Welcome to Bonding
OB: introduction to what bonding is, and
learning to draw (lots of) Lewis Dot Diagrams.
Chemistry is a course about what stuff is, and how this stuff
reacts with other stuff. If a reaction happens, then bonds
were formed, and most likely, other bonds were broken to
let this happen.
There is a lot to chemistry, the energy required or released,
the properties of all that stuff, the conservation of energy,
conservation of mass, and even conservation of CHARGE
(coming soon).
Atoms of the elements are bonded together to make new
stuff. There are a variety of bond types, all with rules and
reasons. We’ll examine lots of them, but please realize that
there is always way more going on - above our heads, and
due to reasons we will not learn about in this course.
Types of bonds we will see:
1. Ionic - are bonds between cations and
anions. The charges of the cations and
anions will always net to zero.
2. Covalent - are bonds between two or
more NON METALS. If there is a metal
in the compound, it must be ionic. If
there are no metals, then it’s covalent.
Types of bonds we will see:
3. Metallic - these are the connection that
metal atoms make with each other when
solid metals exist. They give rise to all
the properties of metals.
One type of metal atom bonds to itself.
4. Intermolecular - these are the kind of
weak attractions between molecules.
There are 3 different kinds of these, but
two are nearly identical.
I like these a lot.
Learning to draw Lewis Dot Diagrams. These will show us
the outermost electrons, the VALENCE ELECTRONS, which
are in the valence orbital. By reminding ourselves about
these electrons it will help us to better understand what
bonding can happen.
5. The outermost electrons are the VALENCE electrons
6. The outermost electron orbital is the VALENCE ORBITAL.
7. Bonds always* form when atoms or ions end up with
full outer orbitals, like the noble gases.
* of course there are exceptions, but not many, and
we’ll get to these exceptions soon.
8. Dots will represent electrons.
9. Lewis dot diagrams will only show valence electrons,
not the inside electrons. The inside electrons do not
participate in the bonding anyway.
Electron Orbitals
10. The first orbital is tiny, it only holds 2 electrons at most.
11. The 2nd orbital is bigger, it can hold only up to 8 electrons
(with a few exceptions!)
We won’t be drawing atoms with more than10 electrons in our
class, but they will be added in college chemistry, so be patient.
12. Together we’ll draw a few atoms, and ions, then YOU will continue these charts which
will run from hydrogen to calcium.
Atom number
Atom symbol
1
H
2
He
3
Li
4
Be
5
B
Lewis Dot
(atom)
Ion Symbol
H
+1
Lewis Dot
(ion)
Atom number
Atom symbol
Lewis Dot
(atom)
Ion Symbol
Lewis Dot
(ion)
1
H
H·
H+1
[H]+1
2
He
He:
---
---
3
Li
:Li
Li+1
[Li]+1
4
Be
:Be:
Be+2
[Be]+2
5
B
:B
Not in our class
Atom number
Atom symbol
6
C
7
N
8
O
9
F
10
Ne
11
Na
Lewis Dot
Ion Symbol
Ion Dot
Atom number
Atom symbol
Lewis Dot
Ion Symbol
Ion Dot
6
C
C
7
N
N
N-3
8
O
O
O-2
9
F
F
F-1
10
Ne
Ne
---
---
11
Na
Na
Na+1
[Na]+1
Not in our class
Atom number
Atom symbol
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
Lewis Dot
Ion Symbol
Ion Dot
Atom number
Atom symbol
Lewis Dot
Ion Symbol
Ion Dot
12
Mg
Mg
Mg+2
[Mg]+2
13
Al
Al
Al+3
[Al]+3
14
Si
Si
Not in our class
15
P
P
P-3
16
S
S
S-2
17
Cl
Cl
Cl-1
Atom number
Atom symbol
18
Ar
19
K
20
Ca
Lewis Dot
Ion Symbol
Ion Dot
Atom number
Atom symbol
Lewis Dot
Ion Symbol
Ion Dot
18
Ar
Ar
---
---
19
K
K
K+1
[K]+1
20
Ca
Ca
Ca+2
[Ca]+2
Lewis Dot diagrams for atoms
show only valence electrons.
Lewis Dot diagrams for ions show
the NEW Valence Orbital Arrangement.
Bonding class #2
OB: Metallic Bonds, More Lewis Dots,
and the Octet Rule.
-----------------------------------------------On the black tables out back:
get a copy of the
Bonding Basics,
Bonding Notes (white and tan),
Bonding I CAN questions,
Bonding HW’s.
14. When sodium chloride forms from sodium metal and
chlorine non metal, the atoms form ions first. To do this,
the sodium TRANSFERS an electron to a chlorine atom .
15. The sodium becomes a sodium cation with a +1 charge
16. The chlorine becomes a chloride anion, with a -1 charge
17. Let’s draw the Lewis dot diagrams for the atoms, the ions, and then the compound.
ATOMS
IONS
COMPOUND
14. When sodium chloride forms from sodium metal and
chlorine non metal, the atoms form ions first. To do this,
the sodium TRANSFERS an electron to a chlorine atom .
