Chapter 8The Periodic Table and

Atomic Structure

Gang Chen, Department of Chemistry, UCF

Lecture 3, 4&5 Review

2

Chapter Objectives• Understand periodic trends in atomic radii, Ionization energy,

and Electron Affinity.

• Predict relative ionic sizes within an isoelectronic series.

• Correlate ionization energies with the chemistry of the elements.

• Master the concept of electron affinity as a measure of the energy required to add an electron to an atom or ion.

• Recognize the inverse relationship of ionization energies and electron affinities.

3

Periodic Trends in Atomic Properties

• Using the understanding of orbitals and atomic structure, it is possible to explain some periodic properties.

• Atomic size(related to size of valence orbitals)

• Ionization energy(affected by attraction between nucleus and valence electron)

• Electron affinity(ability to attract valence electrons)

4

Atomic Size

Covalent atomic radius vs. vdW atomic radius

Nonbonding radius Bonding radius

5

Atomic Size

6

Atomic Size

• The shell in which the valence electrons are found affects atomic size.• The size of the valence orbitals increases with n, so size

must increase from top to bottom for a group.

• The strength of the interaction between the nucleus and the valence electrons affects atomic size.

• The effective nuclear charge increases from left to right across a period, so the interaction between the electrons and the nucleus increases in strength.

• As interaction strength increases, valence electrons are drawn closer to the nucleus, decreasing atomic size.

7

Atomic Size

Zeffective = Z − S

e-

valencecore

Main group Transition

8

Example ProblemQ. Arrange these elements based on their increasing atomic radii .

Br, Ca, Ge, F

F, Br, Group 7A (same group)F < Br

Ca, Ge, Br, Row 4 (same period)Br < Ge < Ca

Down a group, atomic size increases.Across a row (L to R), atomic size decreases.

F < Br < Ge < Ca

9

Ionization Energy• First Ionization energy - the minimal energy required to remove

an electron from a gaseous atom, forming a cation (1+).• Formation of X2+ from X+ would be the second ionization

energy, etc.

• The more strongly held an electron is, the higher the ionization energy must be.

• As valence electrons move further from the nucleus, they become easier to remove and the first ionization energy becomes smaller.

(endothermic)

Mg(g) + 738 kJ/mol → Mg+(g) + e- 1st IE = +738 kJ/mol

Mg+(g) + 1451 kJ/mol → Mg2+

(g) + e- 2nd IE = +1451 kJ/mol

10

Ionization Energy

• Graph of the first ionization energy (in kJ/mol) vs. atomic number for the first 38 elements.

11

Example ProblemQ. Arrange these elements based on their increasing first ionization energies.

Se, Ge, K, S

Se, S, Group 6A (same group)Se < S

Se, Ge, K, Row 4 (same period)K < Ge < Se

Down a group, I.E. decreases.Across a row (L to R), I.E. increases.

K < Ge < Se < S

I.E. Increase

I.E. D

ecre

ase

12

Electron Affinity• Electron affinity - energy change to place an electron on a

gaseous atom, forming an anion.

• Electron affinities may have positive or negative values.• Negative values - energy released • Positive values - energy absorbed

13

Electron Affinity• Electron affinities increase (numerical value becomes more

negative) from left to right for a period and bottom to top for a group.

• An element with a high ionization energy generally has a high affinity for an electron, i.e., EA is largely negative. That is the case for halogens (F, Cl, Br, I), O, and S.

• The greater (more negative) the electron affinity, the more stable the anion will be.

Mg(g) + e- + 231 kJ/mol → Mg-(g) EA = +231 kJ/mol

Br(g) + e- → Br-(g) + 323 kJ/mol EA = -323 kJ/mol

14

Example ProblemQ. Arrange these elements based on their increasing electron affinity (more negative).

Se, Ge, K, S

Se, S, Group 6A (same group)Se < S

Se, Ge, K, Row 4 (same period)K < Ge < Se

Down a group, E.A. decreases (less negative).Across a row (L to R), E.A. increases (more negative).

K < Ge < Se < S

E.A. Increase(more negative)

E.A.

Dec

reas

e

15

Metallic Character• Metallic character is how closely an element’s properties

match the ideal properties of a metal.• Metallic character decreases left to right across a period.• Metallic character increases down the column.

16

Example ProblemQ. Arrange these elements based on their increasing metallic character.

N, Rb, Sb

N, Sb, Group 5A (same group)N < Sb

Rb, Sb, Row 5 (same period)Sb < Rb

Down a group, metallic character increases.Across a row (L to R), metallic character decreases.

N < Sb < Rb

17

Electron Configurations of Cation• Cations form when the atom loses electrons from the valence

shell.• Metals in the main groups form cations with an np6 electronic

configuration.

