Topic 4 - Chemical Bonding & Structure Flashcards
Giant Covalent Structure
Very hard but brittle.
Very high m.p. and b.p.
Do not conduct in any state.
Insoluble.
Giant Ionic Lattice
Hard but brittle.
High m.p. and b.p.
Conduct when molten or aqueous, but not as solids.
Soluble
Giant Metallic Structure
Malleable (not brittle)
M.p. and b.p. dependent on no. of valence e-.
Good conductivity.
Insoluble, (some metals react with water)
Molecular Lattice
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
Different crystalline forms of the same element - Carbon
Bond polarity
Polarity = difference in electronegativity between elements.
The greater the difference, the greater the polarity.
Pi bond (π)
Formed by the sideways overlap of p orbitals
Electron densities = above and below a line drawn through the two nuclei.
Double bonds = one π bond
Triple bonds = two π bonds (perpendicular to each other)
Sigma bond (σ)
Formed by overlap of atomic orbitals from two different atoms along the line drawn through the two nuclei
Electron densities concentrated along the line.
Single, double, triple bonds = one σ bond.
Covalent bond
Bonding by the sharing of electrons.
The electrons are shared and attracted by both nuclei resulting in a bi-directional bond two atoms.
Dative bond (co-ordinate bond)
A bond in which both electrons come from one of the atoms.
Ionic bond
A bond where electrons are transferred from one atom to another - form ions with complete outer shells.
Ionic compound + and – ions = attracted (to each other) by electrostatic force between them
Build up into a strong lattice.
Have high m.p.
Ionic bonds occur between elements with a great difference (>1.8) in electronegativity
Van der Waal’s / London Dispersion forces
Temporary dipole forces due to momentary unevenness in spread of electrons.
Weakest of intermolecular forces.
Conductivity
The extent to which a substance can conduct electricity. Must possess electrons / ions = free to move.
Delocalization
The sharing of one electron pair by more than two atoms.
Dipole-Dipole
Permanent electrostatic forces of attraction between polar molecules.
Stronger than van der Waals’
Weaker than H-Bonding
Hydrogen bonding
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 & van der Waals’
Hybridization
The mixing of atomic orbitals to create new orbitals of the same energy.
Metallic bonding
Valence electrons in metals become detached from the individual atoms = metals consist of a closely packed lattice of + ions in a ‘sea’ of delocalized electrons.
Forces of attraction are between ions + electrons (not between the ions themselves) = metals are malleable and ductile.
Molecular polarity
Depends on both the bond polarity and the symmetry.
Resonance Structures
Structures that arise from the possibility to draw multiple bonds in different positions equivalently.
(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.
