1. Unit 3.3: Covalent Bonds and Intermolecular Forces
Keep until June 2011!
Unit 3.3: Covalent Bonds and Intermolecular Forces

2. Vocabulary:
Covalent bond: “co-” = with; “-valent-” = valence (electrons); bond formed when two atoms share valence electrons
Molecule: a neutral group of atoms joined by covalent bonds.
Molecular formula: a chemical formula that shows the actual number of each type of atom that participates in the compound.
Structural formula: a drawing that represents the bonds of a molecular compound as lines connecting the atoms.
Nonpolar: a molecule that has no regions of partial charge
Polar: a molecule that has regions of partial charge
Dipole: attraction between two polar molecules
Hydrogen bond: bond formed between the slightly positively charged hydrogen on one compound to an unshared pair of electrons on a separate molecule

3. I. Molecular Compounds
A molecule (molecular compound) is a neutral group of atoms joined by covalent bonds.
A diatomic molecule is composed of two atoms of the same element sharing valence electrons.
Br I N Cl H O F series…
A molecular compound is a compound formed when two or more different elements share valence electrons.
Often, refer to both as simply “molecules”; ex: a molecule of oxygen, a molecule of water

4. The molecular formula of a compound is the chemical formula that shows the numbers of each type of atom participating in the compound.
*Subscripts are not necessarily the smallest whole number ratio like ionic compounds
Ex: C6H12O6: 6 atoms of C, 12 atoms of H, and 6 atoms of O, not CH2O
A molecular formula does not show arrangement of atoms.

5. The structural formula of a compound shows the arrangement of the atoms that are participating in the compound in relation to each other.

6. II. Covalent Bonds
A covalent bond forms when two atoms share valence electrons to achieve the stability of a noble gas
“co-” = with; “valent” refers to valence electrons
Satisfies the octet rule for both atoms
(Ionic bonds involve the transfer of an electron between atoms, creating charged atoms that bond through electrostatic attraction.)

7. **Covalent bonds usually form between nonmetals in groups 4A7A**

8. When atoms share a single pair of electrons, forms a single covalent bond:
Represented by a single line in a structural formula
Or by a pair of dots in an electron dot structure

9. When atoms share two pairs of electrons, forms a double covalent bond:
Represented by a double line in a structural formula
(Or by 2 pairs of dots in an electron dot structure)
Stronger than a single bond, pulls atoms closer together b/c greater interaction between nuclei

10. When atoms share three pairs of electrons, forms a triple covalent bond:
Represented by a triple line in a structural formula
(Or by 3 pairs of dots in an electron dot structure)
Even stronger than a double bond, pulls atoms even closer together b/c greater interaction between nuclei

11. Resonance: when there is more than one possible valid dot structure for a compound, the true structure “resonates” between all valid possibilities
Ex: In a molecule of ozone (O3), a central oxygen molecule shares a single bond with one oxygen but a double bond with a third oxygen. The double bond is constantly shifting between both external oxygen molecules.

12. III. Polarity of Bonds
Nonpolar covalent bond: occurs when covalent electrons are shared equally in a bond
All diatomic molecules are nonpolar (H2, O2, Cl2, etc…)
Some compounds are nonpolar as well: CO2, CH4

13. Polar covalent bonds: electrons are not shared equally between atoms, based on electronegativity
Electrons are “pulled” towards more “electronegative” elements, which become slightly negatively charged.
Electrons are “pulled” away from less “electronegative” elements, which become slightly positively charged.
Partial charges in a molecule are denoted with Greek lower-case “delta” (δ)
Can also represent polarity by an arrow pointing to the more electronegative element in a polar molecule

14. Electronegativity predicts which bond will form:
If the difference in electronegativity between the two elements is:
0.0 to 0.4, the bond is nonpolar covalent
0.41 to 1.0, the bond is moderately polar covalent
, the bond is very polar covalent
≥ 2.01, the bond is ionic

15. Properties of polar molecules:
Compounds with polar bonds are often polar (but not always), have two poles, so they are called “dipoles”
Analogy: north/south poles of a magnet – two poles
B/c they have regions of charge, they can interact with other charged particles.
But if the bonding atoms lie in the same plane, the polar bonds can “cancel” each other out, so the compound itself is nonpolar.
Ex: CCl4: Cl is more electronegative, pulls e-’s away from C, but in opposite directions equally
Keep until June 2011! Z

16. IV. Intermolecular Forces (forces between molecules)
Van der Waals forces: weakest forces, two different types:
Dispersion forces: can even occur between nonpolar molecules b/c as electrons move, they may induce a momentary region of charge that can attract a nearby molecule
Sometimes called London forces
Dipole-dipole interactions: polar molecules are attracted to each other b/c of opposite partial charges
Similar to ionic interactions, but much weaker
Dipole Force
Dispersion (London) Force

17. Hydrogen Bonds: dipole interactions that always involve hydrogen atoms
Hydrogen bonded to very electronegative atom can bond weakly to unshared pair of electrons of a separate molecule
Hydrogen bonds are responsible for keeping water a liquid at room temperature.
Hydrogen bonds are critical in many biological molecules
Hold the “rungs” of the DNA ladder together
Strongest of all intermolecular forces

18. Intermolecular forces
Comparing Bonds
Ionic > covalent > hydrogen bonds > dipole> dispersion forces
Intramolecular forces
Intermolecular forces
Characteristic
Ionic Compounds
Covalent Compounds
Representative Unit
Formula unit
Molecule
Bond formation
Transfer of one or more electrons between atoms
Sharing of electron pairs between atoms
Types of elements
Metals with nonmetals
Nonmetals only
Physical state of matter
Solid
Solid, liquid, or gas
Melting Point
High (usually greater than 300° C)
Low (usually below 300° C)
Solubility in water
Usually high
High to low
Electrical conductivity when dissolved in water
Good conductor
Poor or nonconducting

19. Electronegativity Values for Selected ElementsH
2.1 Li
1.0 Be
1.5 B
2.0 C
2.5 N
3.0 O
3.5 F
4.0 Na
0.9 Mg
1.2 Al Si
1.8 P S Cl K
0.8 Ca Ga
1.6 Ge As Se
2.4 Br
2.8 Rb Sr In
1.7 Sn Sb
1.9 Te I Cs
0.7 Ba Tl Pb Bi