7-1
COVALENT COMPOUNDS – ACIDS – MOECULAR GEOMETRY – INTERMOLECULAR FORCES
WHY DO ATOMS SHARE ELECTRONS? (pp. 359 - 363)
1. In the last unit we learned that some metals and nonmetals react to form binary ionic compounds. Electrons are transferred and the resulting ions have noble gas configurations. Compounds are then formed because the ions are attracted to one another.
2. Sometimes two atoms that both need to gain valence electrons to become stable have a similar attraction for electrons.
3. Sharing electrons is another way these atoms can acquire the electron configuration of a noble gas, even though it will be on a part-time basis.
4. In a ______, atoms do not lose or gain electrons. Instead, they share pairs of electrons to achieve stability, often by filling their outer energy levels so they have stable octets.
5. A ______is formed when two or more atoms bond covalently. They are often called ______.
Forces of electric attraction make a covalent bond
1. An attractive force exists between the outer electrons on one atom and the nucleus of a neighboring atom.
2. The force of attraction brings the atoms together until the force of repulsion between the nuclei and between the outer electrons forces the atoms apart.
3. If the forces of attraction are greater than the forces of repulsion, then a covalent bond forms between the atoms.
4. Besides the comparative strengths of the attractive and repulsive forces, another reason the attractive forces can be stronger is that a pair of electrons shared between atoms in a stable covalent bond have opposite spins and occupy less space than a pair of electrons in an orbital or only one atom.
5. The bond is not rigid. It is much like a spring where the atoms vibrate back and forth at some average distance where the attractive and repulsive forces are balanced.
Sharing More Than Two Electrons
1. Covalent bonds between atoms can involve sharing more than two electrons.
2. When a single pair of electrons (2 electrons) is shared; this is known as a ______.
3. When two pairs of electrons (4 electrons) are shared, this is known as a ______.
4. When three pairs of electrons (6 electrons) are shared, this is known as a ______.
Bond Length and Bond Energy are Inversely Related
1. The average distance that separates the atoms in a bond is known as the ______.
2. Bond lengths are never really fixed values because the atoms vibrate. They can also vary depending on the other bonds present in the molecule.
3. ______is the energy required to break a chemical bond to produce individual atoms, each keeping its own electrons.
4. Bond length and bond energy are inversely related.
5. A short bond length requires higher bond energy to break it while a long bond length requires less energy to break it.
Bond Properties
1. Few chemical bonds are either totally ______or totally ______.
2. The bonds in many compounds have some features of both types of bonds.
3. The electrons in a bond are not necessarily shared equally. To determine whether this uneven sharing will be very small or very large, one compares the ability of each atom to pull electrons toward itself.
4. This is called ______.
5. The electronegativity table is used to provide numbers for comparison.
6. The greater the differences in electronegativities between two atoms, the more unequal the sharing and the more ionic character the bond will have.
7. A covalent bond formed between two atoms with equally shared bonding electrons is said to be a ______.
8. Examples are:
9. When atoms of different elements bond, the sharing of electrons can never be truly equal.
10. When a covalent bond is formed between two atoms in which the bonding electrons are more strongly attracted to one atom over the other is said to be a ______.
11. Examples are:
12. To determine bond polarity, we make use of Pauling’s electronegativity values:
13. Look up the electronegativity value of each atom in the bond and then subtract the smaller value from the larger value.
14. The difference is always ______.
15. Use this table to determine the bond polarity:
16. The uneven sharing causes the more electronegative atom to have a partial negative charge while the less electronegative atom will have a partial positive charge.
