Covalent Bonding
The forces that hold two nonmetals together are called covalent bonds.
A covalent bond consists of an electron pair shared between two atoms.
Example: H2 H : H H-H
The stability of H2 depends on the internuclear distance of the H atoms and the corresponding energy.
The attractive forces that bring about stability in H2 include:
(1) The electrostatic energy of the system is lowered by locating two electrons between the two protons in H2. The attractive e- - p+
interactions are greater than the repulsive
p+ - p+ or e- - e interactions.
(2) The two 1s orbitals overlap creating a new bonding orbital with a greater volume for the electrons, which lowers the KE of the electron.
The above is an accurate view of electron density. Depth of color is used to indicate probability of finding an electron.
Lewis Structures; Octet Rule
G.N. Lewis was the first chemist to suggest that nonmetallic atoms can acquire a stable noble gas configuration by sharing electrons to form an electron-pair bond.
Example
H + F ------à H F
1s1 + 1s2, 2s2, 2p5
Lewis structures are used to show valence electrons (the outer orbit electrons that are involved in covalent bonding) as dots. Lewis structures do not show core electrons.
X = symbol for the element
∙ = valence electron
to the right of X = s orbital
above, below, and to the left of X = p orbitals
Examples
Cl B Xe
For main group elements, the number of valence electrons is equal to the last digit of the group number in the periodic table (group 18 = 8 valence electrons, group 17 = 7 valence electrons). If the traditional numbering system is used it is equal to the group # (group VIIA = 7 valence e-).
For a molecule or a polyatomic ion, valence electrons usually occur in pairs. There are 2 kinds of electron pairs:
(1) A pair of electrons shared between two atoms is a covalent bond, shown as a straight line between bonded atoms.
(2) An unshared pair (lone pair) of electrons is owned entirely by one atom and is shown as a pair of dots.
Examples
OH- H2O NH4+
Bonded atoms can share one pair of electrons (single bond) or more.
A double bond occurs when bonded atoms share two pairs of electrons. A triple bond involves sharing three pairs of electrons.
Examples
C2H4 C2H2
The octet rule states that atoms bonding covalently tend to have noble-gas electron configurations.
Nonmetals except hydrogen (duet) achieve this by sharing in an octet (eight) of electrons.
Rules – Writing Lewis Structures (Molecules and Polyatomic Ions)
(1) Count the total number of electrons. Don’t forget to add electrons (equal to charge) for anions and subtract electrons for cations. Brackets with a charge outside should be added to the final structure for all ions.
(2) Draw a skeleton structure for the species. The central atom is usually the first one written in the formula and is placed at the center of the molecule or ion.
(3) Determine the total number of electrons needed to satisfy the octet (or duet for H) rule without sharing.
(4) # of e- needed w/o sharing - # of e- you have = total number shared
(5) Place the shared electrons between the atoms in the structure. Use remaining electrons to complete the octet.
Example
Draw Lewis structures of (a) OCl- and (b) C2H6
Example
Draw Lewis Structures for (a) SO2 and (b) N2
Resonance Forms
In some cases a Lewis structure does not adequately describe the properties of the ion or molecule it represents.
If a Lewis structure implies that there are two or more different kinds of bonds possible (when there is only one kind of bond) resonance structures are used to show that the true form is an intermediate between the alternate forms.
Resonance forms are separated by double-headed arrows.
Resonance can be anticipated any time it is possible to write 2 or more Lewis structures that are plausible without changing the arrangement of atoms (only the distribution of electrons changes).
Examples
Example
Write two resonance structures for the NO2- ion.
Formal Charge
The concept of formal charge can be used to determine which Lewis structure is correct when two structures differ in their arrangement of atoms.
Example
CH4O
The formal charge is the difference between the number of valence electrons in the free atom and the number assigned (unshared electrons owned by the atom and ½ the bonding electrons shared by the atom) to that atom in the Lewis structure.
Cf = X – (Y + Z/2)
The more likely Lewis structure is the one in which:
(1) the formal charges are as close to zero as possible
(2) any negative formal charge is located on the most strongly electronegative atom.
Example
Calculate the formal charges of C and O in the two structures for methyl alcohol.
Exceptions to the Octet Rule
(1) Electron-Deficient Molecules
For some odd electron species (free radicals) it is impossible to write a Lewis structure that follows the octet rule. They contain an odd number of valence electrons.
Examples
NO (11 valence e-) and NO2 (17 valence e-).
In some cases the central atom may have fewer than 4 electron pairs.
Examples
BeF2 and BF3
(2) Expanded Octets
Some molecules violate the octet rule by having a central atom with more than 4 pairs of valence electrons.
Examples
PCl5 and SF6
In most cases the terminal atoms are halogens (sometimes oxygen). The central atom is a nonmetal in the third, fourth, or fifth period of the periodic table (P, As, Sb, S, Se, Te, Cl, Br, I, Kr, Xe).
All of these atoms have d orbitals available for bonding.
If you find that you have an excess of electrons when drawing a Lewis structure, an expanded octet is involved. The extra electrons (2 or 4) are distributed around the central atom as unshared pairs.
Example
Draw a Lewis structure for XeF4.
Molecular Geometry
The geometry of a diatomic molecule is linear.
Cl – Cl or H – Cl
When molecules contain 3 or more atoms, the geometry is not obvious and the angles between bonds (bond angles) must be considered.
Molecular geometry can be predicted based on electron pair repulsion.
The VSERP (valence shell electron pair repulsion) model states that the valence electron pairs surrounding an atom repel each other. In turn, the orbitals housing those electron pairs are oriented as far apart as possible.
Molecular Geometries with Two to Six Electron Pairs on the Central Atom
In many cases, species have a central atom (A) bonded to 2-6 terminal atoms (X) with single bonds and there are no unshared pairs around A.
