define and draw resonance structures.
4. How to write stable electron configurations for ions.
88. Compare and contrast ionic and covalent compounds.
89. Determine whether the following bonds are ionic or covalent. SHOW YOUR WORK.
a. Mg – O
b. Ba – Cl
c. K – I
d. H – C
e. Cl – Cl
37. For each of the following molecules or ions, indicate the bond angle expected between the central
atom and two adjacent hydrogen atoms. H2O bent, 105o
tetrahedral 109o
38. Draw the structures of the following, state bond angle, some may be double bonds.
H2S, SiF4, C2H4, C3H8,
Chapter 11: Gases

Pressure, Volume, Temperature, Amount

STP

Absolute Zero

Laws:
o
Boyles
o
Charles
o
Lussacs
o
Combined
o
Dalton’s Law
o
Avogadro’s Law
o
Ideal Gas Equation

Molar Mass of Gas





Gas Stoichiometry
4. The combined gas law is written
a. P1V1/T1 = P2V2/T2 c. P1V1T1 = P2V2T2
b. P1T1/V1 = P1T2/V2 d. T1V1/P1 = T2V2/P2
____ 5. If a balloon containing 3000 L of gas at 39C and 99 kPa rises to an altitude where the
pressure is 45.5




kPa and the temperature is 16C, the volume of the balloon under these new conditions would be
calculated using the following conversion factor ratios: ____.
a. 2678 L c. 6046 L
b. 1489 L d. 3361 L

____ 7. A sample of 0.50 moles of neon gas has a pressure of 120 kPa at 20 deg C. What volume
does this gas
occupy? (R=8.31 kPa-L/mol-K)
a. 0.69 L c. 5.2 L



b. 2.3 L d. 10.1 L
39. What volume does 4.24 g of nitrogen gas occupy at 58.2oC and 2.04 atm? 2.02L
40. If 45.0 mL sample of gas at 26.5 oC is heated to 55.2 oC, what is the new volume of the gas sample
(at constant pressure)? 49.3 mL
41. What pressure (in atm) is required to compress 1.00 L of gas at 760 mmHg pressure to a volume of
50.0 mL? 20 atm
42. Which of the following has the strongest intermolecular attraction? Li2O, F2, SO2,
43. The difference in the strengths of the intermolecular forces among CH4 and NH3 are mainly due
to..hydrogen bonding
44. The vapor pressure of water at 100 oC is equal to standard atmospheric pressure.
Intermolecular Forces
INTERMOLECULAR FORCES PRACTICE TEST
1.
F2 is a gas and I2 is a solid at room temperature, the difference in their phase is due to
A. H-bonding
B. surface tension
C. size
D. polarity
2.
Dry ice (solid CO2) is a molecular solid held together by _____ forces and easily_____ to form a vapor.
A. dipole, melts
B. dispersion, sublimes
C. covalent bonds, condenses
D. pressure, deposits
3.
Of NH3, PH3, CH4, which has the highest vapor pressure?
A. CH4
B. PH3
C. NH3
4.
The energy to change one mole of a substance from a liquid to a vapor at its boiling point is
A. molar heat of fusion
B. molar heat of vaporization
C. molar heat of boiling
D. critical heat
5.
At a substances critical temperature and pressure, the substance will be in equilibrium between
A. solid & liquid
B. solid & vapor
C. liquid & vapor
6.
On a phase diagram the point where all three lines meet is the
A. boiling point
B. melting point
C. exclamation point
D. triple point
7.
Of H2O, NH3, CH4, which is more volatile?
A. H2O
B. NH3
C. CH4
8.
A diagram showing temperature change over time is a ________
A. heating curve
B. phase diagram
C. energy map
D. vapor pressure chart
9.
Vaporization is A. endothermic
B. exothermic
C. isothermic
D. metaphysical
10.
Boiling occurs when a substances vapor pressure equals
A. melting point
B. critical point
C. atmospheric pressure
D. triple pressure
11.
The boiling point of a substance on Mountain Everest is _____ than its normal boiling point.
A. higher
B. lower
C. same