Free Study Guide for
Cracolice • Peters
Introductory Chemistry: An Active Learning Approach
Second Edition
www.brookscole.com/chemistry
Chapter 10
Atomic Theory:
The Quantum Model of the Atom
Chapter 10–Assignment A: Bohr's Model for Hydrogen and Where It Led
We ended Chapter 5 at roughly the midpoint of the 40-year revolution (circa 1890-1930)
that completely changed our view of atoms and the universe. In 1890, physics reflected
Victorian society. There was a place for everything, and everything was in its place. The
universe was composed of matter and energy, and that was that.
Or was it? Rutherford's 1911 model of the atoms placed all the positive charge of the
atom—the protons—in a small, dense nucleus, and all the negative charge—the
electrons—in the space surrounding the nucleus, at relatively great distances. First year
chemistry or physics students knew that like charges repel, and their instructors were
convinced that the Rutherford atom should fly apart. Rutherford's work also showed that
most of the atom was empty space. How could the atoms that make up a solid be empty
space? The solid should collapse. The experiments done by Rutherford's group were
reproducible and so had to be explained.
The Rutherford model of the nuclear atom describes this tiny particle accurately but
incompletely. How exactly are the electrons arranged around the nucleus? People thought
that electrons circled the nucleus in orbits, like planets circle the sun. How large are these
orbits? What are the positions of the electrons with respect to each other? Rutherford's
model did not answer these questions.
It seemed that the rules of classical physics that accurately predicted large-scale behavior
didn't work on the atomic scale. What new rules would replace the classical ones?
In Assignment A the planetary model, as the nuclear atoms is sometimes called, is enlarged
and defined more specifically. The planetary model is then abandoned for the quantum
mechanical model of the atom that conforms more accurately to experimental observations.
Look for these main ideas:
1)
The Bohr model of the hydrogen atom restricts the electron to certain quantized
energy levels. The electron can have certain definite energies, but never may it have
an energy between the quantized values.
2)
The lowest quantized energy level of the atom is called the ground state. All energy
levels above the ground state are called excited states.
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Study Guide for Introductory Chemistry: An Active Learning Approach
3)
The spectrum of a gaseous atom is the result of energy released as electrons in an
excited state drop to a lower energy level.
4)
The quantum mechanical model of the atom identifies principal energy levels
and sublevels within each principal energy level.
5)
Sublevels are divided into orbitals, mathematically described regions of space that
may be occupied by no more than two electrons.
Learning Procedures
Study
Sections 10.1–10.2. Focus on Goals 1–8 as you study.
Strategy
This material is highly abstract in nature. Quantum chemistry can be
difficult because you have no real-world experience with electrons or
quantized energy levels. Don't forget to study the figures and their
captions. These will help you visualize the quantum model. Focus on the
Summary at the end of Section 10.2.
Answer
Questions, Exercises, and Problems 1–18. Check your answers with those
at the end of the chapter.
Workbook
If your instructor recommends the Active Learning Workbook, do
Questions, Exercises, and Problems 1–18.
Chapter 10–Assignment B: Electron Configurations and Lewis Symbols
There is a clear link between the shape of the periodic table, first proposed in 1869 by
Mendeleev and Meyer, and the quantum mechanical model of the atom, proposed almost 60
years later. This assignment helps you to see that link. The main ideas of this assignment
are:
1)
In ground state atoms, electrons fill the lowest energy orbitals first.
2)
Atoms in the same group of the periodic table (same column) have the same highest
occupied sublevel electron configurations, or the same nsxnpy valence electron
configurations. The highest occupied principal energy level value (n) increases as
you go down a group.
3)
Valence electrons are depicted by Lewis symbols, which are also called electron dot
symbols.
4)
Main group elements in the same period have the same highest occupied principal
energy level, given by n. The highest occupied sublevel changes from s to p between
Groups 2A/2 and 3A/3 as you move left to right across a period.
Learning Procedures
Study
Sections 10.3–10.4. Focus on Goals 9–12 as you study.
