1. Chemical bonding II UNIT 8
CHEM1:PS1:MATTER & ITS INTERACTIONS: Use Lewis Dot structures and electronegativity differences to predict the polarities of simple molecules (linear, bent, trigonal planar, trigonal pyramidal, tetrahedral). Construct an argument to explain how electronegativity affects the polarity of basic chemical molecules.

3. Vsepr theory: the five basic shapes
Valence shell electron pair repulsion theory states that electron groups – lone pairs, single bonds, multiple bonds & single electrons – repel one another through coulombic forces.
Repulsions between electron groups on interior atoms of a molecule determine the geometry of the molecule.
The preferred geometry of a molecule is the one in which the electron groups have the maximum separation (minimum repulsion) possible.
For molecules having just one interior atom, molecular geometry depends on 1.) the number of electron groups around the central atom & 2.) how many of those electron groups are bonding groups and how many are lone pairs.

5. Two electron groups: linear geometry
Consider BeCl2: two electron groups about a central atom. Due to repulsion between electron groups and to maximize separation, this molecule has a bond angle of 180° and is said to have linear geometry.
Molecules that form only two single
bonds , with no lone pairs, are rare
because they do no follow the octet
rule.
However, all molecules with two
electron groups around a central atom
with no lone pairs share linear
geometries.

6. Three electron groups: trigonal planar geometry
Consider BF3 – has an incomplete octet and three electron groups around a central atom. To maximize their separation, this molecule assumes a 120° bond angle in a single plane and trigonal planar geometry.
Formaldehyde, CH2O, is also a trigonal planar
molecule. However, its bond angles deviate
slightly from the expected 120° because there
is a double bond between the oxygen and the
carbon.
The two HCO bond angles measure 121.9° while
the HCH bond angle measures 116.2°. Why? The
double bond has higher electron density therefore
greater repulsive force.

7. Four electron groups: tetrahedral geometry
Molecules with 4 or more electron groups around a central atom are three- dimensional.
Tetrahedral molecules have bond angles of 109.5°. A tetrahedron is a geometrical shape with four identical faces – each an equilateral triangle.
Methane is an example, CH4. Written on
paper, it appears that the molecule should
be square planar with 90° bond angles but
because the molecule is in 3-D space, the
electron pairs can get even further away
from each other. Lowering repulsive
energy and stabilizing the molecule.

8. Five electron groups: Trigonal Bipyramidal geometry
Occurs when 5 electron groups are around a central
atom.
Three groups lie in a single plane like trigonal planar.
The other two are positioned above and below this
plane.
The angles between the equatorial positions (the 3
Bonds in the trigonal plane) are 120°.
The angle between the axial positions (the 2 bonds on
Either side of the trigonal plane) and the trigonal plane
Is 90°. Example: PCl5

9. Six electron groups: Octahedral geometry
Six electron groups around a central atom assume an octahedral geometry.
This structure is named after the eight-sided geometrical shape the octahedron – where four of the groups lie in a single plane, with a 5th group above the plane and a 6th below it.
All angles in this geometry
measure 90° - this is a highly
symmetrical molecule.
Sulfur hexafluoride, SF6 is an
example of this molecular
geometry.

10. Vsepr theory: the effect of Lone pairs
Lone pairs repel electron groups.
Lone pair electrons typically exert slightly greater repulsions than bonding electrons.

12. Representing molecular geometries on paper

13. Predicting the shapes of larger molecules
When molecules have more than one interior angle, apply the same principles of molecular geometry to each interior atom.

15. Molecular shape & polarity

16. Valence Bond Theory: Orbital overlap as a chemical bond
The valence electrons of the atoms in a molecule reside in quantum-mechanical atomic orbitals. Which can be s, p, d, or f orbitals or they may be hybrid combinations of these.
A chemical bond results from the overlap of two half-filled orbitals with spin-pairing of the two valence electrons.
The geometry of the overlapping orbitals determines the shape of the molecule.
The interaction energy between atoms in molecules when forming bonds is usually negative (stabilizing) when the interacting atomic orbitals contain a total of two electrons that can spin-pair (orient with opposing spins).

18. Valence bond theory: Hybridization of atomic orbitals
The orbitals in a molecule are not necessarily the same as the orbitals in an atom – hybridization is a math procedure in which the standard atomic orbitals are combined to form new atomic orbitals called hybrid orbitals – which correspond more closely to the actual distribution of electrons in chemical bonds.
Why?? Stabilizing & energy efficient!
(For now) Only interior atoms hybridize and terminal atoms do not.
The number of standard atomic orbitals added together always equals the number of hybrid orbitals formed.
Use electron geometries as determined by VSEPR theory to predict hybridization type.

21. Molecular geometries analysis
shapes_en.html

22. Single, Double, & triple bonds
When p orbitals overlap end to end, the resulting bond is a sigma bond, 𝜎.
When p orbitals overlap side by side, the resulting bond is a pi bond, π.
A double bond in the Lewis model always corresponds to one 𝜎 and one π bond. A triple bond has one sigma bond and two pi bonds.
In pi bonds, electrons are delocalized and electron density is above and below the internuclear axis. For this reason, pi bonds are weaker than sigma bonds – therefore, pi bonds are easier to break (require less energy).
Multiple covalent bonds do not allow free rotation but single covalent bonds do! This free rotation of atoms in covalent bonds can affect the chemical and physical properties of molecules & produces isomers – compounds with the same molecular formula but different structures (spatial arrangement of atoms). Cis- & trans-