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Electrons in Atoms & Periodic Relationships
Chapter 13-14 Assignment & Problem Set
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Electrons in Atoms & Periodic Relationships
Chapter 13-14 Assignment & Problem Set
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Study Guide: Things You Must Know
Vocabulary (know the definition and what it means):
 electronic structure
 Bohr planetary model
 quantized energy level
 bright line or emission spectrum
 ground state
 excited state
 photon
 Quantum theory or wave theory
 atomic orbital
 electron cloud
 electron configuration
 principal quantum number or principal energy
level
 integer value
 energy sublevel or subshell
 sublevel shape
 Aufbau Principle
 Pauli Exclusion Principle
 Hund’s Rule
 parallel spin vs. opposite spin of an electron
 paired vs. unpaired electrons
 orbital box diagram
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electron configuration
valence electrons
metals
nonmetals
metalloids
representative elements (s & p-block)
transition elements (d-block)
inner transition elements (f-block)
alkali metals
alkaline earth metals
halogens
noble or inert gases
valence electron
Periodic trend
Atomic radius
ionization energy
ionic radius
electronegativity
metallic character of an element
polar molecule
Learning Objectives:
 the historical development of atomic theory, especially the contributions of Bohr and Schrodinger
 Bohr’s planetary model of the atom and how this explains the bright line spectrum of hydrogen
 the electron energy levels in an atom are quantized
 electrons absorb energy when moving to higher energy levels and release energy when moving to lower energy
levels
 when electrons move to a lower energy level, the energy is released in the form of light (photons). The energy
of the light is related to the difference in the starting and ending energy levels of the electron.
 in the quantum or wave model of the atom, an atomic orbital is a region in space with high probability of
finding an electron, and each atomic orbital can hold two electrons.
 in the quantum or wave theory, the location of an electron is designated by its principle quantum number
(energy level) and sublevel (subshell).
 the principle quantum number gives the energy level of the electron, with higher energy level electrons located
farther from the nucleus
 the electron sublevel give the shape of the atomic orbital
 how to determine the electron configuration of an atom using the Aufbau principle, Pauli Exclusion Principle,
and Hund’s Rule.
 how to write the electron configuration, shorthand electron configurations, and draw a box diagram for an
atom.
 how to determine electron configurations using the pattern of elements on the Periodic Table.
 how to determine electron configurations as given on the Regents chart.
 how to determine that an element is in the excited state, given the electron configuration
 that atoms become ions by losing or gaining electrons to attain a Noble gas electron configuration.
Electrons in Atoms & Periodic Relationships
Chapter 13-14 Assignment & Problem Set
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 how to explain neon lights, bright line spectra, and flame tests in terms of electron transition between energy
levels.
 how Mendeleev’s Periodic Table arranged the elements and how the modern Periodic Table arranges the
elements
 the information given for each element on the Regents Periodic Table
 special names for elements in Groups 1, 2, 17, 18
 what characterizes the representative, transition, and inner-transition elements
 transition metal salts dissolve in water to give colored solutions
 what is meant by atomic radius, and how to determine atomic radii of elements using Table S
 how to explain the trend in atomic radius across a Period in terms of nuclear charge
 how to explain the trend in atomic radius down a Group in terms of electron energy levels
 the definition of ionization energy, and how to determine ionization energy of elements using Table S
 how to explain trends in ionization energy across a Period and down a Group in terms of the distance the
valence electron are from the nucleus
 cations are smaller than their atom (less electrons) and anions are larger than their atom (more electrons
 the definition of electronegativity, and how to determine electronegativity of elements using Table S
 how to explain trends in electronegativity across a Period and down a Group in terms of the distance the
valence electron are from the nucleus
 the definition of metallic character
 how to explain trends in metallic character across a Period and down a Group in terms of ionization energy
Key Reference Tables
 Periodic Table
 Table S
Electrons in Atoms & Periodic Relationships
Chapter 13-14 Assignment & Problem Set
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•Read Chapter 13 & 14, except skip “Light and Atomic Spectra” p372-375, “The Quantum Concept and the
Photoelectric Effect” pp376-378, and “Quantum Mechanics” pp381-382. Be sure not to skip over “An
Explanation of Atomic Spectra” pp379-380.
•Lab 12: Flame Test
•Lab 13: A Tour of the Periodic Table
•Regents Tables
Table S: Properties of Selected Elements
•Warm-ups and problems will be collected before you take the test. Answer all 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.
Atomic Models
1. List a major contribution of each of these scientists to the understanding of the atom:
a. Dalton
b. Thomson
c. Rutherford
d. Bohr
e. Schrodinger
2. Which subatomic particles did Thomson include in the plum-pudding model of the atom?
3. How did Bohr explain why an electron did not fall into the nucleus?
Quantum Mechanical Model of the Atom
4. What is an atomic orbital?
5. The energies of electrons are said to be quantized. Explain what this means.
6. What is the shape of:
a. s orbital
b. p orbital
Electrons in Atoms & Periodic Relationships
Chapter 13-14 Assignment & Problem Set
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7. Give the name of and describe the three rules that govern the filling of atomic orbitals by electrons?
