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Name:_________________________________________________ Period:_____ Date:___________
Periodic Trends Review
1. Why did Mendeleyev leave spaces in his periodic table?
Mendeleev left spaces in his periodic table because he predicted the existence of undiscovered elements.
2. What is the periodic law and how does is relate to the modern periodic table.
The periodic law says that chemical similarities occur at regular intervals. The periodic table is arranged according to
increasing atomic number and similar chemical and physical properties.
3. Based on their locations on the periodic table, would you expect carbon and silicon to have similar chemical properties?
Why or why not?
Both carbon and silicon are found in group 14, so they both have 4 valence electrons which make them react similarly.
4. Identify each property listed below as more characteristic of a metal or nonmetal.
a.
b.
c.
d.
e.
a gas at room temperature. nm
Brittle nm
Malleable m
Poor conductor nm
Shiny m
5. Which element in each pair has a greater atomic radius? Explain your determination for each pair.
a.
Sr, Mg Sr has a greater atomic radius because it has more energy levels than Mg resulting in greater shielding
effect and a larger radius.
b.
Al, Cl Al has a larger radius because it has fewer protons than Cl and the same number of energy levels, so Al
has less Zeff (effective nuclear charge)than Cl resulting in a larger radius.
6. Which element in each pair has greater first ionization energy? Explain each determination.
a. Li, B B has a greater first ionization energy because it is a smaller element than Li, resulting in electrons being
held more tightly to the nucleus, therefore the energy needed to remove an electron will be greater.
b. S, Te S has a greater first ionization energy because it is a smaller element than Te resulting in electrons being
held more tightly to the nucleus, therefore the energy needed to remove an electron will be greater.
7. Arrange the following groups of elements in order of increasing ionization energy and explain each determination.
a. Be, Mg, Sr Sr is the lowest ionization energy (easiest to remove an electron) followed by Mg then Be because
Be is the smallest element so the electrons are held more tightly to the nucleus and take more energy to
remove.
b.
Bi, Cs, Ba Cs has the lowest ionization energy followed by Ba then Bi because Bi is the smallest element so the
electrons are held more tightly to the nucleus and take more energy to remove.
8. How does the ionic radius of a typical metallic atom compare with its atomic radius? Explain.
The ionic radius of a typical metallic ion is smaller than the element’s atomic radius because when metals become ions,
they lose electrons and lose an entire energy level, reducing shielding effect and resulting in a smaller size.
9. Which particle has a larger radius in each atom/ion pair? Explain your determination for each.
a.
S, S2- S-2 has a larger radius because the Zeff stays the same but 2 more electrons are added to the energy
level resulting in a greater electrostatic repulsion between electrons and a larger radius than S.
b.
Al, Al3+ Al has a larger radius because when metals become ions, they lose electrons and “lose” their outer
energy levels resulting in an ion with a smaller atomic radius than the original atom.
10. For which of the following properties does lithium have a larger value than potassium? Explain your determination.
a. first ionization energy Li
b. atomic radius K
c. electronegativity Li
d. ionic radius K
Lithium has a larger first ionization energy and a larger electronegativity than K because Lithium is smaller than K, so it takes
more energy to remove an electron (electrons are held more tightly in a smaller atom) and it also is more advantageous for
an electron to be placed in a smaller atom so electronegativity is also higher because it is easier for a smaller atom to
attract electrons than larger atoms.
11. Which would be more reactive out of each pair and explain WHY?
a.
K, Na K is more reactive b/c metals are likely to lose electrons when forming compounds; since K is larger than
Na, K’s electrons are not bonded as tightly and are easier to remove, making K more reactive.
b.
Cl, P Cl is more reactive than P b/c non-metals are likely to gain electrons. It is more advantageous for
electrons to be placed in smaller atoms because electrons are placed closer to the nucleus and feel a greater
Zeff.
12. Write the symbol of the element or elements that fit each description.
a.
a nonmetal in group 14 C
b.
The inner transition metal with the lowest atomic number. La
c.
All of the nonmetals for which the atomic number is a multiple of five. P, Ne, Br
d.
A metal in group 15. Bi
13. In which pair of elements are the chemical properties of the elements most similar? Explain.
a.
b.
