Types of Bonds
A. Electronegativity - The ability of an atom to
attract electrons to itself in a bond
1. Periodic Trends (link to size)
Metals – Low Electronegativity
Non- Metals – High Electroneg
The smaller the atom, the higher the
electronegativity
Types of Bonds
a. Molecular Compounds (non-metals)
• Sharing of electrons
• Covalent or polar covalent bonds
• Ex: H2
b. Ionic Compounds (metal + non-metal)
• Exchange (stealing) of electrons
• Ex: NaCl
1
Types of Bonds
3. Types of bonds
Electronegatvity Difference
Types of Bonds
Are the following bonds ionic, polar covalent or
covalent?
Ba – Br
C–N
Be – F
B–H
O–H
Be – Cl
P–H
C–H
O–O
Example: Na-F
Ionic
Ionic – Electron is completely removed from
one atom and goes to another
a. Metal – Low e-neg (loses e-)
b. Non-metal – high e-neg (gains e-)
Common Charges
IA
IIA
IIIA
IVA
VA
VIA
VIIA
Carbide
C4-
Nitride
N3-
Oxide
O2-
Fluoride
F-
Phosphide
P3-
Sulfide
S2-
Chloride
Cl-
Hydride
HLithium
Li+
Beryllium
Be2+
Sodium
Na+
Magnesium
Mg2+
Potassium
K+
Calcium
Ca2+
Selenide
Se2-
Bromide
Br-
Rubidium
Rb+
Strontium
Sr2+
Telluride
Te2-
Iodide
I-
Cesium
Cs+
Barium
Ba2+
Aluminum
Al3+
2
Complete the following chart:
Formula
Lewis Dot
K and I
Sr and F
Ba and Cl
Sr and N
Ba and O
Ionic Solids
Complete the following chart:
Formula
K and S
Lewis Dot
1. Not Separate molecules
2. Crystal Lattice – Regular 3D pattern in an
ionic solid
Al and O
Na and Br
Ca and O
Ionic Solids
at lattice points
ions at lattice points
3
Ions in Water
1. Water (molecule) dissolves salt (ionic crystal):
Molecules
Other Salts
Many ionic compounds are called salts
CaCl2
NaHCO3
NaF
CaCO3
NaOH
MgSO4
Calcium Chloride (Quik-Joe)
Sodium bicarbonate(Baking soda)
Sodium Fluoride (toothpaste)
Calcium Carbonate(Chalk, antacid)
Sodium Hydroxide (Drano)
Magnesium Sulfate(Epsom Salts)
All big clumps (crystals) of Ions
Quick Review:
What is a Molecule?
• Group of atoms held together by
covalent and polar covalent
bonds (SHARING)
• Usually composed of non-metals
• Stronger than Ionic bonds
• Separate (discrete) groups of
atoms
1. Draw the Lewis Structure of Na2S
2. Is the compound C2H4 ionic or molecular?
Why?
3. Are molecules or ions at the lattice points of
ice crystals?
4. Are molecules or ions at the lattice points of
MgSO4 crystals?
5. Which end of a water molecule is attracted
to Cl- in solution? Why?
4
Old School Lewis Dots
A. Single Bonds (LD and Stick)
CH4
H2O
(bonded versus lone pairs)
Old School Lewis Dots
N and H
P and F
Cl and O
NH2CH3
The Lone
Pear(Pair)
rides again!