***** Determine the bond polarity of the following bonds:
C – H
N – H
C - O
NAMING COVALENT COMPOUNDS (pp 94 - 97)
1. Naming covalent (molecular) compounds is similar to naming ionic compounds.
2. One can use either the Stock naming system or one that makes use of prefixes, roots, and suffixes.
3. The latter system is known simply as the prefix naming system.
4. The root comes from the name of the element and then prefixes and suffixes are added.
5. An example is:
CO
CO2
6. The first element named is usually the one with the lower electronegativity value.
7. If there is only one atom of the first element, then ______prefix is used.
8. The ending ______is used as it was in naming ionic compounds.
9. The common prefixes are:
***** Name the following compounds:
P2O5CCl4
As2O3SO2
NF3
10. The Stock system can be used to name covalent compounds.
P2O5CCl4NF3
As2O3SO2
Writing Formulas for Molecular Compounds
1. If the compound is named using the prefix system, simply translate the prefixes as written into subscripts.
***** Write the formula for each of the following molecular compounds:
dinitrogen tetroxidephosphorus trichloride
disulfur trioxide
NAMING ACIDS (pp 104 - 105)
1. Water solutions of some molecules are acidic and are named as acids.
2. A binary acid contains:
But, no ______.
3. When naming a binary acid, use the prefix ______to name the hydrogen part of the compound.
4. The rest of the name consists of a form of the root of the second element, or polyatomic ion, plus the suffix ______.
5. The word ______is then added.
HBr in a water solution is known as:
***** Name the following acids:
HFHCN
HIH2S
Naming Oxyacids
1. Any acid that contains hydrogen and an oxyanion is called an ______.
2. To name it, first identify the anion present.
3. The name of the oxyacid consists of the root of the anion, a suffix and the word acid.
4. If the anion suffix is ______, change it to ______.
5. If the anion suffix is ______, change it to ______.
Examples:
HNO2
HClO3
HClO4
HC2H3O2
6. Notice there is ______use of the hydro.
7. The following oxyacids were named before the rules went into effect, so they must be memorized:
H2SO3
H2SO4
H3PO4
MOLECULAR GEOMETRY – LEWIS STRUCTURES (pp. 371 - 381)
1. In order to predict the arrangement of atoms in a molecule, a model is used.
2. The nuclei and inner-shell electrons are represented by the element’s atomic symbol. The valence electrons are represented as dots placed around each side of the symbol, up to two per side.
3. Bonds between atoms are represented by either pairs of dots or by lines between the atoms involved in the bond.
4. Unshared pairs are represented by pairs of dots placed around the appropriate atoms.
DRAWING LEWIS STRUCTURES
1. Determine the total number of valence electrons in the compound by adding up all the valence electrons of the atoms in the compound.
2. Arrange all the element symbols according to which element can form more than one bond and those that can only form one bond.
** Atoms that can only form one bond are: H and F. Cl, Br, and I will normally form only one bond unless outnumbered by either O or F.
** Atoms that “love” being in the middle of things are: B, C, N, O, Si, P, S, As, Se, Sb
3. Draw single lines between all the atoms that are bonded together.
4. Count each line, multiply by two, and subtract that number from the total number of valence electrons.
5. This gives you the electrons left to distribute to all the elements still needing electrons so each can have an octet.
6. If there are not enough electrons available to distribute, then one or more of the single bonds may have to made into double or triple bonds.
***** Draw Lewis structure for the following molecules:
a. Iodomethane(CH3I)
b. Methanol, (CH3OH)
c. Dinitrogen difluoride (N2F2)
d. Formaldehyde (H2CO)
e. Hydroxide ion (OH-)
f. Ammonium ion (NH4+)
RESONANCE STRUCTURES
1. Sometimes a single Lewis structure is not enough to accurately depict a molecule.
2. When this happens, more than one equivalent structure is used to represent the molecule.
3. When more than one equivalent Lewis structure can be drawn for a molecule, the molecule is said to be a ______hybrid.
An example is SO3
All of these structures are equivalent and only the double bond moves to different oxygens around the sulfur.
***** Draw the three resonance structures for dinitrogen monoxide (N2O)
EXCEPTIONS TO THE OCTET RULE
1. There are three types of ions or molecules that do not follow the octet rule:
(a) Ions or molecules with an odd number of electrons (NO2)
(b) Ions or molecules with less that an octet around the central atom (BF3)
(c). Ions or molecules with more than eight valence electrons around the central atom (an expanded octet)
PCl5
SF6
IF5
2. It is thought that the extra electrons go into empty “d” orbitals, thus permitting the central atom to exceed the octet rule.
