Ch. 4 Covalent Compounds, Formulas, Structures

Lewis Structures

Covalent Bonds: Electrons are shared, physically attached to each other in molecules
Lewis electron-dot structures used to represent sharing

one electron shared from each atom $\rightarrow$ one covalent bond
two/three pairs shared $\rightarrow$ double/triple bond
octet rule: noble-gas configuration achieved when Lewis structure shows eight electrons around each atom
- except $\ce{H}$ , only requiring two
unpaired electrons generate magnetic field $\rightarrow$ make substance paramagnetic
all electrons paired, cancelling most magnetic field $\rightarrow$ make substance diamagnetic, small magnetic field

Drawing Lewis Structures

Determine central atom
I. $\ce{C}$ atom(s) are usually central
II. $\ce{H}$ is never central
III. Halogens are usually not
IV. Oxygen is usually not, though may link $\ce{C}$
V. atom appearing only once may be central
Put atoms/groups around central atom to form the skeleton
Find the total number of valence electrons (take into account charge if ion)
Add single bonds between each ion
Fill octets on outer atoms and put remaining in central (nonbonding or lone pairs)
Change some outer atom lone pairs to double bonds if central atom does not have a full octet unless it is $\ce{B}$ (may have more, if period 3-7)
Check formal charges

Some structures may have multiple possibilities, called resonance structures
- actual molecule is like a hybrid of resonance structures, imagine "1.5 bonds"
Some structures have odd number of electrons $\rightarrow$ $\to$ free radicals
- may form dimers to pair them, e.g. $\ce{2NO2->N2O4}$

Formal Charge

technique to analytically determine if a Lewis structure is reasonable

For each atom, count all nonbonding electrons
Count half the bonding electrons (e.g. one bond counts as 1 electron)
Add the counts from step 1 and 2
Compare to number of valence electrons

low formal charges are generally better
atoms with higher electronegativity generally have more negative formal charge

left is better because of all 0 compared to a -1 and +1

Covalent Bond Polarity

based on electronegativity difference
polar if different, range from slightly polar to very polar
nonpolar if no difference

Dipole Moments

$\text{Dipole moment}=q\times r$
where $q$ is the difference in charge (polar molecules will have a slightly negative and slightly positive atom) and $r$ is the distance in meters

unit is debye, where $\pu{1 debye}=\pu{3.34\cdot10^{-30}C\cdot m}$

Ionic Character

largest is $\ce{FrF}$ at $\Delta EN=3.3$
$\Delta EN=1.7$ is 50% ionic and 50% covalent in character

$\Delta EN=$ difference in electronegativity

Ionic Character	$\Delta EN$
Nonpolar Covalent	$0$
Polar Covalent	$<1.7$
Ionic	$>1.7$

Characteristics of Bonds

Bond Order

average number of bonds an atom makes
single bond is 1, double bond is 2, triple bond is 3
in $\ce{SO3}$ where there are 4 bonds and 3 $\ce{O}$ atoms, the bond order is $4/3$

Bond Strength/Energy

triple > double > single in bond strength
bonded atoms vibrate, and the frequency is related to bond energy (strength) from $E=hv$
low energy molecules release more energy upon combustion

Bond Length

in general, single > double > triple

Nomenclature

include numeric prefix (mono, di, tri, etc.) (except first atom never gets mono), end second with -ide
- e.g. $\ce{N2O3}$ is dinitrogen trioxide, $\ce{NO}$ is nitrogen monoxide

Molecular Geometry

VSEPR theory (valence-shell electron-pair repulsion): negatively charged electron pairs repel each other as far apart as possible
bonding electron pairs occupy bonding domain
nonbonding electron pairs occupy nonbonding domain

Molecular Geometry Table

$\ce{A}$ is central atom, $\ce{X}$ and $\ce{E}$ are bonding and nonbonding pairs

Notation	Shape	Example	Angle
$\ce{AX}$	Linear	$\ce{HBr}$
$\ce{AX2}$	Linear	$\ce{CS2}$	$180^\degree$
$\ce{AXE}$	Linear	$\ce{CN-}$
$\ce{AX3}$	Planar Triangle	$\ce{BCl3}$	$120^\degree$
$\ce{AX2E}$	Bent	$\ce{SnCl2}$	$120^\degree$
$\ce{AX4}$	Tetrahedron	$\ce{CCl4}$	$109.5^\degree$
$\ce{AX3E}$	Triangular Pyramid	$\ce{NH3}$	$109.5^\degree$
$\ce{AX2E2}$	Bent	$\ce{H2O}$	$109.5^\degree$
$\ce{AX5}$	Trigonal Bipyramid	$\ce{PCl5}$	$120^\degree$ , $90^\degree$
$\ce{AX4E}$	Distorted Tetrahedron	$\ce{SF4}$	$120^\degree$ , $90^\degree$
$\ce{AX3E2}$	T-shape	$\ce{ICl3}$	$90^\degree$
$\ce{AX2E3}$	Linear	$\ce{I3^-}$	$180^\degree$
$\ce{AX6}$	Octahedron	$\ce{XeF6}$	$90^\degree$
$\ce{AX5E}$	Square Pyramid	$\ce{IF5}$	$90^\degree$
$\ce{AX4E2}$	Square Planar	$\ce{XeF4}$	$90^\degree$

Molecular Polarity

symmetrical molecules are nonpolar
nonsymmetrical molecules are polar if bonds are polar
molecule with >1 type of atom attached to the central atom is often nonsymmetrical, so is polar
central atom with nonbonding electron pairs are often nonsymmetrical, so is polar

one end is positive, other is negative

Covalent Bond Formation

valence bond theory (VB theory): paired electron spins overlaps atomic orbitals, creating covalent bond
molecular orbital theory (MO theory): molecule also has distinct energy levels with electrons
results are the same
an overlap between 2 s orbitals, an s orbital and a p orbital, or 2 p orbitals form sigma bond ( $\sigma$ $σ$ )
- every bond has only 1 sigma bond
sideways overlap of 2 p orbitals forms a pi bond ( $\pi$ $π$ )
- can be 0, 1, or 2 to make single, double, or triple bonds in combination with the sigma bond
In short:
- single bond: $\sigma$
- double bond: $\sigma+\pi$
- triple bond: $\sigma+2\pi$

Hybridization

Hybrid orbitals explain complex geometries and covalent bonds
basically count the number of bonds and lone pairs on the central atom
- 2 total: sp
- 3 total: sp $^2$
- 4 total: sp $^3$