Name ____________________________________________________ Period ____

Unit 1 Lesson 2: Periodicity Summative Assessment

Goal: 1.2.1.a I can describe how the number of energy levels in an atom affects the approximate distance between the nucleus and the outermost electrons.

Goal 1.2.1.b I can describe the relationship between the number of protons in a nucleus and the attraction for the electrons that surround it.

Goal 1.2.1.c I can describe how the Coulombic force depends on the charge of the nucleus, the charge of the electrons and the distance between them.

Goal 1.2.2.a I can explain how the force of attraction of the nucleus for its electrons affects the atomic radius.

Goal 1.2.2.b I can explain how the force of attraction of the nucleus for its electrons affects the ionic radius.

Goal 1.2.2.c I can explain how the force of attraction of the nucleus for its electrons affects the ionization energy.

Goal 1.2.2.d I can explain how the force of attraction of the nucleus for its electrons affects the electron affinity.

Goal 1.2.2.e I can explain how the force of attraction of the nucleus for its electrons affects the electronegativity.

Goal 1.2.2.f I can explain how the force of attraction of the nucleus for its electrons affects the reactivity.

Goal: 1.2.3.a I can explain what region of the table houses the metals and how the  ionization energy trend affects the conductivity of the metals (including alkali metals & alkaline Earth metals).

Goal: 1.2.3.b I can explain what region of the table houses the nonmetals and how the  ionization energy trend affects the conductivity and the number of valence electrons affects reactivity of nonmetals (including halogens & noble gases).

Multiple Choice

1.  The presence of inner-shell electrons cause the outermost electrons of an atom to be less attracted to the nucleus.  This is called electron _____.

a. reduction

b. shielding

c. dampening

d. canceling

2.  A single 3s electron in a magnesium atom is shielded by:

a. the 1s and 2s electrons only

b. the 2s and 2p electrons only

c. the 1s, 2s, and 2p electrons

d. the 1s, 2s, and 2p electrons, along with the other 3s electron

3.  Explain why the ionization energy of aluminum is lower than the ionization energy of boron.

a. Aluminum has a greater effective nuclear charge.

b. The electron to be removed from the aluminum atom comes from a different sublevel.

c. The electron to be removed from the aluminum atom is slightly repelled by another electron in the same orbital.

d. There is a greater amount of electron shielding in aluminum.

4.  Why is the ionization energy of potassium lower than the ionization energy of calcium?

a. Calcium has a greater effective nuclear charge.

b. The electron to be removed from the calcium atom comes from a different sublevel.

c. The electron to be removed from the calcium atom is slightly repelled by another electron in the same orbital.

d. There is a greater amount of electron shielding in calcium.

5. The effective nuclear charge of a potassium atom is:

a. 19

b. 18

c. 9

d. 1

6. The effective nuclear charge of a K+ ion is:

a. 19

b. 18

c. 9

d. 1

7. Using Coulomb’s Law, Fe = kq1q2/r2,   as a proton and an electron move farther apart, the force between them

a. increases

b. decreases

Defining the Trends

8. The energy required to remove an electron from an atom.

a. Atomic Radius

b. Ionic Radius

c. Ionization Energy

d. Electron Affinity

e. Electronegativity

9. One half the distance between the nucleii of two adjacent atoms

a. Atomic Radius

b. Ionic Radius

c. Ionization Energy

d. Electron Affinity

e. Electronegativity

10. A measure of the ability of an atom to attract shared electrons when the atom is part of a compound.

a. Atomic Radius

b. Ionic Radius

c. Ionization Energy

d. Electron Affinity

e. Electronegativity

11. One half the distance between the nucleii of two ions in an ionic lattice crystal

a. Atomic Radius

b. Ionic Radius

c. Ionization Energy

d. Electron Affinity

e. Electronegativity

12. Amount of energy released when an atom gains an electron.

a. Atomic Radius

b. Ionic Radius

c. Ionization Energy

d. Electron Affinity

e. Electronegativity

13.  Atomic radius generally _____ from left to right across a period of the periodic table.

a. increases

b. decreases

c. stays the same

d. increases and decreases randomly

14.  Compared to the atomic radius of a potassium atom, the atomic radius of a calcium atom is smaller.  This is primarily a result of the calcium atom having:

