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Periodic Trends
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Printed: June 26, 2015
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C HAPTER
Chapter 1. Periodic Trends
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Periodic Trends
Key Concept
The structure of the periodic table is such that certain trends exist in the physical properties of the elements. Atomic
radii decrease left to right across periods and increase top to bottom down groups. Ionization energy, the energy
required to remove an electron from an atom, increases across periods and decreases down groups. Removing
multiple electrons always takes more energy, especially once the noble gas core is reached. Electron affinity is
the energy change that occurs when an atom gains an electron. Cations are always smaller than the parent atom,
while anions are always larger. Electronegativity, the ability to attract shared electrons, increases across periods and
decreases down groups. The most reactive metals are in the lower left portion of the table, while the most reactive
nonmetals are in the upper right portion.
Standards
Lesson Objectives
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Learn the periodic trends for atomic radius.
Know the relationship between group number and valence electrons.
Describe how ions are formed.
Learn the periodic trends for ionization energy.
Explain how multiple ionization energies are related to noble gas electron configurations.
Describe electron affinity.
Predict the effect that ion formation has on the size of an atom.
Learn the periodic trends for electronegativity.
Lesson Vocabulary
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anion: A negatively charged ion.
atomic radius: One-half the distance between the nuclei of identical atoms that are bonded together.
cation: A positively charged ion.
electron affinity: The energy change that occurs when a neutral atom gains an electron.
electronegativity: A measure of the ability of an atom to attract the electrons when the atom is part of a
compound.
• ion: An atom or group of bonded atoms that has a positive or negative charge.
• ionization energy: The energy required to remove an electron from an atom.
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Teaching Strategies
Introducing the Lesson
Discuss the meaning of trend and emphasize that trends can have exceptions. The key to understanding the trends
is to know the meaning of nuclear charge and to remember that successive principal energy levels consist of larger
orbitals. Use visuals on the board to emphasize atomic radii differences relating to nuclear charge and number of
energy levels. For this purpose, atoms can be drawn in a Bohr-model fashion.
Make sure that students understand ion formation before discussing ionization energy and ionic sizes. Stress that
ions will be very important in future study of chemical compounds.
Demonstration
The link below is for a lab activity that shows the reactivity of the halogen group. The goal is to demonstrate that
chlorine is more reactive than bromine, which is more reactive than iodine. The specifics of the lab involve being
able to write equations for single replacement reactions, which is covered in a later chapter. However, the activity
can be adapted to a demonstration here. If so, perform the displacement reactions in test tubes rather than well
plates. The demonstration reinforces the information from the end of the lesson that the more reactive nonmetals are
located at the top of a given group.
http://www.nuffieldfoundation.org/practical-chemistry/reactions-aqueous-solutions-halogens
Common Misconceptions
Students tend to have difficulty with the concept of electron shielding and its effect on the properties discussed in
the lesson. Pick a student from the back of the room and call him/her the valence electron. Call yourself the nucleus.
There is an attraction between the two of you. However, all the other students are the inner electrons and are getting
in the way of that attraction. Therefore, the “valence electron” student will find it much easier to leave (be ionized).
Taking it Further
The concept of effective nuclear charge (Ze f f ) can be used in concert with electron shielding. In simplified form the
effective nuclear charge is equal to the atomic number (Z) minus the number of inner or non-valence electrons. For
example, lithium has 3 protons, 2 inner electrons, and 1 valence electron. For lithium, the Ze f f = 3 –2 = 1.
The group trend for effective nuclear charge is that it is constant within a group. The element cesium is in the same
group as lithium. Its effective nuclear charge is Ze f f = 55 − 54 = 1. This is the same value as lithium. Despite the
huge increase in nuclear charge, the inner electrons “cancel out” most of that attractive force and leave the valence
electron of cesium very susceptible to removal: even more so than for lithium. The low Ze f f of the alkali metals is
why they have the lowest ionization energies of any group.
The period trend for effective nuclear charge is that it increases by one across a period. For example, the element
nitrogen is in the same period as lithium. It has 7 protons, 2 inner electrons, and 5 valence electrons. The Ze f f = 7 –2
= 5. The higher effective nuclear charge means that nitrogen holds on to its valence electrons more tightly and thus
has a higher ionization energy. Ionization energy increases from left to right across a period because the effective
nuclear charge increases.
Effective nuclear charge is more complicated than what has been discussed here, and takes into account sublevel
differences as well. However, the use of the concept may help students understand electron shielding.
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Chapter 1. Periodic Trends
Science Inquiry
Students can gain insight into the trends by making graphs of the various properties as a function of atomic number.
Since these graphs are shown and discussed in the FlexBook® student edition, it is best to do this activity as an
introduction to the chapter. The link below contains directions for making four graphs. If time is limited, do only
atomic radius and ionization energy.
Download the resource: http://mcs.monet.k12.ca.us/schools/TeacherWebsite/7-12/Durham.K/Worksheets%20preAP/C
hapter%204/Periodic%20Trends%20graphing%20lab.doc
Differentiated Instruction
Have students practice their knowledge of trends by giving them trios of elements, either all in the same group or
all in the same period. For example, Na, S, and Al or P, Sb, and N. Have them order the elements by increasing
atomic radius, increasing ionization energy, increasing electronegativity. When they have mastered that, challenge
them with more elements that they can still figure out just by knowing the trends.
Enrichment
The Alien Periodic Table is an entertaining activity for all students. It allows them to use their knowledge of the
periodic table to construct one from clues given to them about so-called alien elements.
http://www.nclark.net/alienperiodictable___kulis.pdf
Reinforce and Review
Lesson Worksheets
Copy and distribute the Lesson 6.3 worksheets in CK-12 Chemistry –Intermediate Workbook. Ask students to
complete the worksheets to reinforce lesson content.
Lesson Review Questions
Have students answer the Lesson Review Questions at the end of Lesson 6.3 in CK-12 Chemistry –Intermediate
FlexBook® resource.
Points to Consider
Compounds result from the chemical combinations of elements. The nature of chemical compounds depends on the
types of elements that are combining.
• What type of compounds results when a metal reacts with a nonmetal?
• What type of compound results when two nonmetals react with each other?
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