12Mg 1s22s22p63s2 → 12Mg2+ 1s22s22p6 + 2e-

19K 1s22s22p63s23p64s1 → 19K+ 1s22s22p63s23p6 + e-

13Al 1s22s22p63s23p1 → 13Al3+ 1s22s22p6 + 3e-

Group 1A

Group 2A

Group 3A

[Ar]

[Ne]

[Ne]

13Al3+ 1s22s22p6 + 3e- → 13Al 1s22s22p63s23p1

18

Electron Configurations of Cation• Transition metals can form cations with more than one

possible charge.

• Transition metals first lose electrons from the s subshell.

• Additional electrons are lost from the partially filled dorbitals.

• A half filled d orbital set is a fairly stable arrangement.

• Fe2+ and Fe3+ ions are both stable.

26Fe [Ar] 4s23d6 → 26Fe2+ [Ar]3d6 + 2e-

26Fe2+ [Ar]3d6 → 26Fe3+ [Ar]3d5 + e-

19

Electron Configurations of Anion• Nonmetals have negative electron affinities and generally

form anions with an np6 electronic configuration.

• Anions are larger than their corresponding neutral atoms.• Gaining electrons increases electron-electron repulsion.• Valence electrons are less tightly bound to the nucleus.

16S 1s22s23s23p4 + 2e- → 16S2- 1s22s23s23p6

9F 1s22s22p5 + e- → 9F- 1s22s22p6

15P 1s22s23s23p3 + 3e- → 15P 3- 1s22s23s23p6

Group 7A

Group 6A

Group 5A

[Ne]

[Ar]

[Ar]

20

Magnetic Properties• Electron configurations that result in unpaired electrons

mean that the atom or ion will have a net magnetic field; it is paramagnetic.

• Electron configurations that result in all paired electrons mean that the atom or ion will have no magnetic field; this is called diamagnetic.

6C ↑↓ ↑↓ ↑ ↑ 1s2 2s2 2p2

4Be ↑↓ ↑↓ 1s2 2s2

Total spin = (+½ , +½) = 1

Total spin = (+½ , -½) = 0

21

Ionic Radius

• Sizes of ions compared to corresponding neutral atoms.1 pm = 10-12 m

22

Ionic Radius• Cations (positive) are always smaller than their respective

neutral atoms. Cations decrease in size from left to right across the periodic table.• Cations in a period have the same electronic configuration.

• Anions (negative) are always larger than their neutral atoms. Anions decrease in size from left to right across the periodic table.• Anions in a period have the same electronic configuration.

• Number of protons increases from left to right across a period.

• Electrons are held more tightly from left to right across a period, resulting in smaller ions.

Br− > Br

Na > Na+

23

Ionic Radius

Isoelectronic species have the same number of electrons. Examples show number of (protons, Z) and charge (+), (−) or (0).

The more positively charged the isoelectronic species, the smaller the radius.

10e− species (1s2 2s2 2p6 = [Ne]):7N3− > 8O2− > 9F− > 10Ne > 11Na+ > 12Mg2+ > 13Al3+

18e− species (1s2 2s2 2p6 3s2 3p6 = [Ar]):16S2− > 17Cl− > 18Ar > 19K+ > 20Ca2+ > 21Sc3+

Chapter 8The Periodic Table and

Atomic Structure

Gang Chen, Department of Chemistry, UCF

Lecture 1&2 Review

25

Quantum numbers

nlmlms

Principle quantum number - shellAngular momentum quantum number - subshell

Magnetic quantum number – subshell directionSpin quantum number – spin direction

n = 1, 2, 3, ….l = 0, 1, 2, n-1ml = -l, …0, … lms = ± 1/2

• An electron is specified by a set of four quantum numbers.

26

Aufbau Principle: order of electron filling1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

Electron Configurations

s subshell: 1 orbital, 2 electronsp subshell: 3 orbitals, 6 electronsd subshell: 5 orbitals, 10 electronsf subshell: 7 orbitals, 14 electrons

Pauli Exclusion Principle:Each orbital can hold 2 electrons with opposite spins

Orbitals that are in the same subshell have the same energy.

27

Electron Configurations

Quantum number of electrons in 3s subshell:

[ ] [ ] 212 3s Ne Ne Mg ↑↓ (1s22s22p63s2)

(3, 0, 0, +1/2), (3, 0, 0, -1/2)

n = 1, 2, 3, ….l = 0, 1, 2, n-1ml = -l, …0, … lms = ± 1/2

28

Valence Electrons and Core Electrons

Valence electrons in shell with highest n, i.e., the outermost electrons, those beyond the core electrons

11Na: 1s2 2s2 2p6 3s1

4Be: 1s2 2s2

14Si: 1s2 2s2 2p6 3s2 3p2

36Kr: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

27Co: 1s2 2s2 2p6 3s2 3p6 4s2 3d7

They determine the chemical properties of an element. For the representative elements, they are the ns and np electrons; for transition elements they are the ns and (n−1)d electrons.

[Ne]3s1

[He]2s2

[Ne]3s23p2

[Kr]4s23d7