Molecular Shapes

Two ways atoms increase stability. 1. Sharing electrons- covalent bond. 2. Transferring electrons- ionic bond. Lewis diagrams– valence electrons represented by dots.

Types of Covalent Bonds Single bond-shares 2 electrons. Ex. H2, F2 Double bond-shares 4 electrons. Ex. O2 Triple bond-shares 6 electrons. Ex. N2

Molecular shape describes the arrangement of atoms not of the electron pairs. Shared pair of electrons: electrons that are between two atoms. Unshared pair of electrons: electrons that are around the atom.

Rules of Dot Diagrams for Molecules Count up the total valence electrons. 1. Sum valence electrons of atoms present. 2. For polyatomic anion, add one electron for each negative charge. 3. For polyatomic cation, subtract one electron for each positive charge.

Draw skeleton structure, joining atoms by single bonds. Central atom is usually written first. Hydrogen, oxygen, and halogens are terminal. In oxyacids, H is always bonded to O which is bonded to nonmetal.

Deduct two valence electrons for each single bond written. Distribute remaining electrons as unshared pairs to give each atom eight electrons. If there are too many electrons placed, look for double or triple bonds. If there are extra electrons, placed as unshared pairs around the central atom.

Exceptions Extended octets. Valence shell can hold more than 8 electrons. Total number of electrons are odd. Molecule has fewer than eight electrons. C or N never exceed octet.

Coordinate covalent bond One atom furnishes both of the shared electrons in the bond. There are donors and acceptors.

Resonance Resonance structures are two or more valid electron dot structures that can be written for a molecule.

VSEPR Valence shell electron pair repulsion theory. In a small molecule, the pairs of valence electrons are arranged as far apart from each other as possible. Shape of the molecule is determined by the electrons around the CENTRAL ATOM.

Total electron pairs Shared pairs Unshared pairs Shape Example 1 Linear H2 2 CS2 3 Trigonal planar BCl3 Bent NO2-

4 Tetrahedral CH4 3 1 Trigonal pyramid NH3 2 Bent H2O 5 Trigonal bipyramidal PCl5 Unsym. Tetrahedron SF4

5 3 2 T-shaped ClF3 Linear I3- 6 Octahedral SF6 1 Sq. pyramid BrF5 4 Sq. planar XeF4

Shapes In double or triple bonds, all electron pairs must stay together between two atoms. We can treat double or triple bonds just as we do single bonds when predicting molecular shape. Count them as ONE shared pair when predicting shape.

                                  

Hybridization Hybrid orbital is an orbital of electrons in a bond which is a combination of the shapes and properties of the original atomic orbital. Rearrangement of electrons within a valence orbitals of atoms during a chemical reaction. S & p orbitals from the same electron shell can form hybrids.

Types of Hybrids sp hybrid Starts as s2. One electron moves up to a p orbital. s + p = sp Shape is linear. sp2 hybrid Starts as a s2p. 3 hybrid orbitals formed by combining an s and 2 p’s from the same electron shell. s2 + p = sp2 Shape is trigonal planar. sp3 hybrid Starts as a s2p2. 4 hybrids from the same electron shell. Very common. s2 + p2 = sp3 The shape is tetrahedron.

Other types sp3d trigonal bipyramidal sp3d2 octahedral

Polarity Polarity: an uneven sharing of electrons. Electronegativity: measure of the ability of an atom of an element that is chemically combined with another atom to attract electrons to itself. Differences in electronegativity predicts type of bond.

Small difference (less than 1.7), the bond is covalent. Large difference(greater than 1.7), the bond is ionic. If the difference in electronegativity is not zero, the bond is polar. Dipole:the center of positive and negative charge is not the same.

The shape of a molecule and the polarity of its bonds together determine whether a molecule is polar or nonpolar. Even though a bond maybe polar, the shape plays a role. Symmetrical distribution of polar bonds cancel effect. If there is an even distribution of polar bonds, the molecule is nonpolar.