Since atoms are always in motion, when we talk about shapes of molecules or polyatomic ions, we mean the shape based on the average location of the atoms

the size and shape of a molecule of a particular substance, together with the strength and polarity of its bonds, largely determine the physical and chemical properties of that substance.

The shapes of different AX n molecules or ions depend on the number of electron domains surrounding the central atom. In a overwhelming majority of cases, the shapes of the molecules are related to five basic geometries:

Shapes of Ax n molecules where A is a representative element (one of the elements from the s block or p block of the periodic table) are determined by a theory called VSEPR (often read as "vesper"), which stands for Valence Shell Electron Pair Repulsion.(≈90% accuracy) VSEPR theory states the shapes of different AX n molecules or ions depends on the number of “electron domains” surrounding the central “A” atom

The Lewis structure of NH 3 has four electron domains around the central nitrogen atom. (3 bonding pairs & one lone pair)

The number of electron pairs surrounding an atom, both bonding and nonbonding, is called its steric number. Steric number = (number of lone pairs on central atom) + (number of atoms bonded to central atom)

The "AXE method" of electron counting is commonly used when applying the VSEPR theory A represents the central atom and always has an implied subscript one The X represents the number of bonds between the central atoms and outside atoms Note:Multiple covalent bonds (double,triple, etc) count as one X. E represents the number of lone electron pairs surrounding the central atom

In geometry, a tetrahedron (plural: tetrahedra) is a polyhedron composed of four triangular faces, three of which meet at each vertex. A regular tetrahedron is one in which the four triangles are regular, or "equilateral",

Draw the Lewis dot structure for the molecule and count the total number of electron densities or steric number (single bonds + multiple bonds + lone pair electrons + unpaired electrons) For molecules or ions that have resonance structures, you may use any one of the resonance structures. Determine electron-density geometry by arranging electron densities to minimize repulsions Using only bonding pairs from the electron-density geometry, determine molecular geometry

VESPR model can be extended to predict the shapes of complex molecules

While VESPR helps us to determine the ideal bond angles found in a molecule’s molecular geometry, the actual bond angles will deviate from the ideal values for the following reasons:

Triple bonds repel other bonding-electrons more strongly than double bonds. Double bonds repel other bonding- electrons more strongly than single bonds.

The VESPR model is a simple tool for predicting the shapes of molecules but does not explain why bonds exist between atoms Valence-bond theory pictures individual atoms, each with its own orbitals and electrons coming together and forming covalent bonds of the molecule.

Electron orbital notation identifies the valence orbitals

Bonding involves the overlap of valence orbitals on the central atom with those of the surrounding atoms.

sigma ( σ ) bonds are single covalent bonds. They result from the overlap of two s orbitals, an s and a p orbital, or two head-to-head p orbitals.

pi ( π ) bonds result from the sideways overlap of p orbitals, and pi orbitals are defined by the region above and below an imaginary line connecting the nuclei of the two atoms pi bonds never occur unless a sigma bond has formed first, and are always part of a double or triple bond.

Whenever there is a double bond it is made up of one sigma (direct orbital overlap) bond and one pi (lateral orbital overlap) bond.

Triple bonds have two pi bonds arranged at 90º to one another brought about by the lateral overlap of one pair of py orbitals and one pair of pz orbitals.

The experimental shapes and bond angles do not correspond to the theoretical expectation The 'p' orbitals are oriented at 90º to one another and yet there are few molecules that show a bond angle of 90º The valence electrons in the s-orbital cannot bond without interfering with the p- orbitals.

Hybridization is a model that allows us to combine the atomic orbitals and then produce four degenerate orbitals to be used for bonding.

Hybridization is the ‘mixing’ or ‘blending’ of standard atomic orbitals to form new bonding orbitals that will accommodate the spatial requirements in a molecule. The hybrid orbitals are formed by combining s, p and d orbitals

Example: methane (CH 4 ), 1 Carbon binds with 4 Hydrogens. The carbon atom itself has only 2 electrons available for bonding in the 2p subshell. In order for 4 hydrogens to bind there need to be 4 electrons available for bonding, which cannot be achieved with the current orbital arrangement.

This leads to the creation of a new ‘hybridized orbital’, called sp3. The pull of a hydrogen nucleus results in an electron being excited from the 2s subshell into the 2p subshell, where it is available for bonding.

Sigma bonds are generally formed through hybridization: Pi bonds are formed by unhybridized p-orbitals

sp 2 refers to a hybrid orbital being constructed from one s-orbital and two p-orbitals. hybridized orbitals are notated as sp, sp 2, sp 3, sp x, etc.

Valence-bond theory will correspond with the VESPR geometry