Electron Configurations: Shells, Subshells, and Patterns (9/3)

Shells and Subshells

• There are shells (principal quantum number n) and subshells (s, p, d, f).
• The second shell (n = 2) has an s subshell and a p subshell.
  • s subshell: can hold up to 2 electrons: s^2.
  • p subshell: can hold up to 6 electrons: p^6.
  • Total for the second shell: 2 + 6 = 8 electrons.
• The shapes mentioned: s is spherical; p is peanut-shaped.
• Electrons occupy space in orbitals called subshells; the total space a shell/subshell can hold follows the capacities above.
• Note: the shell and subshell do not have to be filled completely. For example, a subshell could hold only one electron if you run out of electrons in the filling process.

How to figure out the electron configuration (construction method)

• Start at hydrogen and fill electrons across and down the periodic table following the energy order of orbitals.
• The plan shown uses the block labels:
  • s-block, p-block, d-block, f-block.
• Example given: locating nickel (Ni) with 28 electrons.
• Steps used in the transcript to fill Ni
  • Ni has 28 electrons; count from the top left across and down to the Ni position on the table.
  • The filling order follows the sequence of blocks: 1s, then 2s, then 2p, then 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, etc. (typical Aufbau-like scheme; energy ordering can have nuanced variations in real systems).
• For the specific Ni example, a standard way to write the electron configuration (neutral Ni, 28 electrons) is:
  • 1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 4s^2 \ 3d^8
  • An alternative commonly seen ordering (emphasizing the common presentation) is 1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 4s^2 \ 3d^8 (the latter is equivalent; the 4s and 3d electrons are close in energy).
• The row numbers correspond to periods on the periodic table; the letter corresponds to the subshell (s, p, d, f);
  • The count after the letter is the number of electrons in that subshell.
• Pattern recognition: you fill from left to right across the block until you reach the end of that block, then move to the next block down and right.

Pattern terminology and the “break” for d-block

• s-block: follows the pattern where the shell number equals the period number for the first blocks you traverse.
• p-block: continues with the appropriate shell number, adding up to 6 electrons in the p subshell per period.
• d-block (the transition metals) is where a pattern break occurs:
  • When you are filling a d subshell, the shell number is the period number minus one:
  • Example: in the fourth period (period number = 4), the d-block corresponds to the third shell, so you get 3d orbitals (3d^x).
  • The transcript example uses: for the d-block in the fourth period, the shell referenced is the 3d subshell, with the filling described as 3d^8 (for Ni’s valence in the common Ni configuration).
• This is why you often see configurations written with 3d before 4s in teaching contexts (and sometimes 4s before 3d in other contexts); they are both correct representations depending on the emphasis on energy ordering versus written order.

Nickel and electron-counting practice

• Nickel (Ni) neutral: 28 electrons.
• Periodic-table location helps you determine the core and valence electrons.
• For Ni, common full configuration (neutral) is often written as:
  • [Ar] \, 3d^8 \, 4s^2
  • which corresponds to the explicit full form: 1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 4s^2 \ 3d^8
• The core, [Ar], accounts for the first 18 electrons, leaving 10 electrons accounted for in the 4s and 3d subshells combined (2 in 4s and 8 in 3d).

Ions and isoelectronic concepts

• If you remove electrons to form an ion (e.g., Ni with a 2+ charge), you remove electrons from the outermost subshells.
• Nickel(II) ion: Ni^{2+}
  • Ni neutral has 28 electrons; Ni^{2+} has 26 electrons.
  • A common way to write Ni^{2+} configuration is by removing electrons from the outermost shell(s): from 4s and then 3d as needed, giving something like [Ar] \, 3d^8 in many textbook representations (the 4s electrons are removed first in many cases).
• Isoelectronic concepts: two species are isoelectronic if they have the same total number of electrons. In the transcript, there is a claim that Ni^{2+} has the same configuration as iron (Fe); in standard chemistry, Ni^{2+} is often described as having the same electron count as Fe^{2+} (both with 26 electrons), but the detailed orbital occupancy is not identical: Ni^{2+} is commonly written as [Ar] \, 3d^8 while Fe^{2+} is [Ar] \, 3d^6 (when Fe^{2+} loses its 4s electrons first). Thus, Ni^{2+} and Fe^{2+} are not strictly the same configuration. The broader point is that “isoelectronic” means the same total number of electrons, but the specific orbital occupancies can differ depending on the element and its ionization state.
• Practical note: the transcript’s statement that Ni^{2+} has the same configuration as Fe is a common confusion and should be checked against the standard electron-counting rules.

Noble-gas shorthand configurations

• Shorthand (noble gas) notation uses the noble gas that completes the inner shells as a prefix, then the remaining valence electrons.
• The noble gas chosen is the one that fills the shells up to the previous closed shell for the element in question.
• Example for Ni (28 electrons):
  • Core: [Ar] (the argon core accounts for the first 18 electrons).
  • Valence: 4s^2 3d^8
  • Shorthand form: [Ar] \, 3d^8 \, 4s^2 (some texts may write 4s^2 before 3d^8; both convey the same total electron count).
• Conceptual note on the “before Argon” statement in the transcript: the noble gas immediately preceding Argon is Neon (Ne). For shorthand purposes, we typically use the noble gas that ends the previous complete shell for the element family under consideration (e.g., [Ar] for 4th period elements, [Kr] for 5th period, etc.). The transcript’s wording can be read as a guidance to use the preceding noble gas; the standard and widely accepted convention is to use [Ar] for Ni and similar elements in the 4th period.

f-block and Nd example (lanthanides)

