Overview of Electron Configuration

Definition of Electron Configuration: The distribution of electrons in an atom's orbitals. It dictates an atom's chemical behavior, reactivity, and position on the periodic table. It provides insight into an atom's stability, magnetic properties, and the types of bonds it can form.

Aufbau Principle

Order of Filling: Electrons occupy the lowest energy orbitals first before moving to higher energy levels. This principle is based on the concept that electrons will fill the lowest available energy states before occupying higher ones, leading to the most stable electron configuration (ground state).
- Orbital Diagram: Orbital diagrams can use various representations ($1s, 2s$, etc.) to show electron configuration, illustrating the order of filling. The order can be visualized using an energy diagram or the ( $n+l$ ) rule, where orbitals with lower ( $n+l$ ) values fill first. If two orbitals have the same ( $n+l$ ) value, the one with the lower 'n' value fills first.
- The general order is $1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p$ .

Hund's Rule

Description: This rule states that for a set of degenerate orbitals (orbitals of the same energy, such as the three p orbitals or five d orbitals), electrons will first occupy each orbital singly with parallel spins, before pairing electrons in an orbital. This minimizes electron-electron repulsion and leads to a more stable configuration.
- Example: Filling of 2p orbitals - for Nitrogen (atomic number 7), the 2p orbitals are filled as $2p<em>x^1 2p</em>y^1 2p_z^1$ before any orbital hosts two electrons.

Orbital Types and Capacities

Each orbital represents a specific region of space around the nucleus where an electron is most likely to be found. Orbitals are characterized by their shapes and energy levels.
s Orbitals: Spherical shape, only one s orbital per energy level. Can hold a maximum of 2 electrons.
p Orbitals: Dumbbell shape, three p orbitals ( $p<em>x, p</em>y, p_z$ ) per energy level (starting from n=2). Each set can hold a maximum of 6 electrons (2 per orbital for 3 orbitals).
d Orbitals: More complex shapes (cloverleaf), five d orbitals per energy level (starting from n=3). Each set can hold a maximum of 10 electrons (2 per orbital for 5 orbitals).
f Orbitals: Even more complex shapes, seven f orbitals per energy level (starting from n=4). Each set can hold a maximum of 14 electrons (2 per orbital for 7 orbitals).

Writing Electron Configurations

This notation concisely summarizes where all electrons in an atom are located and their probable energy states.
Full Configuration Example: For Aluminum (13 electrons): $1s^2 2s^2 2p^6 3s^2 3p^1$ . This shows the sequential filling of orbitals according to the Aufbau principle.
Condensed Notation (Noble Gas Notation): Uses the symbol of the noble gas that precedes the element in question in square brackets, followed by the configuration of the remaining electrons. This highlights the valence electrons.
- Example: For Aluminum, Neon ( $[Ne]$ ) has 10 electrons ( $1s^2 2s^2 2p^6$ ), so Aluminum's condensed notation is $[Ne] 3s^2 3p^1$ .
Exceptions to Aufbau: Some elements, particularly transition metals like Chromium (Cr) and Copper (Cu), exhibit exceptions to the Aufbau principle due to the added stability of half-filled or fully-filled d-subshells.
- Cr (Z=24): Expected is $[Ar] 4s^2 3d^4$ , but observed is $[Ar] 4s^1 3d^5$ .
- Cu (Z=29): Expected is $[Ar] 4s^2 3d^9$ , but observed is $[Ar] 4s^1 3d^{10}$ .

Periodic Table and Electron Configuration

The periodic table is structured based on electron configuration, particularly the highest energy electrons.
Blocks of the Periodic Table: The periodic table is divided into blocks corresponding to the type of orbital being filled.
- s-block: Groups 1 and 2 (alkali and alkaline earth metals) fill s orbitals.
- p-block: Groups 13-18 fill p orbitals.
- d-block: Groups 3-12 (transition metals) fill d orbitals. The principal quantum number for d orbitals is one less than the period number (e.g., $4s$ then $3d$ ).
- f-block: Lanthanides and Actinides fill f orbitals. The principal quantum number for f orbitals is two less than the period number (e.g., $6s$ then $4f$ ).
Period Concept: Periods denote the highest principal quantum number ( $n$ ) of the valence electrons. For example, elements in the 6th period begin filling the $6s$ orbital.
Transition Metals: Their electron configurations often involve d and f block filling, complicating general rules. For transition metal ions, electrons are typically removed from the $s$ orbital before the $d$ orbital, even if the $s$ orbital filled first. For example, $Fe ([Ar] 4s^2 3d^6)$ forms $Fe^{2+} ([Ar] 3d^6)$ by losing its $4s$ electrons.

