Objectives
Chemistry 30
Chapter 6 - Chemical Bonding
1.
Use quantum theory to explain why atoms have distinctive electron configurations (chapter 4).
2.
Be able to draw electron configurations for atoms and ions (chapter 4).
3.
Discuss the meaning of the term valence electron and why they are important (chapter 4).
4.
Be able to draw orbital diagrams for atoms and ions (chapter 4).
5.
Use the concept of valence electrons draw Lewis diagrams for atoms and ions (chapter 6).
6.
Use the concept of Lewis diagrams to explain bonding and non-bonding electrons.
7.
Explain how and why elements from groups 2, 13 and 14 undergo electron promotion.
8.
Explain how and why elements from groups 15, 16, 17 and 18 undergo valence level expansion.
9.
Define chemical bond; differentiate between covalent and ionic bonds. (6-1)
10.
Explain why most chemical bonding is neither purely ionic nor purely covalent. (6-1)
11.
Classify bonding type according to electronegativity differences. (6-1)
12.
Explain bonding in terms of energy and the octet rule. (6-2)
13.
Be able to draw Lewis structures and structural diagrams for covalently bonded molecules. (6-2)
14.
Be able to use Lewis structures and structural diagrams to show double and triple bonds. (6-2)
15.
Explain what is meant by resonance structures. (6-2)
16.
Explain what is meant by a coordinate covalent bond.
17.
Be able to write a Lewis structure for a polyatomic ion.
18.
Explain the importance of isomerism in the properties of chemical compounds.
19.
List and compare the distinctive properties of ionic and molecular compounds. (6-3)
20.
Describe the electron-sea model of metallic bonding, and explain why metals are good electrical
conductors. (6-4)
21.
Explain why metal surfaces are shiny. (6-4)
22.
Explain why metals are malleable and ductile but ionic-crystalline compounds are not. (6-4)
23.
Explain VSEPR theory. (6-5)
24.
Predict the shapes of molecules or polyatomic ions using VSEPR theory. (6-5)
25.
Explain how the shapes of molecules are accounted for by hybridization theory. (6-5)
26.
Describe dipole-dipole forces, hydrogen bonding, induced dipoles and London dispersion forces. (6-5)
27.
Explain what determines molecular polarity and be able to predict the polarity of a given molecule. (6-5)
28.
Discuss the relationship between polarity and physical properties such as boiling point and freezing point.
29.
Compare the physical properties that result from the various intermolecular forces.
Vocabulary
angular molecule
electric dipole
isomer
polar molecule
tetravalent
amorphous solid
crystal
evaporation
hydrogen bond
phase diagram
vapor pressure
bonding electron
electronegativity
linear molecule
pyramidal molecule
trivalent
boiling point
crystalline solid
glass
intermolecular
sublimation
viscosity
covalent bond
electron promotion
lone pair electron
structural isomer
univalent
condensation
deposition
heat of fusion
intramolecular
surface tension
divalent
ionic bond
planar molecule
tetrahedral molecule
covalent-network solid
dipole
heat of vaporization
network solid
vaporization
Notes



Chemical bonding is all about the interactions of atoms involving their valence electrons.
a bond is a force of attraction between two atoms
the source of the force are attractions and repulsions between electrons and the atomic nucleii
Before we discuss bonding we must review some material from grade 11 chemistry:
Quantum Mechanics and Atomic Orbitals
o
o
o
o
o
Quantum mechanics is a mathematical treatment into which both the wave and particle nature of matter could
be incorporated.
since the electron is both a wave and a particle it is impossible to give it’s location or speed with certainty.
gives a probability density map of where an electron has a certain statistical likelihood of being at any given
instant in time.
The equations which describe the wave function of electrons requires three quantum numbers:
o Principal quantum number, n.
 This relates to the energy of the electron
 As n becomes larger, the atom becomes larger and the electron is further from the
nucleus.
 This is directly related to the period of the atom on the Periodic Table
o Angular momentum quantum number, l.
 This quantum number defines the shape of the orbital.
 There are 4 shapes:
 s - begins at n = 1
 p - begins at n = 2
 d - begins at n = 3
 f - begins at n = 4
 Theoretical g, h, i, etc. orbitals exist, but no atoms have been created to use them.
o Magnetic quantum number, ml.
 Magnetic quantum numbers give the three-dimensional orientation of each orbital.
 s - has 1 orientation
 p - has 3 orientations
 d - has 5 orientations
 f - has 7 orientations
2
The s Orbitals
o
o
All s orbitals are spherical.
As n increases, the s orbitals get larger.
The p Orbitals
o
o
o
There are three p orbitals: px, py and pz.
They look like propellers
The letters correspond to allowed the values of ml of –1, 0, and +1.
The d and f Orbitals
o
o
o
o
o
There are five d and seven f orbitals.
Three of the d orbitals lie in a plane bisecting the x-, y-, and z-axes.
Two of the d orbitals lie in a plane aligned along the x-, y-, and z-axes.
Four of the d orbitals have four lobes each.
One d orbital has two lobes and a collar.
o
Spin Quantum Number, ms
 electrons have spin, which creates a magnetic field
 there are two spin states possible, +1/2 and -1/2
 a single orbital can hold a maximum of two electrons, which must have opposite spin.
3
Electron Configuration




