Electron Configuration and Reactivity

Electron Configurations and Orbital Theory (Notes from Transcript)

• Subshell capacity and organization

  • s subshell: 1 orbital, can hold up to 2 electrons
  • p subshell: 3 orbitals (px, py, pz), can hold up to 6 electrons total
  • d subshell: 5 orbitals, can hold up to 10 electrons total
  • Orbital diagrams show all orbitals and their spin states; electron configurations show the principal quantum number n and the azimuthal quantum number l (the “types” of orbitals) without explicit ml and ms labels
  • In orbital diagrams you will see ml (which orbital: px, py, pz) and ms (spin: up or down)

• Notation and visualization differences

  • Electron configuration: focuses on n and l (e.g., 1s, 2s, 2p, 3s, 3p, 4s, 3d, …)
  • Orbital diagram: shows all orbitals in a subshell with arrows indicating electrons and their spins
  • You can use either half-arrows (↑ for spin up, ↓ for spin down) or double-headed arrows; both convey the same information

• Hund’s rule and electron filling order

  • Hund’s rule: electrons fill degenerate orbitals singly first, with parallel spins, before any pairing occurs
  • After the singly occupied stage, electrons pair up in orbitals as needed
  • The filling is done in order of increasing energy ( Aufbau principle ), recognizing that energy gaps exist between shells and subshells

• Energy ordering and the “gaps” between shells

  • Early energy gaps (e.g., between 1s, 2s, and 2p) are relatively large, so moving an electron from one shell to another is energetically significant
  • As you move farther from the nucleus, the energy differences between successive shells/subshells become smaller, so adding electrons in higher shells has less dramatic energy changes
  • Classic filling sequence (simplified):
  • 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < …
  • Beyond 4p the ordering can become less intuitive and “scramble” in practice due to interactions and penetration effects
  • The visual analogies (e.g., bakery smell) illustrate why energy gaps shrink as distance from the nucleus increases

• Building full electron configurations as a blueprint

  • Writing out the full configuration helps identify which electrons are core versus valence
  • Core electrons are those in completely filled shells closer to the nucleus and are generally nonreactive on the chemistry scale
  • Valence electrons are in the outermost shell (the highest n value) and are primarily responsible for chemical reactivity
  • It is not always the last orbital filled that constitutes the valence shell (d-block roles and penetration effects can shift what counts as valence)

• Valence vs core electrons: practical rules of thumb

  • Valence shell: the highest n value in the electron configuration; electrons in this shell are most accessible for bonding/reactivity
  • All electrons in shells with a lower n value are treated as core electrons for many reactivity considerations
  • Example: Group 2 elements (Be, Mg, Ca, Sr, …) all end in ns^2 in their neutral state, and thus have the same number of valence electrons (2) within the same group, leading to similar chemistry
  • Real-world reminder: common anecdotes (e.g., sodium in water) illustrate the highly reactive nature of elements with unstable or highly reducing valence electrons

• Noble-gas shorthand (core abbreviation)

  • Use the noble gas prior to the element of interest in brackets, then continue the remaining valence electrons outside the brackets
  • Be and Mg examples:
  • Be: 1s^2 2s^2 → [ ext{He}] 2s^2
  • Mg: 1s^2 2s^2 2p^6 3s^2 → [ ext{Ne}] 3s^2
  • Strontium (Sr, Z=38): noble gas before Sr is Kr (Z=36)
  • Sr: [ ext{Kr}] 5s^2
  • Rubidium (Rb, Z=37): [ ext{Kr}] 5s^1
  • Yttrium (Y, Z=39): [ ext{Kr}] 4d^1 5s^2
  • Antimony (Sb, Z=51): [ ext{Kr}] 4d^{10} 5s^2 5p^3
  • General rule: valence electrons for p-block elements are in the n= corresponding outer shell (e.g., 5s and 5p for Sb) and are counted as the valence electrons (five in Sb: 5s^2 + 5p^3)

• Valence electrons and reactivity: practical implications

  • For main-group elements, valence electrons largely determine bonding behavior and chemical reactivity
  • In transition metals, d-orbital occupancy influences chemical properties and can lead to special stability patterns (see chromium and copper below)
  • When comparing elements in the same group (e.g., Be, Mg, Ca, Sr), the valence electron count is the same, predicting similar chemistry and trends in reactivity
  • A historical/beverage-like analogy: as you move down a group, the outer electrons are more shielded and spatially extended, altering reactivity and bonding tendencies

