Chapter 9: Periodic Properties of the Elements – Structured Study Notes

Periodic Properties of the Elements

Filling Orbitals with Electrons

Electron arrangement in atoms is governed by three fundamental rules that determine how electrons occupy atomic orbitals. These rules ensure the most stable and unique configuration for each atom.

• Aufbau Principle: Electrons fill the lowest energy (most stable) orbitals first before occupying higher energy orbitals.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers; electrons in the same orbital must have opposite spins.

• Hund's Rule: For degenerate orbitals (orbitals of equal energy), electrons fill each orbital singly with parallel spins before pairing up.

Example: Nitrogen's electron configuration is $1s^2 2s^2 2p^3$.

Electron Configurations

Each element has a unique number of electrons and a ground-state electron configuration, representing the lowest energy arrangement of electrons.

• Nitrogen (7 electrons): $1s^2 2s^2 2p^3$

• Fluorine (9 electrons): $1s^2 2s^2 2p^5$

• Phosphorus (15 electrons): $1s^2 2s^2 2p^6 3s^2 3p^3$

Electron configurations can be represented using orbital diagrams and shorthand notation (using noble gas cores).

Table: Electron Configurations of Several Lighter Elements

Element	Total Electrons	Orbital Diagram	Electron Configuration
Li	3	1s: ↑↓, 2s: ↑	1s22s1
Be	4	1s: ↑↓, 2s: ↑↓	1s22s2
B	5	1s: ↑↓, 2s: ↑↓, 2p: ↑	1s22s22p1
C	6	1s: ↑↓, 2s: ↑↓, 2p: ↑ ↑	1s22s22p2
N	7	1s: ↑↓, 2s: ↑↓, 2p: ↑ ↑ ↑	1s22s22p3
Ne	10	1s: ↑↓, 2s: ↑↓, 2p: ↑↓ ↑↓ ↑↓	1s22s22p6
Na	11	1s: ↑↓, 2s: ↑↓, 2p: ↑↓ ↑↓ ↑↓, 3s: ↑	1s22s22p63s1

Example: Vanadium Electron Configuration

Vanadium (23 electrons) has the configuration:

• $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^3$

• Shorthand: $[Ar] 4s^2 3d^3$

Degenerate Orbitals

Orbitals with the same energy are called degenerate. Near-degenerate orbitals can cause exceptions to the expected filling order, especially in transition metals.

• Chromium: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^5$

• Copper: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^{10}$

More About Electron Configurations

• Cations: Formed by removing electrons (usually from the outer shell).

• Anions: Formed by adding electrons to the highest energy level.

• Isoelectronic Species: Different atoms/ions with the same number of electrons.

• Excited States: Same number of electrons, but electrons occupy higher energy orbitals.

The Periodic Table

The periodic table organizes elements by increasing atomic number and recurring chemical properties. Electron configurations correspond to the table's structure, with blocks for s, p, d, and f orbitals.

Effective Nuclear Charge

Electrons in an atom experience an effective nuclear charge (Zeff) due to shielding by core electrons. Valence electrons are less attracted to the nucleus because of this shielding.

• Core electrons: Close to the nucleus, complete shells.

• Valence electrons: Farthest from the nucleus, involved in bonding.

• Shielding: Core electrons reduce the nuclear attraction felt by valence electrons.

Graph: Probability distributions show that 1s electrons are closest to the nucleus and provide the most shielding.

Electron Shells and Atomic Size

Atoms do not have a sharply defined boundary, but atomic size can be estimated by measuring how close atoms approach each other during collisions or bonding.

• Nonbonding atomic radius: Half the distance between nuclei in a collision.

• Bonding atomic radius: Half the distance between nuclei in a bonded pair.

Bonding Atomic Radii

Bonding atomic radii are typically smaller than nonbonding radii due to the attractive forces in chemical bonds.

• Calculated as half the distance between nuclei of identical bonded atoms.

• Other radii can be inferred by difference.

Average Bonding Atomic Radii

Bonding radii generally increase down a group and decrease across a period.

Cation/Anion Sizes

The size of ions differs from their parent atoms due to changes in electron number and repulsion.

• Cations: Smaller than parent atoms (loss of electrons reduces repulsion).

• Anions: Larger than parent atoms (gain of electrons increases repulsion).

Ionization Energy

Ionization energy (IE) is the energy required to remove an electron from a gaseous atom or ion. Successive ionization energies increase as more electrons are removed.

