Electron Configurations and Periodic Trends: Study Notes

Electron Configurations

Introduction to Electron Configurations

Electron configurations describe the arrangement of electrons in an atom's energy levels, sublevels, and orbitals. Understanding electron configurations is essential for predicting chemical properties and reactivity.

• Electron configuration: The distribution of electrons among the atom's energy levels and sublevels.

• Energy levels: Principal shells around the nucleus, denoted by n = 1, 2, 3, etc.

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

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

Bohr Model of the Atom

The Bohr model proposed that electrons orbit the nucleus in fixed energy levels. Each energy level corresponds to a specific amount of energy.

• Energy levels are quantized; electrons can only occupy certain allowed levels.

• Electrons in higher energy levels are farther from the nucleus.

Example: The hydrogen atom has electrons in the n = 1 energy level.

Energy Levels and Sublevels

Each energy level contains one or more sublevels, which are designated as s, p, d, and f. The number of sublevels in an energy level equals the principal quantum number (n).

• n = 1: 1 sublevel (1s)

• n = 2: 2 sublevels (2s, 2p)

• n = 3: 3 sublevels (3s, 3p, 3d)

• n = 4: 4 sublevels (4s, 4p, 4d, 4f)

Additional info: Sublevels are further divided into orbitals.

Sublevel Electron Capacity

Each sublevel can hold a specific maximum number of electrons:

• s sublevel: 2 electrons

• p sublevel: 6 electrons

• d sublevel: 10 electrons

• f sublevel: 14 electrons

Formula: $\text{Maximum electrons per energy level} = 2n^2$

Electrons per Energy Level

The following table summarizes the distribution of electrons in each energy level:

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

Electron Configuration Notation

Electron configurations are written using a shorthand notation that lists the energy level, sublevel, and number of electrons in each sublevel.

• Format: nsublevelnumber of electrons

• Example: $1s^2 2s^2 2p^6$

Filling Order of Sublevels

Electrons fill sublevels in order of increasing energy, following the Aufbau principle. The order is determined by the diagram below:

• Order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

Additional info: The filling order can be memorized using the diagonal rule diagram.

Blocks of Elements in the Periodic Table

The periodic table is divided into blocks based on the type of sublevel 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

Additional info: Blocks are useful for predicting electron configurations.

Valence Electrons and Ionic Charges

Valence Electrons

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

• Valence electrons are found in the highest principal energy level.

• Only valence electrons are involved in chemical reactions.

Example: Sodium (Na) has 1 valence electron in the 3s sublevel.

Predicting Valence Electrons

The number of valence electrons can be predicted from the group number in the periodic table:

• Group 1: 1 valence electron

• Group 2: 2 valence electrons

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

Predicting Ionic Charges

Atoms gain or lose electrons to achieve a stable electron configuration, often resembling the nearest noble gas. The resulting charge is called the ionic charge.

• Metals tend to lose electrons, forming positive ions (cations).

• Nonmetals tend to gain electrons, forming negative ions (anions).

• Group 1: +1 charge, Group 2: +2 charge, Group 17: -1 charge, Group 16: -2 charge

Group	Valence Electrons	Common Ionic Charge
1 (IA)	1	+1
2 (IIA)	2	+2
16 (VIA)	6	-2
17 (VIIA)	7	-1
18 (VIIIA)	8	0

Electron Configurations of Ions

When atoms form ions, their electron configurations change to reflect the loss or gain of electrons.

• For cations (positive ions), remove electrons from the highest energy sublevel first.

• For anions (negative ions), add electrons to the next available sublevel.

Example: $\text{Na}: 1s^2 2s^2 2p^6 3s^1 \rightarrow \text{Na}^+: 1s^2 2s^2 2p^6$

Periodic Trends

Atomic Radius Trend

Atomic radius is the distance from the nucleus to the outermost electron shell. The periodic table shows predictable trends in atomic radius.

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

• Atomic radius increases from top to bottom within a group (due to addition of energy levels).

Example: Sodium (Na) has a larger atomic radius than chlorine (Cl) in the same period.

Summary Table: Atomic Radius Trend

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

Applications and Practice

Understanding electron configurations and periodic trends is essential for predicting chemical behavior, bonding, and reactivity. Practice problems often involve writing electron configurations, identifying valence electrons, predicting ionic charges, and comparing atomic radii.

• Write electron configurations for given elements and ions.

• Use the periodic table to determine valence electrons and ionic charges.

• Compare atomic radii based on position in the periodic table.

Additional info: These concepts are foundational for further study in chemistry and biochemistry.