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)
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:
Characteristics of Covalent Compounds low melting point do not conduct electricity usually brittle solids, liquids, or
gases
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
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!
Polarity of Bonds polar: having opposite ends
polar bond: e- shared unequally caused by difference in electronegativity
nonpolar bond: e- shared equally
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
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