COVALENT COMPOUNDS, FORMULAS, AND STRUCTURE
Covalent Molecules
1. Formation of ____________________ represents another way in which elements may
combine to form compounds and at the same time attain a noble gas electron
configuration.
2. Covalent bonding occurs when electrons are shared between two or more atoms.
3. In covalent compounds the atoms are physically attracted to one another. In ionic
compounds they are electrostatically attracted to one another.
4. Compounds composed of covalently bonded groups of atoms are called
___________________.
5. To show the sharing of electrons, chemists draw structures of covalent molecules using
_____________________________.
6. In the Lewis representation, the outermost s and p electrons (valence electrons) are shown
as dots arranged around the atomic symbol.
7. When electrons are shared between atoms, one atom donates one electron and the other
atom donates the second electron. The shared pair of electrons represents a
____________________.
8. Two shared pairs is a double bond and three shared pairs is a triple bond.
9. The atoms are trying to attain a noble gas electron configuration, which is to have eight
electrons in the valence shell. These eight electrons represent the octet of the
____________________.
10. In molecules, each atom considers the shared electrons as its own.
11. Magnetic properties of atoms, ions, and molecules arise from unpaired electrons.
12. The electron spin generates a small magnetic field. When electrons are paired, these
magnetic fields cancel. When unpaired, they impart magnetic properties to the substance.
13. Substances that have unpaired electrons are said to be ___________________. They are
attracted toward an external magnet, and the force of attraction is proportional to the
number of unpaired electrons.
14. ____________________ substances have all of their electrons paired, and they exhibit no
magnetic properties.
Lewis Structures of Molecules
1. Add the valence electrons together.
2. If the substance is a polyatomic ion, we must take into account the electrons used to form
the ion. For anions, the charge represents electrons that must be added to the total
valence electrons. For cations, the charge represents missing electrons that must be
subtracted from the valence electrons.
3. Determine the skeleton.
4. Carbon is usually a central atom I the structure. In compounds with more than one
carbon atom, the carbon atoms are joined in a chain to start the skeleton.
5. Hydrogen is never a central atom because it can form only one covalent bond.
6. Halogens form only a single bond when oxygen is not present, and therefore a halogen
will generally not be a central atom.
7. Oxygen forms only two covalent bonds and is rarely a central atom. However, it may
link to carbon atoms in a chain.
8. In simpler molecules, the atom that appears only once in the formula will often be the
central atom.
9. A pair of electrons is placed between two atoms in the skeletal structure to represent a
covalent bond. The pair is represented by a line. These electrons are called a
____________________.
10. The remaining electrons are used to complete the octets of all outer atoms in the skeleton.
These electrons are called ____________________ or ____________________.
11. If any electrons are left over, they are added in pairs to the central atom. These electrons
are also nonbonding, or lone, pairs.
12. When all electrons have been placed, the outer atoms will have octets.
13. If the central atom has an octet, the structure is complete.
14. It is all right if the central atom has fewer than eight electrons, provided that atom is
boron.
15. For other atoms, double or triple bonds may be needed to obtain an octet.
16. The central atom may have more than eight electrons only if it is from periods 3 – 7 of
the periodic table.
***** Draw the Lewis structure of the following molecules or ions.
(a) CH2Cl2
(b) NH3
(c) HClO4
(d) SF6
(e) NH4+
(f) SO42-
(g) CH3CN
(h) SO3
17. These three equally probable structures are called ____________________.
Lewis Structures of Odd Electron Compounds
1. Some compounds have formulas in which the total number of valence electrons is an odd
number.
2. One cannot construct a Lewis structure with an octet for each atom.
3. Molecules that have Lewis structures with an unpaired electron are often called
____________________. The unpaired electron makes the molecule unusually reactive.
Formal Charge
1. Experimental evidence, such as bond length, is the best verification of whether a Lewis
structure is reasonable.
2. Without experimental data, calculating ____________________ on each atom is one
technique that can be used.
3. The formal charge is the difference between the number of electrons an atom has in a
Lewis structure and its number of valence electrons.
4. Based on the electronegativities of atoms, the element with the greater electronegativity
will be the atom with a negative charge. If the formal charges show elements with large
electronegativities as positive compared to other atoms in the structure, we may question
the validity of the structure.
5. To calculate the formal charge on each atom in a Lewis structure, he following steps are
taken:
(a). For each atom count all electrons not used for bonding by the atom.
(b). Count half of the atom’s bonding electron.
(c). Add steps (a) and (b) to obtain the electrons assigned to that atom.
(d). Subtract the assigned electrons from the valence electrons to obtain the formal charge.
