2.2.1, 2.2.2- Electron Structure, Bonding & Structure

Cards (39)

  • Atomic Orbitals
    A region around the nucleus that can hold up to two electrons with opposite spins.
  • s- orbitals: spherical shaped
  • p- orbitals: dumbbell shaped
  • No. of electron pairs in orbitals
    s : 1
    p : 3
    d : 5
    f : 7
  • Ionic bonding
    Electrostatic attraction between positive and negative ions
  • Giant ionic lattices: Oppositely charged ions strongly attracted in all directions
  • High melting and boiling points
    Requires more energy to overcome strong forces of electrostatic attraction.
  • Electrical conductivity
    Solid: No free mobile electrons; ions held in place by strong electrostatic forces of attraction
    Molten: Ions are free to move; free mobile charge carriers.
    Aqueous: Ions disassociate in water; free mobile charge carriers.
  • Covalent bonds
    Strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atom.
  • Single covalent bonding: 2 atoms contribute to and share a pair of electrons
  • Multiple covalent bonding: Atoms contribute and share multiple pairs of electrons between them
  • Dative covalent (coordinate) bonding: A bond in which one atom singly shares all electrons with another atom
  • Average bond Enthalphy: The mean energy required to break 1 mole of bonds in gaseous molecules.
    • A measurement of covalent bond strength
  • Bond angles
    Trigonal planar: 120
    Tetrahedral: 109.5
    Trigonal bipyramidal:
    • Axial: 90
    • Equatorial: 120
    Octahedral: 90
  • Electronegativity
    The ability of an atom to attract the bonding electrons in a covalent bond.
  • Polar bond and permanent dipole
    Within molecules containing covalently-bonded atoms with different electronegativities
  • Polar molecules require polar bonds with dipoles that do not cancel due to their direction.
  • Permanent and induced dipole-dipole forces can both be referred to as Van Der Waal‘s forces.
  • Induced dipole-dipole forces = London Dispersion Forces
  • Hydrogen bonding
    Intermolecular bonding between molecules containing N. O or F and the H atom of -NH, -OH, or HF.
  • Electronegativity
    Ability of an atom to attract bonding electrons in a covalent bond.
    • Increases to F in the periodic table
  • Electron Pair Repulsion Theory
    Pairs of electrons around a nucleus repel each other so the shape that a molecule adopts has the electrons positioned as far apart as possible.
    • Strength of Repulsion
    • LP - LP -> LP - BP -> BP - BP
  • Orbitals are filled in order of increasing energy, with orbitals of the same energy occupied singly before pairing.
  • Energy level
    The shell that an electron is in.
  • Shell
    The orbit that an orbital is in around the nucleus of an atom.
  • Permanent Dipole
    A permanent uneven distribution of charge
  • Polar bond
    A covalent bond that has a permanent dipole due to the different electronegativities of the atoms that make up the bond.
  • Non-linear (bent)
    Central atom has 2 bonding pairs and 2 lone pairs
  • Octahedral
    Central atom has 6 bonding pairs.
  • London (Dispersion) Forces
    Induced dipole-dipole interactions caused when the random movement of electrons creates a temporary dipole in one molecule inducing a dipole in a neighbouring molecule.
  • Macroscopic Properties
    Properties of a bulk material rather than the individual; atoms/molecules that make up the material
  • Polar Molecule
    A molecule that contains polar bonds with dipoles that do not cancel out due to their direction (asymmetrical)
  • Pyramidal
    Central atom has 3 bonding pairs and 1 lone pair
  • Simple Molecular Lattice
    Solid structure made of covalently bonded molecules attracted by intermolecular force.
  • Tetrahedral
    Central atom has 4 bonding pairs
  • Trigonal bipyramidal
    Central atom has 5 bonding pairs.
  • Trigonal Planar
    Central atom has 3 bonding pairs
  • Strength of metallic bonds
    • directly proportional to the number of delocalised electrons
    • Inversely proportional to the radius of the atoms.
  • Boiling points are based on strength of van der Waals’ interactions.
    • Affected by the number of electrons
    Greater opportunity for more extensive temporary dipoles.