4.2 COVALENT BONDS
Covalent bond  the result of electron sharing
Pairs of valence electrons are attracted to two nuclei.
The shared electrons orbit each nucleus.
Electronegativity values must be similar (EN < 1.7) for atoms to share electrons.
Non-metals bond to each other by covalent bonds.
Atoms share electrons so that both atoms appear to gain an electron and complete their valence shell.
It is important to differentiate between the electron pairs that form bonds and those that do not.
Bonding  electrons that are shared by the two atoms
Non-bonding - electrons that only orbit one electron LONE PAIRS
Co-ördinate Covalent Bond (14Å"4)  a covalent bond in which one atom provided BOTH bonding electrons.
A lone pair of ONE atom becomes a bonding pair for both atoms.
Bonding Diagrams for Covalent Bonds
A line represents a pair of electrons.
Draw double-headed arrows ("!) between single electrons of two atoms.
Co-ördinate Draw a single arrow from a lone pair to the bonding area between the atoms
Read 14Å"1 to 14Å"4
Å" Å"
Å" Å"
Å" Å"
Copy bonding diagrams for the following molecules into their notes:
Elemental Molecules
Single Bonds 1 pair of bonding electrons (2 e-)
1. H HÅ" "! Å"H Ä… H - H each H: 1bonding pair 0 lone pairs 2 e- total
2
2. F :FÅ" "! Å"F: Ä… :F-F: or | F-F | each F: 1bonding pair 3 lone pairs 8 e- total
2
Double Bonds 2 pairs of bonding electrons (4 e-)
3. O :OÅ" "! Å"O: Ä… O = O or O = O each O: 2 bonding pair 2 lone pairs 8 e- total
2
Triple Bonds 3 pairs of bonding electrons (6 e-)
4. N :NÅ" "! Å"N: Ä… :N a" N: or | N a" N | each N: 3 bonding pairs 1 lone pair 8 e- total
2
Common Molecular Compounds Common Polyatomic Ions (contain covalent bonds)
+
methane, CH ammonium ion, NH
4 4
2
ammonia, NH carbonate ion, CO
3 3

water, H O nitrate ion, NO
2 3
3
carbon dioxide, CO phosphate ion, PO
2 4
2
carbon monoxide, CO sulfate ion, SO
4
ethanoic acid, CH COOH ethanoate ion, CH COO
3 3
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Shape and Bond Angle
Example Charge Centres Bonding Pairs Non-Bonding Pairs Shape Bond Angle
methane, CH
4
+ 4 4 0 tetrahedral 109.5°
ammonium ion, NH
4
ammonia, NH
3
4 3 1 pyramidal 107°
hydronium ion, H O+
3
water, H O 4 2 2 bent 105°
2
sulfur trioxide, SO
3
boron trifluoride, BF 3 4 (S=O + 2 SO) 0 trigonal planar 120°
3
ethene, C H H C=CH
2 4 2 2
carbon dioxide, CO
2
2 4 (2 C=O) 0 linear 180°
ethyne, C H HCa"CH
2 2
Valence Shell Electron Pair Repulsion (VSEPR) theory
" valence electrons repel each other and tend to distribute themselves as far away from each other as possible
" non-bonding (lone pairs) repel more than bonding electrons resulting in reduced bonds angles in ammonia and water
Bond Length and Bond Strength
Bond Length decreases as the number of bonds increases between two atoms
" single bonds are longer than double bonds; triple bonds are shorter than double bonds
C C > C=C > Ca"C
Bond Strength increases as the number of bonds increase between two atoms
more energy is required to break three bonds of a triple bond than both bonds of a double bond
C C < C=C < Ca"C
Bond Types
Electronegativity difference (EN) between a pair of bonding atoms determines the bond type.
Bonding
Bond Type EN Charges Intermolecular forces
Electrons
electrons formal charges
Ionic 1.8 to 4.0 ionic
transferred (+ and  )
partial charges dipole-dipole possible
Polar unequal
0.5 to 1.7 (´+ and ´ ) H-bonding possible
Covalent sharing
permanent dipoles dispersion if symmetrical
Nonpolar no charges
0 to 0.4 equal sharing dispersion ONLY
Covalent no permanent dipoles
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Molecule Types
Non-polar molecules X NO NET DIPOLE
Electrons are evenly distributed to atoms around a central atom
" All bonds are non-polar covalent bonds
o hydrocarbons (CXHY), elemental molecules (N2, Cl2)
" The polar covalent bonds have symmetry in all directions around the central atom
o carbon dioxide (CO2), carbon tetrachloride (CCl4), sulfur trioxide (SO3)
Polar molecules " NET DIPOLE = polar bonds and no symmetry in all directions
Electrons tend to move toward ONE end of the molecule
Half of area around a central atom is ´+ and the other half is ´
All polar molecules have dipole-dipole interactions
Intermolecular Forces
Giant Covalent HUGE molecular crystals from atoms sharing valence electrons (covalent bonds)
Carbon, Cn n carbon atoms bond to each other  each C bonds to 4 other carbons
diamond  3-dimensional crystal lattice
graphite  2 dimensional sheets with weak C-C bonds between sheets
fullerene  hollow spheres and microtubules
Silicon, Sin semiconductor, computer chips
Silica, SiO2 each silicon atom binds to 4 oxygen atoms in a repeating crystalline pattern
sand, glass
Allotrope different forms of the same element
have different bonding and different physical and chemical properties
Ionic Bonding electrostatic attraction between opppositely charged ions
" electrons must have been transferred (not shared) between atoms (EN>1.7)
" polyatomic ions (having covelent bonds) that have gained or lost electron(s)
" NaCl, MgF2
Hydrogen Bonding electrostatic attraction between H´+ of one molecule and a ´ lone pair on another
" OH, NH, and HF bonds will produce hydrogen bonding
" water H2O, ammonia NH3, alcohols, sugars, amino acids, carboxylic acids
Dipole-Dipole Interactions electrostatic attraction opposite permanent dipoles on two polar molecules
" between ´+ of one molecule and ´ on another
" carbon monoxide CO, chloroform CHCl3
Dispersion Forces the weakest intermolecular force
" electrostatic attraction between temporary dipoles caused by electron motion
o increase in strength with molecular mass (number of electrons)
Br2 > Cl2 > F2 C8H18 > C3H8 > CH4
o increase with the density (closeness) of the molecules
saturated fats > unsaturated oils
van der Waals forces the weak intermolecular forces = dipole-dipole and dispersion forces
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