15. The sodium becomes a sodium cation with a +1 charge
16. The chlorine becomes a chloride anion, with a -1 charge
17. Let’s draw the Lewis dot diagrams for the atoms, the ions, and then the compound.
ATOMS
IONS
COMPOUND
18. It’s important to note here, the sodium atom at 2-8-1 electron configuration
becomes 2-8 as it loses one electron, becoming isoelectric to neon.
19. It loses enough electrons to get a perfect outer orbital, as defined by
noble gases having the most perfect, or most stable electron orbitals of all.
20. The chlorine atom has a 2-8-7 configuration, gains one electron, and
becomes 2-8-8, making it isoelectric to argon.
21. Both ions end up with perfect outer orbitals,
both end up isoelectric to a noble gas.
22. Almost all ions follow the octet rule.
23. The octet rule is that when bonding all ions
will end up with eight outer most electrons,
and when bonding, all non-metals bonding
together with other nonmetals in covalent
bonds, will end up with 8 electrons in the
outermost orbitals.
24. This is a rule, but not a law. There are exceptions:
some ions are too small, like Li.
Some atoms can squeeze 10 electrons, we love exceptions!
25. Fill in this chart!
Compound
name
Compound
Formula
Cation
Anion
Magnesium
oxide
MgO
Mg+2
O-2
LiF
CaCl2
Lewis Dot Diagram
Copy this table BIG, leave enough room for the dot diagrams!
Compound
name
Compou
nd
Formula
Cation
Anion
Magnesium
oxide
MgO
Mg+2
O-2
Lithium
fluoride
LiF
Li+1
F-1
Calcium
chloride
CaCl2
Ca+2
Cl-1
Lewis Dot Diagram
Compound
name
Sodium…
Cesium
oxide
Compound
Formula
Cation
Anion
S-2
Lewis Dot Diagram
Compound
name
Compound
Formula
Cation
Anion
Sodium…
Na2S
Na+1
S-2
Cesium
oxide
Cs2O
Cs+1
O-2
Lewis Dot Diagram
26. Why is the formula for aluminum oxide
Al2O3 and not some other ratio?
Each
metal
atom is
2-8-3
and
needs to
become
2-8
a +3
cation.
Follow
the octet
rule!
Al
O
O
Al
O
Each
nonmetal
atom is
2-6
and
needs to
become
2-8
a -2
anion.
Follow
the
octet
rule!
Why is the formula for aluminum oxide Al2O3
and not some other ratio?
Al
O
O
Al
O
A PERFECT TRANSFER OF ELECTRONS, 6 FROM Al + 6 INTO OXGYEN
27. Draw the UGLY Lewis dot diagram for
Magnesium Nitride
Aluminum Oxide
27. Draw the UGLY Lewis dot diagram for
Magnesium Nitride
Aluminum Oxide
28. This kind of bonding is to explain how metal atoms stick
together to form solid metals. Literally, how does
6.02 x 1023 atoms of copper stick together so you can
weigh 64 grams of copper on the scale?
For the same reason that these atoms can stick together,
nearly all of the properties of metals can be explained at
the same time.
As usual, it’s all about the electrons, where they are, what
they’re doing, and how fast they can move.
First, let’s name a few properties of metals…
29.
Metals are (you better learn the definitions of these ASAP)
Malleable
Ductile
Conduct electricity
Form cations
Have higher densities than non metals
Have low Specific Heat Capacities than non metals
etc.
These main properties can be explained by how we “understand” the metals to be
bonded together.
Draw this diagram quickly…
30. Metals are understood to exist as packed
cations, surrounded by loose valence
electrons. These valence electrons can
move quickly (near the speed of light)
if they have to.
The positives balance the negatives since
they are all atoms. Protons = electrons.
31.
Imagine smashing the metal with a hammer to make the metal exhibit its
malleable nature. The cations will be crushed closer together, and would
repel, but the loose valence electrons flow to offset this excess positive charge.
Same when you squish it into a wire.
32.
Imagine a flow of electrons (electricity) in from the left side. As electrons flow
into the metal, there are too many negative electrons for the cations, so the
excess electrons flow out the other side (the flow of electrons is electricity!).
33.
The cations are awash in a sea of loose valence electrons.
Bonding Class #3
OB: introduction to covalent bonding
34. Covalent bonding is when 2 or more
nonmetals share their valence electrons
to bond.
35. They do not transfer them like ionic
compounds do.
36. With ionic bonding, there is a
TRANSFER OF ELECTRONS FROM METAL → NONMETAL
They still follow the octet rule (mostly).
Ionic bonds require a metal to be first in the formula.
Ionic bonds make formula units (FU’s).
37. In Covalent Bonding, there is
A SHARING OF THE VALENCE ELECTRONS,
AND THE ATOMS WILL FOLLOW THE
OCTET RULE.
38. NO METALS in any covalent bonds.
39. Covalent bonds form molecules.
40. Molecules form with covalent bonds (sharing electrons)
by following the octet rule almost every time.
41.
Let’s draw Lewis Dot Diagrams
H2
F2
41.