You try:
As and H
N and Br
CH3CH2OH
NH2CH2SiCl3
CH3OCH3
As and Cl
C and Br
CH3CH2NH2
CH3OCH2CH3
5
Old School Lewis Dots
Old School Lewis Dots
You Try:
HCN
N2
AlN
O2
B. Multiple Bonds
CO2
C2H2
Lewis Dots
Rules
1. Sum all valence electrons, including
charges
2. Single Bonds
3. Outer atoms get an octet except H
4. Center gets rest even if it violates the octet
5. Double/triple bonds if center atom still
does not have an octet
You try:
SF4
KrF4
Cl2O
ClF3
H2SO4
Lewis Dots
NH3
NCl3
SF6
Br2O
ClF5
Warm-Up:
CH3CH2NH2
SeF4
KrCl4
H2O2
BaCl2 (this is an ionic cmpd)
6
Lewis Dots
CO2
HCN
CNICl4NO+
CO32H3O+
NH4+
PO43SO2
Less Than an Octet
• Hydrogen only makes one bond
• Gr I, II and III
• Especially if with a halogen
BeCl2
PCl5
BCl3
AlCl3
More Than an Octet
• Non-Metals starting with P
P
S
Cl
Ar
As
Se
Br
Kr
Te
I
Xe
Si2H4
AsO43CS32H3S+
PH4+
7
Resonance Structures
Warm-Up:
O3
Resonance Structures
SPECIAL NOTE:
Definition – When a molecule can exist in
more than one arrangement of electrons
1. Atoms don’t move
2. Only the electrons (double bonds) move
3. Also called “delocalized bonding”
Oxygen rarely makes three bonds for
resonance structures. Those structures are
not common.
Examples
1. NO2-, CO32-, CHO2-, HNO3
2. Which needs resonance, SO3 or
SO32-
PO3HSO3+
Shapes of Molecules
• Valence Shell Electron Pair Repulsion Theory
(VSEPR)
• VSEPR – Valence electron pairs on the central
atom push as far apart as possible.
8
Shapes of Molecules
Shapes of Molecules
1. Linear (180o)
BeH2
CO2
Shapes of Molecules
Shapes of Molecules
2. Trigonal Planar (120o)
NO3-
Shapes of Molecules
Shapes of Molecules
3. Tetrahedral (109.5o)
CH4
9
Shapes of Molecules
Shapes of Molecules
4. Trigonal Pyramidal (~107o)
NH3
Shapes of Molecules
Shapes of Molecules
5. Bent (Type I, ~104.5o)
H2 O
10
Shapes of Molecules
H
O
Shapes of Molecules
H
Bent (Type II, 120o)
SO2
•Atoms are bent
•Bond angle = 120o
Shapes of Molecules
Shapes of Molecules
6. Trigonal Bipyramid (120o, 90o)
PCl5
11
Shapes of Molecules
Shapes of Molecules
7. Octahedral (90o)
SF6
Shapes of Molecules
DRAW LEWIS DOT STRUCTURES AND PREDICT
THE BOND ANGLES FOR:
PH3
GeO2
PCl5
NCl3
H2S
SiH4
CF4
SO32SeCl6
SO3
Shapes of Molecules
Ex: Multiple Bonds:
N2
HCN
CO2
H2CO
Ex: All single bonds:
PH3
H 2S
SeCl6
AsF5
CH3F
HF
SO22+
SO22-
SO3
SF6
SiH4
BeCl2
BF3
SO2
12
Shapes of Molecules
Shapes and Formula?
1. Examples:
BeH2
H2O
BH3
NH3
2. There is no relationship between formula
and shape
PO2-
PO43-
PO3-
PO33-
XeF4
13
Household Molecules
H 2O
HCl
NH3
H2SO4
H2
CH4
Polar Molecules
1. Polar molecule – Overall, the
electrons are attracted more to one
end of an entire molecule
2. Non-Polar Molecule – The electrons
are spread out evenly over the entire
molecule
3. d-/ d+ Partial (not full) charges
H2 O
H2
H2 O
CH4
H2CO
H2CO (C is the center)
BeCl2
NH3
CO2
CSO
HCN (triple bond)
14
CCl4
CH3Cl
CH2Cl2
CHCl3
HBeCl
HBr
CH3Cl
CHCl3
Lewis Dot
Shape
Polar or Nonpolar molecule?
