Download Bonding Notes

Table of Contents
I. MetallicA. Definition: sea of mobile electrons
B. Properties
C. Crystal Lattices
1. BCC, FCC, HCP
2. Alloys
a. substitutional
b. interstitial
II. IonicA. Structures and Formulas
B. Assigning Oxidation Numbers
C. Polyatomic Ions
D. Naming Rules
1. binary
2. multiple ox #
3. PAI - Table E
4. Hydrates
E. Properties of ionic compounds
III. CovalentA. Polar covalent bonds (dipoles)
1. Polar molecules
B. Nonpolar covalent bonds
1. Diatomics
2. Nonpolar molecules from polar bonds (symmetrical)
C. Coordinate covalent bonds
D. Naming Rule
E. Properties of Molecular compounds
F. Network Solids
IV. Drawing Molecules / PAI's
A. Formal Charge
B. Drawing Technique / Resonance structures
C. VSEPR Theory
D. Sigma / PI bonds
E. Orbital Hybridization
V. Intermolecular Forces (IMF's)
A. VanderWaals Forces/ London Disperson
B. Dipole/dipole
C. H-bonding
D. Molecule-Ion Attractions
Resources:
http://www.chemguide.co.uk/atoms/structsmenu.html#top
BONDING
There are three types of intermolecular bonding:
I. Ionic
II. Covalent
III. Metallic
Each one differs in how they deal with electrons.
I.
METALLIC BONDSA.
Definition: positive nuclei surrounded by a sea of mobile electrons
B.
Properties
C.
1.
relatively high melting points
2.
ductile
3.
malleable
4.
conduct heat and electricity
Metallic Crystal Lattices:
1.
Body Centered Cube (BCC):
one central atom surrounded by 8 neighbors
4 -1- 4 configuration
a.
2.
Examples: Na, K, Cr, Fe, W
Face Centered Cubic (FCC):
every atom has 12 neighbors
5 - 4 - 5 configuration
a.
3.
Examples: Cu, Ag, Au, Al, Pb
Hexagonal Close Packed (HCP):
every atom has 12 neighbors
7 - 3 - 7 configuration
a.
D.
Alloys
Examples: Mg, Ca, Zn
1.
Substitutional- an alloy composed when metallic atoms of relatively the same size
replace some of the original atoms.
a.
2.
Examples: Brass (Cu & Zn), Bronze (Cu & Sn), Sterling Silver (92%Ag, 8%Cu)
Interstitial - an alloy composed when smaller atoms fit between the larger ones in the
crystal lattice
a.
II.
Examples: Cast Iron, Steel
Ionic- involves the transfer of electrons.
These are typically formed between metals(givers, +) and nonmetals(takers, -) Electrons are
generally gained or lost to fulfill the OCTET Rule!
These are strong electrostatic forces
Compounds are always neutral so the +'s = -"s.
A.
Structure and Formula
Structure
Formula
Ionic bonds must always be represented using:
The molecular formulas
represent
BRACKETS
DOTS
CHARGES
1.
Examples:
2.
Shortcut to formula writing
a.
If the oxidation numbers (charges) balance:
THE RATIO OF ELEMENTS IS 1:1
b.
If the charges are multiples of each other
REDUCE THEN SWITCH THEM
c.
If the charges are not multiples of each other
SWITCH THEM
the number and type of atom.
B.
Assigning Oxidation Number
1.
There are times when you must determine the oxidation number of an element. To do
so:
a.
Remember what you know:
(1)
Any atom "by itself" has an oxidation number of ZERO
ex. O2, Mg, S8
(2)
b.
Oxygen is always -2 in a compound except H2O2 and OF2
Make sure all individual charges add up to overall charge
ex: MnO2 , FeCl2, CO3-2
2.
Practice: what are the oxidation numbers of the following atoms?
1. KMnO4, N2O5, Na2C2O4, Mg(ClO)2, LiH
2. MgY, X2O, H2Z2O3
C.
Polyatmic Ions -"many atom" ions.
These structures re large members of the ion family but act in a bond as a single unit.
"Table E is my friend"
D.
Naming Rules
1.
Binary- (m/nm) only two elements
Rule: metal first, nonmetal second with an "ide" ending
2.
Multiple Ox. #'sRule: "specify" the ox. # using roman numerals
ex: FeO, Fe2O3, CuO, Chromium VI sulfide, cobalt III fluoride
3.
Polyatomic IonsRule: Naming directly using Table E
NO2-
SO3-2
ClO-
NO3-
SO4-2
ClO2ClO3ClO4-
Practice Problems:
titanium IV thiocyanate
scandium thiosulfate
ammonium acetate
MnCl5
RbPO4
Cu(OH)2
(NH4)2 Cr2O7
more practice
4.
Hydrates- many salts are hygroscopic (water attracting)
Rule:
-The ionic compound must be named using the proper rules for naming ionic
compounds.
-The correct Greek prefixes must be added to the term "hydrate" as to
indicate the number of water molecules per formula unit. Example:
Mono - 1
Di - 2
Tri - 3
Tetra - 4
Penta - 5
Hexa - 6
Hepta - 7
Octa - 8
Nona - 9
Deca - 10
If there is no water, the prefix is anhydrous.
Example(s):[2]
ZnCl2 • 6H2O is zinc chloride hexahydrate
Ba(OH)2 • 8H2O is barium hydroxide octahydrate
CaSO4·2H2O is calcium sulfate dihydrate (gypsum=drywall)
E.
Properties (strong IMF's)
1.
Brittle crystalline solids at STP
2.
High m.p.
3.
Most are soluble in water
4.