is the home address of the electrons about an atomic nucleus
is dependent on the quantum numbers
Electrons tend to occupy the lowest available orbital.
o The simplest atom, hydrogen has 1 electron.
o In its’ lowest, or ground state, this electron will occupy the 1s orbital, the lowest energy orbital
available (see chart, page 105)
o The next element, helium, has two electrons, both of which will occupy the 1s orbital.
o Element three, lithium, has three electrons. The first two will fill the 1s orbital while the third must
move up to the next energy level, 2s.
Thus the electron configuration of an atom is the arrangement of the electrons from the lowest energy
level to the highest.
Electron configurations of period two elements
Representative
Element
Electron configuration
Group

1
lithium
1s22s1
2
beryllium
1s22s2
13
boron
1s22s22p1
14
carbon
1s22s22p2
15
nitrogen
1s22s22p3
16
oxygen
1s22s22p4
17
fluorine
1s22s22p5
18
neon
1s22s22p6
Note that the order of the orbitals does not always follow numerically with the energy level or the period.
For instance, the 3d level comes between the 4s and the 4p. The reason for this has to do with quantum
theory and is not important here. It is important, however, that the order be observed when giving the
electron configuration of an element. For instance, the electron configuration of selenium (34 protons, 34
electrons) is:


1s22s22p63s23p64s23d104p4
Electron configuration is simply based on the total number of electrons. A cation is formed if an atom
loses electrons. An anion is formed if an atom gains electrons. In either case, just determine the total
number of electrons and write the appropriate electron configuration.
Electron Promotion

Atoms enter into chemical bonds in order to fill their outer ‘s’ and ‘p’ orbitals. These are the electrons on
the outside of the atom, also called valence electrons.
o since the ‘s’ and ‘p’ orbitals comprise 4 pairs of 2 electrons, there are 8 electrons possible (an
octet) in the valence shell.
o paired electrons are non-bonding (lone-pair electrons)
o orbitals with only 1 electron seek to find another electron to pair with. These are called bonding
electrons.

Electron promotion (or hybridization) involves adding energy to an electron in an ‘s’ orbital so that it can
move up to a ‘p’ orbital in the same quantum.
This is done so that a non-bonding pair of ‘s’ electrons can both become unpaired bonding electrons, one
in the ‘s’ orbital and the other in the ‘p’.

4

If you look at the orbital energy table you were given or the one on page 105 you can see the energy
differences between electrons in different orbitals.
o The valence electrons in the ‘s’ and ‘p’ orbitals for each energy level have a very small difference in
energy, so it is possible for an electron in an ‘s’ orbital to receive a small input of energy and move
up into the adjacent p orbital.

