Polar & Non-Polar
• Electrical Polarity
– Ionic vs Covalent
• Separation of charge
– Basis of polarity
• Degree of polarization
– Molecules with uneven charge distribution
• Effects of polarity
– Like dissolves like
– Movement of ions
Magnetic polarity is similar to electrical
with an energy field between the poles
Dipoles
Include Antennas, Magnets, electric fields,
anything with separation of charge
3
Electrical Polarity
• Net charge on a neutral molecule is zero
– But charge can be unequally distributed
– Degree of asymmetry  polarity
• Ions are charged particles
– Conduct electricity (just like electrons)
– Flow to opposite charge pole
– Electron and ion motion deflected by magnets
• Ionic salts form ionic solutions
– “Electrolytes” in energy drinks
– Conductive path to “ground” in your home
4
Electrical Dipole
Opposite charges at a distance (antenna)
Double-ended charged particles (molecule)
5
Electrical Polarity & Dipoles
• Electronegativity
– Halogens are most electronegative atoms
• Attract electrons, exhibit negative (-) charge
– Metals are least electronegative atoms
• Shed electrons, exhibit (+) charge
– Non-metals are in-between
• Combinations are often “polar”
– One end positive, the other negative
• H-Cl would have H as (+) end, Cl as (-) end
– Polarized molecule has electrical & chemical behavior
• Will align to electric field
• Charged ends attracted to opposites on other molecules
6
Polar Bonds due to Electronegativity
• “Electronegativity” is basis for polar molecules
– Highly electronegative atoms (Fluorine) attracts electrons
• Electron spends more of shared time with electronegative elements
• Net effect is appearance of charge separation
– Polar Molecules more subject to external forces
• Electrical field will orient polar molecules
• Basis for “electret” or electrostatic microphones
7
Fluorine most electronegative, Francium the least negative.
Pauling’s scale maximum is 4, also indicated by column height
Covalent Bond Properties
• Covalent bonds can have “Polarity”
– One end more negative than the other
• Due to unequal distribution of charge
– Charge separation makes a difference
• Chemical reactivity
• “mechanical” properties,
– solubility, boiling point, vapor pressure
• Electrical properties,
– Insulators versus dielectrics,
– Orientation in devices (electrets, capacitors)
9
Polar Molecule HF
Electrons spend more time around F than H
F considered as “more electronegative” than H
10
Water is a polar molecule
Hydrogen loses ownership of its electron, becomes positive
Oxygen acquires 2 additional electrons, becomes negative
Molecular polarity
Opposite charges attract
Ionic vs Covalent Bond
14
Polarized molecules
Align with applied electrical fields
Charged ends attracted to opposite charge
15
“in between” Bond Polarity
Based on Electronegativity
• Not all atoms the same
– Strength of electron attraction varies considerably
• Fluorine has stronger electron attraction than Iodine
• Leads to separation of charge for covalent bonded elements
– Polar Bonds are those with uneven charge distribution
• Electron spends more time around some elements than others
• Unequal probability of electron location means charge separation
– Most bonds are some blend of ionic and covalent
• H-Cl is 83% covalent (and highly polar)
• Na-Cl is 80% ionic
• Non-Polar used to to describe electrically symmetrical molecules.
16
Ionic versus Covalent bonding
17
Shape and Polarity
• “Like dissolves Like”
– Polar solvents dissolve polar solutes
• Salty water, energy (electrolyte) drinks
– Non polar solvents dissolve non-polar solutes
• Oil and gasoline, paint thinner and grease
• Shape of molecules can effect polarity
– Water is “bent” molecule
– Asymmetry leads to polarity
– Water (polar) a great solvent for salts (polar)
– Water would be non-polar if linear
• Life on the planet would be very different
19
Water Shape
• Electronic Structure of water is Tetrahedral
– Oxygen has 4 electron “arms” as SP3 bonds
– Two arms attached to hydrogen + 2 lone pairs
• Molecular Structure of water is Bent
– 2 hydrogen atoms engage two arms of tetrahedron
– 2 lone pairs occupy other 2 arms … felt but unseen
20
Ionic and Covalent comparison
• Ionic solids have bulk properties
– Raise temperature of the entire crystal
• One big “heat sink”
– Local attraction strong between elements
• Electrostatics keep the solid together
• A 3-D network of oppositely charged ions
• Covalent bonding has molecular features
• Individual molecules can react independently
• Minimal attraction betw. covalent bond molecules
• Often gasses, low density, easily melted
21
Covalent Bond Features
• Covalent bond helps complete a shell
• One electron can be shared between two atoms
• Both atoms see combination of shared electrons
• Shared electron does “double duty”
– Sharing augments filling shells of 2 atoms
– Covalent bond has definite length
• Balance of attractive and repulsive forces
• “sweet spot” of lowest energy defines bond length
• Atom bond lengths add to form molecules
– Numerical value only an approximation
22
Arrow Convention
Head of arrow points to negative end(s) of molecule
23
Symmetry negates moment
Molecule with symmetrical charge separation
is NOT effected by external fields
24
Symbol δ indicates partial charge
Charge separation not 100%, and amount is measurable
25
Examples of net dipole moment
26
Polarity and Dipole Moment
27
Literature Example
28
Now to the Experiment
• Solubility in Water (polar) 8 solids to try
– Observe if they dissolve or not
– Dip test tube in hot water bath if not sure
• Solubility in Hexane (non-polar), same 8
– Non-polar solutes soluble in non-polar solvents
– Gentle hot bath heating to verify if in doubt
• A few materials dissolve both ways
– Soap is polar on one end, non-polar the other
– Dissolves grease, but soluble in water too
Classes of materials (page3)
• Ionic (salt, acids, bases)
– Polar materials, dissolve in water (also polar)
– Ionic salts form (+) and (-) ions in solutions
– Ions conduct electrical current
• Polar Covalent (sugar)
– Polar material dissolves in water
– Does NOT form ions
– NO electrical current (or extremely small)
• Non-Polar Covalent (oils, waxes, hydrocarbons)
– Does NOT dissolve in water
– Does NOT conduct electricity
Solubility and Temperature
• Solubility of a solid
– Hot solvents dissolve more solids & faster
• Why we wash with hot water
– Observe rate of KMnO4 dissolving hot water vs cold
• Darker color means more is dissolved
• Solubility of a gas
– Reverse of solids, less soluble in hot solution
– Soda water (Perrier, Pepsi) has acidic pH when cold
• CO2 + H2O  H2CO3 when cold, under pressure
– Goes “flat” if heated, pH increases, gas escapes
• CO2 + H2O  H2CO3 when hot, no pressure
• Don’t forget to add the pH indicator
• Look for color change on hot plate
Solubility of most solids increase with
temperature of water as solvent
This is point of permanganate test
Gases are onlyslightly soluble in water
solubility decreases with temperature increase, so
what happens when we heat carbonated water?