Let’s draw Lewis Dot Diagrams
H2
HH
F2
Br Br
42. In covalent bonds, all atoms get to share enough electrons so
that they get full valence orbitals at least some of the time.
43. These bonds previous are all SINGLE BONDS because they only share a
single pair of electrons (one electron from each atom).
44. They are also NONPOLAR bonds because there is
NO DIFFERENCE IN electronegativity value
between the atoms.
F2 + H2 have
SINGLE NONPOLAR COVALENT
bonds
45. Draw the Lewis Dot Diagrams
HCl
H20
Let’s see if we can draw the Lewis Dot Diagrams for HCl and then, water.
H Cl
HO
H
The hydrogen atom in black has one valence
electron. The chlorine, in red, has 7 valence
electrons. Together the hydrogen gets to borrow
one electron from chlorine to fill its tiny orbital, and
chlorine gets to borrow one electron from hydrogen
to fill its larger orbital (octet rule).
Here, the hydrogen are black again, and need to
borrow one electron from oxygen each, to fill
their tiny orbitals. Oxygen borrows one electron
from each of the hydrogen atoms, to fill up its
larger orbital (octet rule). Water is bent, don’t
forget! We’ll learn why soon enough!
The red/black colors are not important, “just for
seeing” it better as we learn at the beginning.
H Cl
The bond between H and Cl is
between atoms with 2 different
electronegativity values. What
are their EN Values?
46.
HO
H
Is this a polar bond, or
a nonpolar bond?
How about here? What are the
EN Values for H and for O?
47.
Is this a polar or
nonpolar bonds?
H Cl
48.
NAME THIS BOND
_____________________________
HO
H
49.
There are 2 identical bonds here
(both H-O).
Name THESE BONDS
There are 2…
___________________________
H Cl
48.
NAME THIS BOND
SINGLE POLAR COVALENT
HO
H
49.
There are 2 identical bonds here
(both H-O).
Name THESE BONDS
There are 2…
SINGLE POLAR COVALENT
H Cl
HO
H
50. Another way to draw this, with a
lot less dots, is called a structural
diagram. With a structural diagram,
we only show the bonds, with short
lines indicating shared electrons. A
single dash represents a single
covalent bond.
Draw both of these molecules without
dots, with structural diagrams.
Structural Diagrams
H Cl
H―Cl
HO
H
This is a bit turned, but molecules move in
3 dimensions. It’s fine this way, or pointing
in any other way.
51. Draw the Lewis Dot Diagram for AMMONIA
(NH3), then the structural diagram.
NAME THESE 3 BONDS TOO.
Think first:
N
Nitrogen has 5 valence electrons, and they will be paired up in a
Lewis dot diagram (and real life) because this is more stable.
To bond, one pair will have to open up to connect with
3 hydrogen atoms.
N
Bring in the 3 hydrogen atoms…
H
H
H
H NH
H
H―N―H
H
Ammonia as Lewis Dots, and as a structural diagram.
Checking the electronegativity values, we see that H has a 2.2 while N has a 3.0
These bonds are all
single polar covalent.
52. Draw Lewis Dot Diagram, and
Structural Diagram for Methane, CH4
Determine exactly what types of bonds
are present in this molecule.
Draw Lewis Dot Diagram, and Structural Diagram for
Methane, CH4
Determine exactly what types of bonds are present
in this molecule.
H
H
HCH
H
H―C―H
H
Electronegativity values of 2.2 for H, and
2.6 for N, so there are
4 single polar covalent bonds in a molecule of CH4
53.
The greater the difference in electronegativity values between two
atoms, the greater the polarity of the bond. This works like little +/magnets. Some magnets are stronger (greater EN difference) and some
magnets are weaker (lesser EN difference).
Fill in this chart, and then RANK from the
greatest polarity of the bond (1), to the weakest (5).
Polarity
rank
Molecule/
name
EN
#1
EN
#2
EN
diff
Structural diagrams
H2
hydrogen
2.2
2.2
0
H―H
PCl3
OF2
HBr
HI
Polarity
rank
5
1
Molecule/
name
H2
hydrogen
EN
#1
EN
#2
EN diff
zero
2.2
2.2
this is a
nonpolar
bond
H―H
2.2
3.2
1.0
Cl―P―Cl
PCl3
phosphorus
trichloride
Structural
diagrams
Cl
3
OF2
oxygen
difluoride
3.4
4.0
0.6
2
HBr
hydrogen
bromide
2.2
3.0
0.8
H―Br
4
HI
hydrogen
iodide
2.2
2.7
0.5
H―I
O
F
F
54. How do 2 oxygen atoms stay together in O2?
Let’s draw two atoms Lewis Dot to start our
thinking.
O
O
55. How many electrons does EACH atom
of oxygen need to complete the
octet? Can they do this for each other?
Hint: move the bottom pairs of electrons to the open sides.
O
O
Squeeze them together now (this requires you redraw)
In order to both get an octet, the oxygen atoms
must share 2 pairs of electrons with each other.
OO
56. This gives the oxygen molecule a
DOUBLE COVALENT BOND.
