300 Bonding Notes
Chemical bond = a strong attraction between atoms or ions
Bond energy = the energy required to break a bond

Breaking bonds requires energy (is endothermic), forming bonds releases energy
(is exothermic)

The more the bond energy, the stronger the bond
Motivation for bonding… to achieve stable, noble gas electron configuration (filled outermost s and p orbitals)
Octet rule = atoms will form bonds to achieve a stable, noble gas electron configuration
(i.e., the electron configuration of whatever noble gas is closest… recall that noble gases except for helium have 8 valence electrons, hence the name “octet”)
3 types of bonds: ionic, covalent, metallic
Bond Type
Types of elements involved
Metallic
Metals
Relative strength Strong
Common states of matter of compounds at room temperature
Solids
Conductivity Extremely conductive
Ionic
Metals and nonmetals, polyatomic ions
Strong
Solids
Covalent
Nonmetals
Weak
Liquids or gases
Can conduct when molten or in aqueous solution
High
Nonconductive
Melting/boiling points
High
Physical properties Hard but shapeable, malleable, ductile
Hard, brittle
Low
Flow easily (usually liquids or gases)
Metallic Bonding
Properties are confusing: hard, high melting points, yet easily shaped… difficult to separate atoms but easy to slide them around past each other as long as they stay in contact with each other
“Electron sea model” is the explanation – a regular array of metal cations in a “sea” of valence electrons (mobile electrons can conduct heat and electricity, ions can move around as metal is hammered into a sheet or pulled into a wire)
Alloy = substance that contains a mixture of elements and has metallic properties

Ex: pure iron very soft, ductile, malleable; add carbon (to get steel) and properties change- harder, stronger, less ductile; add other elements (to get steel alloys) to further tune properties (ex: add chromium to get stainless steel that resists corrosion)
Ionic Bonding
Ionic bond = strong attraction between closely packed, oppositely charged ions
Ionic bonds result when an atom that loses electrons relatively easily (metal or cation) reacts with one that has a high affinity (attraction) for electrons (nonmetal or anion) (the two elements involved have a high electronegativity difference, usually > 1.9
If you subtract the electronegativities of the two elements involved and it is >1.9, the bond will be IONIC
Ionic bonds involve a TRANSFER of electrons
How do Lewis structures show ionic bonding?
(examples…. NaCl, CaF
2
, Mg
3
N
2
)
Now… think about how this relates to how we wrote formulas for ionic compounds!!
(same examples)
Covalent Bonding
Covalent bond = sharing of an electron pair between two nuclei
Covalent bonds result between atoms that have similar affinities for electrons (both nonmetals)
This sharing of electron pairs can be even or uneven
Nonpolar covalent bond = covalent bond with an even sharing of electrons
(electronegativity difference between elements is essentially zero)
Polar covalent bond = covalent bond with an uneven sharing of electrons
(electronegativity difference between elements is relatively small but not zero, 0.5 to about 1.8)

With an uneven sharing, the electron pair will be pulled closer to one atom or the other

The end of the bond with the electrons closer to it will be more negative in charge, the other end of the bond will be more positive in charge… this is a dipole
If you subtract the electronegativities of the two elements and get 0 or almost 0, the bond is NON POLAR COVALENT… if it is between 0.5 and 1.8, the bond is POLAR
COVALENT

(tug of war analogy: nonpolar, polar, ionic)
A special type of covalent bonding…
Coordinate covalent bond = covalent bond where both electrons in a bond are contributed by the same atom (common examples: ammonium, hydronium… more on that later)
How do Lewis structures show covalent bonding?

Recall, a Lewis structure for an element shows how the valence electrons are arranged around a single atom

A Lewis structure for a compound shows how valence electrons are arranged among the atoms in a molecule

A shared pair of electrons (e.g., bond) between two atoms is shown as a 2 dots or a single line (1 line (2 dots) = single bond, 2 lines (4 dots) = double bond, 3 lines
(6 dots) = triple bond)

Unshared pairs of electrons can be shown as two dots on one side of the chemical symbol
Foolproof Steps to Drawing Lewis Structures for Covalently Bonded Substances
1.
Start with the formula for the compound
2.
Count up the number of valence electrons for each atom using the Periodic Table
3.
Add up the total number of valence electrons for each atom in the formula to get the total number of valence electrons available a.
Note: If the substance is an ion, add or subtract electrons to this total depending on the charge (+ charge means subtract electrons, - charge means add electrons)
4.
Use a pair of electrons to form a bond between each pair of bound atoms (show these single bonds either as a pair of dots or a line)
5.
Distribute the remaining electrons making sure that the octet/duet rule is satisfied and that the total number of valence electrons used is exactly equal to what was available (this may take some trial and error)
6.
Distribute any extra electrons as lone (i.e., nonbonding) pairs around the central atom
7.
Count up all electrons and make sure the octet/duet rule is satisfied for each atom and that the total number of electrons used is equal to what was available originally… if it doesn’t add up, you’ll need to revise and try again (trial and error) possibly adding double or triple bond(s)!
Resonance = when a single Lewis structure does not adequately reflect the properties of a substance (i.e., when you can draw several equivalent Lewis structures)
(examples, including coordinate covalent: HF, H
2
S, CH
4
, O
2
, N
2
, HCN, CN
-
, CO, SO
4
CO
3
2-
, O
3
, BF
3
, BeCl
2
, NH
3
, H
3
O
+
, NH
4
+
)
2-
,