62
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Chapter 10
Atomic Theory: The Quantum Model of the Atom
Strategy
Electron configurations are best learned by learning the information in
Figure 10.8. Learn to read electron configurations from the periodic table.
Section 10.4 is relatively easy, but learn it well because it is preparing you
for an important concept to come.
Answer
Questions, Exercises, and Problems 19–30. Check your answers with
those at the end of the chapter.
Workbook
If your instructor recommends the Active Learning Workbook, do
Questions, Exercises, and Problems 19–30.
Chapter 10–Assignment C: Trends in the Periodic Table
In Assignment B, you saw the linkage between the periodic table and the quantum
mechanical model of the atom. In this assignment, you will take another look at the periodic
table, using the quantum mechanical model for the deeper insights it can provide.
Look for these big ideas in this assignment:
1)
First ionization energy, the energy required to remove an electron from a neutral
atom, decreases as you move down a group and increases as you move across a
period.
2)
Groups of elements in the periodic table exhibit similar behavior and are called
chemical families.
3)
You will need to identify the following chemical families: alkali metals, alkaline
earths, halogens, and noble gases.
4)
The sizes of atoms in the periodic table increase as you move down a group, but
decrease as you move left to right across a period.
5)
The elements in the periodic table are classified as metals or nonmetals, based on
their chemical behavior.
Learning Procedures
Study
Section 10.5. Focus on Goals 13–18 as you study.
Strategy
Learn the reasons for the trends in the periodic table and not just the pattern
for those trends. It takes a little more time to learn NOW rather than just
memorize a few trends, but understanding these concepts will pay off in the
long term. Don't neglect the figures when you are studying the textbook.
Answer
Questions, Exercises, and Problems 31–45. Check your answers with
those at the end of the chapter.
Workbook
If your instructor recommends the Active Learning Workbook, do
Questions, Exercises, and Problems 31–45.
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No part of this work may be reproduced without the written permission of the publisher.
Study Guide for Introductory Chemistry: An Active Learning Approach
Chapter 10–Assignment D: Summary and Review
When learning how to write electron configurations, remember that all the rules already
exist in picture form in the periodic table. The easiest way to master the material in Chapter
10 is to keep a “clean” periodic table next to you as you study, and use it often.
To place electrons in their proper orbitals, convince yourself that the periodic table is a grid
of the elements. The position of each element on that grid tells you the electron
configuration of that element. The period gives you the highest value of n. For the main
group elements, which are the elements in the U.S. A groups (IUPAC Groups 1–2, 13–18),
the group gives the highest occupied energy sublevel and the number of electrons in that
sublevel.
Some students make unneeded problems for themselves in this chapter by careless
handwriting. The lower case letters s, p, d, and f are symbols for subshells; upper case S, P,
D, and F are not. Also, the number of electrons in a subshell is given by a superscript, like
3p2 ; it is never given by a subscript, like 3p2 .
Learning Procedures
Review
your lecture and textbook notes.
the Chapter in Review and the Key Terms and Concepts, and read the Study
Hints and Pitfalls to Avoid.
Answer
Concept-Linking Exercises 1–9. Check your answers with those at the end
of the chapter.
Questions, Exercises, and Problems 46–47. Include Questions 48–55 if
assigned by your instructor. Check your answers with those at the end of the
chapter.
Workbook
If your instructor recommends the Active Learning Workbook, do
Questions, Exercises, and Problems 46. Include Questions 47–54 if
assigned by your instructor.
Take
the chapter summary test that follows. Check your answers with those at the
end of this assignment.
Chapter 10 Sample Test
Instructions: You may use a “clean” periodic table on this test.
1)
The Bohr model of the atom provides all of the following except
a) an explanation of the line spectra of the elements.
b) a description of electron behavior in many-electron atoms.
c) evidence that electron energies are quantized.
d) calculation of energies of known lines in the hydrogen spectrum.
e) the radius of the electron orbit in a hydrogen atom.
