Electron Configurations and Periodic Trends: Study Notes

Chapter 11: Electron Configurations

Introduction to Electron Configurations

Electron configurations describe the arrangement of electrons in atoms, which is fundamental to understanding chemical properties and periodic trends. This topic is essential for chemistry and related life sciences, including Anatomy & Physiology, as it underpins atomic interactions and molecular structure.

Bohr Model of the Atom

Energy Levels and Subshells

The Bohr model proposes that electrons orbit the nucleus in fixed energy levels or shells. Each shell contains subshells, which are further divided into orbitals.

• Energy Level (n): The principal quantum number indicating the shell's distance from the nucleus.

• Subshells: Subdivisions within energy levels, labeled as s, p, d, and f.

• Orbitals: Regions within subshells where electrons are likely to be found.

• Example: The second energy level (n=2) contains s and p subshells.

Energy Levels and Subshells

Subshells and Their Order

Each energy level contains a number of subshells equal to its principal quantum number (n). The order of subshells is s, p, d, f.

• First energy level (n=1): 1 subshell (1s)

• Second energy level (n=2): 2 subshells (2s, 2p)

• Third energy level (n=3): 3 subshells (3s, 3p, 3d)

• Fourth energy level (n=4): 4 subshells (4s, 4p, 4d, 4f)

Maximum Electrons per Subshell

The number of electrons each subshell can hold is determined by the type of subshell:

• s subshell: 2 electrons

• p subshell: 6 electrons

• d subshell: 10 electrons

• f subshell: 14 electrons

Electrons per Energy Level

The total number of electrons in each energy level is the sum of electrons in all its subshells.

Energy Level (n)	Subshells	Max Electrons
1	1s	2
2	2s, 2p	8
3	3s, 3p, 3d	18
4	4s, 4p, 4d, 4f	32

Electron Configurations

Writing Electron Configurations

Electron configurations are written by listing subshells in order of increasing energy and indicating the number of electrons in each subshell as a superscript.

• Example: Carbon: 1s2 2s2 2p2

• Order of filling: Subshells are filled in the order of increasing energy, following the Aufbau principle.

Filling Diagram for Sublevels: The order of subshell filling is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

Blocks of Elements

The periodic table is divided into blocks based on the type of subshell being filled:

• s-block: Groups 1 and 2, plus Helium

• p-block: Groups 13-18

• d-block: Transition metals (Groups 3-12)

• f-block: Lanthanides and actinides

Valence Electrons

Definition and Identification

Valence electrons are electrons in the outermost energy level (shell) of an atom. They are crucial in determining chemical reactivity and bonding.

• Group number: Indicates the number of valence electrons for main group elements.

• Example: Oxygen (Group 16) has 6 valence electrons.

Predicting Valence Electrons

Use the periodic table to predict the number of valence electrons:

• Group 1: 1 valence electron

• Group 2: 2 valence electrons

• Groups 13-18: Number of valence electrons equals the group number minus 10

Predicting Ionic Charges

Formation of Ions

Atoms gain or lose electrons to achieve a stable electron configuration, often resembling the nearest noble gas.

• Metals: Tend to lose electrons, forming positive ions (cations)

• Nonmetals: Tend to gain electrons, forming negative ions (anions)

• Example: Sodium (Na) loses one electron to form Na+; Chlorine (Cl) gains one electron to form Cl-

Predicting Ionic Charges by Group

Group	Typical Ionic Charge
Group 1 (IA)	+1
Group 2 (IIA)	+2
Group 13 (IIIA)	+3
Group 15 (VA)	-3
Group 16 (VIA)	-2
Group 17 (VIIA)	-1
Group 18 (VIIIA)	0

Periodic Trends

Atomic Radius

The atomic radius is the distance from the nucleus to the outermost electron shell. It varies predictably across the periodic table.

• Trend across a period: Atomic radius decreases from left to right due to increased nuclear charge pulling electrons closer.

• Trend down a group: Atomic radius increases due to addition of energy levels.

• Example: Sodium (Na) is larger than Chlorine (Cl) in the same period.

Summary Table: Atomic Radius Trend

Direction	Trend
Across a period (left to right)	Decreases
Down a group	Increases

Additional info:

• Electron configurations and periodic trends are foundational for understanding chemical interactions in biological systems, such as ion transport and molecular bonding in Anatomy & Physiology.

• Knowledge of valence electrons and ionic charges is essential for predicting how atoms interact in physiological processes.