8. What is the significance of the boundary of an electron cloud?
9. What is meant by 3p3?
10. How many orbitals are in the following sublevels?
a. 3p sublevel
b. 2s sublevel
c. 4f sublevel
11. An atom of an element has two electrons in the first energy level and five electrons in the second energy level.
Write the electron configuration for this atom and name the element. How many unpaired electrons does an
atom of this element have?
12. Give the symbol and names of the elements that correspond to these configurations.
a. 1s22s2 2p6 3s1
b. 1s2 2s2 2p6 3s2 3p2
c. 1s2 2s2 2p4
d. 1s2 2s2 2p6 3s2 3p6 3d2 4s2
13. Write the complete quantum mechanical and Regents electron configuration for each atom.
a. lithium
b. fluorine
c. potassium
14. Write the complete quantum mechanical and Regents electron configuration for the elements that are identified
only by these atomic numbers.
a. 15
b. 9
c. 18
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15. What is the maximum number of electrons that can go into each of the following sublevels?
a. 2s
b. 3d
c. 4p
d. 4f
16. How many electrons are in the second energy level of an atom of each element?
a. chlorine
b. phosphorus
c. potassium
17. Write the complete quantum mechanical and Regents electron configuration for atoms of these elements.
a. selenium
b. vanadium
c. calcium
18. How many valence electrons (electrons in the highest occupied energy level) do these atoms have?
a. barium
b. sodium
c. aluminum
d. nickel
Atomic Spectra
19. Compare the ground state and the excited state of an electron.
20. Explain the origin of the atomic emission spectrum (bright line spectrum) of an element.
21. List the colors of the visible spectrum in order of increasing wavelength.
The Periodic Table
22. What is the difference between the way Mendeleev organized the Periodic Table and the way it is done today?
23. Explain how an element's outer electron configuration is related to its position in the periodic table.
Electrons in Atoms & Periodic Relationships
Chapter 13-14 Assignment & Problem Set
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24. Why do the elements potassium and sodium have similar chemical and physical properties?
25. What are the representative elements?
26. Which of the following are representative elements? Na, Mg, Fe, Ni, Cl?
27. Categorize each element as a representative element, a transition metal, or a noble gas.
a. 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6
b. 1s2 2s2 2p6 3s2 3p6 3d10 4s2
c. 1s2 2s2 2p6 3s2 3p2
28. Which of the following are transition metals: Cu, Sr, Fe, Au, Al, Ge, Co?
29. What are the symbols for all the elements that have an outer configurations of s2 p5?
Trends in Atomic and Ionic Size
30. What is the trend in atomic radius going left to right across any period? Explain.
31. Indicate which element in each pair has the greater atomic radius.
a. sodium, lithium
b. strontium, magnesium
c. carbon, germanium
d. selenium, oxygen
32. Arrange these elements in order of decreasing atomic size: sulfur, chlorine, aluminum, and sodium.
33. How does the ionic radius of a typical metallic atom compare with its atomic radius?
34. How does the ionic radius of a typical anion compare with the radius for the corresponding neutral atom?
Electrons in Atoms & Periodic Relationships
Chapter 13-14 Assignment & Problem Set
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35. Which particle has the larger radius in each in each atom/ion pair?
a. Na, Na+
b. S, S-2
c. I, Id. Al, Al+3
36. The Mg+2 and Na+ ions each have ten electrons surrounding the nucleus. Which ion would you expect to have
the smaller radius? Why?
37. Atoms and ions with the same number of electrons are called isoelectronic. Name a cation and an anion that are
isoelectronic with krypton.
Trends in Ionization Energy
38. Define Ionization Energy.
39. What is the trend in ionization energy going left to right across any period?
40. Which element in each pair has the greater first ionization energy?
a. sodium, potassium
b. magnesium, phosphorus
c. lithium, boron
d. magnesium, strontium
e. cesium, aluminum
41. Would you expect metals or nonmetals to have higher ionization energies? Why?
42. There is a large jump between the second and third ionization energies of magnesium. The corresponding large
jump is between the third and fourth ionization energies of aluminum. Explain.
Electrons in Atoms & Periodic Relationships
Chapter 13-14 Assignment & Problem Set
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Trends in Electronegativity
43. Define electronegativity.
44. Which element has the greatest electronegativity? What is its value?
45. What is the trend in electronegativity going down any group? Why?
Metallic Character
46. What is meant by metallic character?
47. Select the more metallic element in each pair.
a. Na or Cs
b. Na or Al
c. Si or S
d. Sr or F
Review
48. Give the number of protons and electrons in each of the following:
a. Cs
b. Ag+
c. Se-2
49. Name or give the formula for the following, and state whether each is molecular (M) or ionic (I).
a. N2O4
b. Cs2O
c. Tin(II) sulfate
d. Ammonium carbonate
e. Pb3(PO4)2
50. Convert the following quantities.
a. 3.7 mg to kg
b. 132 km to mm
c. 62 m/s to km/hr
51. Chlorine has two isotopes: Cl-35 which is 75.8 % abundant and Cl-37 which is 24.2% abundant. Calculate the
atomic mass of chlorine?