H2O
NH3
CH3Cl
Polar Molecules
CH2Cl2
HBeCl
HBr
Hydrogen Bonding
a. Water Beading
b. SurfaceTension
c. Ice
d. DNA
e. Miscibility (NH3, ethanol)
f. Immiscibility (Water and Oil)
CH3CH2OH
SO2
Water Beading
15
Surface Tension
Surface Tension
Ice
Ice
16
MoleculesBoiling Point
DNA
• Generally
increases with
increasing molar
mass
• H2O unusually
high - H-bonding
DNA
DNA is TWO
molecules
that are
hydrogen
bonded (like a
zipper)
DNA
Human
Genome
Project
Miscibility
• DNA in one cell = 1 meter
• DNA in all your cells = 93,000,000 miles
•“Like dissolves like.”
•Polar dissolves Polar
•Water, ammonia
17
Miscibility
Miscibility
Miscibility
Miscibility
Miscibility
•Water, alcohol
•Water, alcohol
18
Miscibility
•Water, alcohol
•Water, alcohol
Miscibility
•Non-Polar dissolves Non-Polar
•oil paint, thinner
Miscibility
Would acetone (shown below) dissolve in
water?
:O:
||
CH3CCH3
Acetone
19
Molecules
London Dispersion Forces
• Very weak force
• Caused by temporary imbalances in electrons
S2
O2
• More electrons, more chance for temporary
dipole
Boiling Point Table
Halogen
Molar
Mass
BP(oC)
Noble
Gas
Atomic
Mass
BP(oC)
F2 (g)
38.0
-188
He
4.0
-268
Cl2 (g)
71.0
-35
Ne
20.2
-246
Br2 (l)
159.8
59
Ar
39.9
-186
I2 (s)
253.8
185
Kr
83.8
-152
London Forces: Organic Compounds
Molecules
Consider the following molecules:
Te2
London Forces: Inorganic Molecules
Molecules
Se2
a. Rank them from weakest to strongest London
forces
b. Which should have the highest boiling point?
Which should have the lowest?
c. Would it take more energy to melt a sample
of carbon or silicon?
• The longer the carbon chain, the higher the
London Dispersion Forces (the higher the melting
point and boiling point)
CH4
C2H6
C3H8
C4H10
BP(oC)
-161.6
-88.63
-42.07
-0.5
Which should have a higher boiling point:
Explain the following trends in solubility in water
C6H14 or C12H26
Using you knowledge of “like
dissolves
Alcohol
Solubility
in Hlike”,
2O
(mol/100
g H2O at 20oC)
explain the following trends
in solubility.
CH3OH
∞
CH3CH2OH
∞
CH3CH2CH2OH
∞
CH3CH2CH2CH2OH
0.11
CH3CH2CH2CH2CH2OH
0.030
CH3CH2CH2CH2CH2CH2OH
0.0058
CH3CH2CH2CH2CH2CH2CH2OH
0.0008
20
Emulsifying agents
– Mayonaise
– Soap
The Key to the Universe
Polar Molecules
You Try:
CCl4
CH3Cl
CH2Cl2
NO3H2CO
H2S
KrCl4
Si2H4
BN
CH3SiH2PH2
SiS2
H2SO4
O
O
S
O
H
O
H
21
9. (c) Sc2+
(d) I2- (e) As215. Rb+ Ba2+ Te2- N316. Al3+, Mg2+, Na+, F-, O2-, N320.RbF
Rb2Te
Rb3P
MgF2
MgTe
Mg3P2
CrF3
Cr2Te3
CrP
22. BeS
Cs2S
Ga2S3
8+
3+
24.a) Ir
b) Sc
c) Pt6+
3+
+
e) Tc
f) Ag
27. a) HI
49. a) PC
50. a) PC
52. a) N-H
octahedral
tetrahedral
linear
trigonal planar
bent
trigonal planar
Linear
c) SiBr4
c) PC
d) Ionic
c) S-F
d) H3As
d) Ionic
d) PC
P-Cl
SrS
d) Co3+
54.a)
c)
e)
56.a)
a)
b)
c)
d)
e)
f)
g)
b) SeBr2
b) Cov
b) PC
b) Si-O
h) tetrahedral
i) trigonal bipyramid
j) trigonal pyramid
k) octahedral
l) linear
m) linear
n) linear
trigonal planar
linear
trigonal planar
H2O
c) NH4+
b) trigonal pyramid
d) trigonal planar
Assessing the Objectives (page 278)
a) 2, 3
b) Less than 109.5o
c) Trigonal Planar
d) Bent
e) Linear, 180o
bent, 120o
22
Questions
1. Rank the three liquids from weakest to the
strongest forces between their molecules.