Na and Cl
N and P because both are non-metals in the
c. B and O
same group
d. C and Pb
14. Explain why fluorine has a smaller atomic radius that both oxygen and chlorine.
F is smaller than O b/c F has a greater Zeff, pulling the electrons closer to the nucleus. F is smaller than Cl b/c Cl has
more e- energy levels resulting in a larger shielding effect and larger radius.
15. Would you expect metals or nonmetals in the same period to have higher ionization energies? Explain.
Non-metals would have higher ionization energies than metals in the same period because non-metals will have a smaller
radius due to greater Zeff and will hold onto electrons more tightly making it more difficult to remove electrons and
resulting in a greater ionization energy.
16. In each pair, which ion is larger? Explain your reasoning for each.
a. Ca2+, Mg2+ Ca+2 b/c larger atoms form ions with larger radii (more shielding effect)
b.
Cl-, P3- P-3 is larger because it has less Zeff and greater e- repulsion due to 3 electrons.
c.
Cu+, Cu2+ Cu+ is larger b/c both have the same Zeff, but Cu+2 has less electrostatic repulsion because it has one
less electron.
17. Explain what an element’s valence configuration indicates about the element’s position on the periodic table.
The electron configuration indicates an element’s period and group on the table.
18. The bar graphs below shows the relationship between ionic and atomic radii for group 1 and group II elements.
a. Describe and explain the trend in atomic radii within the groups.
As you go down a group on the periodic table, the atomic radii increases due to increased shielding effect. As you
go across the periodic table, the size of the atomic radii decreases due to increased Z eff.
b. Explain the difference between the size of the atoms and the size of the ions.
When metal atoms form ions, the ions they form are smaller than the size of the original atom because outer shell
electrons are lost resulting in the loss of an entire energy level.
19. The graphs below show the relationship between the electronegativities and first ionization energies for period 2 and 3
elements.
a. Based on the data for these two periods, what is the general trend for each of these two values?
As atomic number increases across a period, ionization energy also increases. As Atomic number increases across a period,
electronegativity also increases. (harder to remove an electron from a smaller atom; more advantageous for an electron to
be placed in a smaller atom.)
b. Describe and explain the trend found in the noble gases Noble gases have the highest ionization energies because
they are the most stable elements on the periodic table, making it more difficult to remove an electron. Noble gases
do not have an electronegativity value because they are stable and do not attract electrons readily.
20. Explain why it takes more energy to remove a 4s electron from zinc than from calcium.
Zinc has more protons (greater Zeff) and therefore holds the 4s electron more tightly.
21. Which would have a smaller radius; a Mg2+ ion or a Na+ ion? Explain.
Mg+2 would have a smaller radius because it has more protons (greater Zeff) and the same number of electrons as Na +
22. Write a symbol for both a cation and for an anion that are isoelectronic with krypton and label which one is the cation
and which is the anion.
Cation: Rb1+ , Sr2+
Anion: Se2- , Br1-
23. Is it possible for a cation to be isoelectronic with an anion from the same period? Explain.
No. Since cations lose electrons and anions gain electrons, cations must always come from a higher period than anions in
order to be isoelectronic with the same element.
24. There is a large jump between the second and third ionization energies of magnesium. There is a large jump between
the third and fourth ionization energies of aluminum. Explain this observed difference.
Both large jumps occur when electrons are removed from inner energy levels instead of outer energy levels. It takes much
more energy to remove inner energy level electrons.
25. Explain each of the following comparisons.
a. Calcium has a smaller second ionization energy than does potassium.
Removing 2 electrons from Ca still involves only removing outer shell electrons whereas removing the 2 nd electron
from K involves removing an inner shell electron which takes much more energy.
b. Lithium has a larger first ionization energy than does Cesium.
Lithium has a larger first ionization energy than Cs because Li is a smaller atom, so electrons are held more tightly,
making electrons harder to remove which takes more energy.
c. Magnesium has a larger third ionization energy than does aluminum.
Removing a third electron from magnesium takes more energy than removing a third electron from aluminum
because the third electron from magnesium is an inner energy level electron.