3. When it is necessary to exceed the octet rule for one of several third row (or higher) elements, assume the extra electrons should be placed on the central atom.
MOLECULAR GEOMETRY – DETERMINING THE SHAPE OF THE MOLECULE
(pp 381 - 392)
1. Shape is an important factor in determining the chemical properties of a molecule.
2. One example is the difference between normal, healthy red blood cells and the shape of a sickle cell.
3. The theory used to predict shapes is called the ______or (______) theory.
4. VSEPR is based on the idea that electron pairs surrounding the central atom will arrange themselves to be as far apart as possible.
5. Shapes cannot be predicted from the molecular formula. One must know:
(a)
(b)
(c)
6. Once the bonding pairs and nonbonding pairs attached to the central atom have been determined, use this chart to determine the molecular geometry (shape) of the molecule:
TOTAL PAIRS / BONDING PAIRS / NONBONDING PAIRS / MOLECULAR GEOMETRY / EXAMPLE2 / 2 / 0 / Linear / CO2
3 / 3 / 0 / Trigonal planar / BF3
3 / 2 / 1 / Bent / NO2-
4 / 4 / 0 / Tetrahedral / CH4
4 / 3 / 1 / Trigonal pyramidal / NH3
4 / 2 / 2 / Bent / H2O
5 / 5 / 0 / Trigonal bipyramidal / PCl5
5 / 4 / 1 / See-saw / SF4
5 / 3 / 2 / T-shaped / ClF3
5 / 2 / 3 / Linear / XeF2
6 / 6 / 0 / Octahedral / SF6
6 / 5 / 1 / Square pyramidal / BrF5
6 / 4 / 2 / Square planar / XeF4
***** Determine the shape of each of the following:
a. Ammonia (NH3)
b. Water (H2O)
c. Dichlorodifluoromethane (CF2Cl2)
d. Arsenic pentafluoride (AsF5)
e. Selenium hexabromide (SeBr6)
MOLECULAR DIPOLE
1. Remember that each bond within a molecule can be either nonpolar or polar.
2. If the bond is polar, then one end of the bond appears to have a slight positive charge while the other appears to have a slight negative charge.
3. This creates a dipole.
4. If the molecule has polar bonds and the shape of the molecule causes the polarities to cancel, then the molecule is ______.
5. If the molecule has polar bonds and the shape of the molecule does not cause the polarities to cancel, then the molecule will be ______.
6. If the molecule has all nonpolar bonds, then no matter what the shape is, the molecule will be ______.
***** Determine the polarity of each of the following molecules:
Ammonia (NH3)
Water (H2O)
Dichlordifluoromethane (CF2Cl2)
Arsenic pentafluoride (AsF5)
Selenium hexabromide (SeBr6)
INTERMOLECULAR FORCES (pp. 442 - 444)
1. All atoms and molecules attract each other. But, these forces of attraction are not as strong as the forces of attraction between atoms in the bonding process (ionic and covalent bonds).
2. ______are forces that cause attractions between molecules.
3. Intermolecular forces of attraction between molecules, or between atoms and molecules, do not involve either the transferring or sharing of electrons.
4. The weakest intermolecular forces are ______. These are the forces of attraction between nonpolar substances.
Examples are:Br2, I2, N2, Cl2, H2, O2, F2, CH4, and the noble gases
5. Next in increasing strength are ______. These intermolecular forces occur between polar molecules because there is a dipole in the molecule. The positive end of one molecule attracts the negative end of a nearby molecule.
Examples are: CO, NO
6. The last type of intermolecular force is known as ______. This especially strong force of attraction occurs between molecules containing hydrogen bonded directly to a highly electronegative atom such as ______, ______or ______.