6. The formal charge calculations may be quickly checked since the formal charges on the
atoms in a molecule must add up to zero, obeying the law of electroneutrality.
7. For a polyatomic ion, the formal charges must add up to the charge on the ion.
8. A molecule or polyatomic ion with the lowest possible formal charge on each atom is
usually judged to be a more probable structure than the one where the formal charge is
larger.
***** Look at the sulfate ion:
These formal charges are large, so an alternative structure should be sought.
This is the preferred structure since it has the minimum formal charges.
9. Formal charges can be used to deduce the proper structure for a compound that has many
possible Lewis structures.
Resonance Structures
1. At times we can construct several Lewis structures for a substance that are totally
equivalent, even down to the formal charges on the atoms. These structures are called
____________________.
2. Most resonance structures are very similar for a given substance, usually differing only
in the geometry of the molecule or ion.
3. It is found experimentally that none of the resonance structures properly describes the
molecule.
4. The true properties of the substance are found to be a blend of all the resonance structures
together.
5.
In a variety of experiments it is found that all of the sulfur-oxygen bonds are identical.
The measured properties of these bonds indicate that they are neither purely single bonds
nor purely double bonds.
6. Benzene is an organic molecule that exhibits resonance.
Abbreviated as:
Covalent Bond Polarity and Electronegativity
1. Electrons are shared equally only in a covalent bond between two identical atoms.
2.
If the electrons are not shared equally by the two atoms, they will spend more time
localized near on e atom or the other. The result is that the atom that attracts the
electrons will be relatively more negative than the other atom.
3. The bond between the atoms is __________ with a positive end and a negative end.
4. The concept of ____________________ was developed to numerically represent the
ability of an atom to attract electrons.
5. Electronegativity increases from left to right across a period and decreases from top to
bottom down a group.
6. Electronegativity increases from the lower left corner of the periodic table to the upper
right corner. This is known as a ____________________.
7. Using trend information, one can determine bond polarity.
8. Polarities are indicated by the symbols _____ and _____ for partially positive and
partially negative.
***** Indicate the positive and negative ends of each of the following bonds:
S–O
Si - O
C–N
H - Br
S–P
H-O
C–F
9. The electronegativity table can be used to evaluate the magnitude of bond polarity.
The larger the value of ∆EN, the greater the polarity of the bond.
10. Without the electronegativity table, it is possible to determine which of two bonds is the
more polar if the bonds have one atom in common. The more polar bond will be the one
where the second atom is furthest in the periodic table from the common atom.
***** For each of the following pairs determine which bond is more polar.
C – N or C – O
H – Cl or H – S
S – O or S – Cl
H – N or H – O
P – Br or S – Br
Dipole Moments
1. A better measure of how polar a bond is to use the ____________________.
2. The dipole moment is a measure of the difference in charge, q, on two covalently bonded
atoms and the distance, r, between the two nuclei.
3. The units for dipole moment are Coulomb – meters, and a common unit for dipole
moment is the debeye.
4. Dipole moments a mathematically treated as vectors.
Bond Order
1. ____________________ is a term that refers to the average number of bonds that an
atom makes in all of its bonds to other atoms.
***** Determine the bond order of the central atom in each of the following compounds:
(a). CH3Cl
(b). CS2
(c). SO3 resonance structure.
2. Bond order is an index of bond strength. The larger the bond order, the stronger the
bond.
MOLECULAR GEOMETRY
1. Once a valid Lewis structure has been determined, the overall geometry of a simple
molecule with one central atom can be established.
2. The overall geometry of a molecule is extremely important in understanding the
properties of chemical compounds.
3. The Valence – Shell electron – Pair Repulsion (VSEPR) Theory allows us to determine
the three-dimensional shapes of covalently bonded molecules with a minimum of
information.
4. This theory states that the geometry around each atom will depend on the repulsion of the
valence-shell electrons (bonding electrons and nonbonding electrons) away from each
other.
Structures
1. For basic structures, to determine the three dimensional geometry around a central atom
_____ all one needs to know is how many atoms, _____, are covalently bonded to it.
2. Nonbonding electron pairs, _____, on the central atom take up space just as an atom
does. A nonbonding pair actually takes up more space.
3. Nonbonding pairs cannot be seen on x-ray crystallography. We can only see the result of
the repulsions.
4. To determine the shape of the molecule, one must draw the Lewis structure and then
count the total electron pairs, the bonding electron pairs, and the nonbonding (lone) pairs
attached to the central atom.