• $M(g) \rightarrow M^+(g) + e^-$ — First IE

• $M^+(g) \rightarrow M^{2+}(g) + e^-$ — Second IE

• $M^{2+}(g) \rightarrow M^{3+}(g) + e^-$ — Third IE

Table: Successive Ionization Energies (kJ/mol)

Element	IE1	IE2	IE3	IE4	IE5	IE6	IE7
Na	496	4560	6910	9540	13350	16600	20800
Mg	738	1450	7730	10500	13600	18000	21700
Al	578	1820	2750	11600	14800	18300	23200
Si	786	1570	3230	4360	16000	19800	24300
P	1012	1900	2910	4960	6270	21200	25400
S	1000	2250	3400	4550	7000	8490	27100
Cl	1251	2290	3820	5150	6540	10400	12900
Ar	1521	2665	3930	5770	7840	8700	12000

First Ionization Energy Trends

• Increases across a period (left to right): Due to increasing effective nuclear charge and decreasing atomic size.

• Decreases down a group: Due to increasing atomic size and electron shielding.

Electron Affinity

Electron affinity (EA) is the energy change when an atom gains an electron. A more negative EA means the atom more readily accepts an electron.

• $A(g) + e^- \rightarrow A^-(g)$ — First EA

• $A^-(g) + e^- \rightarrow A^{2-}(g)$ — Second EA

Table: Electron Affinities (kJ/mol)

Element	EA (kJ/mol)
H	-73
Li	-60
Be	>0
B	-27
C	-122
N	>0
O	-141
F	-328
Ne	>0
Na	-53
Mg	>0
Al	-44
Si	-134
P	-72
S	-200
Cl	-349
Ar	>0

Metals, Nonmetals, and Metalloids

Elements are classified based on their physical and chemical properties:

• Metals: Good conductors, malleable, ductile, tend to lose electrons.

• Nonmetals: Poor conductors, brittle, tend to gain electrons.

• Metalloids: Intermediate properties, semiconductors.

Metallic character increases down a group and decreases across a period.

Summary of Periodic Trends

Property	Trend Down a Column	Reason Down	Trend Across a Row	Reason Across
Atomic Radii	Increasing	Size of outermost occupied orbital increases	Decreasing	Effective nuclear charge increases
First Ionization Energy	Decreasing	Outermost electrons farther from nucleus	Increasing	Effective nuclear charge increases
Electron Affinity	No definite trend	Ionization energy decreases	Decreasing (more negative)	Effective nuclear charge increases
Metallic Character	Increasing	Ionization energy decreases	Decreasing	Ionization energy increases

Group Trends

Alkali Metals

Element	Electron Configuration	Melting Point (°C)	Density (g/cm³)	Atomic Radius (Å)	IE1 (kJ/mol)
Lithium	[He]2s1	181	0.53	1.28	520
Sodium	[Ne]3s1	98	0.97	1.66	496
Potassium	[Ar]4s1	63	0.86	2.03	419
Rubidium	[Kr]5s1	39	1.53	2.20	403
Cesium	[Xe]6s1	28	1.88	2.44	376

Alkaline Earth Metals

Element	Electron Configuration	Melting Point (°C)	Density (g/cm³)	Atomic Radius (Å)	IE1 (kJ/mol)
Beryllium	[He]2s2	1287	1.85	0.96	899
Magnesium	[Ne]3s2	650	1.74	1.41	738
Calcium	[Ar]4s2	842	1.55	1.76	590
Strontium	[Kr]5s2	777	2.63	1.95	549
Barium	[Xe]6s2	727	3.51	2.15	503

Group 6A Elements

Element	Electron Configuration	Melting Point (°C)	Density (g/cm³)	Atomic Radius (Å)	IE1 (kJ/mol)
Oxygen	[He]2s22p4	-218	1.43	0.66	1314
Sulfur	[Ne]3s23p4	115	1.96	1.05	1000
Selenium	[Ar]3d104s24p4	221	4.82	1.20	941
Tellurium	[Kr]4d105s25p4	450	6.24	1.38	869
Polonium	[Xe]4f145d106s26p4	254	9.20	1.40	812

Halogens

Element	Electron Configuration	Melting Point (°C)	Density (g/cm³)	Atomic Radius (Å)	IE1 (kJ/mol)
Fluorine	[He]2s22p5	-220	1.69	0.57	1681
Chlorine	[Ne]3s23p5	-102	3.12	1.02	1251
Bromine	[Ar]3d104s24p5	-7.3	3.12	1.20	1140
Iodine	[Kr]4d105s25p5	114	4.94	1.39	1008

Noble Gases

Element	Electron Configuration	Boiling Point (K)	Density (g/L)	Atomic Radius (Å)	IE1 (kJ/mol)
Helium	1s2	4.2	0.18	0.28	2372
Neon	[He]2s22p6	27.1	0.90	0.58	2081
Argon	[Ne]3s23p6	87.3	1.78	1.06	1521
Krypton	[Ar]3d104s24p6	120	3.75	1.16	1351
Xenon	[Kr]4d105s25p6	165	5.46	1.40	1170
Radon	[Xe]4f145d106s26p6	211	9.73	1.50	1037

Additional info: These notes expand on the original slides by providing definitions, explanations, and context for each concept, as well as recreating and summarizing all tables for clarity and completeness.