COVALENT COMPOUNDS
§ Formed when two atoms share electrons
o Usually involve 2 or more non-metal atoms
o Similar electronegativity values (electronegativity difference is less than 1.7).
§ Can be pure covalent – equal sharing of e (<.3)
§ Polar covalent – uneven sharing of electrons (.3-1.7)
o Called molecules
o There are more covalent compounds than ionic compounds.
§ Binary covalent – include 2 elements only
§ Organic molecules (which include carbon) are covalent
· Glucose C6H12O6
· Sucrose C12H22O11
· Hydrocarbon C6H6 etc.
§ The shared valence electrons are found in a molecular orbital, formed by overlapping atomic orbitals of the atoms involved in bonding.
§ Recall the properties of covalent molecules listed previously in chapter 5 notes (flow chart).
NAMING COVALENT COMPOUNDS
§ Covalent molecules are NOT named like ionic compounds!!!
Steps to naming binary covalent molecules: (only include 2 elements)
1. the first element (least electronegative) is named first
2. the second element has the ending –ide
3. Prefixes are used before the element name to indicate the number of each atom in the molecule (the subscript)
# of atoms / Prefix1 /
Mono
2 / Di3 / Tri
4 / Tetra
5 / Penta
6 / Hexa
7 / Septa
8 /
Octa
9 / Nona10 / Deca
Do NOT use a prefix for the first atom when there is only one present.
When the element name begins with a vowel, the (a) and (o) are dropped from the end of the prefix
Examples:
CCl4
carbon tetrachloride
CO
carbon monoxide
(eliminate the o from the prefix)
P2Cl5
diphosphorus pentachloride (more than one of the first element)
N2O3
dinitrogen trioxide
Si3N4
trisilicon tetranitride
H2O
dihydrogen monoxide
Steps for writing covalent formulas from the name.
1. the elements appear in the same order as in the name
2. the prefix indicates the subscript in the chemical formula
Examples:
Boron trifluoride BF3
Dinitrogen monoxide N2O
Dinitrogen tetroxide N2O4
ENERGY and STABILITY of COVALENT BONDS
· Most atoms have relatively LOW STABILITY and HIGH POTENTIAL ENERGY.
· When a compound forms, the atoms become MORE STABLE and the potential energy is at a MINIMUM.
Example:
The potential energy curve for hydrogen.
http://www.usm.maine.edu/~newton/Chy251_253/Lectures/LewisStructures/Dihydrogen.html
· When the nuclei are farthest apart, the potential energy is ZERO.
· As they get closer, the energy DECREASES.
· When the potential energy is LOWEST (at -436 kJ//mole), the atoms bond.
· The distance between the two nuclei at this lowest energy is called the BOND LENGTH and is 75 pm for the H2 molecule.
· When the REPULSION of the two atoms perfectly balance the ATTRACTIVE FORCES between the two nuclei, a COVALENT BOND forms.
· Since the potential energy DECREASED, energy has been RELEASED. For H2, the energy released is -436 kj/mol.
· The energy released when forming the bond is the exact amount of energy that would be needed to break the bond. Energy needed = +436 kj/mol
· The energy required to break a bond is known as the BOND ENERGY.
BOND ENERGY
and
BOND LENGTH
Bond energy(kj/ mol) / Bond length (pm) / Electronegativity Difference
HF / 570 / 92 / 1.8
CF / 552 / 138
OO / 498 / 121
HH / 436 / 75
HCl / 432 / 127 / 1.0
CCl / 397 / 177
HBr / 366 / 141 / 0.8
HI / 299 / 161 / 0.5
· BOND ENERGY (STRENGTH) is inversely related to the BOND LENGTH. The higher the bond energy, the SHORTER the bond.
· Bond length is actually an AVERAGE DISTANCE between the two nuclei since the distance is constantly changing due to the bond being able to vibrate and bend.
