Topic 4: Intermolecular forces and metallic bonding Flashcards
Structure, giant covalent
Very hard but brittle. Very high melting point and boiling point. Do not
conduct in any state. Insoluble.
Structure, giant ionic
Hard but brittle. High melting point and boiling point. Conduct when molten or
aqueous, but not as solids.
Structure, giant metallic
Malleable, not brittle. Melting point and boiling point dependent on no. of
valence e-. Good conductivity.
Structure, molecular covalent
Usually soft and malleable unless hydrogen bonded.
Low melting point and boiling point. 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.
Bond polarity
A polarity caused by a difference in electronegativity between the
elements. The greater the difference, the greater the polarity.
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.
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.