Structure
Notation
Total
electron
pairs
AX2
2
Bonding
Pairs
2
Lone
Pairs
Geometry
Hybridization Example
0
Linear
sp
CO2
sp2
BF3
AX3
3
3
0
Trigonal
planar
AX2E
3
2
1
Bent
sp2
NO2-
AX4
4
4
0
Tetrahedral
sp3
CH4
AX3E
4
3
1
Trigonal
pyramidal
sp3
NH3
AX2E2
4
2
2
Bent
sp3
H2O
AX5
5
5
0
Trigonal
bipyramidal
sp3d
PCl5
AX4E
5
4
1
Seesaw
sp3d
SF4
AX3E2
5
3
2
T-shape
sp3d
ClF3
AX2E3
5
2
3
Linear
sp3d
XeF2
AX6
6
6
0
Octahedral
sp3d2
SF6
sp3d2
BrF5
sp3d2
XeF4
AX5E
6
5
1
Square
pyramidal
AX4E2
6
4
2
Square
planar
Complex Structures
1. Geometries of more complex molecules are constructed by determining the geometry
around each atom in sequence and then stringing the geometries together.
***** Predict the geometry around each of the carbon atoms in this molecule:
Molecular Polarity
1. Bond polarities depend on the electronegativities of the two elements bonded together.
2. For most molecules, more than one bond must be considered in determining the polarity
of the molecule as a whole.
3. Even if a bond is polar, the molecule as a whole may or may not be polar.
4. Four general rules for determining molecular polarity are:
(a). A molecule that is symmetrical is nonpolar. It does not matter how polar the individual
bonds are.
(b). A nonsymmetrical bond is polar if the bonds are polar.
(c). A molecule with more than one type of atom attached to the central atom is often
nonsymmetrical and therefore polar.
(d). A central atom with nonbonding electron pairs is often nonsymmetrical and polar.
***** Draw the Lewis structure and then predict the geometry and the polarity of the
molecule or ion.
(a). IF2+
(b). OSF4
(c). SiF62-
(d). IF5
COVALENT BOND FORMATION
Wave Mechanics and Covalent Bond Formation
1. The valence bond theory considers a covalent bond to be the overlapping of two atomic
orbitals when the electron spins are paired.
2. The molecular orbital theory considers that a molecule is similar to an atom in that they
both have distinct energy levels that can be populated by electrons.
Valence Bond Theory
1.
In VB theory two atoms approach one another and interact with an overlap of atomic
orbitals with electrons of opposing spin. As the bond is formed, the electrons spread out
over the molecule to form the final electron cloud surrounding the nuclei.
2. The overlap of two s orbitals forms a sigma (σ) bond.
3. Sigma bonds may also be formed by the overlap of an s orbital and a p orbital, or by the
overlap of two p orbitals.
Orbital Overlap Model (pi Bonds)
1. Every covalent bond has one, and only one, sigma bond. If a compound has either a
double or triple bond, then additional overlap of orbitals is needed.
2. Such a bond, called a pi (π) bond, is formed by the sideways overlap of two p orbitals.
Its electron density is arranged in two clouds, one above and one below the internuclear axis.
3. A double bond involves one sigma and one pi bond. A triple bond involves one sigma
and two pi bonds.
4. Al of the bonds are arranged so that their electron clouds do not interfere with each
other.
Hybrid Orbital Model
1.
If all covalent molecules were formed from the overlap of s and p orbitals, we would
expect all covalent molecules to have 900 bond angles. This is hardly the case.
2. Even a simple molecule, such as methane, cannot be explained adequately by the overlap
model.
sp3 Hybrid Orbitals
1. A ____________________ may be defined as a set of orbitals with identical properties
formed from the combination of two or more different orbitals with different energies.
2.
In an sp3 hybrid, one s and three p orbitals combine to form four hybrid orbitals.
3. Any molecule whose basic structure is the tetrahedron will have sp3 hybrid orbitals. This
includes:
sp2 Hybrid Orbitals
1.
It looks like:
2.
It is usually associated with trigonal planar shapes and double bonds.
sp Hybrid Orbitals
1.
It looks like:
2. Usually associated with linear geometries, double bonds such as those in CO2, and triple
bonds such as those in CN-.
sp3d Hybrid Orbitals
1.
It looks like:
2. The five electrons in the sp3d hybrid will form five sigma bonds in a covalent compound.
The basic structure is trigonal bipyramidal. It also includes seesaw, T-shape, and linear.
sp3d2 Hybrid Orbitals
1.
It looks like:
2. The hybrid allows six covalent bonds to form. The basic structure is the octahedron and
also includes square pyramidal and square planar.