Honors Chemistry
Chapter 8: Covalent Bonding
Molecular Compounds
• Molecular or covalent compounds are those where electrons are
shared between atoms, so as for all atoms to “see” the electrons they
need to achieve a noble gas configuration
• The diatomic elements are a good set of examples of covalent
bonding, with 5 having single bonds (hydrogen, fluorine, chlorine,
bromine, iodine), 1 having a double bond (oxygen), and 1 having a
triple bond (nitrogen)
• These compounds tend to have relatively lower MPs and BPs than
ionic compounds
• Molecular formula: the number of atoms in a molecule, such as two
hydrogens and one oxygen in a molecule of water!
The Nature of Covalent Bonding
• Octet rule: again generally similar to ionic bonding…noble gas
configuration; useful but may be violated…more later
• Single covalent bonds: two electrons shared
• Does not include any unshared pairs of electrons…these remain on the atom
where they originated…such as the unshared pair (aka lone pair) on nitrogen
in ammonia (NH3) or the two lone pairs on the oxygen in water
• How to draw: I like Hank Green’s method…lines for bonds, pairs of dots for
lone pairs…more distinctive about what is going on, but I won’t fault any of
you if you use only pairs of dots or only lines…main point is to know where
the electrons are, right?
The Nature of Covalent Bonding, continued
• Double and Triple Covalent Bonds
• Yes, we are trying to follow the octet rule for satisfying the electron needs of
atoms. As we have learned about models, it may not always be correct…our
book calls us on one: the oxygen model. If you think in terms of a double
bond between the two atoms, that is wrong but okay in terms of knowing
that they are sharing the electrons. If you want the right answer, visit this
site: http://courses.chem.psu.edu/chem210/mol-gallery/oxygen/oxygen.html
but expect to not really understand what they are saying…it is well beyond
the scope of Honors Chem in NC. Suffice it to say, as we learned from Hank
Green, that models do fail, but are still useful. For nitrogen molecules, the
triple bond follows the rule well and accounts for nitrogen’s relative
“inertness”
The Nature of Covalent Bonding, continued
• Coordinate Covalent Bonds
• Same as covalent bonds, in general, but that both electrons shared originated
at one of the bonding atoms
• An example I may have already drawn is the ammonium ion, whereby a hydrogen ion
attaches to the lone pair of the central nitrogen atom of an ammonia molecule, forming
the NH4+ ion…both electrons came from the nitrogen atom
• Covalent compounds
• Please look at page 224 for examples of some common compounds: H2O2,
SO2, SO3, NO, NO2, N2O, HCN, HF, HCl, and let’s not forget H20!
Bond Dissociation Energies
• This is the energy required to break the bond between two covalently
bonded atoms…the larger the amount of energy, the stronger the
bond was
• More on this later…or you can read in chapters 17-18 now, if it suits
you, but I think you would prefer waiting
Resonance
• Resonance is the concept where there may be several structures
capable of being drawn to represent the locations of the bonding
electrons, and since we cannot typically distinguish between them,
and since experimental evidence indicates that the actual bonds are
usually a hybrid of the two/more structures, we feel that the
resonance structure does a better job of representing what is
happening
• We usually use double headed arrows (↔) to indicate this type of
structure
• These can be either “straight line” molecules or cyclic ones, such as
benzene
Resonance: example: Carbonate ion CO3-2
Resonance: example: ozone O3
Resonance: example: benzene C6H6
Exceptions to Octet Rule
Just like irregular verbs in Spanish, here come some exceptions
• Duet rule: 2 electrons: H, He, Li, Be, B
• More than 8: P, S, Cl, Br, I, Kr, Xe
• Caveat: for this to occur, these atoms are the central atom of the molecule!
This is where the more unusual shapes are modeled: the trigonal bipyramid, octahedron,
pentagonal bipyramid.
Exceptions to Octet Rule, continued
• Yes, there are more than we would like…always keep an eye open to
lone pairs, too!
Bonding Theories
• Molecular orbitals
• This is the theory that assumes that bonding atoms have their atomic orbitals
overlap to form molecular orbitals…this is the new region of probability for
the electrons to inhabit. This means that it would be occupied by two
electrons.