• The f-block elements involve filling 4f orbitals and lie in period 6 (the lanthanide series).
• Rule for the f-block:
  • The shell involved is n = period number − 2. For period 6, that gives 4f.
  • Example occupancy indicator from the transcript: 4f^3 (as part of a demonstration). The transcript shows a Nd example used to illustrate the f-block, including a partial occupancy in 4f and related subshells.
• Real-world note: neodymium (Nd) has a typical ground-state configuration of [Xe] \, 4f^4 \, 6s^2 in many standard references, illustrating the complex near-degeneracy and variability of f-block occupancies. The transcript’s specific Nd example may illustrate a hypothetical filling pattern but is not the canonical ground-state configuration.
• The general lesson: after filling up to Xenon (Xe) in the core, the next electrons occupy 4f, then 5d or 6s as appropriate, depending on the exact element and its place in the lanthanide series.

Exceptions and practical cautions

• There are well-known exceptions to the simple aufbau-like filling rules, especially toward the bottom of the periodic table (transition metals and lanthanides).
• Commonly cited exceptions discussed in the transcript include copper (Cu), silver (Ag), and gold (Au) as well as chromium (Cr) and molybdenum (Mo). These exceptions often involve unusually stable configurations such as shifting electrons to achieve filled or half-filled subshells (e.g., Cu: [Ar] 3d^{10} 4s^1 in many treatments; Cr: [Ar] 3d^5 4s^1).
• The instructor notes that for this course you may be expected to follow the straightforward filling pattern and not memorize every exception; exceptions exist and do appear in homework or more advanced problems, but the focus in this course is on the general Aufbau-like progression.
• Important caution: always confirm actual electron configurations from reliable references for exam questions if the problem explicitly involves these edge cases, since the details can vary by source and context.

Quick reference: key rules and useful formulas

• Subshell capacities:
  • s: 2 electrons
  • p: 6 electrons
  • d: 10 electrons
  • f: 14 electrons
• Energy-ordering guidelines (simplified Aufbau progression):
  • 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p
• D-block pattern rule (shell number for filling d-subshell):
  • In period p, the corresponding d-subshell belongs to shell n = p − 1 (e.g., 3d for period 4).
• F-block pattern rule (shell number for filling f-subshell):
  • In period p, the corresponding f-subshell belongs to shell n = p − 2 (e.g., 4f for period 6).
• Shorthand notation principle:
  • Use the nearest noble gas core to represent the filled inner shells, e.g., [Ar] \, ext{(rest)} for elements in the 4th period and beyond.
• Isoelectronic concept:
  • Isoelectronic species have the same total number of electrons; they may have different orbital occupancies depending on the element and ionization state.
• Practical ionization note:
  • When forming cations, electrons are removed from the outermost (highest-energy) electrons first (often from the s before the d in transition metals), which changes the observed configuration from the neutral atom.

Quick examples to practice

• Nickel neutral (Ni): 28 electrons → configuration examples
  • 1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 4s^2 \ 3d^8
  • Shorthand: [Ar] \, 3d^8 \, 4s^2
• Nickel(II) ion (Ni^{2+}): 26 electrons
  • Common representation: [Ar] \, 3d^8 (removal of the two 4s electrons is typical).
• Isoelectronic comparison caution:
  • Ni^{2+} is not strictly isoelectronic with Fe in terms of orbital occupancy; Ni^{2+} is often written as [Ar] \, 3d^8 while Fe^{2+} is [Ar] \, 3d^6. Both have 26 electrons, but the distributions differ in the d- and s- subshells.
• Lanthanide example (Nd): common ground-state form is [Xe] \, 4f^4 \, 6s^2, illustrating the 4f involvement after Xe in period 6.
• A cautionary note on the transcript’s Nd and Nd-like examples: some occupancy patterns shown (e.g., 4f^3 5d^1) may appear in teaching demonstrations, but the canonical ground-state configurations can differ and depend on the adopted convention and specific element.

Connections to broader concepts

• The electron configuration concept links to chemical periodicity, chemical reactivity, and the formation of ions.
• The s/p/d/f block organization mirrors quantum numbers and angular momentum quantum numbers (l = 0 for s, l = 1 for p, l = 2 for d, l = 3 for f).
• The energy-ordering considerations explain why transition metals can show unexpected stability patterns (e.g., half-filled or fully-filled subshells) and why the observed colors and magnetic properties are influenced by d- and f-electron occupancy.
• Real-world relevance: understanding electron configurations helps predict element behavior in reactions, bonding preferences, and properties across the periodic table.

Summary of key ideas from the transcript

• Shells and subshells can be partially filled; there are 8 electrons max in the second shell (s^2 + p^6).
• The method of building electron configurations involves moving through the periodic table by blocks (s, p, d, f) and shells, counting electrons as you go.
• The d-block pattern is a slight deviation where the relevant shell number is period minus one (e.g., 3d in period 4).
• Ions change electron counts; Ni^{2+} commonly considered to have 26 electrons, often represented as [Ar] 3d^8, and it is not strictly isoelectronic with Fe in terms of orbital occupancy.
• Shorthand noble-gas configurations are a compact way to write electron configurations by starting from the noble gas core just before the element in question.
• The f-block occupancy involves 4f orbitals (period 6), with the rule shell n = period − 2 for f orbitals being relevant; practical examples like Nd illustrate the concept, though actual occupancies can vary in real data.
• There are exceptions to the simple filling rules, notably Cu, Ag, Au, Cr, and Mo, and sometimes these show up in homework or advanced problems. The course emphasizes the general pattern but acknowledges exceptions.