CHAPTER 7 - Cengage · Web viewGiven an orbital diagram or electron configuration, decide whether...

CHAPTER 8

Electron Configurations and Periodicity

CHAPTER TERMS AND DEFINITIONSNumbers in parentheses after definitions give the text sections in which the terms are explained. Starred terms are italicized in the text. Where a term does not fall directly under a text section heading, additional information is given for you to locate it.

periodic table* arrangement of the elemental symbols in rows and columns (chapter introduction)

electron configuration particular distribution of electrons among available subshells (8.1)

orbital diagram diagram showing how the orbitals of a subshell are occupied by electrons; orbital can hold at most two electrons and then only if the electrons have opposite spins (8.1)

Pauli exclusion principle no two electrons in an atom can have the same four quantum numbers (8.1)

nuclear magnetic resonance (NMR)* condition wherein an atomic nucleus with a net spin and magnetism aligned with an external magnetic field is caused to change or “flip” to an alignment against the applied field (Instrumental Methods: Nuclear Magnetic Resonance [NMR])

magnetic resonance imaging (MRI)* medical diagnostic tool based on nuclear magnetic resonance (Instrumental Methods: Nuclear Magnetic Resonance [NMR])

ground state* quantum-mechanical state in which an atom is at its lowest energy level (8.2)

excited states* quantum-mechanical states in which an atom is at energy levels other than the lowest (8.2)

building-up (Aufbau) principle scheme used to reproduce the electron configuration of the ground states of atoms by successively filling subshells with electrons in a specific order (the building-up order) (8.2)

building-up order* order in which electrons successively fill the subshells: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d (8.2)

noble gases* Group VIIIA elements (8.2)

alkaline-earth metals* Group IIA elements (8.2)

noble-gas core inner-shell electron configuration corresponding to one of the noble gases (8.2)

pseudo-noble-gas core noble-gas core together with (n 1)d10 electrons (8.2)

valence electron electron in an atom outside the noble-gas or pseudo-noble-gas core (8.2)

valence-shell configurations* arrangements of electrons in the outer ns and np subshells (8.2)

main-group (representative) elements* elements in the A columns of the periodic table, in which an outer s or p subshell is filling (8.2)

d-block transition (transition) elements* ten columns of elements inserted between Groups IIA and IIIA in the periodic table, in which d subshells are filling (8.2)

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166 Chapter 8: Electron Configurations and Periodicity

f-block transition (inner-transition) elements* two rows of elements, each with 14 columns, at the bottom of the periodic table, in which f subshells are filling; they fit between Groups IIIB and IVB of Periods 6 and 7 of the periodic table (8.2)

x-ray spectroscopy* analysis of x rays emitted from a target hit by an electron beam (Instrumental Methods: X Rays, Atomic Numbers, and Orbital Structure [Photoelectron Spectroscopy])

x-ray photoelectron spectroscopy* analysis of the kinetic energies of electrons ejected from a target irradiated with x rays (Instrumental Methods: X Rays, Atomic Numbers, and Orbital Structure [Photoelectron Spectroscopy])

photoelectric effect* ejection of electrons from the surface of a metal when light of the proper frequency shines on it (Instrumental Methods: X Rays, Atomic Numbers, and Orbital Structure [Photoelectron Spectroscopy])

ionization energy* energy necessary to remove an electron from an atom (Instrumental Methods: X Rays, Atomic Numbers, and Orbital Structure [Photoelectron Spectroscopy])

Hund’s rule the lowest-energy arrangement of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before pairing electrons (8.4)

paramagnetic substance substance that is weakly attracted by a magnetic field; generally due to unpaired electrons (8.4)

diamagnetic substance substance that is not attracted by a magnetic field or is weakly repelled by such a field; generally means the presence of only paired electrons (8.4)

ferromagnetism* strong, permanent magnetism of iron objects owing to the cooperative alignment of electron spins in many iron atoms (8.4, marginal note)

eka* Sanskrit word meaning “first” (8.5)

melting point* temperature at which a solid substance changes to a liquid (8.5)

boiling point* temperature at which the vapor pressure of a liquid equals the external pressure (8.5)

periodic law when the elements are arranged by atomic number, their physical and chemical properties vary periodically (8.6)

covalent radii* lengths of atomic radii obtained from measurements of distances between the nuclei of atoms in the chemical bonds of molecular substances (8.6)

effective nuclear charge positive charge that an electron experiences from the nucleus, equal to the nuclear charge but reduced by any shielding or screening from any intervening electron distribution (8.6)

first ionization energy (first ionization potential) minimum energy needed to remove the highest-energy (the outermost) electron from a neutral atom in the gaseous state (8.6)

alkali metals* Group IA elements (8.6)

electron affinity energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion (8.6)

basic oxide oxide that reacts with acids (8.7)

acidic oxide oxide that reacts with bases (8.7)

amphoteric oxide oxide that has both acidic and basic properties (8.7)

alloys* metallic mixtures (8.7, marginal note)


