COVALENT BONDING
A. Single Covalent bonds -- a bond in which two atoms share one pair of electrons
1. Atoms share electrons to obtain a noble gas configuration
H` + ,H --> H:H
a. Represent shared pair of electrons with a dash
H--H
1. Structural formula -- chemical formulas that show the arrangement of atoms in molecules and polyatomic ions
2. Refresher --> ionic formula units represent lowest whole # ratio of ions while covalent molecular formulas represent # of atoms
3. Most covalent compounds are combinations of two or more nonmetals
a. Diatomic atoms follow these rules since they need only one electron
1. Group 7A atoms (need 1 electron)
b. Pairs of electrons not involved in bonding are called unshared pairs (unshared electrons)
B. Double and Triple bonds
1. Double bond -- a bond in which two atoms share two pairs of electrons to obtain an octet
a. Group 6A elements (need 2 electrons)
Ex. --> O2, CO2
2. Triple bond -- a bond in which two atoms share three pairs of electrons to obtain an octet
a. Group 5A elements (need 3 electrons)
Ex. --> N2
C. Coordinate covalent bonds -- a bond in which one atom contributes a pair of electrons to form a bond
1. If one atom is stable and the other is not, then the stable one donates a pair of electrons
Ex. --> CO
a. Represent coordinate covalent bonds with an arrow pointing towards atoms receiving pair of electrons
b. most polyatomic ions contain covalent and coordinate covalent bonds
D. Bond Dissociation Energies -- total energy required to break the covalent bond between atoms
1. When covalent bonds form, energy is released
a. Therefore, energy is needed to break bonds
Ex. --> H2 needs 435 kJ/mol to break the bond
E. Resonance structures -- structures that occur when it is possible to write two or more valid electron dot formulas
1. Electron pairs flip back and forth (resonate) between different structures
a. Not totally correct
2. bond lengths are equal
a. Actual bonding is an average of all structures (hybrid or mixture)
Ex. --> O3
F. Exceptions to the Octet Rule
1. Happens when total # of valence electrons is an odd # (Ex. --> NO2)
2. Consider electrons as spinning charge
a. Spinning charge creates a magnetic field
b. Pairs spin opposite directions and therefore cancel each other
1. Diamagnetic -- substance in which all electrons are paired
2. These compounds are not affected by outside magnetic forces
c. Single electrons create magnetic field since they are not cancelled
1. Paramagnetic -- substance in which one or more electrons are unpaired
2. These compounds are affected by outside magnetic forces
3. Mass appears greater in magnetic field
4. Reason for O2 not following octet rule
II. Bonding Theories
A. Molecular Orbitals -- orbitals that apply to the entire molecule (quantum mechanical model)
1. Atomic orbitals overlap when atoms bond to make molecular orbitals
a. Occupied by only two electrons max
2. # of molecular orbitals = # of overlapping atomic orbitals
a. Called bonding and antibonding orbitals
1. Bonding Orbitals -- molecular orbital with lower energy than the atomic orbitals
2. Antibonding Orbitals -- molecular orbital with higher energy than the atomic orbitals
b. Electrons go into bonding orbital first since it has less energy (more stable)
c. If Antibonding orbital is left empty, then the bond occurs
Ex. -- H2, 1s atomic orbitals overlap to form a Sigma bond
1. sigma bond -- electrons in molecular orbitals that are symmetrically along the axis connecting the two atomic orbitals (happen when s orbitals overlap with s orbitals)
3. Covalent bonding is really imbalance between attractive and repulsive forces
a. If attractive forces of nuclei for electrons is greater, electrons go into bonding orbitals and form a stabel molecule
b. If repulsive forces of nuclei for each other is greater, electrons go into antibonding orbital and do not form a molecule (stay separated)
4. p-orbitals can overlap also