If a species has one or more unshared pairs of electrons around the central atom the electron pair geometry is generally the same as when only single bonds are involved, but bond angles are generally smaller.
The molecular geometry (which refers to the positions of bonded atoms only) is quite different when one or more unshared pairs are present. The locations of unshared pairs are not specified.
In general, unshared pairs will occupy slightly more space than bonding pairs. This forces the bonding pairs closer together and reduces the bond angle.
Geometries can be predicted using the VSEPR model for expanded-octet molecules.
The model can also be expanded to cases where double and triple bonds exist. For molecular geometry, a multiple bond (which occupies the same region of space as a single bond) behaves like a single bond.
The VSEPR model also applies where there is no single central atom. Example: H – C = C - H
To determine geometry:
(1) Count the number of terminal atoms bonded to the central atom (single, double, or triple bonds do not matter).
(2) Count the number of unshared pairs, E, of electrons around the central atom.
Type / Shape / Geometry† / Geometry‡ / ExamplesAX1 / Linear / / / HF, O2
AX2 / Linear / / / BeCl2, HgCl2, CO2
AX2E1 / Bent / / / NO2−, SO2, O3
AX2E2 / Bent / / / H2O, OF2
AX2E3 / Linear / / / XeF2, I3−
AX3 / Trigonal or triangular planar / / / BF3, CO32−, NO3−, SO3
AX3E1 / Trigonal or triangular pyramidal / / / NH3, PCl3
AX3E2 / T-shaped / / / ClF3, BrF3
AX4 / Tetrahedral / / / CH4, PO43−, SO42−, ClO4−
AX4E1 / See-saw / / / SF4
AX4E2 / Square Planar / / / XeF4
AX5 / Trigonal or triangular Bipyramidal / / / PCl5
AX5E1 / Square Pyramidal / / / ClF5, BrF5
AX6 / Octahedral / / / SF6
Example
Predict the geometry of (a) NH4+ (b) GeF2 (c)PF3
Example
Predict the geometries of the following given the Lewis structures.
(a) (b) (c)
Polarity of Molecules
Covalent bonds and molecules held together by such bonds may be:
(1) Polar – An unsymmetrical distribution of electrons causes the bond or molecule to contain a positive and negative pole (dipole).
(2) Nonpolar - A symmetrical distribution produces a bond or molecule with no + or – poles.
Nonpolar covalent bonds are formed when two atoms are joined identically. Ex – H2 or F2
Polar covalent bonds occur when the density of the electron cloud is greater about one atom (unsymmetrical density). Ex – HF
The degree of polarity is related to the differences in electronegativities of the bonded atoms. The larger the difference in EN the greater the polarity.
A polar molecule is one that contains positive and negative poles. The positive pole on a polar molecule will align itself with an external negative charge. A nonpolar molecule does not.
The extent to which molecules align themselves in an electrical field is a measure of their dipole moment.
Nonpolar molecules have a dipole moment of zero.
Diatomic molecules have the same polarity as their bond.
If a molecule contains more than two atoms you must consider bond polarity and molecular geometry.
If the polar A-X bonds in a molecule AXmEn are arranged symmetrically around the central atom A, the molecule is nonpolar.
Examples
H ─│------à F
The arrow points to the more negative end.
CCl4 & BeF2 – nonpolar H2O & CH3Cl - polar
Example
Determine whether each of the following is polar or nonpolar (a) SO2 (b) BF3 (c) CO2
Example
Match:
(1) CO2
(2) CH2Cl2
(3) XeF2
(4) BF3
(e) polar, bent
(f) nonpolar, triangular planar
(g) nonpolar, linear
(h) nonpolar, triangular pyramid
(i) polar, tetrahedral
(j) polar, triangular pyramid
(1) (2) (3) (4)
Atomic Orbitals and Hybridization
Linus Pauling developed the atomic orbital or valence bond model, which states that a covalent bond consists of a pair of electrons of opposed spin within an orbital.
According to this model, atoms should form a number of bonds equal to the number of unpaired electrons. This works for H, the halogens, and noble gases.
To explain other cases (like Be, C, or B) the valence bond theory must be modified to include the concept of hybrid orbitals.
In order for hybrid orbitals to form, orbitals in isolated atoms are mixed (hybridized) as the atoms approach.
1 s (atomic orbital) + 1 p (atomic orbital) ---à 2 sp (hybrid orbitals) (AX2)
1 s + 2 p --à 3 sp2 (AX3, AX2E)
1 s + 3 p --à 4 sp3 (AX4, AX3E, AX2E2)
Unshared as well as shared electron pairs can be located in hybrid orbitals.
The extra electron pairs in an expanded octet are accommodated by using d orbitals.
1 s + 3 p + 1 d ----à 5 sp3d (AX5, AX4E, AX3E2)
1 s + 3 p + 2 d ----à 6 sp3d2 (AX6, AX5E, AX4E2)
The number of hybrid orbitals formed is always equal to the number of atomic orbitals mixed.
Example
Give the hybridization of (a) C in CH3Cl (b) P in PH3 (c) S in SF4
Multiple Bonds
The extra electron pairs in a multiple bond are not located in hybrid orbitals.
A sigma (σ) bond consists of a hybridized electron pair occupying a sigma bonding orbital (one lobe).
A ------B
All single bonds are sigma bonds.
A pi (π) bond consists of unhybridized electron pairs (associated with multiple bonds) occupying pi bonding orbitals (two lobes).
A ------B
One of the electron pairs in a multiple bond is a sigma bond; the others are pi bonds.
Example
State the hybridization of N in (a) NH3 (b) NO2-
(c) N2
Example
Give the number of pi and sigma bonds in (a) NH3 (b) NO2- (c) N2