Atomic Structure and Periodicity

Electronic Structure and the Periodic Table

The lecture aims to enhance understanding of the periodic table and its relation to electronic structure, reducing the need for memorization.
The periodic table will be available during exams for reference.

The lecture builds upon previous discussions about orbitals and electron placement in elements.
The focus is on the outermost orbitals and valence electrons, which determine elemental behavior.
The periodic table is structured into blocks (S, P, D, F) based on the azimuthal quantum number, representing the shape of orbitals.
The principal quantum number (n) corresponds to the period (row) number in the periodic table.
Groups 1 and 2 have valence electrons in the S orbitals.
Transition metals in periods 3-6 have 10 columns due to the five orbitals in the D subshell, each holding two electrons.
The P block has six columns due to the three orbitals in the P subshell, each holding two electrons.

Energy diagrams show the increasing energy of orbitals from 1s to 2s, 2p, 3s, 3p, 4s, and 3d.
The Alf Bauer filling principle helps remember the order of orbital filling.
The S subshell has one orbital (two electrons max), the P subshell has three orbitals (six electrons max), and the D subshell has five orbitals (10 electrons max).

Different notations serve various purposes in representing electronic configurations.
Spectroscopic notation and orbital box notation are shorthand representations of energy diagrams.
Example: Carbon (atomic number 6) has the electronic configuration 1s^22s^22p^2.
Hund's rule dictates that electrons in degenerate orbitals (same energy level) remain unpaired for lower energy.
Spectroscopic notation (e.g., 1s^22s^22p^2) is a universal shorthand for electronic configuration.
Orbital box notation shows spin pairing explicitly, representing electrons as arrows in boxes. Hund's rule is shown by placing single electrons in each box before pairing.

The lecture provides a summary of the first 11 elements, showing element names, symbols, spectroscopic notation, and orbital box notation.
Increasing energy is represented from left to right in both notations.
Orbital box notation shows degenerate orbitals and electron pairing.
Electrons fill in order of increasing energy, following Hund's rule.

Noble gas notation is used as a shorthand for elements with higher atomic numbers.
Group 18 elements (noble gases) have filled shells.
Core electrons represent filled shells, while valence electrons participate in chemical reactions.
Noble gas notation focuses on valence electrons, simplifying the representation.
Example: Sodium (11 electrons) can be written as [Ne]3s^1, where [Ne] represents the neon core (10 electrons).
Elements in the same group have similar valence electron configurations.

Chromium (atomic number 24) is an exception: [Ar]4s^13d^5 instead of [Ar]4s^23d^4.
This anomaly is due to Hund's rule, favoring a half-filled d subshell for lower energy.
Anomalies occur near half-filled subshells, especially in transition metals.

Electronic structure significantly influences chemical behavior.
Filled electron shells in noble gases result in low reactivity.
Elements in the same group have similar valence electron configurations.
Group 1 (alkali metals): ns^1
Group 2 (alkaline earth metals): ns^2
Group 17 (halogens): ns^2np^5
Group 18 (noble gases): ns^2np^6
Unpaired electrons lead to high reactivity, while paired electrons result in low reactivity.
Group 14 (carbon family) shows variation in properties, from non-metal carbon to metal lead.

Effective nuclear charge: the net positive charge experienced by an electron in a multi-electron atom. Shielded by core electrons, valence electrons experience a reduced positive charge from the nucleus.
The valence electron in sodium does not experience the full +11 charge due to shielding by 10 core electrons.
A lower effective nuclear charge makes the valence electron more likely to participate in chemical reactions.

Atomic size is determined by measuring the distance between nuclei in bonded atoms (bond length).
The bonding atomic radius is half the bond length.
Example: Diatomic bromine molecule has a bond length of 228 picometers, so the bonding atomic radius is 114 picometers.

Atomic radius generally increases down a group due to increasing principal quantum number and more electrons.
Atomic radius generally decreases across a period due to increasing effective nuclear charge.
Group 1 metals have high atomic radii, while Group 17 elements have smaller radii.

Ions are formed when atoms gain or lose valence electrons.
Cations are positively charged ions formed by donating electrons.
Anions are negatively charged ions formed by receiving electrons.
Cations are smaller than their parent atoms because outermost electrons are removed, and repulsions are reduced.
Anions are larger than their parent atoms because extra valence electrons increase repulsions and occupy larger orbitals.
Ionic size increases down a group due to increasing nuclear size and valence electrons extending further out.

Sodium loses its 3s^1 electron to form Na^+, becoming smaller.
Sulfur gains two electrons to form S^{2-}, filling the 3p subshell and becoming larger.
Ion formation is driven by achieving a stable noble gas configuration.

Fluorine gains an electron to form F^-, achieving the same configuration as neon.
Sodium donates an electron to form Na^+, also achieving the neon configuration.
Ions are often more stable than their neutral atoms, and reactions will proceed to form them.

Ionization energy is the energy needed to remove an electron from an atom, trends generally increasing across the periodic table and decreasing down columns, but there are complexities and exceptions.
Removing the first electron (first ionization energy) is easier than removing subsequent electrons from core electrons.
Noble gases have the highest ionization energies, and Group 1 elements have low ionization energies.
Half-filled valence shells exhibit some additional stability due to Hund's rule, leading to anomalies in ionization energy.

Electron affinity is the energy change when an electron is added to an atom.
Chlorine has a high electron affinity, indicating it strongly wants an extra electron.
Group 1 elements have some electron affinity, while Group 2 and 18 elements have very low electron affinities.
Some anomalies occur due to electron configurations; for instance, beryllium formation is not very favorable.

Eelectronegativity is the ability of atoms in a molecule to attract electrons. It's a dimensionless scale ranging from 0.79 (cesium) to 4 (fluorine).
Electronegativity helps predict which elements will react with each other.
It increases across a period and decreases down a group, similar to electron affinity.