57.
Since each atom of oxygen has the same
electronegativity value, it’s proper to call this a:
OO
Double Non-Polar
Covalent bond
O O
Take Out HW 1 and 2
Bonding Class #4
OB: become masterful with both the Double
and the Triple Covalent Bonds, plus some practice
drawing structural diagrams for larger molecules
58. O2 Oxygen How does it bond? (review)
O
O
Each oxygen needs to gain
2 electrons to fill it’s valence
orbital. Each oxygen must lend
2 electrons to the other.
O
O
Will become…
O O
Drawn structurally… this way:
O=O
Each O has a 3.4 electronegativity value, so this is a
Double Non-Polar Covalent Bond
59. Looking at the HONClBrIF twins, in order, let’s figure out
the kinds of bonds that they all have… (draw dots or structural's to think)
H2
O2
N2
Cl2
Br2
I2
F2
Looking at the HONClBrIF twins, in order, let’s figure out the
kinds of bonds that they all have…
H2 Single non-polar covalent H―H
O2 Double non-polar covalent O=O
N2 ???
Cl2 Single non-polar covalent Cl―Cl
Br2 Single non-polar covalent Br―Br
I2 Single non-polar covalent I―I
F2 Single non-polar covalent F―F
Time for
some
Thinking!
N
N
60.
We’ll need two nitrogen
atoms, which both need
to follow the octet rule.
How many electrons do
they need to borrow from
each other?
To do that, we’ll
have to rearrange
the electrons so
that they can
share them with
each other.
N
N
Will shift →
N N
And structurally this will become:
N N
61. Nitrogen shares 3 pairs of electrons, it makes a
triple nonpolar covalent bond
62. Covalent bonds are between 2 or
more nonmetals, and usually follow
the octet rule.
63. Covalent bonds can be:
SINGLE, DOUBLE, or TRIPLE
64. Covalent bonds can also be:
POLAR or NON-POLAR
65. Let’s draw electron dot diagrams and then STRUCTURAL diagrams for these compounds.
DOTS
C2H6
C2H4
Structural
Let’s draw electron dot diagrams and then STRUCTURAL diagrams for these compounds.
DOTS
C2H6
Structural
H
H
C C
H
C2H4
H
H
H
C C
H
H H
H
H
H
H―C―C―H
H H
H H
C C
H H
66. Let’s draw electron dot diagrams and then STRUCTURAL diagrams for these
compounds. Name each bond type.
C2H2
C3H8
DOTS
Structural
Let’s draw electron dot diagrams and then STRUCTURAL diagrams for these compounds.
DOTS
Structural
C2H2
H
C C
H
H―C C―H
H-C bond is single polar covalent
C-C bond is triple nonpolar covalent
C3H8
H
H
H
C C C
H
H
H H H
H
H
H
H-C bond is single polar covalent
C-C bond is single nonpolar covalent
H―C―C―C―H
H H H
67. Let’s draw electron dot diagrams and then STRUCTURAL diagrams for these compounds.
Name each bond type.
DOTS
CO2
AsCl3
Structural
DOTS
Structural
CO2
O C O
O = C bond is a double polar covalent bond
AsCl3
Cl As Cl
Cl
As – Cl bond is a single polar covalent bond
O=C=O
Carbon dioxide is a
STRAIGHT molecule
Cl―As―Cl
Cl
Propane goes camping with some people in blue
tanks, to run the stoves. It’s a repeating type
molecule, a little chain really. The formula is C3H8.
68. Draw the Structural Diagrams now.
69. Name both kinds of bonds in this molecule
Propane formula is C3H8.
H H H
H―C―C―C―H
H H H
The C―C bond is single nonpolar covalent
and the C―H bond is single polar covalent
70. Draw Dot diagrams and structural
diagrams for oxygen dibromide.
71. Name the bonds between Br + O.
Structurally, this becomes…
O Br
Br
O Br
Br
These bonds are SINGLE POLAR COVALENT bonds.
EN diff of 3.4 - 3.0 = 0.4
72. Draw structural diagrams for
carbon tetrachloride.
73. Name the bonds between carbon and chlorine.
CCl4 The Electronegativity difference between Cl - C is
3.2 – 2.6 = of C is 0.6, These are all
single polar covalent bonds.
Cl
Cl C Cl
Cl
Cl
Cl―C―Cl
Cl
Bonding Class #5
OB: more practice with Lewis Dot diagrams,
Structural Diagrams, and we get to meet the
weird hybrid bonds of ozone, carbon
monoxide, PCl5, and NO2
Fresh minds, periodic tables at the ready!
ALLOYS
Alloys are MIXTURES of either 2 or more metals, OR metals + nonmetals. They’re not chemically bonded!
They get mixed most often by melting them together since metals are solids and don’t mix well otherwise.
The resulting “stuff” is not a new substance, it’s a mixture of the original substances.
The alloy has different properties than the original substances because they pack together mixed up.
Most alloys are made for strength, non-corrosiveness, or beauty.
