Download Chapter 9 Notes - UIC Department of Chemistry

Chemistry 112
Kotz Chapter 9
Fetzer Gislason
Drawing Lewis Structures for Ionic and for Covalent Substances
Electron dot symbols are sometimes called Lewis symbols: each dot represents 1 valence electron.
# valence electrons = the group number
Group #
I
Na
II
III
IV
Mg
Al
Si
V
P
VI
VII
VIII
S
Cl
Ar
Lewis Structures for Ionic Compounds
Ionic compounds contain both a cation and an anion held together by strong electrostatic forces
(originating in their charges.) Lewis structures for ionic compounds are written by putting two Lewis
ionic symbols together.
F−
Lewis structures for ions: Mg2+
Ionic compounds-put the ions together as in a formula: MgF2
Lewis Structures for Covalent Compounds
1) Nonmetals tend to share electrons in order to acquire a stable octet (8) of electrons.
2) Shared electrons are counted by both "partners" as contributing to their octet.
3) Shared electrons between two elements constitutes a covalent bond between them.
4) Molecules are held together by covalent bonds.
Examples: H2O
CO2
CH4
N2
Unequal sharing of electrons leads to partial charges on some of the atoms in a molecule. Determination
of the formal charges on atoms in a molecule can serve as a check on the reasonableness of the Lewis
structure we have drawn. A “good” Lewis structure has:
1) Formal charges as close as possible to zero
2) Negative formal charges reside on the more electronegative atom. (F is most electronegative.)
Formal charge of atom = group number - # electrons in lone pairs - 1/2 (# of shared electrons)
= group number - # electrons in lone pairs − #bonds to atom
Draw Lewis structures for the following. Check by determining the formal charge on each atom in the
structure. Summation of all the formal charges must equal the charge on the overall structure.
[ClO3]-
POCl3
1
Chemistry 112
Kotz Chapter 9
Fetzer Gislason
Resonance Structures Sometimes the experimental measurement of bond length is not close to either what we
expect for a single bond nor for a double bond but rather for something in between these values. The bonds in O3,
for example have equal bond lengths between that of a single and of a double bond length. The Lewis picture fails
here, but we can modify it by saying that the true structure is a resonance hybrid of two Lewis structures. How do
we draw resonance structures?
O3
[CO3]2- has three resonance structures. Draw them.
Lewis structures that differ in the placement of single and double bonds are called resonance structures. We
understand that neither drawing is accurate but that the molecule has bond lengths between a single and a double
bond. We call this a resonance hybrid; a cross between a single and a double bond.
VESPR Valence electron shell pair-repulsion. This is a method
of predicting the shape of the molecule by counting electron
domains. A domain is either a lone pair or the bond between two
atoms.
Electron domains: 1) each lone pair is a domain
2) each bond is a domain;
a single bond = 1 domain; a double bond = 1 domain; a triple bond = 1 domain
Main idea: Electron domains repel one another; the lowest energy (and therefore the most stable)
electron arrangement in a molecule results when the electron domains are as far away from each
other as possible.
We will always start with a central atom and consider molecules with 2-6 atoms bonded to the central
atom.
Shapes of Molecules.
Electron Pairs
Electronic Geometry
Molecular Geometry
total
bonding lone
2
2
0
linear
linear, AX2
3
3
0
trigonal planar
trigonal planar AX3
2
1
bent, AX2
4
4
0
tetrahedral
tetrahedral, AX4
3
1
trigonal pyramid, AX3
2
2
bent, AX2
5
5
0
trigonal bipyramidal
trigonal bipyramid, AX5
4
1
seesaw, AX4
3
2
T-shaped, AX3
2
3
linear, AX2
6
6
0
octahedral
octahedral, AX6
5
1
square pyramid, AX5
4
2
square planar, AX4
2
Chemistry 112
Kotz Chapter 9
Fetzer Gislason
Practice: Find the electronic geometry for
IF4−
BrO4−
SeOF4
SO2
Properties of a Chemical Bond: length, enthalpy, strength, and polarity
C−C
Bond length
1.54 Å
C=C
C≡C
1.34 Å
1.20 Å
Bond enthalpy
Bond strength
348 kJ
614 kJ
839 kJ
Strength of chemical bonds: (as measured by bond enthalpy)
weak chemical bonds
up to 200 kJ/mol
average chemical bonds about 500 kJ/mol
strong chemical bonds
above 800 kJ/mol
Judging the strength of covalent bonds.
Breaking bonds
requires energy
Making bonds
releases energy
sign of ∆H is ??
sign of ∆H is???
Bond enthalpy is the energy needed to break a bond in one mole of gaseous substance.
Electronegativity and bond polarity: Electronegativity measures the attraction for shared electrons in a
bond. We use electronegativity differences to predict the type of bonding.
• Find electronegativity value for both elements in table.
• Subtract.
• If the difference is
Zero
pure covalent bond
<2
polar covalent bond
≥2
ionic bond
Examples:
P −O
S−F
Br − Br
O − Cl
3