Periodic Trends Atomic Radis, Ionic Radius, Ionization Energy, and Electron Affinity

Overview

John Flood Chemistry connects electron configurations and valence electrons to periodic trends in atomic radius, ionic radius, ionization energy, and electron affinity. The lesson explains orbital-filling rules and half-filled or fully filled subshell stability, then uses effective nuclear charge, electron shielding, and examples such as chromium, copper, tellurium, iodine, and the isoelectronic series O²⁻/Ne/Mg²⁺ to predict relative sizes and energy changes.

Key takeaways

Across a period, core-electron shielding changes little while proton number rises, so effective nuclear charge increases and atomic radius generally decreases.
Down a family, valence electrons occupy higher principal energy levels farther from the nucleus, so atomic radius and metallic character generally increase.
Adding electrons to form an anion generally expands the electron cloud, while removing electrons to form a cation generally contracts it.
For isoelectronic species, the species with more protons is smaller: among O²⁻, Ne, and Mg²⁺, Mg²⁺ is smallest and O²⁻ is largest.
Ionization energy generally decreases as the electron’s distance or subshell energy increases, but half-filled and fully filled subshell stability can make specific electrons harder to remove.
Chromium and molybdenum favor half-filled d subshells, while copper and silver favor fully filled d subshells, producing their characteristic anomalous electron configurations.

Chapters

0:00 Orbital Diagrams, Subshells, and Quantum-Mechanical Nodes

An electron configuration lists occupied orbitals, while an orbital diagram shows their relative energies and electron spins.
Energy shells contain subshells—s, p, and d—with allowed angular momentum values limited by the principal quantum number.
Nodes are regions of zero electron probability; they can be spherical or angular, and p and d orbitals have one and two angular nodes, respectively.

10:00 Aufbau, Pauli Exclusion, and Hund’s Rule in Orbital Filling

The Aufbau principle fills lower-energy orbitals before higher-energy orbitals, though the energy ordering can cross between shells.
The Pauli exclusion principle prevents two electrons in an atom from sharing all four quantum numbers.
Hund’s rule places electrons singly with aligned spins in orbitals of the same subshell before pairing them.

17:00 Why Chromium and Molybdenum Have Anomalous Configurations

Chromium is [Ar] 4s¹ 3d⁵ rather than the simple predicted [Ar] 4s² 3d⁴ configuration.
Molybdenum follows the analogous pattern, [Kr] 5s¹ 4d⁵, producing two half-filled subshells.
John Flood Chemistry attributes these exceptions to the greater stability of half-filled or fully filled subshells when nearby orbital energies make electron rearrangement possible.

21:00 Copper and Silver: Stable Fully Filled d Subshells

Copper is [Ar] 4s¹ 3d¹⁰ rather than the predicted [Ar] 4s² 3d⁹.
Moving one 4s electron into 3d gives copper a half-filled 4s subshell and a fully filled 3d subshell.
Silver displays the corresponding anomalous pattern, while the lesson identifies chromium/molybdenum and copper/silver as the key families to recognize.

27:00 Carbon and Silicon Orbital Diagrams: Shared Valence Patterns

Students apply Aufbau, Pauli, and Hund’s rules to draw orbital diagrams for carbon’s 6 electrons and silicon’s 14 electrons.
Carbon and silicon each have two unpaired electrons and the same outer-shell arrangement: s²p².
The 2s and 2p energy gap is too large for carbon to gain stability by shifting an s electron into p, unlike certain transition-metal configurations.

34:00 Valence Electrons and Why Periodic Families Resemble Each Other

An orbital is described as the region containing about 95% of an electron’s probability density, with a specific size, shape, and orientation.
Carbon and silicon each have four valence electrons in their outermost s and p subshells, helping explain their shared chemical patterns.
For main-group elements, valence electrons are the outer-shell s and p electrons; for transition metals, the lesson includes relevant s, d, and f electrons.

38:00 Iodine’s Seven Valence Electrons Versus Its Full Configuration

Iodine’s configuration is [Kr] 5s² 4d¹⁰ 5p⁵, but its main-group valence count is seven: 5s² plus 5p⁵.
The 4d¹⁰ electrons belong in iodine’s full configuration but are not counted as its main-group valence electrons.
The group pattern links fluorine, chlorine, bromine, and iodine through their seven valence electrons.

47:03 Metallic Character and Atomic Size Increase Down a Family

Metallic character generally increases down a periodic-table family: carbon is a nonmetal, silicon and germanium are metalloids, and tin is a metal.
Atomic size also increases down a family because valence electrons occupy successively higher energy shells farther from the nucleus.
The nitrogen family illustrates the metallic trend as it progresses from nonmetal nitrogen to metallic bismuth.

55:21 Covalent Radius: Estimating Atomic Size from Bond Length

An atom’s electron cloud has no sharp outer edge, so its radius cannot be measured as if it were a solid sphere.
For a bond between two identical atoms, covalent radius is estimated as half the covalent bond length.
The F–F and I–I bonds illustrate how bond lengths provide estimates of the smaller fluorine and larger iodine radii.