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Chapter 10
Atomic Theory: The Quantum Model of the Atom
2)
Identify the false statement about electron energies:
a) Electron energies are quantized in excited states, but not quantized in the ground
state.
b) Light spectra of the elements are experimental evidence of the quantization of
electron energies.
c) Lines in the spectrum of an element are produced by electrons dropping from a
high energy level to a lower energy level.
d) Energy must be absorbed to raise an electron from ground state to an excited
state.
e) Electrons cannot possess an energy between two quantized energy levels.
3)
Identify the incorrect statement among the following:
a) Except for n = 1, each principal energy level has three p orbitals.
b) There are three sublevels when n = 3.
c) Electrons in the 5d orbitals have higher energies than electrons in the 5p orbitals.
d) The n = 4 sublevels are at higher overall energies than the corresponding n = 5
sublevels.
e) A 3s orbital is at higher energy than a 1s orbital.
4)
Identify the incorrect statement about electron orbitals:
a) An orbital may be occupied by no more than two electrons.
b) All energy sublevels contain the same number of orbitals.
c) For a given atom, the 3p orbitals are larger than the 2p orbitals, but smaller than
the 4p orbitals.
d) At a given sublevel, the maximum number of d electrons is 10.
e) The orbital sketched below is a p orbital.
Z
X
Y
5)
Using the symbolism nsxnpy, where x and y are whole numbers, write the general
electron configuration that is responsible for the family properties of the alkali
metals.
6)
Write the Lewis symbol for oxygen, and then for any alkaline earth element.
7)
Write the electron configuration for an atom of chlorine.
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Study Guide for Introductory Chemistry: An Active Learning Approach
8)
Write the electron configuration for an atom of titanium, Z = 22.
9)
The halogens chlorine and bromine form the insoluble silver chloride, AgCl, and
silver bromide, AgBr, upon reaction with silver. These halogens exhibit these
reactions with silver because:
a) Elemental silver has its outermost electrons in the n = 5 quantum level.
b) Chlorine and bromine are in the same period of the periodic table.
c) Silver is a reactive metal.
d) Chlorine and bromine have similar electron configurations in their highest
energy occupied orbitals.
e) Silver and these two halogens all have similar electron configurations in their
highest energy occupied orbitals.
10)
List the symbols of the following elements in order of increasing first ionization
energy: helium, lithium, boron, oxygen, neon.
11)
Which trio of atomic numbers is arranged in order of increasing atomic size?
a) 54–18–2
b) 52–33–14
c) 8–9–17
d) 17–35–34
e) 12–19–20
12)
Consider atoms with atomic numbers 7, 11, 16, 20, 43, and 53, then. . .
a) write the atomic numbers of the metals listed above and
b) write the atomic numbers of the nonmetals listed above
Answers to Chapter 10 Sample Test
1) b
5)
2) a
3) d
4) b
ns1 . The alkali metals share an ns1 highest occupied energy level electron
configuration. In this description, n is the principal quantum number of the highest
occupied energy level when the atom is in its ground state. For lithium, n = 2; for
sodium, n = 3; for potassium, n = 4, and so forth. The s refers to the s sublevel in
the highest occupied energy level. The p, when needed, refers to the p sublevel in the
same energy level. Superscript numbers identify the number of electrons in those
sublevels. The ns1 indicates that the highest occupied energy level of a ground state
alkali metal is occupied by one electron in the s orbital.
66
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Chapter 10
6)
The Lewis symbol for oxygen is
Atomic Theory: The Quantum Model of the Atom
O
The Lewis symbol for an alkaline earth element is X :
X can be Be, Mg, Ca, Sr, Ba, or Ra.
7)
Cl: 1s2 2s2 2p6 3s2 3p5 or [Ne]3s2 3p5
8)
Ti: 1s2 2s2 2p6 3s2 3p6 4s2 3d2 or [Ar]4s2 3d2
9)
d
10)
Li < B < O < Ne < He
11)
d
12)
a) The metals are Z = 11 (Na), 20 (Ca), and 43 (Tc)
b) The nonmetals are Z = 7 (N), 16 (S), and 53 (I)
67
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No part of this work may be reproduced without the written permission of the publisher.