Chapter 8: Electron Configurations and Periodicity 167

chalogens* Group VIA elements (8.7)

halogens* reactive nonmetals with the general molecular formula X2, where X symbolizes a halogen (8.7)

CHAPTER DIAGNOSTIC TEST1. Which of the following electron configurations is(are) incorrect? Give the correct one(s).

a. Al: [Ne] 3s23p1

b. Co: [Ar] 3d84s1

c. Br: [Ar] 4d105s25p5

d. Cs: [Xe] 6s1

e. S: [Ne] 3s23p4

2. Atoms in the same group of the periodic table have the same _____________________________

______________________________________________________________________________

3. Silicon has the following number of valence electrons:

a. two.

b. five.

c. one.

d. eight.

e. none of the above.

4. Determine whether each of the following statements is true or false. If a statement is false, change it so that it is true.

a. The orbital diagram of the nitrogen atom in the ground state is [He] . True/False: __________________________________________________________________________________________________________________________________________

b. Paramagnetism arises when a species has unpaired electrons and is demonstrated experimentally when a sample of the species is attracted into a magnetic field. True/False:___________________________________________________________________________________________________________________________________________

c. According to Hund’s rule, the lowest energy state of an atom will have the maximum number of parallel spins for a given n and l designation. True/False: ____________________________________________________________________________________________________________________________________________________

d. The correct order in increasing size of atomic radii is Na, S, Cs, Ba. True/False: ____________________________________________________________________________________________________________________________________________________

5. Write electron configurations for the following:

a. Sr ________________________________________________________________________

b. Ni _______________________________________________________________________

c. Cl _______________________________________________________________________



6. Choose the answer (or answers) that will complete the following phrase. The Pauli exclusion principle

a. states that no two electrons in an atom can have the same four quantum numbers.

b. states that the ground-state electron configuration of an atom will have the maximum number of parallel spins.

c. implies that the energy ordering of orbitals depends on the numerical values of n and l.

d. excludes the possibility of having two electrons with opposed spins.

e. determines the number of electrons in each orbital.

7. Following is a labeled drawing of the periodic table:

A

B

C

D

Which of the following sets correctly identifies the labeled parts?

a. A—main group, B—nonmetal, C—transition, D—lanthanide elements

b. A—transition, B—main group, C—inner transition, D—actinide elements

c. A—inner transition, B—transition, C—main group, D—lanthanide and actinide elements

d. A—main group, B—transition, C—main group, D—inner-transition elements

e. None of the above is correct.

8. Of the five atoms Rb, O, P, Sr, and Se, which should exhibit the lowest first ionization potential?

a. Rb

b. O

c. P

d. Sr

e. Se



ANSWERS TO CHAPTER DIAGNOSTIC TESTIf you missed an answer, study the text section and problem-solving skill (PS Sk.) given in parentheses after the answer.

1.

a. Should be [Ar] 3d74s2

b. Should be [Ar] 3d104s24p5 (8.3, PS Sk. 2)

2. Number of valence electrons and similar chemical and physical properties (8.5, 8.6, PS Sk. 3)

3. e (8.3, PS Sk. 3)

4.

a. False. [He] (8.1, 8.4, PS Sk. 1, 4) 2s 2p

b. True. (8.4)

c. True. (8.4)

d. False. The correct order is S, Na, Ba, Cs. (8.6, PS Sk. 5)

5.

a. 1s22s22p63s23p63d104s24p65s2, or [Kr] 5s2

b. 1s22s22p63s23p63d84s2, or [Ar] 3d84s2

c. 1s22s22p63s23p6, or [Ar] (8.3, PS Sk. 2)

6. a and e (8.1)

7. e (8.2)

8. a (8.6)

SUMMARY OF CHAPTER TOPICS

8.1 Electron Spin and the Pauli Exclusion Principle

Learning Objectives Define electron configuration and orbital diagram.

State the Pauli exclusion principle.