Examples:
Sterling Silver made from SILVER + COPPER for strength
Cast Iron made from IRON and CARBON for strength and non-corrosiveness
Stainless steel made from IRON and CHROMIUM for strength and non-corrosiveness
Brass is made from ZINC and COPPER for durability and beauty
First we draw 2 diagrams, (structural or dots), and
determine the types of bonds present.
calcium oxide + butane C4H10.
calcium oxide
+2
[Ca] [O]
butane C4H10.
-2
This is an IONIC bond, due to
the transfer of electrons
from calcium to oxygen.
H H H H
H-C-C-C-C-H
H H H H
This has single non polar
covalent bonds at the C-C
locations, and single polar
covalent bonds at the C-H
locations
Make sure you don’t forget that metals will only make IONIC bonds!
Now, let’s look over the model of NaCl crystal (and the diagram here).
Each Na+1 cation is surrounded by 6 Cl-1 ions.
The reverse is also true,
Each Cl-1 ion is surrounded by 6 Na+1 cations
76.
We can say that the
COORDINATION NUMBER
for chloride is 6, and that
the COORDINATION
NUMBER for the sodium
cations is also 6.
Big deal, right?
77.
Well, the coordination number, at the ion level, will give rise to specific shapes
to the salt crystals when they grow up to be big enough to see with your eyes.
Uric Acid crystals look like this:
Cute little swords, really.
They cause gout, which is what can
happen to your teacher from time to
time.
The last time my sister in law Andrea
made Portobello Mushroom appetizers
for Christmas Eve 3 years back, I
woke up at 2 AM in pain and was
unable drive home the next day,
because of those crystals
(all making a home in my big toe!).
78. Let’s work on carbon monoxide now,
how does it bond together?
CO2 is so important, straight line, two double
polar covalent bonds, but what about it’s little
cousin, CO?
C
O
This will be an exceptional bond, with a cool name
too. Here goes…
C
O
Can we share 4 + 6 to get two octets? This is tricky!
CO
6
8
Looks like a
polar double
bond is in
order, but
carbon has no
octet yet.
CO
CO
~
C=O
So, oxygen will
“lend” 2 of it’s unshared
electrons to the bonding
“mix”, and it keeps an octet,
and although they are not
bonded in the same way,
carbon “gets” an octet too.
78.
It will form a double polar covalent
bond and also form what’s called a
COORDINATE COVALENT BOND.
The oxygen electrons coordinate this
situation so that carbon “gets” an octet in a
sort-of cheating way. Weird, but it happens!
79.
My shorthand for this type of bond
80. Phosphorus pentachloride is up next, a real
but weirdo molecule…
It’s used as part of fertilizer preparation, and
other chemical reactions. Draw the dot diagram.
Phosphorus Pentachloride
Cl
PCl5
Cl
Cl
P
Cl
81.
Cl
This compound breaks the octet rule, how many electrons does
phosphorus end up with? How is this possible? Isn’t it a rule???
The compound called 2-pentene has five carbons in a chain,
similar to octane (8) that we drew yesterday. The #2 means that
the one double bond fits between the 2nd and 3rd carbon in the
chain. Only hydrogen atoms are bonded to this set of five carbon
atoms.
82.
Can you draw the dots here, then the structural diagram?
(I dare you to try, no talking 2 minutes! My kids can do anything!)
The compound called 2-pentene has five carbons in a chain,
similar to octane (8) that we drew yesterday. The #2 means that
the one double bond fits between the 2nd and 3rd carbon in the
chain. Only hydrogen atoms are bonded to this set of five carbon
atoms. Can you draw the dots, and the structural diagram?
(I dare you to try, no talking 2 minutes! My kids can do anything!)
H
H
H
H―C ―C C ―C―C―H
H
H H
H
H
C5H10
John Adams High School,
in Ozone Park, New York
City at left. I graduated in
1978 on that sidewalk out
in front of the school. We
had 2 graduations that
day, 975 kids were too
many for one ceremony!
Both my sister and my brother
worked here, and they had the best
crumb buns in the world. And the
cookies were fab too.
Those stairs go up to the A Train to
Manhattan. It’s only the subway in
Manhattan, most of Queens the
subway runs on the el.
84.
Ozone and Oxygen, both pure oxygen, but one you breathe to live,
one you breathe to die. (sorry).
O3 vs. O2
85. You already know that oxygen has a double, nonpolar,
covalent bond. No need to review (right?).
86. Ozone is an ALLOTROPE of oxygen.
87. Allotropes are pure forms of an element but due to
different bonding, they have different properties. Other
allotropes are carbon in the graphite mode and carbon in
the diamond mode!
88. Try to bond 3 oxygen atoms…
88
O
O
89.
O
O
O
O
With ozone (and other molecules, like NO2, the electrons
can’t add up to full octets all around. In this case, the
oxygen atoms become “most” stable by making a
double bond and a single bond, which RESONATES,
back and forth.
It’s a resonating bond.
O
O
90.
O
O
O
O
In reality, this switching back and forth is constant, and
becomes, 2 “one and a half bonds” all the time.