a. a greater number of principal energy levels

b. fewer principal energy levels

c. a greater nuclear charge

d. less nuclear charge

15.  Compared to the atomic radius of an aluminum atom, the atomic radius of a boron atom is smaller because of an decrease in:

a. the number of principal energy levels

b. the number of valence electrons

c. the nuclear charge

d. the number of neutrons

16.  The removal of one or more electrons from a neutral atom results in a _____.

a.negatively-charged anion

b.negatively-charged cation

c. positively-charged anion

d. positively-charged cation

17.  Anions are always _____ than the parent atom from which they were formed. Cations are always _____ than the parent atom from which they were formed.  

a. larger; larger

b. larger; smaller

c. smaller; larger

d. larger; larger

18.  Which electron is the first to be removed from any atom?

a. an s electron

b. a p electron

c. a valence electron

d. an electron from the first principal energy level

19.  Ionization energy generally _____ from top to bottom down a group of the periodic table.

a. increases

b. decreases

c. stays the same

d. increases and decreases randomly

20.  Compared to the ionization energy of a lithium atom, the ionization energy of a beryllium atom is larger.  This is primarily a result of the beryllium atom having:

a. more energy sublevels

b. greater nuclear charge

c. greater electron-electron repulsion

d. more principal energy levels

21.  Compared to the ionization energy of a magnesium atom, the ionization energy of a calcium atom is smaller.  This is primarily because the calcium atom:

a. has fewer valence electrons than the magnesium atom

b. is larger than the magnesium atom

c. has more electron shielding than the magnesium atom

d. has a greater nuclear charge than the magnesium atom

22.  The removal of an electron from a neutral atom _____ energy, while the addition of an electron to a neutral atom usually _____ energy.

a.absorbs; absorbs

b. releases; absorbs

c. absorbs; releases

d. releases; releases

23.  Which equation shows the electron affinity of flourine?

a. F - → F + e- + energy

b. F + e- → F - + energy

c. F + e- + energy → F -

d. F - + energy → F + e-

24.  Which group has the highest electron affinities?

a. noble gases

b. alkali metals

c. group 13

d. halogens

25.  There are frequent exceptions to the period and group trends for electron affinity.

a. True

b. False

26.  Chemical bonds involve the transfer or sharing of _____.

a. valence electrons only

b. all electrons

c. inner-shell electrons only

d. one electron only

27.  Metals tend to have generally _____ electronegativities, while nonmetals tend to have generally _____ electronegativities.

a. high; high

b. high; low

c. low; high

d. low; low

28.  Which group has the highest electronegativities?

a. alkali metals

b. alkaline earth metals

c. halogens

d. noble gases

29.  Which group has the lowest electronegativities?

a. alkali metals

b. alkaline earth metals

c. halogens

d. noble gases

30.  Metals are the largest of the three classes of elements.

a. True

b. False

31.  Properties of most metals include

a. high melting point.

b. ability to conduct heat.

c. shiny appearance.

d. all of the above

32.  Some metals are gases at room temperature.

a. True

b. False

33.  The properties of metals depend mainly on their number and arrangement of neutrons.

a. True

b. False

34.  Examples of nonmetals include

a. carbon.

b. phosphorus.

c. sulfur.

d. all of the above

35.  Properties of nonmetals include

a. high boiling point.

b. ability to conduct heat.

c. dull appearance.

d. none of the above

36.  All nonmetals are very reactive.

a. True

b. False

37.  Nonmetals are located on the right side of the periodic table.

a. True

b. False

Short Answer

38.  Why are metals good conductors of electricity? Explain in terms of a periodic trend.

39.  The following ions all have the same number of electrons.  Arrange them in order from smallest to largest and explain why this is the order: Mg2+, F−, N3−, Na+, Al3+, O2−

40.  Arrange the following atoms in order from smallest to largest and explain why:  Ge, Br, K, Se, Cs

41.  Which atom/ion of each pair below is smaller? Why?

Cl−/P3− 

O2−/O

B3+/B