Apply the Pauli exclusion principle. (Example 8.1)

Problem-Solving Skill1. Applying the Pauli exclusion principle. Given an orbital diagram or electron configuration,

decide whether it is possible or not, according to the Pauli exclusion principle (Example 8.1).

You will need to remember the essence of the Pauli exclusion principle to work problems. In short, it is that no two electrons in an orbital can have the same spin.



Exercise 8.1Look at the following orbital diagrams and electron configurations. Which are possible and which are not, according to the Pauli exclusion principle? Explain.

a.

1s 2s 2p

b.

1s 2s 2p

c.

1s 2s 2p

d. 1s22s22p4

e. 1s22s42p2

f. 1s22s22p63s23p103d10

Known: Pauli exclusion principle (from text)

Solution:

a. Possible.

b. Possible.

c. Impossible; two electrons in a 2p orbital have the same spin.

d. Possible.

e. Impossible; there are four electrons in the 2s subshell.

f. Impossible; there are ten electrons in the 3p subshell.

8.2 Building-Up Principle and the Periodic Table

Learning Objectives Define building-up principle.

Define noble-gas core, pseudo-noble-gas core, and valence electron.

Define main-group element and (d-block and f-block) transition element.



A memory device often called the diagonal rule is shown below. It is a very useful aid for writing the arrangement of electrons in atoms.

The diagonal rule:

To obtain the order in which orbitals are filled, start at the right of the bottom arrow and follow it to its point. Then begin at the top right of the next arrow and follow it to its point, and so forth.

It is a good idea to memorize the following exceptions to the building-up principle and their configurations: Cr and Cu in Period 4 and Ag in Period 5. There are others as well. See Appendix D.

8.3 Writing Electron Configurations Using the Periodic Table

Learning Objectives Determine the configuration of an atom using the building-up principle. (Example 8.2)

Determine the configuration of an atom using the period and group numbers. (Example 8.3)

Problem-Solving Skills2. Determining the configuration of an atom using the building-up principle. Given the atomic

number of an atom, write the complete electron configuration for the ground state, according to the building-up principle (Example 8.2).

3. Determining the configuration of an atom using the period and group numbers. Given the period and group for an element, write the configuration of the outer electrons (Example 8.3).

Mendeleev and Meyer found that when the elements were ordered by atomic weight, properties of the main-group elements recurred every eighth element. (The noble gases were not discovered until around 1900.) They arranged the elements in horizontal rows with like elements under one another. The reason for the recurring similarities, we believe, is the recurrence of a similar electron configuration. In our periodic table, arranged by atomic number, all elements in Group (column) IA have one electron in the s subshell of the highest occupied quantum level, n. All elements in Group IIA have two electrons in the s subshell of the highest occupied quantum level, etc. Thus the Group A number is the number of valence electrons. Moreover, each period of the



table corresponds to a major quantum level. In Period 1, for example, quantum level 1 is the quantum level that fills.

Thus it is the electrons with their particular energies that are the stars of the chemical drama. In Groups IA and IIA, the s subshell (and orbitals) are filling. In Groups IIIA to VIIIA, the p subshell (and orbitals) are filling, except for helium, in which the 1s sublevel is filled. In the transition metals, the d subshell (and orbitals) are filling, and in the inner-transition metals, the f subshell (and orbitals) are filling.

Once you are comfortable with this quantum level, subshell, and orbital correlation with periodic table groups and sections, you can read the building-up order directly from the table. Begin at the upper left with hydrogen, and read across to helium: 1s1 and 1s2. Then begin at the left again with lithium, 2s1, and continue across to neon, 2p6. Skipping down to potassium, which begins Period 4, you have 4s1, 4s2, then 3d1, etc. (remember the exceptions, chromium and copper). Practice doing this until you can write the electron configurations of main-group elements and at least the 3d transition elements directly from the table.

If you look into the subject, you will find that there are several forms of the periodic table. The one in your text is called the long form. The short form takes less space but is much more difficult to read.

In some texts you will see the lanthanides called lanthanoids and the actinides called actinoids. According to the currently accepted rules in naming compounds, an ionic compound of two elements is a salt and ends in -ide. Recall that NaCl is sodium chloride. Some authors feel that the group names should reflect these rules accurately, and since these substances are elements, they should not be named as salts. Other authors, including your textbook authors, believe that history should be preserved and call them by their original names.

Exercise 8.2Use the building-up principle to obtain the electron configuration for the ground state of the manganese atom (Z = 25).

Known: The diagonal rule memory aid; Z, the atomic number = number of protons; number of electrons in the neutral atom equals the number of protons = 25; 2electrons fill an s subshell, 6 fill a p subshell, and 10 fill a d subshell.

