Chapter 9

Covalent Bonding

Covalent bond

• Sharing of electrons– Nonmetal- nonmetal– electronegativity

difference less than 1.7

Molecule

• Base unit of a covalent bond

Diatomic molecule

• Naturally formed 2 atom groups – H2, N2, 02, F2, Cl2, Br2, I2

2 types of covalent bonds

• Polar covalent- bonding electrons are shared unequally

• Nonpolar covalent- bonding electrons are shared equally– Balanced distribution of charge

How do you tell the difference in polar and nonpolar?

• If the 2 atoms have a difference in electronegativity the bond is polar, if no difference then the bond is nonpolar

• If polar bonds draw the electron to one side of the molecule and the other side has none the molecule is polar

Bond length

• Distance between two bonded atoms at their minimum potential energy– Average distance between bonded atoms

Bond energy

• Energy required to break a bond

Coordinate covalent bond

• One atom donates both electrons to the bond

Nomenclature for covalent bonds

• (nonmetal- nonmetal)

• 2 systems: – Stock- name(+), (+) oxidation # in roman

numerals, name(-)– Classical- use prefixes except for a single (+)

ion • *Never use mono first*

Prefixes

• Mono• Di • Tri• Tetra• Penta• Hexa• Hepto• Octa• Nano• Deca

Naming acids

• Binary: hydro + root + ic acid

• Ternary: polyatomic, drop –ate, add –ic acid

Oxyacids

• +1 0 per- ate per- ic

• Memory -ate - ic• -1 0 -ite -ous• -2 0 - hypo-ite

hypo-ous

Lewis structures

• Use electron dot diagrams to show bonding and electron arrangement

Symbols uses in molecular structural formulas

unshared pair- (lone pair) pair of electrons that is not involved in bonding- belongs to one atoms

single bond- 2e- shared

double bond- 4e- shared

triple bond- 6e- shared

Structural formula

• Indicates kind, number, arrangement, and bond type in a molecule

Sigma bond

• - electrons are shared along the bond axis

Pi bond

– electrons are shared above and below axis

Resonance

• More than one valid Lewis structure can be drawn

VSEPR theory

• Valance shell electron pair repulsion– Model for molecular geometry– Bond angles– Arrangement minimizes repulsion of e- around

the central atom– Molecules adjust their shape, so that valence

e- are as far apart as possible

3 types of repulsion

1. Unshared-unshared

2. Unshared- shared

3. Shared-shared

hybridization

• A process in which atomic orbitals are mixed to form new identical hybrid orbitals

Hybrid orbitals

• Orbital of equal energy produced by the combination of two or more atomic orbitals

• (sp, sp2, sp3)

Intramolecular forces

• Forces within a molecule that hold atoms together

Intermolecular forces

• Forces of attraction between molecules

(Van der Waals)

Dipole

• Molecule that has two poles (polar)

Dipole moment

• Measure of the strength of the dipole and is a property that results from the asymmetrical charge distribution in a polar molecule- depends on size and distance Qd

Van der Waals forces

• Groups of intermolecular forces

Dipole- dipole force

• Between polar molecules

Induce dipole

• A normally nonpolar molecule is transformed into a dipole

Hydrogen bonding

• Intermoleculer (Van der Waals) force in which H bonded to a highly electronegative atom is attracted to an unshared pair of an electronegative atom in a nearby molecule

London Dispersion forces

• (dispersion force) result from the constant motion of e- ’s and the creation of instanteous dipoles– Force generated in a temporary dipole

interaction– Most important Van der Waals force– Proposed by Fritz London in 1930– Strength increases with the number of e- in

the interacting atoms