Scientists know this because they can measure the bond
lengths. Single bonds are longer than double bonds.
These “resonating” bonds are 1½ sized all of the time.
Getting the ozone bonding, the RESONATING bonding correct
on the regents is a very personal thing for me. If you have any
kindness in your heart for me, please remember this bond and
how it is a “hybrid” bond (abnormal, like me!)
Bonding Class #7
OBJECTIVE:
91. Defining the 3 kinds of Intermolecular Bonding:
the weak attractions between molecules,
much weaker than ionic or covalent
bonds, but they are important and have a
real effect on the compounds
Quick review….
92. Ionic bonds form between metals (that
lose electrons) and nonmetals (that gain
electrons). The transfer of electrons result in
the formation of neutral ionically bonded
compounds, such as NaCl, MgO, or CuCl2
93. Covalent bonds form between 2 or more
nonmetals (no metals ever) by sharing electrons.
The molecules that form will have single, double or
triple bonds, and atoms follow the octet rule.
Examples are water, CH4, and CO2.
94. These bonds are all inside the compound.
95.
There are three kinds of INTERMOLECULAR BONDS, bonds
formed by the molecules with each other. These are all MUCH
WEAKER that inside the compound bonds, but they are important.
96.
Weakest to strongest they are: electron dispersion force,
dipole interaction, and hydrogen bonding.
97. When I was in college there were only 2 kinds, electron dispersion
and dipole interaction. Hydrogen bonding is very similar to dipole
interaction, and we’ll see how they work today.
98. The weakest is the electron dispersion force.
It’s created by the movement of electrons.
99.
Electron Dispersion forces.
Example one: fluorine F2
Each of these F2 molecules has a 2-7
doubled electron configuration. Each
atom has 9 electrons, the molecules
have 18 electrons.
100.
When these electrons all
“move” to one side, for a nanosecond,
there will be a temporary dipole
created, a positive side, and a
negative side of the molecule.
This allows for the weakest of
temporary attractions to exist.
F2 is a gas at STP, because the
kinetic energy at 273 Kelvin
exceeds the attractive force
of the electron dispersion
forces, so it’s a
GAS.
Electron Dispersion forces.
Example two: Chlorine Cl2
101.
Each of these Cl2 molecules
has a 2-8-7 doubled electron
configuration. Each atom has 17
electrons, the molecules have 34
electrons.
When these electrons all “move” to
one side, for a nanosecond, there
will again be temporary dipoles. This
happens more often than with
fluorine, but not often enough to
make a difference at 273 Kelvin.
This allows Cl2 to be a gas at STP,
because the kinetic energy at 273
Kelvin exceeds the attractive force
of the electron dispersion
forces, so chlorine is also a
GAS.
Electron Dispersion forces.
Example 3: Bromine Br2
102.
Each of these Br2 molecules
has a 2-8-18-7 doubled electron
configuration. Each atom has 35
electrons, the molecules have 70
electrons.
These electrons all “move” to one side,
so many times per second that there
will be a dipole created, a positive
side, and a negative side of the
molecule. This happens often enough
that these attractive forces make Br2
a liquid!
The weak but constant intermolecular
attractions accumulate.
The 273 Kelvin kinetic energy cannot
overcome the intermolecular attractions,
so bromine becomes a
liquid.
Electron Dispersion forces.
Example 4: Iodine I2
103.
Each of these I2 molecules
has a 2-8-18-8-7 doubled electron
configuration. Each atom has 53
electrons, the molecules have 106
electrons.
The electrons move so much, that a
near constant dipole exists due to
these electron dispersions.
This allows for the weakest of
temporary attractions to exist at all
times, which makes I2 a solid at STP.
The kinetic energy at 273 Kelvin
DOES NOT exceed the attractive
force of the electron dispersion
forces, so iodine is a
SOLID
104. The halogens clearly show how electron dispersion
forces accumulate and then affect the molecules.
105. When there are dipoles, that means a positive and a negative side to a
molecule (or a bond). Here, there are near permanent dipoles created
by polar bonds but ONLY IN POLAR MOLECULES.
H
H C H
S
Cl
Cl
H
We have seen already that BONDS can
be polar, due to either having a
difference in electronegativity (they
share electrons unevenly), or if they are
IONIC and transfer electrons to bond
(forming positive and negative ions)
We’re about to start talking
MOLECULAR POLARITY, is the
whole molecule polar (different than
bond polarity).
106. Molecular polarity is based upon
SHAPE OF THE MOLECULE.
107. If the molecule is “balanced” it will
be nonpolar.
108. The balance, or SYMMETRY we’re
looking for is called RADIAL SYMMETRY
109. There are other symmetries,
but they DO NOT matter.
110.
In SCl2, the bonds are single polar covalent. The
molecule itself is polar because it does not have
radial symmetry. So, the sulfur will become
positively charged most of the time, and the
chlorine atoms will be negative most of the time.
S
Cl
Cl
SCl2 had
what sort of
symmetry?
Gingerbread
Man
symmetry?
111. Methane, which has polar bonds too,
also has radial symmetry.
