Periodicity and Ionic Bonding Lecture Notes

Structure and Bonding

  • Structure + Bonding = Arrangement in space + Forces holding atoms together.

Types of Bonding

  • Ionic: Electron transfer.
  • Covalent: Electron sharing.
  • Metallic: Mobile valence electrons.
  • Most materials exhibit a blend of these bonding types.

Valence Electrons

  • Valence electrons are outer shell electrons that participate in chemical bonding.
  • Core electrons are all non-valence electrons.
  • Valence electrons determine chemical properties.
  • Main group elements' valence electrons = "A" group number.

Atomic and Ionic Sizes

  • Determined by electronic structure and nucleus-electron interactions.
  • Electrons (- charge) outside, Protons (+ charge) inside nucleus.
  • Electrostatic principles govern interactions: Opposites attract, likes repel; greater charge = greater force; shorter distance = stronger force.

Effective Nuclear Charge (Zeff)

  • ZeffZ_{eff} is the net positive charge experienced by an electron.
  • ZeffZ_{eff} < (number of protons) due to shielding by core electrons.
  • Zeff=Z−SZ_{eff} = Z - S where S is the shielding constant.

Trends in Atomic Radius

  • Decreases from left to right across a period due to increasing ZeffZ_{eff}.
  • Increases down a group due to increasing principle quantum number (n).

Ionic Radius

  • Cations are smaller than their neutral atoms.
  • Anions are larger than their neutral atoms due to electron-electron repulsion.

Isoelectronic Species

  • Isoelectronic species have the same number of electrons.
  • For isoelectronic ions, radius decreases with increasing number of protons.

Ionization Energy (IE)

  • Energy required to remove an electron from a gaseous atom: Na(g)→Na+(g)+e−Na(g) \rightarrow Na^+(g) + e^-
  • Decreases as atoms get larger (smaller ZeffZ_{eff}).
  • IE2 > IE1 always because it's harder to remove subsequent electrons.

Exceptions to IE Trend

  • New subshells: IE1(Be) > IE1(B) because 2s is lower energy than 2p.
  • Paired electrons: IE1(N) > IE1(O) because paired electrons in O repel each other.

Electron Affinity (EA)

  • Energy change when an electron is added to a gaseous atom: Cl(g)+e−→Cl−(g)Cl(g) + e^- \rightarrow Cl^-(g)
  • Negative EA: exothermic (energy released).
  • Positive EA: endothermic (energy required).

Periodic Trend Summary

  • Atomic radius: Increases down, decreases across.
  • Ionization Energy: Decreases down, increases across.
  • Electron Affinity: Becomes less exothermic down, more exothermic across.

Ionic Bonding

  • Metal atoms transfer electrons to nonmetal atoms, forming ions.
  • Ions attract due to opposite charges, forming ionic bonds.
  • Atoms gain/lose electrons to achieve noble gas configurations.

Ionic Lattice

  • Ions form 3D lattice structures, not individual pairs.
  • Chemical formula indicates simplest ion ratio.
  • Held together by electrostatic attractions.

Lattice Energy

  • Energy required to separate one mole of solid into gaseous ions: MX(s)→Mn+(g)+Xn−(g)MX(s) \rightarrow M^{n+}(g) + X^{n-}(g)
  • Decreases as ion size increases.
  • Increases as ion charge increases.
  • E=k(q1q2)/dE = k(q1q2)/d