Book: Ch. 10

Shows symbols for atoms, the bonding electron pairs as lines, and the lone electron pairs that fill each atom’s outer level as pairs of dots.

Indicated with two-headed arrow (↔), molecules with two Lewis structures, that have the same relative placement of atoms but different locations of bonding and lone electron pairs.

Something each atom has. The charge it would have if the bonding electrons were shared equally. Formal charge of atom = no. of valence e- - no. of unshared valence e- - 1/2 no. of shared valence e

In gaseous molecules containing either beryllium or boron as the central atom, that atom is often electron deficient: it has fewer than eight electrons around it (an incomplete octet).

Species that contain a lone (unpaired) electron, which makes them paramagnetic.

Occur only with nonmetals from period 3 or higher because they have d orbitals available. A central atom expands its valence shell to form more bonds.

A basic principle that says that, in order to minimize repulsions, each group of valence electrons around a central atom is located as far as possible from the others.

The three-dimensional arrangement of nuclei joined by the bonding groups.

The angle formed by the bonds joining the nuclei of two surrounding atoms to the nucleus of the central atom, which is at the vertex.

Two electron groups attached to a central atom point in opposite directions; they have a linear arrangement. This results in a molecule of linear shape with a bond angle of 180º.

Given by bond angles of 120º and three electron groups around a central atom pointing to the corners of an equilateral triangle.

Molecular shape given by three electron groups in a trigonal plane with a lone pair at any one of the triangle’s corners. Bond angle then goes to considerably less than 120º, because the lone pair repels the other two bonds far more.

All molecules or ions with four electron groups around a central atom adopt the tetrahedral arrangement. Bond angle of AX₄ is 109.5º.

AX₃E. Tetrahedral electron shape with one lone pair. Gives a bond angle of 107.3º.

Five mutually repelling electron groups gives this. You have two different positions for electron groups and two ideal bond angles: 120º and 90º.

For a trigonal bipyramidal arrangement, there are three equatorial groups that lie in a trigonal plane; these have 120º bond angles. The other two are the axial groups, which lie above and below this plane with a 90º angle separating the axial from the equatorial plane.

Occurs for AX₄E, where there’s a trigonal bipyramidal shape with one lone pair. Bond angles are 101.5º and 86.8º.

AX₃E₂. Trigonal bipyramidal shape with two lone pairs. Bond angles are 86.2º.

AX₆. Gives 90º bond angles between six electron groups.

AX₅E. Reduces the bond angle to 81.9º.

AX₄E₂. You get two opposing lone pair electron groups and a plane of bonds with angles of 90º to each other.

Determined by shape and bond polarity, this refers to an uneven distribution of charge over the whole molecule or a large portion of it.

A measure of molecular polarity in units of debye (D), which is given by: 1 D = 3.34×10^-30 Cm