Solution: 1s22s22p63s23p64s23d5. To be consistent with the text, we order the shells with 3d before 4s: 1s22s22p63s23p63d54s2, which also could be written [Ar]3d54s2.

Exercise 8.3Using the periodic table on the inside front cover of the text, write the valence-shell configuration of arsenic (As).

Known: Definition of valence-shell configuration; the period gives the number of the highest occupied quantum level; the Group A number gives the number of valence electrons.

Solution: Arsenic is in Period 4 and Group VA. Its valence-shell configuration is 4s24p3.

Exercise 8.4The lead atom has the ground-state configuration [Xe] 4f145d106s26p2. Find the period and group for this element. From its position in the periodic table, would you classify lead as a main-group, a transition, or an inner-transition element?



Known: Valence electrons are in the highest occupied quantum level (shell); the quantum number of this level is the period number; the number of valence electrons gives the column (Group) A number.

Solution: The highest occupied quantum level is 6, so lead is in Period 6. There are 4 valence electrons, 2 in the 6s subshell and 2 in the 6p subshell, indicating GroupIVA. In this position, lead is a main-group element.

8.4 Orbital Diagrams of Atoms; Hund’s Rule

Learning Objectives State Hund’s rule.

Apply Hund’s rule. (Example 8.4)

Define paramagnetic substance and diamagnetic substance.

Problem-Solving Skill4. Applying Hund’s rule. Given the electron configuration for the ground state of an atom, write

the orbital diagram (Example 8.4).

You will need to remember the essence of Hund’s rule to work problems: One electron goes into each orbital of a subshell before two electrons occupy any of them, and all electrons in the singly occupied orbitals have the same spin.

Exercise 8.5Write an orbital diagram for the ground state of the phosphorus atom (Z = 15). Write all orbitals.

Known: There are 15 electrons to place in the filling order; Hund’s rule; Pauli exclusion principle (no more than 2 electrons per orbital; spins must be opposite).

Solution: First write the electron configuration:

1s22s22p63s23p3

Then write the orbital diagram:

1s 2s 2p 3s 3p

8.5 Mendeleev’s Predictions from the Periodic Table

Learning Objective Describe how Mendeleev predicted the properties of undiscovered elements.



8.6 Some Periodic Properties

Learning Objectives State the periodic law.

State the general periodic trends in size of atomic radii.

Define effective nuclear charge.

Determine relative atomic sizes from periodic trends. (Example 8.5)

State the general periodic trends in ionization energy.

Define first ionization energy.

Determine relative ionization energies from periodic trends. (Example 8.6)

Define electron affinity.

State the broad general trend in electron affinity across any period.

Problem-Solving Skill5. Applying periodic trends. Using the known trends and referring to a periodic table, arrange a

series of elements in order by atomic radius (Example 8.5) or ionization energy (Example 8.6).

Exercise 8.6Using a periodic table, arrange the following in order of increasing atomic radius: Na, Be, Mg.

Known: Size increases down a group (column), decreases across a period; we can get the order from the periodic table.

Solution: Be < Mg < Na.

Exercise 8.7The first ionization energy of the chlorine atom is 1251 kJ/mol. Without looking at Figure 8.18, state which of the following values would be the more likely ionization energy for the iodine atom. Explain.

a. 1000 kJ/mol

b. 1400 kJ/mol

Known: Ionization energies decrease going down a column; the periodic table shows that iodine is below chlorine.

Solution: Iodine should have a lower ionization energy than 1251 kJ/mol, so (a) is the more likely value.

Electron affinity is the energy change when an electron is added to a neutral atom in the gaseous state to form a negative ion. Note that the associated sign indicates whether the energy is lost () or gained (+). The more stable the ion, the more energy is released on its formation and the higher is the electron affinity. The less stable the ion, the more positive is the electron affinity. Across a period, electron affinity generally increases. This trend is the same as the trend for ionization energy. (Recall, however, that ionization energy is the energy needed to remove an electron from a neutral atom in the gaseous state.)



Exercise 8.8Without looking at text Table 8.4 but using the general comments in this section, decide which has the larger negative electron affinity, C or F.

Known: Definition of electron affinity; the more stable the anion formed by the addition of an electron, the greater is the electron affinity; anions with half-filled p subshells are quite stable; anions with filled p subshells, and thus noble-gas configurations, are more stable; the valence-electron configuration gives the number and arrangement of electrons in the subshells.