Valence Electrons vs Core Electrons

Valence Electrons: These are the electrons in the outermost principal energy level or the highest energy orbitals (often s and p orbitals in the outermost shell, or sometimes d orbitals for transition metals). They are involved in chemical bonding and determine an atom's reactivity and chemical properties. For main group elements, the number of valence electrons is typically the same as the group number (e.g., Group 1 has 1 valence electron, Group 17 has 7).
Core Electrons: These are the inner-shell electrons that are not involved in bonding. They are found in filled principal energy levels beneath the valence shell and are represented by the noble gas configuration in condensed notation.

Trends in Atomic Properties

Atomic Radius

Definition: Distance from the nucleus to the boundary of the surrounding cloud of electrons, providing a measure of how large an atom is.
Effect of Principal Quantum Number (Moving Down a Group): As you go down a group, the atomic radius increases. This is because electrons are added to new principal energy levels (shells), which are progressively further from the nucleus. The increase in shielding by inner electrons also reduces the effective nuclear charge experienced by the outermost electrons, allowing them to spread out more.
- Example: Lithium (Li, $1s^2 2s^1$ ) vs Sodium (Na, $1s^2 2s^2 2p^6 3s^1$ ) - Li's outer electron is a $2s$ while Na's is a $3s$ , meaning Na has an additional electron shell, making it larger.
Effect of Nuclear Charge (Moving Across a Period): As you go across a period from left to right, the atomic radius generally decreases. Although electrons are added, they are added to the same principal energy level (or subshell). Simultaneously, the nuclear charge (number of protons) increases, pulling these valence electrons more strongly towards the nucleus without a significant increase in shielding from inner electrons, resulting in a smaller atomic radius.

Ionization Energy

Definition: The minimum energy required to remove one electron from a gaseous atom or ion in its ground state. The first ionization energy (IE1) refers to removing the first electron, IE2 the second, and so on.
Trends:
- Decreases down a group: As atomic radius increases down a group, the outermost electron is further from the nucleus and experiences less effective nuclear charge due to increased shielding. This weaker attraction means less energy is required to remove it.
- Increases across a period: As atomic radius decreases across a period, the valence electrons are held more tightly due to an increasing effective nuclear charge and a decreasing distance from the nucleus. More energy is thus required to remove an electron.
Successive Ionization Energies: Each subsequent ionization energy is significantly higher than the previous one because removing an electron from a positively charged ion is harder than from a neutral atom. Also, as electrons are removed, the remaining electrons are pulled closer to the nucleus, increasing the effective nuclear charge.
Note: A very large jump in successive ionization energies indicates that a core electron is being removed. For example, for an element in Group 1, the second ionization energy will be dramatically higher than the first because it involves removing an electron from a stable noble gas configuration (a core electron).
Exceptions: Minor exceptions exist due to electron-electron repulsion (e.g., removing a paired electron is easier) or orbital stability (e.g., removing from a half-filled p subshell vs. a slightly more filled one).

Electronegativity

Definition: A measure of the tendency of an atom to attract a bonding pair of electrons in a chemical bond. It is a relative value, with Fluorine being the most electronegative element.
Trends:
- Decreases down a group: As atomic size increases, the distance between the nucleus and the valence electrons in a bond increases, reducing the nucleus's attraction for those electrons.
- Increases across a period: As effective nuclear charge increases and atomic size decreases across a period, the nucleus's ability to attract bonding electrons increases.

Electron Affinity

Definition: The energy change that occurs when an electron is added to a gaseous atom to form an anion. A more negative (exothermic) electron affinity indicates a greater tendency to accept an electron.
Trends:
- Generally becomes more negative (more favorable) across a period due to increasing effective nuclear charge, making it easier for the nucleus to attract an additional electron.
- Becomes less negative (less favorable) down a group due to increasing atomic size and shielding, which decreases the nucleus's attraction for an incoming electron.
Exceptions: Some elements (like noble gases, alkaline earth metals, and Nitrogen) have positive (endothermic) electron affinities because adding an electron disrupts a stable electron configuration, requiring energy input.

Trends Summary

Atomic Radius: Increases down a group, decreases across a period.
Ionization Energy: Decreases down a group, increases across a period.
Electronegativity: Decreases down a group, increases across a period.
Electron Affinity: Generally becomes more negative (favorable) across a period and less negative (less favorable) down a group.

Conclusion

Understanding electron configuration and associated periodic trends (such as atomic radius, ionization energy, electronegativity, and electron affinity) is crucial for grasping concepts in atomic structure, chemical bonding, and predicting the chemical behavior and reactivity of elements.