 ```Chapter 11
Section 11.1: Team Learning Worksheet
1. As the wavelength increases, the frequency and energy decrease (frequency and
energy are directly related). The speed is constant (it is the speed of light). These are
explained in the book, but have the students explain them in their own words.
2. This phenomenon is explained in the text, but the students should explain it in their
own words. Because energy levels are quantized, excited electrons that return to
lower energy states can emit photons of only certain energy levels. Because of this,
only certain wavelengths of light are emitted; thus we see only certain colors.
3. White light is a combination of all colors of visible light. In an incandescent light
bulb, the wire in the bulb is heated as current passes through it. As the heat dissipates
from the wire, energy of all wavelengths in the visible spectrum is emitted and we see
white light.
4. An excited state is a state with excess energy. An electron that is in an excited state is
higher in energy and therefore less stable than an electron in the ground state (lowest
energy state). Students should understand the concepts of lowest energy state and
stability from Chapter 10 and know how to apply them to an atom.
5. Light seems to behave both as a stream of discrete particles of energy (called
photons) and as though it consists of waves.
Section 11.2: Team Learning Worksheet
1. Saying that energy levels are quantized means that only certain energy levels exist.
Make sure the students realize that this is counter-intuitive (even to the scientists who
discovered it). It implies that a hydrogen atom, for example, can exist in one state or
another, but not “in between” states. We are not used to thinking of the world
working in this way (we have a continuous view of matter, time, and energy) and the
students need to know this.
2. The fact that excited hydrogen emits only certain colors supports the idea of
quantized energy levels (as do flame tests, which the students may do as a laboratory
activity). The students should be able to explain this evidence.
3. We do not know how electrons move in an atom, but we are sure that they do not
move in circular orbits. Our model for describing an electron in an atom is based on
probability. We are likely to find the electrons close to the nucleus (although the
electrons also repel each other and so we should expect them to “spread out” to some
degree). Students should understand that a universe system of the atom (with the
nucleus as the sun and electrons as planets) is simply not correct.
4. The firefly analogy shows us that although we can’t tell exactly where the firefly is
going next, or predict its “flight pattern,” we do have reason to believe where the
firefly is likely to be. The same holds true for electrons—we do not know exactly
how they move in an atom, and we cannot predict with certainty their “flight pattern”;
however, we have an idea of where they are likely to be. Students should be able to
explain this analogy in their own words. This problem should also help the students
become more active and critical readers of the text.
5. A good way to think about an orbital is to think of a probability map of an electron
(or electrons) around the nucleus. Although the name is unfortunate, an orbital does
not imply an “orbit” like a planet around the sun. An orbital is also not a physical
thing—it is a region of high probability of finding an electron.
Section 11.3: Team Learning Worksheet
1. An orbital is a probability map of one or more electrons around the nucleus. The
probability of finding an electron decreases the further we move from the nucleus, but
it never becomes zero. The orbital, then, does not have a size that can be exactly
defined, so we arbitrarily choose a size that contains a 90% probability of finding the
electron.
2. The energy level can also be thought of as the energy state. Atoms have discrete
energy levels. The energy levels also have sublevels, where we designate the
different orbitals. For example, the second energy level has two sublevels, which are
constituted by the s orbital and the p orbitals (there are three distinct p orbitals). As
we go from energy level, to sublevel, to orbital, we are further specifying the
probable location of an electron.
3. If energy is added to a sample of hydrogen, the hydrogen atoms become excited and
the electrons move to higher energy levels. Thus, although a “ground state” hydrogen
atom has the electron in the 1s orbital, the electron can be excited to any level and
orbital.
4. Each p orbital can hold two electrons. For each energy level (except n = 1), there are
three p orbitals, so a total of six electrons can be in the p orbitals.
5. The 1s and 2s orbitals are both spherically shaped, but the 2s orbital is larger than the
1s orbital; in addition, the 2s orbital is higher in energy than the 1s orbital. A 2p
orbital has a different shape from the 2s orbital; for the hydrogen atom these two
orbitals are equal in energy. For elements other than hydrogen, the 2s orbital is lower
in energy than the 2p orbitals (each of the three 2p orbitals is equal in energy to the
others).
Section 11.4: Team Learning Worksheet
1. This is explained in the text, but the students should be able to explain it in their own
words. The students should understand that certain “blocks” of the periodic table
correspond to different orbitals of the last electron added and that the rows of the
periodic table correspond to different energy levels.
2. Students often memorize trends for ionization energy and atomic radius, but it is also
important that they can explain them. This question shows the students that we cannot
merely look at the trends in terms of additional protons and electrons. Students should
know that we need to consider the number of protons and the energy level of the
valence electrons. If we compare the atomic radius of lithium and fluorine, the
valence electrons are in the same energy level. However, fluorine has additional
protons, so the “pull” from the positive nucleus to the negative electrons is greater
and the electrons are held closer (smaller radius). In comparing fluorine to iodine, it is
true that iodine has additional protons (and thus a stronger “pull” from the nucleus).
However, the valence electrons of iodine are at a higher energy level than those of
fluorine. Thus, the electrons are further from the nucleus for iodine than fluorine.
3. In writing the shorthand form, the core electrons are represented by the closest noble
gas with fewer electrons. Since the chemically important electrons are the valence
electrons (they are the only ones that interact in a chemical reaction), we need to
know only the number of valence electrons and their orbitals, which we get from the
shorthand form.
4. False. All ionization energies are endothermic because energy is always required to
remove an electron from an atom. Students often overuse the phrase “wants to be like
a noble gas” when discussing the loss or gain of electrons. We can predict the number
of electrons that some metals and nonmetals lose and gain by knowing the number
needed to achieve a noble gas electron configuration. However, this does not mean
that an atom will spontaneously gain or lose one or more electrons.
5. Atom A has the higher ionization energy. The electrons in Atom A are lower in
energy, or more stable. Thus, more energy will be required to remove these electrons
from the atom (which is the definition of ionization energy). Students often get this
backwards; that is, many students will say that less energy is required to remove an
electron that is lower in energy.
6. The correct answer is “a.” Choices that are known to be tempting to the students are
included here to test their understanding. The phrase “wants to be like a noble gas” is
often cited as support for both “b” and “c.” Choice “d” is also popular. The phrase
“an atom will achieve an electron configuration like a noble gas” helps us to predict
that a magnesium ion will have a 2+ charge; however, it does not explain it. An
explanation is beyond the material at this time (it involves lattice energy of ionic
compounds). However, differentiating between predicting and explaining is also
useful.
```