• Notable exceptions: chromium and copper (d-subshell stability)

  • Chromium (Cr) and Copper (Cu) illustrate why you can’t always rely on a simple “fill according to the last shell” rule
  • Cu family vs Cr family terminology refers to the tendency of d-orbitals to favor either a half-filled or fully filled configuration for extra stability
  • Chromium in its neutral state is commonly described as: [ ext{Ar}] 3d^5 4s^1 rather than [ ext{Ar}] 3d^4 4s^2, i.e., one electron from 4s is promoted to 3d to achieve a half-filled d subshell (3d^5) and a more stable arrangement
  • Copper in its neutral state is commonly described as: [ ext{Ar}] 3d^{10} 4s^1, i.e., one electron contributes to completing the d subshell to a full set (3d^{10}) with a single 4s electron left
  • The driving idea: partially filled d subshells tend to confer enhanced stability when arranged as a half-filled or fully filled set; this can cause occasional deviations from the naive Aufbau ordering
  • When forming ions, the typical energy considerations still lead to removing electrons from the valence shell first (often 4s before 3d for many transition metals), but the presence of d-subshell stability can influence which electrons are removed in practice

• Ions: neutral configuration as a starting point

  • For an ion, first write the neutral electron configuration as a baseline
  • Then remove (for cations) or add (for anions) electrons starting from the valence shell
  • Example: Oxygen neutral configuration: 1s^2 2s^2 2p^4
  • If it gains two electrons to form oxide (O^{2-}), its valence shell (2p) becomes filled: 1s^2 2s^2 2p^6
  • In transition metals, removal generally starts with the outermost 4s electrons before 3d electrons, but the exact pattern can be influenced by d-subshell stability and oxidation state
  • Important rule of thumb: focus on the valence shell first when considering reactivity and ion formation, then address inner-shell electrons as needed

• Connecting to broader chemistry concepts

  • Electron configuration underpins many chemical properties: periodic trends, reactivity, oxidation states, and bonding tendencies
  • Understanding valence electrons helps explain why elements in the same group behave similarly and why transition metals exhibit unique chemistry
  • The concept of core vs valence electrons connects to concepts like shielding, effective nuclear charge, and atomic size (to be explored in future units)

• Practical exam-oriented takeaways

  • Be comfortable writing neutral electron configurations first, then adjust for ions
  • Identify valence electrons by locating the outermost shell with electrons (highest n value)
  • Use noble gas shorthand to simplify configurations, especially for heavier elements
  • Remember exceptions (Cr, Cu) and understand the rationale: half-filled or fully filled d subshells confer extra stability
  • For exams, you may be asked to determine valence electrons or to describe the electron arrangement that yields similar chemistry within a group
  • You may be asked to interpret how valence electron distribution affects reactivity, not just to reproduce a full configuration

• Quick reference formulas and examples

  • Subshell capacities: Ns = 2,\, Np = 6,\, N_d = 10
  • General fuller configurations (examples):
  • Be: 1s^2 2s^2 = [\text{He}] 2s^2
  • Mg: 1s^2 2s^2 2p^6 3s^2 = [\text{Ne}] 3s^2
  • Sr: [\text{Kr}] 5s^2
  • Rb: [\text{Kr}] 5s^1
  • Y: [\text{Kr}] 4d^1 5s^2
  • Sb: [\text{Kr}] 4d^{10} 5s^2 5p^3
  • Cr (neutral): [\text{Ar}] 3d^5 4s^1
  • Cu (neutral): [\text{Ar}] 3d^{10} 4s^1
  • O: 1s^2 2s^2 2p^4; O^{2-}: 1s^2 2s^2 2p^6

• Final takeaway

  • Electron configuration is a foundational tool for predicting and understanding chemical behavior
  • Practice writing neutral configurations first, then adapt for ions
  • Always identify valence electrons as the key players in reactivity
  • Be mindful of exceptions (Cr, Cu) and the underlying stability considerations for d-subshells

Note: This set of notes mirrors the main ideas and practical examples given in the transcript, including the emphasis on valence electrons, noble-gas shorthand, Hund’s rule, energy gaps, and the chromium/copper exceptions. Use these as a blueprint for solving electron-configuration problems and for connecting these ideas to chemical reactivity and periodic trends.