Periodic Properties of Elements: Electron Arrangement and Atomic Trends

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Ch. 9: Periodic Properties of Elements

How the Arrangement of Electrons Affects Properties of Elements

The periodic properties of elements are determined by the arrangement of electrons in atoms. Understanding electron configuration and related quantum principles is essential for predicting chemical behavior and trends across the periodic table.

Spin Quantum Number, ms

Definition: The spin quantum number (ms) describes the intrinsic angular momentum (spin) of an electron in an orbital.
Electrons in the same orbital do not have the same energy due to their different spins.
Spin determines the magnetic field, which affects energy.
Allowed values: ms = +1/2 (spin up) or -1/2 (spin down).

Pauli Exclusion Principle

Statement: No two electrons in the same atom can have identical quantum numbers.
Only two electrons (with opposite spins) can occupy the same orbital.
Example: The atomic orbital below at energy level 3 has 2 electrons. Each electron must have a set of 4 unique quantum numbers.

Energies of Orbitals

One-electron atoms (e.g., hydrogen): Orbitals at the same energy level are degenerate (same energy).
Many-electron atoms: Electron repulsions cause orbitals at the same energy level to have different energies (not degenerate).

Electron Configuration: Distribution of Electrons in Orbitals

Notation: Superscript indicates the number of electrons in each orbital (e.g., 1s2).
Ground-state configuration: Electrons fill orbitals starting from the lowest energy level (Aufbau principle).
Maximum electrons per orbital: 2

Order of Filling Electrons in Orbitals

Electrons fill orbitals in order of increasing energy, following the diagonal rule (shown by arrows in diagrams).
Order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p

Orbital Diagram

Electron configuration can be represented using arrows for electron spins.
Electrons in the same subshell are placed one per orbital before pairing (Hund's rule).
Unpaired electrons are easily identified in orbital diagrams.
Magnetic properties:
- Paramagnetic: Atoms with unpaired electrons; affected by magnetic fields.
- Diamagnetic: All electrons are paired; not affected by magnetic fields.

Condensed Electron Configuration

Uses the symbol of the previous noble gas in brackets to represent core electrons (e.g., [Ne] 3s2 3p1).
Valence electrons are those outside the noble gas core.
Example: Si: [Ne] 3s2 3p2

Row 4, 5, and 6 Elements: Filling Order

For Row 4: 4s orbital is filled before 3d due to slightly lower energy.
For Row 5 & 6: 5s is filled before 4d, and 6s before 4f and 5d.
Example: Zr: [Kr] 5s2 4d2

Exceptions in Transition Elements

Some transition elements have electron configurations that differ from the expected order due to stability of half-filled or fully filled d orbitals.

Element	Expected Configuration	Actual Configuration	Reason
Cu	[Ar] 4s2 3d9	[Ar] 4s1 3d10	Full d orbital stability
Cr	[Ar] 4s2 3d4	[Ar] 4s1 3d5	Half-filled d orbital stability
Ag	[Kr] 5s2 4d9	[Kr] 5s1 4d10	Full d orbital stability

Valence Electrons: s, p, d, and f Blocks

Valence electrons are found in the outermost s and p orbitals for main group elements, and in d or f orbitals for transition and inner transition elements.
Periodic table blocks:
- s-block: Groups 1 & 2
- p-block: Groups 13-18
- d-block: Transition metals (Groups 3-12)
- f-block: Lanthanides and actinides

Electron Configuration of Ions

Neutral atoms lose or gain electrons from the highest energy orbital.
Cations: Lose electrons (usually from s orbital first).
Anions: Gain electrons.
Transition metals lose s electrons before d electrons.

Properties of Elements: Explained by Electron Configuration and Valence Electrons

Elements in the same group have similar properties due to similar valence electron configurations.
Elements in the same period show repeating patterns.

Reactivity Trends

Group 1A: Very reactive metals
Group 2A: Reactive metals
Group 7A: Very reactive nonmetals
Group 8A: Chemically unreactive (noble gases)

Size of Atoms and Effective Nuclear Charge (Zeff)

Atomic size: Refers to the volume occupied by an atom.
Size is determined by the attraction between electrons and nucleus, and repulsion between electrons.
Effective nuclear charge: The net positive charge experienced by valence electrons.

Formula:

Z: Number of protons
S: Number of core electrons (screening electrons)
Across a period, Zeff increases as the number of protons increases but core electrons remain the same.

Atomic Radii

Atomic radius decreases from left to right across a period due to increasing Zeff.
Atomic radius increases from top to bottom of a group due to increasing energy levels (distance from nucleus).

Ions: Cations and Anions

Cations: Atoms that have lost electrons; smaller than neutral atoms.
Anions: Atoms that have gained electrons; larger than neutral atoms.
For the same element:

Ionization Energy

Definition: The energy required to remove an electron from an atom in the gas phase.
First ionization energy () removes the first electron.
Down a group: decreases (easier to remove electrons).
Across a period: increases (harder to remove electrons).

Example: For Si: kJ/mol

Ionization Energy and Reactivity

Reactivity of metals increases down a group (lower ).
Reactivity of nonmetals decreases down a group (lower electron affinity).

Electron Affinity (EA)

Definition: The ability of an atom to take electrons.
Group 7A elements have the most exothermic electron affinities (favorable process).
Smaller nonmetals are more reactive than larger nonmetals.
Exothermic EA means energy is released when an electron is added.

Summary Table: Periodic Trends

Property	Trend Across Period	Trend Down Group
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electron Affinity	More negative (exothermic)	Less negative
Effective Nuclear Charge	Increases	Decreases slightly

Practice Problems and Examples

Write ground state electron configurations for N, Na, S, Ar.
Write electron configuration for an excited state aluminum atom.
Use the periodic table to write condensed ground state electron configurations and orbital diagrams for Ca, I, Fe, Pb.
Arrange atoms in order of decreasing radius: S, Cl, Na, Li, Cs, Ge, Sn.
Determine which element has the lowest ionization energy: Li, Na, Mg, Ca, Ba.
Determine which element is more reactive and why: I or Xe, Mg or Na, Mg or Ca, F or Br.

Additional info: These notes expand on the original slides by providing definitions, explanations, and context for quantum numbers, electron configuration, periodic trends, and atomic properties. Tables and diagrams have been described and recreated in text and HTML format for clarity.