Atomic Structure and Valence Summary

Dalton's Atomic Theory

• Elements are composed of very small indivisible particles called atoms; atoms retain identity through chemical reactions.
• Atoms of the same element are alike; atoms of different elements are different.
• Atoms enter into chemical reactions; reactions involve building up or breaking down atoms.
• Atoms are permanent and cannot be decomposed (as stated by Dalton).
• Historical note: Democritus (400 B.C.) suggested matter is made of indivisible particles called atoms.

Modern Atomic Theory

• Isotopes exist: atoms of the same element with different masses.

• Not all atoms of the same element have identical masses; atoms can be decomposed by radiation.

• Structure: nucleus (protons and neutrons) with a diffused electron cloud surrounding it.

• Subatomic particles:

  • Electron: charge $-1$, mass ≈ $\frac{1}{1850}$ of a proton.
  • Proton: charge $+1$, mass ≈ 1 (in proton units).
  • Neutron: charge $0$, mass ≈ 1.

• The nucleus provides the positive charge; the electron cloud provides negative charge; atom is neutral when electrons = protons.

• The fundamental difference between elements arises from the positive charge of the nucleus (atomic number, $Z$).

• Key historical milestones:

  • Rutherford: nucleus exists.
  • Bohr (1913): electrons move in orbits around nucleus with fixed energy levels.
  • Moseley: measured the numerical value of nuclear positive charge via x-ray studies.

Electronic Structure of the Atom

• Electrons occupy orbits around the nucleus, later described as orbitals.
• Shells (groups) and energy levels governed by the principal quantum number $n$ (values: 1, 2, 3, …).
• Shell labels (designation): $n=1,2,3,4,\dots$ correspond to K, L, M, N, O, P, Q, …
• Electrons are characterized by a principal quantum number $n$; energy increases with increasing $n$.
• Each shell is divided into sublevels (subshells): S, P, D, F, etc.
• In a given shell, the number of sublevels equals $n$.
• Orbital capacities:
  • s-sublevel: 1 orbital → max $2$ electrons
  • p-sublevel: 3 orbitals → max $6$ electrons
  • d-sublevel: 5 orbitals → max $10$ electrons
  • f-sublevel: 7 orbitals → max $14$ electrons
• The total maximum electrons in a shell: $2n^2$.
• Example maxima:
  • K-shell ($n=1$): $2$ electrons
  • L-shell ($n=2$): $8$ electrons
• Electrons fill the lowest-energy orbitals first; orbitals within the same sublevel have the same energy, but different orientations.

Orbital Energies and Occupation

• The order of filling follows increasing energy: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p
• Note: 4s fills before 3d because 4s energy is lower than 3d; 3d fills after 4s but before 4p, etc.
• Example electron configurations:
  • Calcium (Z=20): $1s^2\ 2s^2\ 2p^6\ 3s^2\ 3p^6\ 4s^2$
  • Potassium (Z=19): $1s^2\ 2s^2\ 2p^6\ 3s^2\ 3p^6\ 4s^1$
• Alternate shell notation: Mg (A=24, Z=12) can be written as Mg (2, 8, 2) to denote electrons in shells K–L–M.

Atomic Diagrams and Notation

• Atomic number (Z): number of protons (and electrons in neutral atom).
• Atomic weight (A): approximately the sum of protons and neutrons.
• Example Mg: A=24, Z=12 → neutral Mg has 12 electrons.
• Electron distribution for Mg: 1s^2, 2s^2, 2p^6, 3s^2 (often written as 1s^2 2s^2 2p^6 3s^2).
• Example Aluminum: A=27, Z=13 → electrons: 13; distribution: 1s^2, 2s^2, 2p^6, 3s^2, 3p^1.

Valence, Ionic Charge, and Reactivity

• Valence electrons: electrons in the outermost shell (valence shell).
• Outer shell capacity (except the first shell) is typically eight for stability (octet rule).
• Atoms tend to complete their outermost shell by gaining, losing, or sharing electrons, forming ions or bonds.
• Ions: charged atoms formed by gain or loss of electrons.
  • Cations: atoms that lose electrons; positive valence. Example: Magnesium loses 2 electrons → valence +2.
  • Anions: atoms that gain electrons; negative valence. Example: Chlorine gains 1 electron → valence -1.
• Noble gases (inert elements): complete outer shells; generally have no valence in reactions.
• Metalloids (amphoteric elements): have four electrons in outer shell; variable behavior in reactions.
• Transition elements: elements with electrons transitioning between shells (e.g., 3rd to 4th); exhibit variable valences (examples: Fe, Co, Ni, Cr, Mn).

Quick reference formulas and notes

• Maximum electrons in a shell: $2n^2$
• Subshell capacities: $s:2,\, p:6,\, d:10,\, f:14$
• Outer shell stability: typically 8 electrons (octet rule), except the innermost K-shell can have fewer; inner shells sequentially can hold up to $2n^2$ with inner shells filling before outer shells as energy dictates.
• Order of filling (example sequence): 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
• Example notations for electron configurations are often written as: 1s^2 2s^2 2p^6 3s^2 (calcium-like), or as shell groupings like (2, 8, 2).