Electron Configuration and Periodic Trends — Study Notes

Neon and Noble Gas Configuration

• Proper nomenclature starts with the n value (principal quantum number) and then the subshell; example for neon shows how to read across blocks.
• Neon electron configuration (full): $1s^2 \; 2s^2 \; 2p^6$
• How to read the sequence: begin at 1s (n=1, s-block), then move to n=2 (2s, then 2p).
• This gives the full configuration for neon: $1s^2 \; 2s^2 \; 2p^6$
• Noble-gas shorthand and why it matters: if a problem doesn’t specify a full configuration, you can use the noble-gas core to skip the inner transition metals; e.g., for many elements past the noble gas core, write as [Noble Gas] followed by the remaining outer-shell electrons.
• Example: if asked for a 3s^1 configuration (Na-like), you can use the noble gas core: [Ne] 3s^1.
• In practice, the noble-gas approach is especially important for transition metals, where inner d-block electrons become relevant but may be treated separately from the outer valence electrons.

valence electrons and counting with silicon example

• Silicon’s valence electron count is 4 (outer shell configuration is 3s^2 3p^2).
• Silicon example in context: to determine the number of electrons in the outermost shell, identify the highest energy level n and sum its s and p contributions to get the valence electrons.
• For germanium (Ge) and the discussion of the fourth shell, the valence electrons are four: Ge has 4s^2 4p^2 as its valence shell.
• When writing a valence shell for elements in the n=4 shell, you use the “fourth row” conventions: Ge = [Ar] 3d^10 4s^2 4p^2; here the 3d^10 is part of the core in many contexts, but the 4s^2 4p^2 are the valence electrons.
• The general idea: valence electrons are those beyond the previous noble gas core; they determine reactivity and bonding.

Noble-gas core for germanium and the d-block discussion

• Germanium (Ge) example: write in noble-gas form as [Ar] 4s^2 4p^2, but you must include the d-block contribution because 3d^10 is filled in between argon and the 4s/4p level (not part of argon’s configuration).
• Full Ge electron configuration: $[ ext{Ar}]\; 3d^{10} \; 4s^2 \; 4p^2$
• Valence electrons for Ge: 4 (4s^2 + 4p^2).
• The key point: the d^10 core (3d^10) is not counted as valence electrons, but it must be written in the full configuration since it lies between the noble-gas core and the outer p-levels.
• When counting electrons for a neutral atom, the total number of electrons equals the number of protons (Z).

Chromium and copper: transition-metal exceptions and box diagrams

• General idea: transition metals can rearrange electrons between s and d orbitals to reach lower energy (stability) by making half-filled or fully filled subshells.
• Hund’s rule reminder (box diagram): fill with single electrons of parallel spins before pairing.
• Chromium (Cr) example (box-diagram reasoning and exception):
  • If you start from [Ar] with 4s^2 and 3d^4, you notice that moving one electron from 4s to 3d can give a more stable arrangement: 4s^1 3d^5 (two half-full orbitals rather than one full and one incomplete).
  • The actual electron configuration of chromium is: $[ ext{Ar}]\; 4s^1 \; 3d^5$
  • The textbook rationale: half-filled 3d (3d^5) is particularly stable; by moving one electron from 4s, the system lowers its energy.
• Copper (Cu) example: transition metals can also gain stability by achieving a fully filled d-sublevel.
  • Copper’s configuration (as stated in the transcript) is: $[ ext{Ar}]\; 4s^1 \; 3d^{10}$
  • Note: the commonly cited ground-state configuration is often written as $[ ext{Ar}]\; 3d^{10} \; 4s^1$, which is equivalent to the above and reflects the close energy of 4s and 3d orbitals.
• Additional transition-metal connection: moving down a period can lead to analogous patterns; the transcript mentions Mo (molybdenum) as an example: “Mo would be 4d^4” in a hypothetical analogy, illustrating the idea that electrons can rearrange to achieve half-filled or fully filled d subshells across the block.
• Copper’s special case explains why the 4s and 3d orbitals are often near-degenerate in energy for transition metals, enabling such rearrangements.

