Chemical Bonds and Atomic Structure - Quick Reference

Atomic Identity and Isotopes

  • Element: substance composed of only one kind of atom.
  • Atomic number $Z$: number of protons; identifies the element (e.g., Hydrogen $Z=1$).
  • In a neutral atom, number of electrons $E$ equals $Z$: E=Z.E = Z.
  • Mass number $A$: A=Z+NA = Z + N where $N$ is the number of neutrons.
  • Atomic weight (on the table): average of mass numbers across isotopes; varies due to isotope distribution.
  • Isotopes: same $Z$ (same element) but different $N$ and $A$; example: hydrogen isotopes have different $A$ but are all hydrogen.

Atomic Structure and Subatomic Particles

  • Nucleus contains protons (positive, mass ~1 Da) and neutrons (neutral, mass ~1 Da).
  • Electrons: negative, very small mass; orbit around the nucleus in electron shells.
  • A neutral atom has equal numbers of protons and electrons; charges cancel.
  • Elements’ chemical/physical properties depend on protons and electrons distribution.

Electron Configuration and Valence

  • Electron shells around the nucleus:
    • First (innermost) shell capacity: up to 2 electrons.
    • Outer shells (n ≥ 2) capacity: up to 8 electrons each (per shell).
  • Orbitals: regions where electrons are found; each orbital holds a maximum of 2 electrons.
  • Valence shell: outermost shell that contains electrons.
  • Reactivity is driven by unpaired electrons in the valence shell.
  • A full valence shell (octet) tends to be stable and nonreactive (simplified Octet Rule).

Practical Periodic Table Notes

  • Periodic table orientation: atomic number $Z$ is shown above the symbol; atomic weight is shown below.
  • Mass number $A$ is not directly listed on the standard table; it varies with isotopes.
  • For example, oxygen has $Z=8$; a neutral oxygen atom has $E=8$ electrons.

Covalent Bonding and Molecular Structure

  • Chemical bond: an attractive force that links atoms into molecules; atoms may lose, gain, or share electrons.
  • Covalent bonds: atoms share electrons so that their valence shells are filled; often high bond energy.
  • Examples:
    • H–H forms a single covalent bond (H$_2$).
    • CH$_4$ (methane): carbon forms four covalent bonds with four hydrogens, achieving an octet.
  • Multiple bonds: double bonds (two shared electron pairs) and triple bonds (three shared electron pairs).
  • Bond energy: energy required to break a bond; higher bond energy means a stronger bond.

Electronegativity and Bond Character

  • Electronegativity: the nucleus’s attractivity for electrons in a bond; higher value means stronger pull on electrons.
  • Example values:
    • Oxygen: EN(O)≈3.5\text{EN}(O) \approx 3.5
    • Sodium: EN(Na)≈0.9\text{EN}(Na) \approx 0.9
  • Tug-of-war view: in a bond, electrons may be pulled more toward the more electronegative atom.
  • Result: bonds can be polar covalent if there is a significant electronegativity difference.
  • Even highly electronegative atoms can bond with less electronegative ones; bonds may be polar rather than purely covalent.

Quick Concepts to Remember

  • $Z$ identifies the element; $E = Z$ in a neutral atom.
  • $A = Z + N$; isotopes differ by $N$ but share $Z$.
  • Valence electrons determine reactivity; unpaired electrons in the valence shell drive bonding.
  • Covalent bonding involves sharing electrons to fill valence shells (octet rule).
  • Electronegativity differences influence bond polarity and bond type.
  • Practice: use the periodic table to draw Lewis structures and predict bonding and reactivity.

Study Tips

  • Draw Lewis structures to visualize valence electrons and bonding.
  • Create quick tables to compare bond types, bond energies, and electronegativity differences.
  • Use practice problems (e.g., identify bond type for given element pairs) for last-minute prep.