Chapter 12
Valence Electrons: Review
• The outermost electrons are involved in bonding.
– Valence electrons.
• Most atoms are stable with 8 valence electrons.
– This is the octet rule.
– Which only need 2?
• Atoms will lose, gain or share electrons in order to have a filled valence shell.
– The octet rule is broken by some atoms.
Ions and Valence Electrons
• Atoms tend to form ions based on their number of valence electrons.
• Metals tend to form + ions (cations).
• Non-metals tend to form –ions (anions).
• Noble gases have filled valence shells and tend not to form ions.
• We can determine the number of valence electrons in an atom from its group
number.
Ionic Bonds
• Opposite charges attract.
• If a + charged ion comes into contact with a – charged ion, they will attract each
other.
• This attraction holds the ions close together and is call an ionic bond.
• Because metals tend to form + ions and non-metals form – ions, ionic bonds usually
form between metals and non-metals.
– NaCl, CaO, FeBr2
Ionic Compounds
• Compounds held together by ionic bonds are call ionic compound.
• The + and – charges of ionic compounds must balance.
– Mg2+ will bond to two ions of Cl1-.
– MgCl2
• The common name for ionic compound is “salt”.
• Metal and Non-metal: CaCl2.
• May involve polyatomic ions: NaNO3.
Crystals
• Most ionic compounds consist of thousands of + and – ions held together in a crystal.
– The formula of the compound indicates the overall ratio of + and – ions in the
crystal.
– A MgCl2 crystal may have thousands of ions, but the ratio of Cl to Mg in the crystal
is 2:1.
• The shape of the crystal depends on the packing of ions in the crystal.
Ionic vs. Covalent
• When 2 non-metals form a bond, they do not tend to form ions.
• The non-metals will share valence electrons to get a more stable electron
configuration.
• They form a covalent bond.
– CO, H2O, HCl, N2
• Covalently bonded atoms form molecules.
• Non-metal to non-metal.
Example
• Oxygen has 6 valence electrons.
• How do we know that?
– Oxygen needs 2 more electrons to be “happy”.
• Hydrogen has 1 valence electron.
– It needs 1 more electron.
• If an oxygen atom shares electrons with 2 hydrogen atoms, all the atoms’
requirements are satisfied.
• The result is a water molecule H2O.
Covalent vs. Ionic Bonds
•
•
Ionic Bonds
– Attraction between positive and negative ions.
– Metal to non-metal.
– Cation to Anion
Covalent Bonds
– Sharing of electrons.
– Non-metal to non-metal.
Example
•
Label the following bonds as ionic or covalent:
– NaCl
– MgO
– CO
– CCl4
– NH3
Using Lewis Dot Structures
More on Covalent Bonds
•
•
Each covalent bond has 2 electrons.
– Single bond
Some atoms will form two or three bonds with the same atom.
– Double Bond =
– Triple Bond 
More difficult Lewis Structures Ex: SO2
•
Count the number of valence e- available.
•
S = 6, O = 6. So we have 18
•
Count the # of e- needed for octet rule.
•
S & O each need 8. So we need 24.
•
e- needed - e- available = e- shared
•
24 – 18 = 6
•
e shared/2 = # bonds
•
6/2 = 3 bonds
•
Arrange the atoms and insert the bonds.
O-S=O
•
Fill in the remaining valence e.. .. ..
:O – S = O
..
..
Lewis Structures of Polyatomic Ions
•
Each negative charges adds 1 to number of valence e- available. Example: NO3•
Count the number of valence e- available.
•
N = 5, O = 6, charge adds 1. So we have 24
•
Count the # of e- needed for octet rule.
•
N & O each need 8. So we need 32.
•
e needed - e- available = e- shared
•
32 – 24 = 8
•
e shared/2 = # bonds
•
8/2 = 4 bonds
Continued
5. Arrange the atoms and insert the bonds.
O-N=O
|
O
6. Fill in the remaining valence e-, and include the charge.
..
..
:O – N = O
..
| ..
:O:
..
Electrons in Molecules
•
•
•
The valence electrons of a molecule come in pairs.
– Orbitals hold 2 electrons each.
– Bonds have 2 electrons each.
Bonding
Non-bonding (lone pairs).
Shapes of Molecules
•
Depends on:
– Number of atoms bound to central atom.
– Number of lone pairs of electrons.
Molecular Shape VSEPR
•
•
•
The pairs of valence electrons help to set up the shape of the molecule.
However, only bonded atoms count in the final shape.
This is because atoms are much larger and heavier than lone pairs.
Determining Molecular Shape
•
•
•
Count the # of atoms bound to the central atom.
Count the lone pairs of electrons.
Determine the geometry of the molecule.
Shapes to Know
• CH4
– Tetrahedral
– 4 atoms attached to the central atom and no lone pairs.
• NH3
– Trigonal pyramidal
– 3 atoms attached to central atom and 1 lone pair.
• H2O
– bent
– 2 atoms attached to central atom and 2 lone pair.
Polar Molecules
• In a covalent molecule, the electrons are not always distributed evenly.
• Because of:
– Unequal sharing of electrons (see next).
– Shape of molecule (more later).
• Result:
– One side of the molecule is more negative (electrons spend more time there).
– One side is more positive.
– Polar molecule.
Electronegativity
• Electrons in a covalent bond are not always shared equally between atoms.
– Called a “polar covalent bond”.
• Some atoms pull harder on the electrons in a bond.
– They have a higher electronegativity.
– Fluorine has the highest.
– O and Cl are next.
Polar Bonds
• Covalent Bond.
• Unequal sharing of electrons.
• The result of 2 atoms with different electronegativities bonded together.
• The bigger the difference the more polar the bond.
• Is sometimes indicated by an arrow pointed in the direction of the more
electronegative atom.
Shapes of Molecules
•
The shape of a molecule decides whether it is a polar molecule or not.
– Example: water.
– Shape: bent.
O
H
H
Non-Polar Molecules
• If a molecule has no polar bonds, it is non-polar.
– O2
• Some molecules have polar bonds that cancel out because of the shape of the
molecule.
– These molecules are also non-polar.
– CH4
– BF3
Example
• The BF3 molecule has a triangular shape.
• Each F pulls on electrons more than B.
– Fluorine has a higher electronegativity.
• However, due to the shape, the “pulls” cancel out.
• Result: no side of the molecule is more negative than another.
– Non-polar.
Polar vs. Non-Polar
• Polar molecules attract each other.
– High boiling points.
– Mix with other polar molecules.
– Don’t mix well with non-polar molecules.
• Oil and water.
• Non-polar molecules don’t attract other molecules well.
– Low boiling points.
– Mix well with other non-polar molecules.
• Like dissolves like.
Homework
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Ch 12 starts p. 335
Key Concepts
1, 3 – 7, 10
Exercises
5 – 10, 33 –37, 39 – 50, 53 – 56, 67 – 74