Chemical Bonding

Chemical Bonding

Bonds form in 2 main ways atoms share electrons electrons are transferred

between atoms

Type of bond depends on the atom’s electronegativity and electron configuration

3 Main Types of Bonds

1. Ionic Bonds electrostatic force atoms transfer e- to become ions usually between a metal and a

nonmetal

2. Covalent Bonds electrons are shared by atoms usually between two nonmetals

3. Metallic Bonds forces that hold metals

together metals have many freely

moving electrons that attract positive metal ions

Ex. What type of bonding would exist in solid aluminum?

Ionic Bonding

valence electrons: outermost s and p electrons of an atom

Dot Diagrams: show valence electronsExamples:

isoelectronic: having the same electron configuration

Characteristics of Ionic Compounds high melting point able to conduct electricity in

molten state tend to be water soluble crystallize in definite patterns

(crystal lattice)

A closer look at ionic bonding

Naming BINARY Ionic Compounds name the metal first – do not

change ending name nonmetal second – change

ending to –ide Examples:

Writing Formulas for Ionic Compounds Sum of all ion charges MUST equal

ZERO! Use the “criss-cross” method Examples:

Covalent Bonding

molecule: name for a covalently bonded particle

Characteristics of Covalent Compounds low melting point do not conduct electricity usually brittle solids, liquids, or

gases

A closer look at covalent bonding

Naming Covalent Molecules

make sure the bond is covalent (usually 2 nonmetals)

first element’s name does not change second element’s ending becomes –ide Use Greek prefixes to indicate the # of atoms of

each element

mono = 1

di = 2

tri = 3

tetra = 4

penta = 5

hexa = 6

hepta = 7

octa = 8

Writing Formulas for Covalent Molecules Prefixes tell the # of atoms of each

element Examples:

Molecular Geometry

VSEPR Theory

Lewis Structures

Show arrangement of atoms in molecules

Show shared (bonding) and free electrons

Drawing Lewis StructuresUsed for covalently bonded molecules

ONLY!1. Determine the atoms in the molecule2. Count valence electrons for each atom.3. Find total # of valence electrons4. Arrange atoms in skeleton structure.

*Least electronegative atom in center!*5. Add electrons to structure.

The number of covalent bonds normally formed by an atom in a Lewis structure depends on its group in the periodic table. H is expected to form one bond.

F, Cl, Br, I, all in group 17 are expected to form one bond each.

O, S, Se, in group 16, are expected to form two bonds each.

N, P, As, in group 15, are expected to form three bonds each.

C, Si, Ge, in group 14 are expected to form four bonds each.

Octet Rule Atoms try to achieve Noble Gas

configuration (8 outer e-) Hydrogen – forms “duet” instead Some atoms exceed octet – more

than 8 bonding e-

VSEPR Theory

Valence Shell Electron Pair Repulsion Theory

electron groups arranged to minimize repulsion

Molecular Shapes FILL IN SHAPE CHART!

Show relative positions of atomic nuclei MUST determine Lewis structure to

determine shape!

Resonance Equivalent Lewis structures Shows possible locations of double

bonds

Resonance

Polarity of Bonds polar: having opposite ends

polar bond: e- shared unequally caused by difference in electronegativity

nonpolar bond: e- shared equally

ALL COVALENT BONDS ARE POLAR EXCEPT:

1. C – H 2. any atom bonded to

itself

A closer look at polar bonds

Polarity of Molecules Bonds must be polar for molecule to

be polar.

Molecule must have a definite top and bottom with opposite charges in order to be polar.

Intermolecular Forces (Weak Bonds)

Three main types1. Dispersion forces (London, van

der Waals)2. Dipole-Dipole Interactions3. Hydrogen bonding

Dispersion Forces (van der Waals) Very weak Between nonpolar molecules Induces momentary (temporary)

dipole Ex. – occurs in Cl2, CO2, CH4, etc.

Dipole-Dipole Interactions Stronger than dispersion Occur between

molecules with permanent dipoles (aka – polar)

Partially + end of one molecule attracted to partially – end of another

Hydrogen bonding Stronger type of dipole-dipole

interactions Results from H being covalently

bonded to either F, O, or N Stronger because…

H is so small F, O, & N are very EN Partial +/- charges are more

intense

H-bonds > Dipole-dipole > Dispersion

Affect BP, MP, solubility More E required to boil/melt

substances w/ stronger intermolecular forces Why?

Intermolecular Forces

Pop Quiz1. Name CaCO3

2. Write a formula for sodium sulfite.

3. Draw a dot diagram for boron (B).

4. How many valence electrons does carbon have?

5. What is the oxidation number of potassium?