2. Discuss how you chose your rankings from
the results of the experiment.
3. Is isopropyl alcohol or hexane more like
water? Explain, using the results from the
experiments.
4. What information did the salt experiment
provide about the molecules?
Formula
Lewis Dot
Ca and O
Ca and Br
P and H
C and F
N and F
GAK – NEATNESS!!!!!!!
1. Glue
2. Food coloring
3. Borax solution
Consider the following three chemical bonds:
C-O
Cl-Br
Rb-Cl
a. Determine the type of bond present in each
situation
b. In which bond is an electron transferred?
c. Draw the Lewis Dot structure of this ionic
compound.
d. In which bond are electrons unequally
shared?
e. Identify the element in that bond that the
electron spends more time with.
The element chlorine forms compounds with both
strontium and sulfur.
a. Write the formula of the compound that forms
between strontium and chlorine.
b. Draw the Lewis Dot Structure of this compound,
indicating whether it is ionic or molecular.
c. Write the formula of the compound that forms
between sulfur and chlorine.
d. Draw the Lewis Dot Structure of this compound,
indicating whether it is ionic or molecular.
e. Is the structure you drew in (d) linear? Why or why
not?
f. State the bond angle of the structure in (d).
23
Draw the Lewis Dot Structure and predict whether it is a polar or
non-polar molecule
Water
Isopropyl
Alcohol
Hexane
H2O
CH3CH(OH)CH3
CH3CH2CH2CH2CH2CH3
Draw Lewis Structures (Ionic)
BaF2
Li2O
Draw Lewis Structures (Molecular)
ClF3
SiO32SO2
Number 1
a. Ionic
b. Cov
c. Polar Covalent
d. Ionic
e. Ionic
f. Polar Covalent
g. Polar Covalent
h. Ionic
Compound
Central Atom
BeH2
Tan
BH3
Tan
CH4, NH3, H2O
Black
PCl5
Tan
SF6
Yellow or silver
Number 2
a. CaCl2
b. NaF
c. AlF3
d. MgO
e. CsCl
f. SrO
g. BaF2
24
Number 3
a. K+ Brb. Al3+ 3Clc. 2Ga3+ 3O2d. 2Li+ O2e. Ca2+ 2Brf. Na+ OHg. Ca2+ 2NO3h. 2Na+ CO32i. 2Fe3+ 3SO42j. Co4+ 4NO2-
25
26 a) H2Se
e) NCl3
27 a) HI
b) GeH4
f) CBr4
b) SeBr2
c) CLF
d) Cl2O
c) SiBr4
d) AsH3
26
7a. Octahedral
b. Tetrahedral
c. Linear
d. Trigonal Planar
e. Bent I
f. Trigonal Planar
g. Linear
1.
a.
b.
c.
d.
e.
f.
g.
h.
h. Tetrahedral
i. Trigonal bipyramid
j. Trigonal pyramid
k. Octahedral
l. Linear
m. Linear
n. Linear
Covalent, Ionic or Polar Covalent Bonds?
Ionic
Covalent
Polar Covalent
Ionic
Ionic
Polar Covalent
Polar Covalent
Ionic
27
3a. K+ Brb. Al3+ 3Clc. 2Ga3+ 3O2d. 2Li+ O2e. Ca2+ 2Brf. Na+ OHg. Ca2+ 2NO3h. 2Na+ CO32i. 2Fe3+ 3SO42j. Co4+ 4NO2-
http://www.youtube.com/watch?v=WaPxTAeW
5e8
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