112. This offsets that polarity, and the molecule is
nonpolar. SCl2 will be liquid at room temp
while methane would be a gas.
Why???
H
S
Cl
H C H
Cl
H
113. Draw these: All the positive sulfur atoms are nearly permanently
attracted to the negative chlorine atoms. The EN difference in a polar
molecule can create intermolecular bonds called dipole attractions.
S
S
Cl
Cl
S
Cl
Cl
S
Cl
Cl
S
Cl
Cl
Cl
Cl
114.
Draw these. These methane molecules
(nonpolar) have nearly no attraction to each other, so
they will be gas at room temperature. Dipole
attraction is way less powerful than ionic or even
covalent bonding, but it can affect the phase of the
compound. Nonpolar molecules are hardly attractive
to each other.
H
H
H C H
H C H
H
H
H
H C H
H H
H C H
H
Is there ANY attraction here between molecules?
115. Hydrogen bonding is exactly the same as dipole
attraction, but, and it’s a SMALL but, hydrogen has to be
present in the molecule.
116.
H has a much smaller EN value than most other
atoms, so when it’s included, like with water, the dipole it
creates is usually much stronger than when it’s something
like SCl2.
117.
The EN difference between chlorine and sulfur is 3.2 – 2.6 = 0.6
118.
The EN difference between oxygen and hydrogen is 3.4 – 2.2 = 1.2
119.
This greater difference creates a “stronger” dipole. Strong enough that
we now have to give it a new name. Instead of just calling it a strong
dipole attraction, we call it hydrogen bonding.
S
Cl
O
Cl
H
H
120. Draw these now. All of the negative
O
H
O
H
oxygen are magnetically attracted to the positive
hydrogen atoms in nearby molecules. This is an
intermolecular attraction. Hydrogen bonding is
the strongest of the 3 intermolecular attractions.
H O
H
H
O
H H
O
H
H
O
H
H
H
121
Give an example molecule (or formula unit) for each type of bond:
Ionic
Single nonpolar covalent
Single polar covalent
Double nonpolar covalent
Double polar covalent
Triple non polar covalent
Triple polar covalent
Coordinate covalent
Resonant
Ionic + Covalent at the same time
Breaks the octet rule (more than 8e-)
Breaks the octet rule (less than 8e-)
Give an example molecule (or formula unit) for each type of bond:
Ionic……………………………………………………………………………………..NaCl
Single nonpolar covalent……………………………………………………F-F
Single polar covalent………………………………………………………….H-Cl
Double nonpolar covalent………………………………………………….O=O
Double polar covalent………………………………………………………..O=C=O
(both)
Triple non polar covalent…………….…………………………………….NΞN
Triple polar covalent………………………………………………………….NΞC-H
Coordinate covalent…………………………………………………………..carbon monoxide
Resonant……………………………………………………………………………….ozone O3
Ionic + Covalent at the same time………………………………….CuSO4·5H2O*
Breaks the octet rule (more than 8e-)………………………….PCl5
Breaks the octet rule (less than 8e-)…………………….………H-H (too small)
* Also has hydrogen bonding as well. (wow!)
Bonding Class #8
OB: master relative oxidation numbers,
review all bonding
for celebration tomorrow
-----------------
CuSO4·5H2O*
A long, long time ago, in a galaxy, far, far away…
This is going to be great!
122. We learned about oxidation numbers, those little
positive and negative numbers in the corners of the
periodic table, that told us what ratios of atoms to
atoms molecular compounds make.
Time to revisit them.
123. Hydrogen has a +1 and a -1 oxidation number.
Oxygen has only a -2 oxidation number.
To “make” molecules, you have to combine atoms to atoms, so that the sum of the
oxidation number is zero. These are numbers, not ion charges!
Since oxygen is only a -2, it will
take two +1 hydrogen atoms to
make a molecule. That is why the
formula is H2O, and that’s why
H3O or HO is not a real compound.
124.
Let’s determine the relative oxidation numbers in these molecules…
HCl
CH4
CO2
Let’s determine the relative oxidation numbers of the atoms in these molecules…
HCl
+1
H
-1
Cl
CH4
-4
C
+1
H
(-4) + 4x(+1) = 0
CO2
+4
C
-2
O
(+4) + 2x(-2) = 0
(+1) + (-1) = 0
125
Sulfur dioxide
SO2
S+4 O-2 O-2 (0)
Chromate ion
CrO4-2
Cr+6 O-2 O-2 O-2 O-2 (-2)
Permanganate ion
126
NH3
127
NaOH
128
KClO3
129
Carbon monoxide
130
Carbon dioxide
131
Dihydrogen sulfate
132
Nitrate ion
133
Nitrogen dioxide
134
Phosphorus
trichloride
Sulfur dioxide
SO2
S+4 O-2 O-2 (0)
Chromate ion
CrO4-2
Cr+6 O-2 O-2 O-2 O-2 (-2)
Permanganate ion
MnO4-2
Mn+6 O-2 O-2 O-2 O-2 (-2)
ammonia
NH3
N-3 H+1 H+1 H+1 (0)
Sodium hydroxide
NaOH
Potassium chlorate
KClO3
Carbon monoxide
CO
Carbon dioxide
CO2
Dihydrogen sulfate
H2SO4
H+1 H+1 S+6 O-2 O-2 O-2 O-2 (0)
Nitrate ion
NO3-1
N+5 O-2 O-2 O-2 (-1)
Nitrogen dioxide
NO2
Phosphorus
trichloride
PCl3
Na+1 O-2
H+1 (0)
K+1 Cl-5 O-2 O-2 O-2
C+2
C+4
O-2 (0)
O-2 O-2 (0)
N+4 O-2 O-2
P+3
Cl-1
(0)
(0)
Cl-1 Cl-1 (0)
135.