• Sigma bonds: two orbitals combining to form a molecular orbital that is
symmetrical around the axis between the two atoms; this can be either with s
orbitals or p orbitals, but only one set of the p orbitals can form a sigma bond
• Pi bonds: p orbitals overlapping to form side-by-side pi molecular orbitals;
these are in addition to the sigma bond; there can be one or two pi bonds,
hence either a double or triple bond. Pi bonds have less overlap than sigma
bonds, so they tend to be weaker than sigma bonds (but, they are in addition,
so double bonds, overall, are stronger than single, triple stronger than
double)
VSEPR
• The theory that the repulsion between electron pairs is such that they
move as far apart in the molecules as possible, creating the
appropriate lattice of atoms for the molecule
• We see this reflected in the shapes defined by the number of electron
domains around the central atom
• These new orbitals formed are hybrids of the original orbitals present
in the central atoms, and the hybrids formed are dependent upon
how many are needed
• Key point: the new hybrid orbitals are equal in energy, creating a much more
stable compound
VSEPR
VSEPR: hybrid orbitals
• Two electron domains: sp hybrid
• examples: (draw on board)
• ethyne…sigma bond along axis, leaving two sets of p orbitals to form the two pi bonds,
making the triple bond complete
• Magnesium hydride…sigma bonds along axis of molecules, from Mg to each Hydrogen
VSEPR: hybrid orbitals
• Three electron domains: sp2 hybrids
• Examples (draw on board)
• Aluminum trihydride: 3 equivalent orbitals formed, for the trigonal planar molecule
• Ethene…sigma bond along axis, two sigma bonds to hydrogens, leaving one set of p
orbitals to form the one pi bond, making the double bond complete
VSEPR: hybrid orbitals
• Four electron domains: sp3 hybrids
• Examples (draw on board)
• Methane…4 equivalent orbitals formed, making 4 sigma bonds between the central
carbon and the 4 hydrogen atoms
Polar bonds and molecules
• Covalent bonds involve the sharing of electrons between atoms
• If both atoms are of the same element, their attraction for these electrons
will be equal
• If both atoms are different, they will have a somewhat different attraction,
resulting in some degree of polarization in the bond, referred to as a polar
covalent bond, or just polar bond
• This differential attraction is based upon electronegativity, a periodic trend
which we looked at before…the tendency for an atom of a given element to
have a greater attraction for electrons
Polar bonds and molecules, continued
• Reality check: can we have polar bonds and nonpolar molecules?
• Yes!
Polar bonds and molecules, continued
• Implications of polarized molecules, also referred to as dipoles (having two poles)
• Will orient when placed between charged plates
• Will orient when around other molecules/ionic compounds with some degree of polarity
• If the “solvent,” may dissociate ionic compounds…think dissolution of ionic salts in water!
Electronegativity chart: refresher!
Intermolecular attractions
• Subtitle: why do some substances act like they do?
• Intermolecular attractions…gee, attractions between molecules of the same
substance…weaker than either ionic or covalent bonds
• Several types
• Van der Waals forces
• Dipole interactions: slight attractions of partial (+) with partial (-) regions of polar molecules…similar but much
weaker than ionic bonds
• Dispersion forces (aka London dispersion forces): momentary dipole interactions, transient, caused by electron
movement within the molecules; strength dependent upon number of electrons, which is why iodine and
bromine show more cohesion than chlorine or fluorine
• Hydrogen bonds (warning…misnomer…not really a bond like ionic or covalent!)
• Attractions between the hydrogens bonded to one molecule with the lone pairs of another molecule…occurs
because the valence electrons are not shielded by other levels of electrons
• Approximately 5% of the strength of an average covalent bond…the strongest of the intermolecular attractions
• Think about how water tries to form a sphere against gravity…yes, this is very significant!
Intermolecular attractions: implications
• Molecular compounds’ physical properties are largely dependent
upon the intermolecular attractions
• As we have spoken, melting and boiling points of molecular compounds are
lower than those of ionic compounds
• A few exceptions: carbon: diamond! Tetrahedral lattice of bonds…very strong
• Silicon carbide: same deal, two very similar atoms