Periodic Trends — Atomic Radii, Ionic Radii, Ionization Energy, Electron Affinity, and Electronegativity
Atomic Radii Trends
Definition
Atomic radius (bonding atomic radius) describes the size of an atom in a particular context (often comparison within a period or group).
1 Å = 10^{-10} m.
When discussing trends, consider shielding and effective nuclear charge (ENC).
General trends across the periodic table
Within each group (vertical column): bonding atomic radius tends to increase from top to bottom.
Within each period (horizontal row): bonding atomic radius tends to decrease from left to right.
Shielding decreases going up the periodic table; ENC increases moving left to right across a period.
Small summary: large radius in the lower-left; small radius in the upper-right.
Key concepts
Shielding: inner electron shells shield outer electrons from the full nuclear charge.
ENC (effective nuclear charge): effective positive pull felt by valence electrons after accounting for shielding.
As you move down a group, additional electron shells increase shielding, but nuclear charge increases only slightly relative to shielding, so radius grows.
As you move across a period, electrons experience a stronger effective nuclear charge with little additional shielding, so radius shrinks.
Quick practice question (from transcript)
Question: Rank the following atoms in terms of decreasing atomic radius: Na, N, O, Mg, F?
Answer: d) Na, Mg, N, O, F
Rationale: radius is largest in the lower-left and smallest in the upper-right; Na (lower-left) is largest; F (upper-right) is smallest; among Na and Mg, Na is larger than Mg due to its position further left and lower.
Practical implications and connections
Radius trends help predict chemical reactivity and bond lengths: larger atoms typically form bonds with longer bond lengths; smaller atoms tend to hold electrons tighter.
ENC and shielding underpin many other periodic properties (e.g., ionization energy, electronegativity).
Ionic Radii Trends
What determines ionic size?
The size of an ion depends on:
Its nuclear charge (Z)
The number of electrons it possesses (electron count)
General rules
Cations (positively charged) are smaller than their parent atoms.
Anions (negatively charged) are larger than their parent atoms.
Isoelectronic series
Isoelectronic series: a group of ions that all contain the same number of electrons.
In an isoelectronic series, the more positively charged the species, the smaller the radius (because the same electron cloud is pulled tighter by a larger nuclear charge).
Quick practice question (from transcript)
Question: Rank Li+, Be^{2+}, He, H^−, B^{3+} in order of decreasing size.
Answer: H^− > He > Li^+ > Be^{2+} > B^{3+}
Rationale: these species are isoelectronic (2 electrons) except for their differing nuclear charges; the larger the nuclear charge, the smaller the radius for a fixed electron count.
Practical implications
Ion size affects lattice energies, lattice stability, and ionic bonding characteristics.
In biological systems and materials science, ionic radii influence solubility, diffusion, and conductivity.
Ionization Energy (Ionization Energies)
Definition
Ionization energy (IE) is the minimum energy required to remove an electron from the valence shell in the ground state (gas phase).
First ionization energy (I1): energy to remove the first electron.
Second ionization energy (I2): energy to remove a second electron after the first has been removed, and so on.
Notation examples:
ext{Na}(g)
ightarrow ext{Na}^+(g) + e^- ext{ with energy } I_1
ext{Na}^+(g)
ightarrow ext{Na}^{2+}(g) + e^- ext{ with energy } I_2
General trends
I1 < I2 < I3 < …
Across a period (left to right), ionization energies generally increase (harder to remove electrons as ENC increases).
The transition metals show only a slow gradual increase from left to right across a period.
The f-block elements also show only small variation in I1 values.
Sharp rises in IE
A sharp increase occurs when an inner-shell (core) electron is removed, because those electrons are more tightly bound.
Conceptual notes
Higher IE means greater difficulty in removing an electron; chemistry becomes less electropositive.
IE values relate to chemical reactivity, particularly in determining which elements form cations readily.
Practical example (interpretation from transcript)
The progression from removing a valence electron (low IE) to removing a core electron (high IE) reflects increasing stability of filled shells.
Connections and implications
IEs underpin redox chemistry and electrochemical series.
Trends help rationalize periodic reactivity patterns across groups and periods.
Electron Affinity
Definition
Electron affinity (EA) is the change in energy when an electron is added to the valence shell of a neutral atom.
For most atoms, energy is released when an electron is added (negative EA indicates exothermic process).
Trends
The trend is less obvious than ionization energy.
Electron affinities do not change greatly as you move down a group.
There is a slight decrease in EA downs a group, attributed to the enlargement of the atom (valence shell farther from the nucleus), which weakens the attraction for an added electron.
Practical note
EA helps explain tendencies to form anions and the strength of bonding in compounds.
Conceptual recap
IE measures the ease of removing electrons; EA measures the tendency to gain electrons.
Transcript practice context
A demonstration question contrasted IE and EA to identify valence shell behavior and trends across the periodic table.
Electronegativity
Definition
Electronegativity is the ability of an atom in a molecule to attract electrons toward itself.
Scale
The most widely used scale is the Pauling scale, developed by Linus Pauling.
General trends
Electronegativity increases from left to right across a period.
Electronegativity decreases with increasing atomic number down a group.
Practice question (from transcript)
Rank the following in decreasing electronegativity: N, P, Sb, Bi
Answer: N > P > Sb > Bi
Quantitative comparison
Example question: What is the difference in electronegativity between N and F atoms?
EN differences explain bond polarity and the nature of chemical bonds (ionic, covalent, polar covalent).
Connections
EN correlates with dipole moments, bond strength, and reaction energetics in organic and inorganic chemistry.
Summary of Key Concepts and Connections
Shielding and ENC drive periodic trends in atomic size and reactivity.
Atomic radii increase down a group and decrease across a period; ionic radii follow related patterns with charge (cations smaller, anions larger).
Ionization energies generally rise across a period and show a pronounced jump when core electrons are involved.
Electron affinity trends are subtler but generally become less favorable down a group due to increased distance from the nucleus.
Electronegativity rises across a period and falls down a group, influencing bond type and molecular polarity.
Practical implications span reactivity, bonding, materials properties, and biological processes; understanding these trends aids in predicting compound formation and behavior.
Practice Problems (summary answers)
Atomic radii ranking (decreasing): Na > Mg > N > O > F
Ionic radii (decreasing size) in isoelectronic series for 2 electrons: H^- > He > Li^+ > Be^{2+} > B^{3+}
Electron affinities and their trends: generally small downward trend down a group; energy release when adding an electron is common but varies by element.
Electronegativity ranking: N > P > Sb > Bi
Difference in EN between N and F: roughly 0.94 (χ(F) ≈ 3.98, χ(N) ≈ 3.04)