Chemical Bonding and Structure

Chemical Bonding and Structure

• Ionic compounds: Consist of ions held together in lattice structures by ionic bonds.
• Covalent compounds: Form by sharing electrons.
• Lewis structures: Show electron domains in the valence shell, used to predict molecular shape.
• Molecular properties: Result from forces between their molecules.
• Metallic bonds: Involve a lattice of cations with delocalized electrons.
• Bonding systems: Often require sophisticated concepts and theories.
• Hybridization: Results from mixing atomic orbitals to form hybrid orbitals with equivalent energy.
• Chemical bonds: How atoms combine to form stable associations.
• Ionic bonds form: When electrons are transferred between atoms.
• Ionic compounds: Usually solids with lattice structures and are electrically conductive in molten or aqueous forms.
• Metals form cations: By losing valence electrons
• Non-metals form anions: By gaining electrons
• Electron configurations: Determine the number of electrons lost or gained during reactions.

Ionic bonding and structure

• Ions: Charged particles formed by loss or gain of electrons.
• Ionic bond: Electrostatic attraction between oppositely charged ions.
• Ionic compounds properties:
  • Mostly solids
  • High melting and boiling points.
  • Often soluble in water
  • Conduct electricity when molten or dissolved.
• Polyatomic ions: Groups of atoms acting as a single ion (e.g., (NH_4^+), (OH^-), (NO_3^-), (HCO_3^-), (CO_3^{2-}), (SO_4^{2-}), (PO_4^{3-})).
• Predicting ionic charge: Determined by the number of electrons lost or gained from the group number on the Periodic Table.

Covalent Bonding

• Covalent bond: Electrostatic attraction between a shared pair of electrons and the positively charged nuclei of atoms.
• Types of covalent bonds:
  • Single bond: One shared pair of electrons.
  • Double bond: Two shared pairs of electrons.
  • Triple bond: Three shared pairs of electrons.
• Bond length: Decreases as the number of shared electrons increases.
• Bond strength: Increases as the number of electrons shared increases.
• Polarity: Results from the difference in electronegativity of the bonded atoms.

Covalent Structures

• Lewis diagrams: Show the arrangement of electrons in covalent molecules.
• Octet rule Atoms tend to gain a valence shell with eight electrons
• Exceptions to octet rule: Certain elements (e.g., Be, B) form stable compounds with fewer than eight electrons in their valence shell.

Polar bonds

• Polar covalent bonds: Result from unequal sharing of electrons due to differences in electronegativity.
• Dipole moments occur: When bonds are polar, and are not symmetrically arranged.
• Polar molecules: When the bond dipoles in a molecule do not cancel out.
• Non-polar molecules: When the bond dipoles in a molecule cancel out.

Intermolecular Forces

• London dispersion forces: Weak forces between all molecules, caused by temporary dipoles.
• Dipole-dipole forces: Attractive forces between polar molecules.
• Hydrogen bonds: Strong forces between molecules containing H bonded to F, O or N. These are especially strong intermolecular forces.
• Strength comparison: London dispersion < Dipole-dipole < Hydrogen bonds.

Metallic bonding

• Metallic bond: Electrostatic attraction between positively charged metal ions and delocalized valence electrons.
• Properties of metals:
  • High electrical conductivity due to delocalized electrons.
  • High thermal conductivity due to delocalized electrons.
  • Malleability and ductility due to non-directional metallic bonding.
  • High melting points due to strong metallic bonds.

Alloys

• Alloys: Mixtures of two or more metals, often with enhanced properties.
• Examples of alloys:
  • Steel
  • Brass
  • Bronze
  • Duralumin
• Properties of alloys:
  • Increased strength and durability
  • Improved corrosion resistance
  • Altered melting points