Once you can draw the Lewis structure of a compound, there are many other interesting things to consider…. 3D molecular shape/geometry, molecular polarity, and intermolecular forces!
Three Dimensional Shape/Geometry of Molecules
It is useful to be able to predict shape of a molecule in three dimensions
Valence Shell Electron Pair Repulsion (VSEPR) Model:

A molecule’s shape in 3D is determined by minimizing electron-pair repulsions

Bonding and nonbonding pairs will want to be as far apart as possible to minimize repulsions

Repulsion from lone pairs (nonbonding pairs) and double or triple bonds is greater than that from bond pairs
To determine the 3D geometry of a molecule:
1.
Draw the Lewis structure for the molecule
2.
Count the number total electron domains (bond/lone pairs)
3.
Arrange the bonds and lone pairs to minimize repulsions in 3D
4.
Describe the shape of the molecule by memorizing the information in the chart below
Bond angle = angle made by between lines made by connecting nuclei of the two atoms
Examples and shape names/descriptions:
Total electron domains
2
Number of bond regions
2
Number of lone pairs
0
Name of shape
Linear
Picture of shape Bond angle
Examples
180 degrees
CO
2
,
BeCl
2
3 3 0 Trigonal planar
120 degrees
BF
3

Total electron domains
Number of bond regions
Number of lone pairs
Name of shape
Picture of shape Bond angle
Examples
4 4 0 Tetrahedral 109.5 degrees
CH
4
4 3 1 Trigonal pyramidal
About
107 degrees
NH
3
4 2 2 Bent About
104.5 degrees
H
2
O

Total electron domains
Number of bond regions
Number of lone pairs
Name of shape
Picture of shape
(revisit examples)
Molecular Polarity
Polar molecule = molecule that has positive and negative poles
Bond angle
Examples
Polar molecules tend to orient (align) themselves in an electric field and are said to have
“dipole moment”
Nonpolar molecule = molecule that does NOT have positive and negative poles (i.e., is said to have a “dipole moment” of zero)
To determine if a molecule is polar, you must do two things:
1.
Look at the bonds and determine the polarity of each bond
2.
Look at the symmetry of the molecule as a whole a.
If polar bonds are arranged symmetrically in 3D, the molecule is nonpolar b.
If polar bonds are NOT symmetrically in 3D, the molecule is polar
(revisit examples)
Intermolecular Forces
Intra molecular forces = forces WITHIN a single molecule (i.e., BONDS)
Inter molecular forces = forces between molecules that result in physical properties of substances
Comparison of States of Matter and their IMFs
State of Matter
Solids
Liquids
Intermolecular Forces
Very strong
Moderate
Effect of IMFs
Particles are held close together and in place; substances are rigid, incompressible, and do not flow; high melting/boiling points
Particles are held closer together but can still slide

Gases Very weak past each other; substances can flow and be poured but are virtually incompressible; moderately high melting/boiling points
Particles can spread out and move wildly about; substances have no definite shape or volume and are highly compressible; low melting/boiling points
NOTICE from above… low IMFs = low melting/boiling points, strong IMFs = high melting/boiling points!
Several types of IMFs between neutral molecules that are of interest to us: London dispersion, dipole-dipole, and hydrogen bonding (all of which are sometimes called “van der Waals forces”) and ion-dipole forces
London dispersion forces = forces of attraction between molecules based on motion of electrons, these are the only IMFs in nonpolar molecules but can and do exist in any type

Larger molecules, molecules with higher masses have more electrons and are more polarizable (can move electrons around to get an instantaneous dipole)

The more polarizable, the stronger the London forces

So, the bigger/heavier the molecule (the more electrons), the stronger the London forces
Dipole-dipole forces = forces between polar molecules (between positive and negative ends of different molecules)

For molecules with equal masses and sizes, strength of these forces increases with increasing polarity

The more polar, the stronger the dipole forces
Hydrogen bond forces = special type of dipole-dipole force between H atoms in a polar bond and a small, electronegative atom (F, O, or N)

Explains anomalies observed in certain physical properties (ex: abnormally high boiling point of water, ammonia, and hydrogen fluoride compared to other compounds of elements)
Ion-dipole forces = IMF between an ion and the partial charge on the end of a polar molecule

Between a cation and – end of dipole or between an anion and + end of dipole

Exist in aqueous solutions (this is what helps water pull apart the ions in an ionic compound when they dissolve)
(revisit examples)

Comparing/Identifying IMFs

Identifying type: o Everything has at least London disperson o If nonpolar- London dispersion only o If polar – also dipole-dipole o If H bonded to F, O, or N - also hydrogen bonding o If an ionic compound in aqueous solution – ion-dipole (we won’t use too many of these)

London dispersion < Dipole-dipole < Hydrogen bonding

All IMFs are WEAKER than actual bonds

If you are asked to rank substances based on their melting or boiling points, remember… the stronger the IMFs, the higher the melting and boiling points because it is harder to separate the molecules from each other!