58:32 Effective Nuclear Charge Explains Shrinking Across a Period

Effective nuclear charge is introduced as nuclear charge minus shielding from core electrons.
Across a period, proton number increases while the number of core electrons stays approximately constant, strengthening the nucleus’s pull on valence electrons.
For atoms in the same shell, this stronger effective pull contracts the electron cloud and decreases atomic radius from carbon toward fluorine.

1:10:53 Using the Periodic Table to Count Core and Valence Electrons

The periodic table serves as an orbital-filling roadmap as well as an arrangement of elements; its s- and p-block positions show outer-shell electrons.
Silicon has four valence electrons and 10 core electrons, while iodine has seven valence electrons and 46 core electrons in the lesson’s counting scheme.
For main-group elements, core-electron counting must distinguish outer-shell s and p electrons from inner d electrons, as iodine’s 4d¹⁰ demonstrates.

1:16:51 Tellurium and Iodine: Comparing Radii Across Period 5

Tellurium and iodine both have 46 core electrons in the lesson’s model, but their atomic numbers are 52 and 53.
Their effective nuclear charges are therefore estimated as +6 for tellurium and +7 for iodine.
Because iodine pulls more strongly on valence electrons in the same shell, it has a smaller atomic radius than tellurium.

1:20:48 Anions Expand and Cations Contract Relative to Neutral Atoms

Adding an electron to form an anion increases electron–electron competition while proton number stays fixed, generally increasing ionic radius.
Removing an electron to form a cation leaves fewer electrons for the same nuclear charge to attract, generally decreasing radius.
The lesson uses the magnet analogy: greater competition weakens each electron’s effective attraction, while fewer electrons allow a stronger pull.

1:22:06 Isoelectronic Radii: O²⁻, Neon, and Mg²⁺

O²⁻, neon, and Mg²⁺ each have 10 electrons and the configuration 1s² 2s² 2p⁶, making them an isoelectronic series.
Within an isoelectronic series, proton number determines the relative nuclear pull on the same number of electrons.
Mg²⁺ is smallest because magnesium has 12 protons; O²⁻ is largest because oxygen has only 8 protons.

1:24:36 Ionization Energy: Energy Required to Form Cations

Ionization energy is the energy required to remove an electron from an atom and form a cation.
The process is endothermic because energy must be supplied to free an electron from its attraction to the nucleus.
An energy-well diagram represents the bound electron at lower energy than an electron infinitely far from the nucleus.

1:28:32 Successive Ionization Energies Rise as Electrons Are Removed

The first ionization energy, IE₁, is lower than the second, IE₂, in the calcium example.
After calcium loses one electron, the positively charged ion attracts its remaining electrons more strongly.
Removing the next electron therefore requires overcoming a stronger attraction and supplying more energy.

1:33:34 Atomic Radius and Subshell Energy Predict Ionization Energy

Larger atomic radius generally corresponds to lower ionization energy because more distant electrons are easier to remove.
Sodium’s outer electron is farther from the nucleus than lithium’s, so sodium has the lower first ionization energy.
Within the same principal shell, electrons in higher-energy subshells are easier to remove; aluminum loses a 3p electron more readily than magnesium loses a 3s electron.

1:36:42 Subshell Stability Creates Ionization-Energy Exceptions

A partially filled subshell can have lower ionization energy than a particularly stable half-filled or fully filled subshell.
Vanadium’s [Ar] 4s² 3d³ configuration allows removal from a partially filled d subshell more readily than removal that disrupts chromium’s half-filled 3d⁵ subshell.
The class corrects and discusses the comparison: manganese is also [Ar] 4s² 3d⁵, so removing an electron disrupts a half-filled d subshell.

1:39:18 Half-Filled d-Shell Stability and the Chromium Comparison

Chromium’s [Ar] 4s¹ 3d⁵ configuration contains a half-filled d subshell, making it harder to ionize than vanadium in the comparison.
Removing a d electron from vanadium does not break the same half-filled arrangement, so its ionization energy is lower.
The stability argument complements the general trends based on atomic radius and subshell energy.

1:43:37 Electron Affinity: Adding Electrons to Form Anions

Electron affinity describes adding an electron to an atom, corresponding to anion formation.
The process is generally exothermic in the lesson’s energy model, so released energy is represented by a negative value.
A more negative electron-affinity value means more energy is released; the lesson’s main prediction is that greater effective nuclear charge strengthens attraction for an added electron.

1:46:03 Reviewing Periodic-Trend Practice and Previewing Structure and Bonding

The assigned review asks students to compare atomic and ionic sizes, organize isoelectronic species, and explain why ionization is endothermic while electron addition is exothermic.
John Flood Chemistry previews Unit 4 on structure and bonding, which connects to topics already appearing in laboratory work.
The class is preparing for a midterm with two signup windows, reflecting a possible schedule extension into the second week.

Summary, takeaways, and chapters were generated by AI from the video's transcript and may contain errors. The video belongs to its creator, John Flood Chemistry.