BOND STRENGTH and POLARITY
· The HIGHER the electronegativity difference, the STRONGER the bond.
o Ex. HF has a much stronger bond than HI.
o HF is nearly IONIC and HI is almost PURE COVALENT.
LEWIS DOT STRUCTURES
RESONANCE
· when 2 or more equivalent lewis structures exist for a molecule or compound
· all structures should be represented with a double arrow between
· the actual bond strength is an average of all the bonds.
· Examples:
(see back of notes)
MOLECULARGEOMETRY: ARRANGEMENT and SHAPE
· The three dimensional shape of molecules can be predicted using the VALENCE SHELL ELCTRON PAIR REPULSION THEORY.
o maximizes space between electron pairs to predict shape.
o VSEPR
· Arrangement of electrons – based on the number of regions of electron density around the central atom.
· Shape of the molecule – can be predicted by counting the number of bonding pair electrons and lone pair electrons.
· Treat double and triple bonds as single electron pairs when determining VSEPR shape.
once the shape is determined, certain bond ANGLES can be predicted. Bond angles are altered by the present of LONE PAIR ELECTRONS. When present, the bond angles are compressed due to the REPULSION of the.
o lone pairs require more space since they are only controlled by one nucleus.
· The arrangement of electrons and the shape of the molecule with be the same when there are NO LONE PAIR ELECTRONS.
VSEPR and MOLECULAR GEOMETRY
BONDING PAIRS of ELECTRONS / LONE PAIR ELECTRONS / TOTAL NUMBER OF ELECTRON PAIRS / ARRANGEMENT / SHAPE / SKETCH andBOND ANGLES
1 / 0 / 1
2 / 0 / 2
3 / 0 / 3
2 / 1 / 3
4 / 0 / 4
3 / 1 / 4
2 / 2 / 4
5 / 0 / 5
4 / 1 / 5
3 / 2 / 5
2 / 3 / 5
BONDING PAIRS of ELECTRONS / LONE PAIR ELECTRONS / TOTAL NUMBER OF ELECTRON PAIRS / ARRANGEMENT / SHAPE / SKETCH
and
BOND ANGLES
6 / 0 / 6
5 / 1 / 6
4 / 2 / 6
3 / 3 / 6
2 / 4 / 6
HYBRIDIZATION
In order for bonds to have equivalent energy, mixing of orbitals must occur.
Ex. Methane (CH4) predicts 4 equivalent bonds. Which orbitals of C are used?
mix the orbitals used into a new “hybrid orbital” with a new name.
Total number of e pairs around a central atom / Arrangement / Hybridization2 / Linear / sp
3 / Trigonal planar / sp2
4 / Tetrahedral / sp3
5 / Trigonal bipyramidal / sp3d
6 / octahedral / sp3d2
POLARITY and DIPOLE MOMENT
POLAR BONDS:
§ Result from high ELECTRONEGATIVITY DIFFERENCES between atoms of a molecule.
§ Causes PARTIAL POSITIVE and PARTIAL NEGATIVE ends of the molecule (called DIPOLES).
SHAPE AFFECTS POLARITY:
§ DIPOLE MOMENT – overall direction of electron “pull” within a molecule. Show using a molecular model.
§ Sometimes dipoles will CANCEL each other, and the result will be a molecule with NO NET DIPOLE due to the shape.
Molecule / Sketch / Direction of dipole or no net dipoleHBr
BeCl2
BBr3
SeCl2
CO2
H2O
POLARITY AFFECTS PROPERTIES:
§ For example:
o CO2 is non-polar, so the ATTRACTIVE FORCE between CO2 atoms is VERY WEAK.
o This results in a lower MELTING POINT and BOILING POINT.
§ H2O is POLAR, the molecules interact with each other and attractive forces are greater.
§ This results in a higher MELTING POINT and BOILING POINT.
§ A very strong INTERMOLECULAR force exists between water molecules due to its polarity.
o This force is called HYDROGEN BONDING
§ It is not the bond within water….it is a bond between water molecules.
HYBRIDIZATION