Nonconductors as solids, conductors as liquid or in solution (proportional to
concentration)
solids
liquids
gases
III. Covalent- a bond formed by the "sharing of electrons" (nm/nm)
Why do non-metals share electrons instead of transfer?
There are three ways electrons can be shared:
A. Polar - unequally
B. Nonpolar - equally
C. Coordinate - both electrons donated by 1 atom
A.
POLAR COVALENT- electrons are shared unequally (unequal distribution of charge)
1.
When atoms of different e-neg's form a bond, the electrons tend to more attracted to
the more electronegative atom (effective nuclear charge).
2.
The resulting bond is called a dipole, with one side slightly positive and the other
slightly negative.
3.
The bigger the difference in electronegativity the more polar the bond, and therefore
the greater it's ionic character.
4.
This type of bond is called a polar bond and often results in producing a polar
molecule.
a.
Polar Molecule - to be a polar molecule the entire molecule must be a dipole(it must
have a slightly
positive and slightly negative end.)
b.
B.
Examples:
NONPOLAR COVALENT- electrons are shared equally (no unequal distribution of
charge)
When the electronegativities are equal (same pull on electron pair). Usually same element.
1.
Examples: All diatomic molecules.
2.
Nonpolar Molecules from polar bonds.
If a molecule is structurally balanced (symmetrical) the bond dipoles will cancel out.
C.
COORDINATE COVALENT - both electrons in a bond are donated by 1 atom alone
D.
E.
Properties of Molecular Compounds:
1.
Low m.p. / b.p. (gases, liquids and soft solids at STP)
2.
Nonconductors of heat and electricity
3.
Most insoluble in polar solvents
Naming Molecules:
Rule: Use PREFIXES to denote the number of atoms in the molecule.
1.
Examples:
CO2 - carbon dioxide
CO - carbon monoxide
CCl4 - carbon tetrachloride
N2O5 - dinitrogen pentoxide
F.
Network Soilds- huge networks of covalently bonded atoms
Atoms can bond repeatedly with themselves because each makes 4 covalent bonds.
1.
Properties of Network Solids: Covalently bonded nm's with ATYPICAL properties!!!!
a.
extremely hard
b.
high melting point
c.
DO NOT conduct electricity
IV. DRAWING MOLECULESA.
Formal Charge- #original (ve's) - # immediate (ve's)
The structure of the molecule that is most evenly balanced will be preferred in nature. You will
see later
that the goal of making the structure is for the formal charge of each atom to equal zero (or as
close to zero) while adhering to the other rules as well.
1.
Again, FC= # valence electrons as an atom - # immediate valence electrons (all pairs +
1/2 bonding e's)
2.
Examples: 6-6= 0 , 5-5= 0, 6-4= +2
6-6= 0, 5-7= -2
6-5= +1, 5-6= -1
B.
Drawing Technique: (Lewis dot structures)
Step 1 - Determine # of V.E.'s
Step 2 - Organize central atom and ligands
Step 3 - Distribute electron pairs (make bonds, fill octet on ligands, all extra's back on central
atom)
Step 4 - Evaluate (Octet rule and formal charge for PERIODS 1 & 2, Formal charge and octet
rule(maybe) for Period 3+)
Step 5 - Redistribute and recheck.
More Practice On Lewis Dot Structures
C.
Molecular Geometry (VSEPR - valence electron pair repulsion theory)
The number of covalent bonds and additional electron pairs has a direct effect on the geometrical
shape of a molecule.
KEY POINTS:
1. Electron pairs and/or bonds are areas of negative charge that repel each other.
2. E-pairs and/ or bonds will arrange themselves to minimize repulsions.
3. Single, double and triple bonds are considered as ONE bond in VSEPR.
1.
Structures to know:
1. linear
2. bent
3. trigonal planar
4. pyramidal
5. tetrahedral
6. trigonal bipyramidal
7. octahedral
8. square planar
9. T-shaped
D.
Sigma / Pi Bonds
1.
All single bonds are sigms bonds with their overlapping orbitals in the same direction
2.
Double and triple bonds represent pi bonding with p sublevel overlaps
(above and below and left and right of sigma bond)
E.
V.
Orbital Hybridization -
Intermolecular Bonding:
Excluding ionic, metallic and network solids there are three major types of intermolecular forces:
Dipole/dipole (POLAR) or Van derWaals Forces (NONPOLAR) and Molecule/Ion attractions
these are for covalently bonded molecules
A.
Van der Waals Forces: if the molecules are nonpolar they will have minimal
forces of attraction caused by slight,
temporary dipoles (aka LONDON DISPERSION
FORCES).
TEMPORARY DIPOLES
1.
Mol. Wts:
F2 = 38 g/mol
Br2 = 160 g/mol
I2 = 254 g/mol
2.
These forces are temporary slight dipoles 51-49% electron distribution.
B.
Dipole/dipole (PERMANENT) - if the molecules are polar then they will have
"MICROMAGNETISM" forces of attraction. These are relatively strong forces of attraction.
Remember the more polar the bond, the stronger the dipole/ dipole forces. Therefore the higher
the m.p./b.p., surface tension, density, etc.
C.
If the SLIGHTLY (delta) POSITIVE side of the molecule is composed of Hydrogen (the
lowest e-neg nonmetal) the dipole/dipole forces of attraction are called HYDROGEN
BONDING. H-bonding is the Hulk of IMF's because of the great difference in e-neg and
the small size of the hydrogen atom. Even though these are strong they are nothing
compared to metallic, ionic. More info CLICK HERE.
D.
Molecule / Ion Attractions: To have these you must have 1)Molecules and 2)Ions
................go figure!
1.
Example: NaCl(aq)