payment page
-
Bonding electrons
Electrons shared between two atoms in a covalent bond.
-
Bond order
The number of electron pairs involved in a covalent bond. In terms of Molecular Orbital theory, it is the number of bonding electron pairs minus the number of antibonding electron pairs. Greater bond order means greater bond strength and shorter bond length.
-
Carbocation
The phenomenon in which carbon, usually tetravalent, can form a trivalent species with a positive charge. Such a carbon atom is extremely unstable due to its lack of an octet, but its reactivity makes it a source of a great deal of fascinating chemistry that will be discussed in later sections.
-
Covalent bond
Interaction between atoms held together by the sharing of electrons.
-
Curved-arrow formalism
A method of keeping track of the movements of pairs of electrons during chemical reactions. Uses double-headed arrows to denote movement of electrons from source to destination.
-
Delocalization
The phenomenon in which electrons in some molecules are not fixed to specific atoms or bonds but are spread out over several atoms or bonds. Delocalization is an energetically favorable process: by distributing charge over a greater volume, the net energy of the molecule is lowered, resulting in resonance stabilization.
-
Dipolar
The presence, in a bond or molecule, of a positive end and a negative end.
-
Dipole moment
The measure of the polarity of a molecule. Higher dipole moments are accorded to more polar molecules. Not all molecules with dipolar bonds have dipole moments: the dipole moment is dependent on both orientation and magnitude of dipolar forces; it is possible for the dipolar forces of a molecule to cancel each other out, resulting in a molecule with no significant dipole moment.
-
Electronegativity
The relative tendency for an atom to attract electrons to itself. Measured on an arbitrary scale of 4.0, with fluorine being the most electronegative element. Electronegativity increases from left to right across the periodic table and decreases as you move down a group.
-
Formal charge
An accounting scheme that estimates the charge on an atom. The formal charge is calculated by subtracting the number of lone electrons and half the number of bonded electrons from the group number.
-
Lewis Dot structure
A common way to represent molecules, it uses lines to depict bonded electron pairs and dots to represent lone pairs. Inner-shell electrons are not shown.
-
Lone pair
Electrons in the valence shell of an atom that don't participate in bonding.
-
Molecule
A collection of atoms held together by covalent bonds.
-
Octet rule
The cardinal rule of bonding. The octet rule states that atoms gain stability when they have a full complement of 8 electrons in their valence shells.
-
Polar covalent bond
A covalent bond between atoms of differing electronegativities such that one atom has a partial positive charge and the other has a partial negative charge.
-
Resonance hybrid
The weighted average of several resonance structures that gives a composite view of the electronic structure of a molecule.
-
Resonance stabilization
Because resonance allows for delocalization, in which the overall energy of a molecule is lowered since its electrons occupy a greater volume, molecules that experience resonance are more stable than those that do not. These molecules are termed resonance stabilized.
-
Resonance structure
One of several Lewis structures that can be drawn for a given atomic connectivity. Each resonance structure contributes an aspect of the resonance hybrid.
-
Valence
The number of bonds an atom typically forms. Carbon is tetravalent, nitrogen is trivalent, oxygen is divalent, and hydrogen/halogens are monovalent.
-
Valence electron
The electrons in the outermost energy shell of an atom. The configuration of these electrons determine the chemical properties of the element.
-
Valence shell
The highest energy shell in an atom. All interactions between atoms take place through the electrons of the valence shell.