Solution: The valence-electron configuration of C would be 2s22p3. The valence-electron configuration of F would be 2s22p6. Because F has the filled p sublevel and a noble-gas configuration, it will be more stable and thus have the larger electron affinity.

8.7 Periodicity in the Main-Group Elements

Learning Objectives Define basic oxide, acidic oxide, and amphoteric oxide.

State the main group corresponding to an alkali metal, an alkaline-earth metal, a chalcogen, a halogen, and a noble gas.

Describe the change in metallic/nonmetallic character (or reactivities) in going through any main group of elements.

ADDITIONAL PROBLEMS1. Briefly explain how the quantum-mechanical model of the atom helps us to understand the

groupings of the elements in the periodic table.

2. Write the ground-state electron configuration for each of the following species.

a. Br

b. Al

c. Ca2+

d. N3

e. V

f. S2

3. Arrange the following atomic orbitals in order of increasing energy: 5p, 4p, 3p, 4s, and 3d. Explain your ordering.

4. Decide whether each of the following orbital diagrams or electron configurations is possible according to the quantum-mechanical view of the atom, the Pauli exclusion principle, and Hund’s rule. Explain your decisions.

a.

1s 2s 2p

b.

1s 2s 2p



c.

1s 2s 2p

d. 1s22s1

e. 1s22s22p62d103s2

f. 1s22s22p63s23p63d84s2

5. Write an orbital diagram for each of the following electron configurations. Indicate whether each atom is diamagnetic or paramagnetic, and explain why.

a. 1s22s22p63s23p5

b. 1s22s22p3

c. [Ne] 3s2

6. Write the valence-electron configuration for atoms of the element in each of the following positions in the periodic table.

a. Period 4, Group VIIA

b. Period 6, Group VA

c. Period 3, Group IA

d. Period 2, Group VIA

7. Indicate the expected trends (increase or decrease) in atomic radii, cation radii, ionization energy, and electron affinity for the elements as we move from left to right across the periodic table and then as we move down the periodic table.

8. The observed trends in atomic and ionic radii for the halogens and halide ions are as follows, in picometers:

F Cl Br I F Cl Br I

64 99 114 133 119 167 182 206

Discuss these data in terms of the valence-shell electron configuration, nuclear charge, and ionic charge of these species.

9. Arrange the elements in each of the following groups in order of decreasing ionization energy. Use only the periodic table.

a. Zn, K, Br

b. Ca, Be, Ba

c. Sr, Cs, Rb

d. Al, Ca, Cs

10. Show, in a drawing, your idea of what the first four periods of the periodic table would look like if each orbital could have three electrons.



ANSWERS TO ADDITIONAL PROBLEMSIf you missed an answer, study the text section and problem-solving skill (PS Sk.) given in parentheses after the answer.

1. The valence-shell electron configurations of elements appearing in the same column of the periodic table are the same, except that the principal quantum number, n, increases by one as you go down a given column. (8.5)

2.

a. Br: [Ar] 4s23d104p5

b. Al: [Ne] 3s23p1

c. Ca2+: [Ar]

d. N3: [Ne]

e. V: [Ar] 4s23d3

f. S2: [Ar] (8.1, PS Sk. 2)

3. The relative energy ordering of these atomic orbitals is as follows:

3p < 4s < 3d < 4p < 5p

From the building-up principle, we know that the lower the value of the principal quantum number n, the lower the energy of the orbital. For orbitals with the same principal quantum number, the energies of orbitals increase with the l quantum number. The 3d energy is quite close to that of the 4s orbital, and the energy of the atom is lowest when the 4s orbital is filled before the 3d orbitals. (8.2)

4.

a. Not possible; three electrons cannot occupy an orbital (1s), and single electrons in orbitals in the same subshell must have like spins (2p).

b. Not possible; two electrons cannot occupy one orbital when other orbitals in the same subshell are empty (2p).

c. Possible; this atom is in an excited state, because the 2s orbital contains only one electron.

d. Possible; the s subshell can contain two electrons.

e. Not possible; there is no 2d subshell.

f. Possible. No violations. (8.1, PS Sk. 1)

5.

a.

1s 2s 2p 3s 3pThe spin of the one unpaired electron makes the atom paramagnetic.

b.

1s 2s 2pThree unpaired electrons (with like spins) make the atom paramagnetic.



c.

