Chemical Bonding & Lewis Dot Structures
I. Covalent & Ionic bonding
A. chemical bonding –
1. octet rule – chem copounds ten to form so that each atom (gain e-, share e-,
lose e-) has an octet (8) of electrons in its highest occupied E-level.( or
have a noble gas configurations)
[why??? because nature favors a lower PE & atoms have lower PE when bonded--the bond length between 2 bonded atoms is the distance between the atoms when the PE
is the least---- this PE is the bond energy or how much E is required to break the chem
bond chart p.168 pm = picometer = 10-12m]
2. attractions between nuclei & valence e- of dif. atoms that hold atoms together
3. atoms bond & redistribute their valence e- to make the atoms more stable
4. types of bonds
a. ionic bonds
b. covalent bonds —polar covalent & nonpolar covalent
c. metallic bonds
d. hydrogen bonds
5. How to determine if a bond will be ionic, polar covalent, or nonpolar covalent
a. look at the electronegativity chart!!! p. 162 & p.151
1) Bonding between atoms of dif. elements is rarely purely ionic or purely
covalent. It usually falls somewhere in between the 2 extremes.
depending on how strongly the atoms of each attract electrons….their
electronegativity.
2) look at dif. in electroneg. to determine type of bond (p. 151)
a) difference from 0 to 0.3 = nonpolar covalent bond (p.162)
b) difference from 0.3 to 1.7 = polar covalent bond
c) difference from 1.7 to 3.3 = ionic bond
d) the atom with the larger electroneg. will be the more negative atom
B. Covalent bonding – sharing electron pairs between 2 atoms—a molecule is
produced -1. molecules – produced by covalent bonds between atoms
a) diatomic molecules – contain only 2 atoms ex: O2, N2, H2
b) chemical formula – shows #’s & kinds of atoms in a compound using
element symbols & subscripts
2. nonpolar-covalent bond – “pure”covalent bond where the electrons are shared
equally so there is an even distribution of e3. polar covalent – covalent bond where there is an unequal attraction for the
shared electrons….which means one end of the molecule is slightly more
neg. δ- & the other end is slightly more positive δ+
4. octet rule applies to most there are exceptions
a) less than an octet – H (2), Be (4), B (6)
b) more than an octet – SF6, PCl5
c) odd # of electrons – NO (5 + 6 = 11)
5. characteristics of covalent cmpds
a) low melting & boiling pts
b) most are rubbery, powdery, liquids or gases
c) does not conduct electricity
d) polar molecules dissolve in other polar molecules
e) nonpolar molecules dissolve in nonpolar molecules
6. single bond – sharing 1 pair of electrons (short bond length)
7. doulbe bond – sharing 2 pr of e- between 2 atoms (shorter bond length)
8. triple bond – sharing 3 pr of e- between 2 atoms (shortest bond length)
C. Ionic bonding—transfer of electrons from one atom to another –
1. ionic compound -- made of pos & neg ions that are combined so # of pos & neg
charges are equal –
2. formula unit – shows the types & # of ions in an ionic cmpd ( pos ion – cationalways goes first)
3. ions reduce PE be combining in crystal formations
4. to compare bond strengths – lattice E – E released when 1 mole of gaseous ions
form into a solid crystal structure
5. characteristics of ionic cmpds
a) high melting & boiling pts
b) exist as crystals that are brittle, shifting the structure (striking it) will cause
cleaveage (break along patterns)
c) many are soluble in water
d) conduct electricity when dissolved or melted
6. polyatomic ion – charged group of covalently bonded atoms – has both ionic &
molecular char.--- they combine with other ions to form ionic cmpds
D. Metallic bonding—chem bonding between metals that results from the attraction
between metal atoms & the SEA of e- surrounding them – the bonding is the same
in all directions
1. in metals the highest E level has very few e2. the vacant p (& d) orbitals overlap
3. this allows the valence e- to roam freely throughout the entire metal
a) this allows for great conductivity of heat & electricity
b) small distances between orbitals allow for a wide range of photons to be
absorbed easily (& re-emitted easily in the form of light which makes
metals shiny)
c) since bonding is same in all directions metals are malleable (able to be
hammered into sheets) & ductile (able to be pulled into thin wires)
d) very high melting & boiling pts.
4. heat of vaporization – measure of bond strength – the amt of heat required to
vaporize metal into individual metal atoms in the gaseous state
E. Other types of bonds (weaker)
1. Hydrogen bonding – intermolecular attractions that occur between polar
covalent molecules that have a hydrogen atom bonded to an extemely
electroneg. atom (N, O, F)—the H atom will be δ+ and be attracted to the
δ- (unshared pair of e-) of another like molecule
a) this accounts for the high boiling pt of H2O, HF, NH3
2. London Forces – due to the constant motion of e- momentary polarity occurs in
all atoms & molecules (distribution of e- uneven) this can weakly hold
atoms/molecules together
a) the larger the atomic/molar mass the greater the attraction—the greater
London forces can increase the bp
II. Electron Dot Notation
A. electron configuration that shows how the valence e- are arranged around an atom
……shown usually as dots *fill in dots in counterclockwise & in singles
1. ex: p. 170 H, B, Kr, Sn, Bi, C, O, N, P, K
III. Lewis Structures
A. electron dot notation can be used to show electron config for molecules (& ionic
cmpds) p.170
B. Lewis structures keep the dots* to show unshared e- & use a line to show a shared
pair of electrons *later the dots will be dropped*
1. structural formula - shows only the shared pair of e- (p.171)
How to draw Lewis dot structures for molecules
1. determine the type & number of atoms in the molecule
2. write the electron-dot notation for each type of atoms in the molecule
3. find total # of valence e4. find the center atom (C if it’s present if not then it’s the least-electronegative atom
– except for H it’s never the center)
5. connect the atoms by electron-pair bonds (dots or lines)
6. add in unshared prs of electrons so that each atom has an octet (or 2 valence e- for H)
7. count e- to be sure you have the correct number of valence e-
How to draw Lewis dot structures for ionic compounds
1. Draw the Lewis dot structure for the positive ion (with charge) If there is more
than one positive ion, be sure to draw them all (with charge).
2. Draw the Lewis dot structure for the negative ion (with charge) very close by. If
there is more than one negative ion, be sure to draw them all (with charge)
p.180
Examples:
N2, HF, NH3, CH4, CF4, NO+
IV. Molecular Geometry - VSEPR