Review (push yourself)
Name a compound or molecule or formula unit for each type of bond:
Single polar covalent
Double polar covalent
Triple polar covalent
Single nonpolar covalent
Double nonpolar covalent
Triple nonpolar covalent
Ionic
Resonant
Coordinate covalant
Breaks octet rule (too small)
Breaks octet rule (too big)
Review (push yourself)
Name a compound or molecule or formula unit for each type of bond:
Single polar covalent
H-Cl
Double polar covalent
O=C=O
Triple polar covalent
NΞC-H
Single nonpolar covalent
Cl-Cl
Double nonpolar covalent
O=O
Triple nonpolar covalent
Ionic
Resonant
NΞN
KCl MgO
O3 ozone
Coordinate covalant
CO
Breaks octet rule (too small)
H-H
Breaks octet rule (too big) PCL5
Intermolecular bonding system Jeopardy!
136. It keeps ammonia NH3 together as a liquid
what is…
137. It keeps Br2 bromine a liquid, but iodine I2 a solid
what is…
138. It keeps phosphorus trichloride PCl3 together
as a liquid
what is…
Intermolecular bonding system Jeopardy!
It keeps ammonia NH3 together as a liquid
What is hydrogen bonding?
It keeps Br2 bromine a liquid, but iodine I2 a solid
What is the electron dispersion force or
electron dispersion attraction?
It keeps phosphorus trichloride PCl3 together as a liquid
What is the dipole attraction force?
139. In one sentence explain
the difference between bond
polarity and molecular polarity.
Who has the guts to stand and orate
this one?
Bond polarity is when there is a difference in electronegativity value between two
atoms that are bonding. All ionic bonds are polar, but for covalent bonds we have to
check table S.
Molecular polarity has to do with molecular shape.
If a molecule has radial symmetry, it is a nonpolar molecule.
A molecule that doesn’t exhibit radial symmetry is polar.
A polar molecule water
A non polar molecule CCl4
In Queens, especially in Ozone Park, you can get on the A train and go to Brooklyn.
Then you get off, cross the platform, and go back to Ozone Park in Queens. You
can do this over and over all day long, all night long, all for one price.
You can resonate back and forth from Queens to Brooklyn.
The bonds in ozone O3 resonate back and forth, they are exceptional bonds, but
they exist.
In reality the bonds are “both” 1½ sized rather than small doubles and bigger
single bonds.
140.
Once and for all, with a little dot diagram, and one sentence,
explain how carbon monoxide bonds together.
C O
Once and for all, with a little dot diagram, and one sentence, explain how
carbon monoxide bonds together.
C O
It’s called a double polar covalent bond (the bottom 2 pairs
of electrons) and a coordinate covalent bond, which means
oxygen just “lends” 2 electrons into the mix so carbon
“gets” an octet too.
141. True or False?
1. Ionic bonds can be double or single bonds
2. Covalent bonds cannot be nonpolar bonds
3. Oxygen molecules have double polar covalent bonds
4. Nitrogen molecules have double nonpolar covalent bonds
5. Hydrogen atoms can make single or double covalent bonds
6. Oxygen atoms must make double bonds ONLY
7. Water is sometimes a straight line molecule by shape
8. Molecules with polar bonds can never be non polar molecules
9. Molecules with nonpolar bonds only can never be polar molecules
10. The weakest intermolecular bond is the dipole force of attraction
True or False? ALL FALSE!!!
1. Ionic bonds can be double or single bonds No, just magnetic ionic
2. Covalent bonds cannot be nonpolar bonds No, F2 or Cl2 are nonpolar
3. Oxygen molecules have double polar covalent bonds
No, double nonpolar
4. Nitrogen molecules have double nonpolar covalent bonds No, triple nonpolar
5. Hydrogen atoms can make single or double covalent bonds No, only single
6. Oxygen atoms must make double bonds ONLY No, in water they make 2 singles
7. Water is sometimes a straight line molecule by shape No, always, always bent!
8. Molecules with polar bonds can never be non polar molecules No, CO2 or CH4
9. Molecules with nonpolar bonds only can never be polar molecules No, NBr3
10. The weakest intermolecular bond is the dipole force of attraction
No, electron dispersion forces are weakest, watch them in Group 17
Study tonight, and every night.
We celebrate next
Wed+ Thurs
Bonding + Water
2 class periods long
Friday’s NO SCHOOL!