1s 2s 2p 3sThe atom is diamagnetic, because all electron spins are paired. (8.4, PS Sk. 4)

6.

a. 4s24p5

b. 6s26p3

c. 3s1

d. 2s22p4 (8.3, PS Sk. 3)

7. The trends in the observed properties of the elements are as follows:

Trend from Left Trend from Top to Right in to Bottom in

Observed Property Periodic Table Periodic TableAtomic radii decrease increaseCation radii decrease increaseIonization energy increase decreaseElectron affinity increase decrease (8.6)

8. Going down the group from fluorine to iodine, the atoms, as well as their ions, increase in size. This is due to the increased value of the principal quantum number, n, causing the average distance of the valence electrons from the nucleus to increase.

The valence-shell electron configuration of each of the elements is ns2np5, and that for each of their ions is ns2np6 (the configuration of the nearest noble gas in the periodic table).

The nuclear charge of F and that of F are the same, and similarly for Cl and Cl, Br and Br, and I and I. The negative ion in each case has a larger radius than the neutral atom because the electrons (one more than in the neutral atom) of the ion are held less tightly by the nucleus than those of the neutral atom. (8.6)

9.

a. Br > Zn > K

b. Be > Ca > Ba

c. Sr > Rb > Cs

d. Al > Ca > Cs (8.6, PS Sk. 5)



10.

(8.2)

CHAPTER POST-TEST1. For each of the following pairs, indicate which atom or ion should have the larger radius.

a. Au3+ or Au+

b. Ar or Xe

c. Br or Br

d. F or I

e. O or B

2. _________________________ electrons are found in the outermost energy level.

3. Which of the following statements about periodic properties of the elements is (are) incorrect?

a. The ionization energies of the elements in a given period generally increase from left to right.

b. The electron affinities of the elements in a given period generally increase from left to right.

c. Chemical properties of the elements are periodic functions of the atomic number.

d. Atomic radii increase across a period and down a group.

e. The ionization energies of the elements generally increase in going down a given group in the periodic table.

4. The element with atomic number 53 has how many electrons in its valence shell?

a. 7

b. 53

c. 8

d. 2

e. 126



5. Write the abbreviated electron configuration for each of the following in the ground state.

a. potassium, Z = 19

b. titanium, Z = 22

c. aluminum, Z = 13

d. Ag, Z = 47

6. The d subshell can accommodate, for any given principal quantum number, the following number of electrons:

a. 6.

b. 14.

c. 2.

d. 10.

e. none of the above.

7. Which of the following orbital diagrams is (are) incorrect for the respective ground-state electron configuration?

a. 4s23d5 =

b. 4s23d104p4 =

c. 3s23p3 =

d. 5s1 =

e. 4f6 =

8. Which of the following is (are) diamagnetic?

a. Sn2+

b. F

c. H

d. Si

e. K+

9. Arrange the following sets of atoms in order of increasing first ionization potential.

a. K Cs Rb Li

b. Ga As Br Ca

c. O N Se Te

d. B He Li Ne

10. Given the general trend in electron affinities across any period in the periodic table, explain why the electron affinities of C and N are 122 and 0 kJ/mol, respectively.

11. Explain in simple terms how (or why) the Pauli exclusion principle works.



ANSWERS TO CHAPTER POST-TESTIf you missed an answer, study the text section and problem-solving skill (PS Sk.) given in parentheses after the answer.

1.

a. Au+

b. Xe

c. Br

d. I

e. B (8.6, PS Sk. 5)

2. Valence (8.2)

3. d and e (8.6)

4. a (8.3, PS Sk. 3)

5.

a. [Ar] 4s1

b. [Ar] 3d24s2

c. [Ne] 3s23p1

d. [Kr] 4d105s1 (an exception) (8.3, PS Sk. 2)

6. d (8.1)

7. b (8.1, 8.4, PS Sk. 1, 4)

8. a and e (8.4)

9.

a. Cs < Rb < K < Li

b. Ga < Ca < As < Br

c. Te < Se < O < N

d. Li < B < Ne < He (8.6, PS Sk. 5)

10. Electron affinity generally increases across a given period. Nitrogen is an exception, as are all Group VA elements. They have the valence electron configuration ns2np3 with the half-filled p subshell. Adding an additional electron sacrifices the stability attributed to the half-filled subshell and produces an unstable negative ion. (8.6)

11. Electrons are negatively charged particles that repel each other. When degenerate orbitals exist (separate orbitals with the same energy), it is energetically favorable for the electrons to fill the orbitals singly until all available degenerate orbitals are used. At that point, it becomes energetically favorable for electrons to pair rather than to move to a new subshell or to a new quantum level. (8.1)


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