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
Metallic
Ionic
Metals
Metals and nonmetals,
Types of elements involved
polyatomic ions
Covalent
Nonmetals
Relative strength
Common states of matter of
compounds at room temperature
Strong
Solids
Strong
Solids
Weak
Liquids or gas
Conductivity
Extremely conductive
Can conduct when
molten or in aqueous
solution
Nonconductiv
Melting/boiling points
High
High
Low
Physical properties
Hard but shapeable,
malleable, ductile
Hard, brittle
Flow easily (u
liquids or gase
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 changeharder, 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 , Mg N )
2
3
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
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)
-
2-
(examples, including coordinate covalent: HF, H S, CH , O , N , HCN, CN , CO, SO ,
2
2-
+
+
CO , O , BF , BeCl , NH , H O , NH )
3
3
3
2
3
3
4
4
2
2
4
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 a