Box diagrams, Hund’s rule, and implications for periodic trends

• Box diagram practice: for Cr, starting with [Ar] 4s^2 3d^4, you would place electrons in the diagram with parallel spins before pairing, leading to 4s^2 and 3d^4 in the simple fill; but recognizing the stability of 3d^5 with one electron in 4s yields Cr = [Ar] 4s^1 3d^5.
• The key implication: stability of half-filled (d^5) or fully filled subshells drives deviations from a naive Aufbau filling, and explains transitions metals’ chemistry and exceptions to simple filling rules.
• So, the “correct” chromium configuration (as excited by the speaker) is not simply 4s^2 3d^4, but 4s^1 3d^5; the latter is lower in energy due to exchange energy and symmetry considerations.

Implications of half-full and fully filled orbitals for periodic trends

• The stability of half-filled and fully filled subshells helps explain why certain elements have unexpected configurations.
• As a consequence, some elements show deviations from the simple Aufbau progression, especially in the d-block where the energy gap between s and d is small.
• These effects influence chemical behavior and periodic trends (e.g., ionization energies and oxidation states).

Copper and chromium: real-world takeaway for electron configurations

• Copper’s anomalous configuration is tied to achieving a filled d-subshell: a tendency for d^10 stability when possible.
• Chromium’s anomaly is tied to achieving a half-filled d-subshell (d^5) with minimal energy via moving an electron from the s to d orbital.
• In both cases, the result is a set of configurations that differ from a straightforward Aufbau fill and illustrate why transition metals can be exceptions.

Covalent radius, orbitals, and periodic trends

• Covalent radius grows as the principal quantum number n increases (orbital size increases with n).
• Across a period, covalent radius generally decreases from left to right; moving from right to left, covalent radius increases (the transcript’s example mentions Kr, Br, Ge in that order as you move left).
• The smallest atoms are generally in the upper-right corner of the periodic table, while the largest are in the bottom-left (consistent with the standard trend: radii shrink across a period and grow down a group).
• Example trend discussion from the transcript: krypton (Kr) → bromine (Br) → germanium (Ge) illustrates increasing radius as you move from right to left within a row.
• Ion size changes: when you remove electrons (cation formation) or add electrons (anion formation), the effective nuclear charge felt by the remaining electrons changes, altering the radius. For example, removing electrons (fewer electrons) generally reduces shielding and can lead to a smaller radius, while adding electrons (as in Cl^- compared to Cl) increases radius.
• Conceptual note: covalent radius is a practical measure; direct diameter measurements are challenging, so radii are used to compare sizes and trends.

Using electrons to identify an element in neutral atoms

• In neutral atoms, the number of electrons equals the number of protons (Z).
• By summing the electrons in a given configuration, you can determine Z and then identify the element on the periodic table.
• Example given in the transcript: an electron configuration like $1s^2 \; 2s^2 \; 2p^5$ sums to 9 electrons, indicating 9 protons, which corresponds to fluorine.

Quick practical takeaways for exams

• For a neutral atom, total electrons = total protons = atomic number Z.
• Use noble-gas core to simplify electron configurations for heavier elements: write [Xe], [Ar], [Ne], etc., then add the outer-shell electrons.
• Know the common chromium and copper exceptional configurations:
  • Cr: $[ ext{Ar}]\; 4s^1 \; 3d^5$
  • Cu: $[ ext{Ar}]\; 4s^1 \; 3d^{10}$ (often written as [ ext{Ar}] 3d^{10} 4s^1).
• Remember Hund’s rule when filling box diagrams: electrons occupy separate orbitals with parallel spins before pairing.
• Be aware that s and d orbitals can be close in energy in transition metals, which allows electron promotion between s and d to reach more stable configurations.
• Covalent radius trends: covalent radius increases from right to left across a period and increases down a group; smallest atoms tend to be at the top-right, largest at the bottom-left.
• When discussing ions vs. neutral atoms, adding electrons generally increases radius (e.g., Cl^- vs Cl), while removing electrons reduces radius due to decreased shielding and higher effective nuclear charge.