Conductivity
• Polar solutions often good electrical conductors
– Why we don’t use appliances near bathtubs
– “GFI” detects stray currents, to shut off appliances
• LED test is subjective (dim vs bright)
– Pour sample into small test tube
– Insert electrical probes, look for light
• Brighter light means more conductivity (less resistance)
– Be careful, don’t touch probes while running test
– Safe to use, only a 9V battery
• Light Bulb Demo
– Show conductivity of tap water
– Dangers of electrical shock standing in water
Ionic material
forms ions, electrically conductive solutions
Ionic Conductivity
Our experiment includes this test,
we use an LED to monitor conductivity
Non-polar Molecules
Hexane has no ions, no conductivity
Carbon tetrachloride similar behavior
both are COVALENT and NON-IONIC
… and do NOT dissolve in water
Interpreting the water results
• Dissolves in water, bright light (NaCl)
– Must be polar, and highly ionized
– ions are the electrical carriers
• Dissolves in water, dim light
– Polar because it dissolved
– Ionization is low, fewer ions to conduct
• Dissolves in water, NO light (sugar)
– Polar because it dissolves, but NOT ionic
– If not ionic, must be covalent
• Does NOT dissolve in water, NO light (oils)
– Non-Polar, non-ionic, must be covalent (oil, wax)
Interpreting the hexane results
• Dissolves in hexane, NO light
– Must be non-polar, and non-ionized
– ions are the electrical carriers, but none here
• Does NOT dissolve in hexane, NO light
– Polar because it did NOT dissolve in non-polar
– no ions to conduct electrical current in hexane
– Likely to dissolve in water, may conduct current
Los Alamos National Laboratory's Periodic Table
Group**
Period
1
IA
1A
2
3
1.008
3
4
H
Li
Be
6.941
9.012
11
12
Na Mg
22.99
4
5
8
9
10
3
4
5
6
7
11 12
------- VIII IIIB IVB VB VIB VIIB
IB IIB
-----3B
4B 5B 6B
7B
1B 2B
------- 8 ------
20
21
Ca
Sc
39.10
40.08
37
38
Rb
Sr
85.47
87.62
Cs
87
Fr
(223)
56
88
6
7
8
9
B
C
N
O
F
22
23
24
25
26
27
28
29
30
13
14
Al Si
32
Y
40
41
42
44
45
46
47
48
49
50
72
73
74
(98)
75
17
18
Cl
Ar
33
34
35
51
52
53
I
101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9
76
77
78
79
80
81
82
83
84
85
Pt Au Hg Tl Pb Bi Po At
138.9 178.5 180.9 183.9 186.2 190.2 190.2 195.1 197.0 200.5 204.4 207.2 209.0 (210) (210)
107
108
109
86
Rn
(222)
116
118
---
()
()
()
59
60
61
62
63
64
111
Xe
131.3
---
(257) (260) (263) (262) (265) (266)
110
54
114
58
106
83.80
---
Lanthanide
Series*
105
36
Kr
112
(227)
104
39.95
Ra Ac~ Rf Db Sg Bh Hs Mt --- --- --(226)
89
Ne
20.18
S
Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
88.91 91.22 92.91 95.94
57
43
10
16
44.96 47.88 50.94 52.00 54.94 55.85 58.47 58.69 63.55 65.39 69.72 72.59 74.92 78.96 79.90
39
4.003
15
26.98 28.09 30.97 32.07 35.45
31
2
He
P
Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br
Ba La* Hf Ta W Re Os Ir
137.3
5
10.81 12.01 14.01 16.00 19.00
19
132.9
7
24.31
13
14
15 16
17
IIIA IVA VA VIA VIIA
3A 4A 5A 6A 7A
K
55
6
8A
2
IIA
2A
1
1
18
VIIIA
()
()
()
65
66
67
68
69
70
71
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
140.1 140.9 144.2 (147) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0
41
Polar & Non-Polar
• Now to the experiment!
43
Fig. 13-17, p. 384
44
p. 386
45
Fig. 13-18, p. 385