Structure, giant covalent: Very hard but
brittle. Very high m.p. and b.p. Do not conduct in any state. Insoluble.
Structure, giant ionic: Hard but
brittle. High m.p. and b.p. Conduct when molten or aqueous, but not as solids.
Structure, giant metallic: Malleable, not
brittle. M.p. and b.p. dependent on no. of valence e-. Good conductivity.
Structure, molecular covalent: Usually soft and
malleable unless hydrogen bonded. Low m.p. and b.p. Do not conduct in any
state. Often soluble in non-aqueous solvents, unless they can hydrogen bond to
water.
Allotropes: Occur when an element can exist in different
crystalline forms, such as in carbon, which can exist as graphite, fullerene
and diamond.
Diamond is exceptionally hard because there is
no plane of weakness in the molecule made up of sp3 hybridized
carbon atoms. In graphite, the carbon atoms are sp2 hybridized.
Remaining electrons after the three σ bonds, are delocalized, resulting in the
fact that graphite is a good conductor of electricity.
Bond polarity: A polarity caused by a difference in
electronegativity between the elements. The greater the difference, the greater
the polarity.
Bond, π: Pi bond. A bond formed by
the sideways overlap of p orbitals with electron densities concentrated above
and below a line drawn through the two nuclei. Double bonds have one π bond,
while triple bonds have two which are perpendicular to each other.
Bond, σ: Sigma bond. A bond formed by
the head on overlap of atomic orbitals from two different atoms along the line
drawn through the two nuclei, with electron densities concentrated along the
line. Single, double and triple bonds have one σ bond.
Covalent bond: Bonding by the
sharing of electrons. The electrons are shared and attracted by both nuclei
resulting in a directional bond between the two atoms.
Dative bond: A bond in which both electrons come from one of
the atoms. Also known as coordinate bond.
Ionic bond: A bond by which electrons are transferred from
one atom to another to form ions with complete outer shells.
In an ionic compound the + and – ions are
attracted to each other by the electrostatic force between them, and build up
into a strong lattice. Have relatively high m.p. Ionic bonds occur between
elements with a great difference (>1.8) in electronegativity.
Conductivity: The extent to which a substance can conduct
electricity. Must possess electrons or ions that are free to move.
Delocalization: The sharing of
one electron pair by more than two atoms.
Forces, dipole-dipole: Permanent electrostatic
forces of attraction between polar molecules. Stronger than van der Waals’.
Forces, Hydrogen bonding: Occurs when
hydrogen attached to a highly electronegative element (N, F, or O) is bonded to
another highly electronegative element (N, F, or O). Stronger than
dipole:dipole forces.
Forces, van der Waal’s: Temporary dipole
forces due to momentary unevenness in spread of electrons. Weakest of
intermolecular forces. Increase with increasing molar mass.
Hybridization: The mixing of
atomic orbitals to create new orbitals of the same energy.
Metallic bonding: The valence
electrons in metals become detached from the individual atoms so that the
metals consist of a closely packed lattice of + ions in a ‘sea’ of delocalized
electrons. Forces of attraction are between ions and electrons and not between
the ions themselves, which means that metals are malleable and ductile.
Molecular polarity: Depends on both
the bond polarity and the symmetry.
Resonance hybrid: Structures that
arise from the possibility to draw a multiple bond in different positions
equivalently. Can be better explained by delocalization.
Solubility: The extent to which one substance dissolves in
another.
VSEPR theory: Valence Shell
Electron Pair Repulsion theory. States that pairs of electrons arrange
themselves around the central atom so that they are as far apart from each
other as possible. Greater repulsion between lone pair of electrons than bonded
pairs.
Information for this post came directly from http://liakatas.org/chemblog/?page_id=17