a. Overlap end-to-end to form sigma bond (ex. --> F2)
b. Overlap side-to-side to form pi bond (ex. --> CO2)
c. Pi bond -- electrons in "sausage-shaped" regions above and below the axis connecting two atomic orbitals
d. Weaker than sigma bonds
B. VSEPR Theory -- since electron pairs repel, molecular shapes adjust so valence electron pairs are as far apart as possible
1. Used to describe the 3-D shape of molecules
a. electron dot & structural formulas are only 2-D
Ex. --> CH4, molecule is actually a tetrahedron (Pyramid with a triangular base)
each H-C-H bond angle will = 109.5 degrees, most stable arrangement
2. Unshared electron pairs (ammonia, NH3)
a. Unshared electrons are held more tightly, force bonded electrons away
b. H-N-H bond angle = 107 degrees
3. Two unshared electron pairs (water)
a. Force bonded electrons away more
b. H-O-H bond angle = 105 degrees
c. Molecule is planar (flat) & bent
III. Polar bonds and Molecules
A. Bond Polarity
1. Covalent bonds are like a tug-of-war for electrons
a. If atoms have equal electronegativity values, then the bond is nonpolar
1. Nonpolar Covalent bond -- bonding electrons are shared equally
2. Diatomic molecules (H2, O2)
b. If one atom has higher electronegativity than the other, then the electron pair is pulled slightly to that atom
1. Polar Covalent bond (Polar bond)-- bonding electrons are shared unequally
2. Ex. --> HCl, H2O
3. Notation for polar bonds
a. Greek sign delta with charge
b. arrow with a tail, arrow head towards more electronegative atom
Electronegativity Difference Type of Bond
0.0 - 0.4 Nonpolar Covalent
0.4 - 2.0 Polar Covalent
>= 2.0 Ionic
B. Polar Molecules -- A molecule with a slightly positive end and a slightly negative end
1. Polar bonds can create a polar molecule (dipole -- a molecule with two poles)
a. Affected by magnetic fields
b. Ex. --> HCl
2. Polar bonds can cancel each other out based on symmetry
a. Nonpolar molecule
b. Ex. --> CO2
C. Attractions Betwen Molecules
1. Molecules attract to each other by forces that are weaker than ionic or covalent bonds
a. Van der Waals forces -- weakest attractive forces between molecules
1. Dispersion forces -- caused by moving electrons (weakest type of Van der Waals)
a. Increase as # of electrons increase
2. Dipole interactions -- attraction betwen polar molecules (2nd type of Van der Waals)
a. Positive end of one molecule attracts to the negative end of another molecule
b. Ex. --> H2O
c. Similar to ionic
2. Forces are the reason for the state of matter of the molecular compound
a. Halogens go from gas (F, Cl) to liquid (Br), to solid (I, At) due to increased Intermolecular forces
3. Hydrogen bonds -- a covalently bonded H is weakly bonded to an unshared pair of electrons from another molecule
a. H usually makes polar bonds since it is electron deficient
b. Shares a nonbonded pair of electrons with another molecule to compensate
1. Bond is 5% the strength of a covalent bond
D. Intermolecular Attractions and Molecular Properties
Characteristic Ionic Compound Covalent Compound
Representative Unit Formula Unit Molecule
Bond formation Transfer electrons Share pair of electrons
Type of elements Metal with Nonmetal Two nonmetals
Physical State Solid Solid, Liquid, or Gas
Melting Point High (>300 C) Low (<300 C)
Solubility High High to Low
Electrical conductivity Good Poor
in aqueous solution
1. Exceptions to the melting point trend
a. Network solids -- solids in which all atoms are covalently bonded to each other
b. Most covalent solids break intermolecular forces when melted
c. Network solids break covalent bonds
1. Very high melting points (>1000 C)
2. Usually decompose or vaporize before melting occurs
Outline based upon:
Matta, M. S., Staley, D. D., Waterman, E. L., & Wilbraham, A. C. (2000). Chemistry, Addison-Wesley. (5th ed.). Menlo Park, CA: Prentice Hall, pp. 437-468.
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