Why electron promotion happens can be explained in terms of energy.
o All atoms desire to achieve the lowest possible total energy.
o When an atom bonds with another atom it moves to a lower energy level, thus explaining why
atoms bond.
o When an atom promotes an electron, it gains energy, but the extra bonding capacity achieved
allows it to get to a lower total energy than it could without electron promotion.
o See the graph below:



Point 1 represents the energy of an unhybridized atom. Point 2 represents the energy increase needed
to promote an electron. Point 3 represents the net decrease in energy which occurs when a hybridized
atom forms bonds with other atoms.
Electron promotion happens only when it increases the bonding capacity and only happens between
the s and p orbitals of the same energy level.
Following these criteria, hybridization occurs only in atoms of groups 2, 13, and 14:
Electron Configuration of Hybridized and Unhybridized Atoms
Element
Unhybridized
Hybridized
beryllium
1s22s2
1s22s12p1
boron
1s22s22p1
1s22s12p2
carbon
1s22s22p2
1s22s12p3
Note: For atoms of groups 2, 13, and 14, hybridization is the rule. It happens all the time for these elements, so it
must be taken into account when using these elements.
Assignment: Write the electron configuration for the following elements and ions:
Si, S, P, Ca, As, Fe, Br, Kr, At, U, Na1+, F1-, Ne
5
Orbital Diagrams

These give information concerning bonding.

They are best learned by comparison with electron configuration:

eg.

electron configuration:

orbital diagram :
Na (11 protons, 11 electrons)
1s22s22p63s1
1s
2s
↑↓
↑↓
2p
↑↓ ↑↓ ↑↓
3s
↑

Each line represents an orbital.

Note the three lines under the 2p orbital to represent the three sub-orbitals.

An arrow above a space represents an electron occupying an orbital. It can be either up or down, to
represent the spin (+1/2 or -1/2) of the electron.


Two arrows, one up, the other down, represents a situation where the orbital is full.
Note how electron promotion is incorporated using orbital diagrams.
o In the elements from groups 2, 13 and 14 and electron is taken from the ‘s’ orbital of the valence
electrons and placed in the ‘p’, giving each element 2 more bonding electrons than it would
normally.
Group
Electron arrangements and orbital diagrams of period two elements
Representative
Element
Electron configuration
Orbital Diagram
1s
2
1
1
lithium
1s 2s
2
beryllium
1s22s12p1
2
1
↑
↑↓
↑
↑
↑↓
↑
↑ ↑
boron
1s 2s 2p
14
carbon
1s22s12p3
↑↓
↑
↑ ↑ ↑
3
↑↓
↑↓
↑ ↑ ↑
15
nitrogen
1s 2s 2p
16
oxygen
1s22s22p4
↑↓
↑↓
↑↓ ↑ ↑
17
fluorine
1s 2s 2p
5
↑↓
↑↓
↑↓ ↑↓ ↑
18
neon
1s22s22p6
↑↓
↑↓
↑↓ ↑↓ ↑↓
2
2
↑↓
2p
2
13
2
2s
2
Assignment:
Repeat the last assignment, giving the orbital diagrams for the elements and ions.
6
Electron dot (Lewis) diagrams




Are another way to illustrate the position of the electrons in an atom or ion.
It is more compact than the electron configuration, because it gives information only concerning the
valence electrons.
Valence electrons are the electrons on the outside of an atom;
o they are the electrons responsible for bonding and are also the electrons gained or lost when an
atom ionizes.
o Valence electrons are electrons in the s and p orbitals of the highest energy level reached by the
electrons of an atom.
o In this class when valence electrons are mentioned, the only elements concerned are those in
groups 1, 2, and 13 through 18.
An electron-dot diagram (or Lewis diagram) begins with the symbol of the element.
o Dots are drawn around the symbol to represent the valence electrons; one for each electron.
o The placement of the electrons is important; they must be placed in four pairs, left, right, top, and
bottom.
o Pay close attention to the placement of the electrons in the table below:
Electron arrangements, orbital diagrams and Lewis diagrams of period two elements
Group
Representative
Electron
Element
configuration
Valence
electrons
1
lithium
1s22s1
1
2
beryllium
1s22s12p1
2
13
boron
1s22s12p2
3
14
carbon
1s22s12p3
4
15
nitrogen
1s22s22p3
5
16
oxygen
1s22s22p4
6
17
fluorine
1s22s22p5
7
18
neon
1s22s22p6
8
Lewis Diagrams of Ions
When a cation forms it loses its’ valence electrons:
Sodium:
Sodium ion:
7
Lewis
diagram
When an anion forms, it gains enough electrons to fill its’ octet:
Sulfur:
Sulfur ion:
Assignment:

Give the correct Lewis diagram for K, Sr, Al, Se, Ca 2+ and Br1-
This table gives a summary of the electron arrangements characteristic of elements from groups 1, 2 and 13 to
18, the groups with outer electrons in the ‘s’ and ‘p’ orbitals.
o electron promotion is a given for elements in groups 2, 13 and 14
o both the orbital diagram and the Lewis diagram show the bonding capabilities of each atom.
 Where two dots are together or arrows are paired, these represent paired valence
electrons which fill an orbital and cannot be used for bonding.
 These are called lone pair electrons.
 Where a single dot or arrow is found, this represents an unpaired electron in a half-filled
orbital. This electron can be used for bonding and is thus called a bonding electron.
1s
Orbital and Lewis diagrams with Lone Pair and Bonding Electrons
Valence
Electron-dot
Lone Bonding
Bond
Orbital Diagram
electrons
Diagram
Pairs Electrons
Capacity
2s
2p
lithium
↑↓
↑
beryllium
↑↓
↑
↑
2
boron
↑↓
↑
↑ ↑
3
carbon
↑↓
↑
↑ ↑ ↑
4
nitrogen
↑↓
↑↓
↑ ↑ ↑
5
oxygen
↑↓
↑↓
↑↓ ↑ ↑
6
fluorine
↑↓
↑↓
↑↓ ↑↓ ↑
7
neon
↑↓
↑↓
↑↓ ↑↓ ↑↓
8
Element
Li
1
Be
B
C
N
O
F
Ne
8
0
1
1
0
2
2
0
3
3
0
4
4
1
3
3
2
2
2
3
1
1
4
0
0
Valence level expansion

The number of bonding electrons is a good guide to the number of other atoms a given element can interact
with. Number of bonding electrons = Number of chemical bonds

Some compounds occur which cannot be easily explained with the information given so far:
PF5, SF6, ClF7, ArF8



Normally these elements have a bonding capacity much lower than is indicated by these formulas.
o Phosphorus, in group 15, should make 3 bonds. Here it makes 5
o Sulfur, in group 16, should make 2 bonds. Here it makes 6
o Chlorine, in group 17, should make 1 bond. Here it makes 7
o Argon, in group 18, should not bond at all. Here it makes 8
What is happening is that every valence electron becomes a bonding electron. The number of valence orbitals
expand allowing a large increase in bonding capacity.
o These elements promote electrons from their outer ‘s’ and ‘p’ orbitals into the ‘d’ orbital of the
same quantum.
o The reason for this is the same as for hybridization; to achieve a lower total energy.
Elements which expand their valence level include elements of groups 15 to 18, from period 3 down.
o The reason for this is that expansion is into the unused ‘d’ orbital of the same energy level;
periods 1 and 2 do not have a ‘d’ orbital, so there is no valence level expansion for elements of
these periods.
For example, P as it exists normally:
1s
2s
2p
3s
3p
↑↓
↑↓
↑↓ ↑↓ ↑↓
↑↓
↑ ↑ ↑
P with valence level expansion:
1s
2s
2p
3s
3p
↑↓
↑↓
↑↓ ↑↓ ↑↓
↑
↑ ↑ ↑
3d
↑
Please note that valence level expansion is the exception, not the rule. The only time you will be expected
to use v.l.e. in a problem is if I ask about it directly or ask you to explain a molecule like PF 5.
9
Chemical Bonding
When two atoms are joined together chemically, there is a chemical bond between them. Bonds form for reasons
of energy. When atoms combine to form molecules they achieve a lower energy state than they could as individual
atoms (see figure 6-5, page 165). There are two main reasons why elements will combine:
1)
To share electrons - All elements want to have the same number of valence electrons as the noble
gases, and will either give electrons away or steal electrons in order to achieve this. Sometimes,
however, this is not possible. The next best thing is to get close to another atom with the same
problem. (see figure 6-7, page 167). For example, oxygen has 6 valence electrons. It would like to
have 8 in order to have the same number of valence electrons as Neon. If it cannot steal two electrons
it will get close to another oxygen atom. Each atom will put forward its 2 bonding electrons and their
electron clouds will mesh, forming the compound O 2. This type of bonding is called covalent bonding.
2)
To balance charge - When atoms gain or lose electrons they become ions. These ions tend to be
attracted to ions of opposite charge and repelled by ions of the same charge, much like the poles of a
magnet. These ions will combine in proportions which completely balance the charges and form a
compound which is electrically neutral You used this principle when you studied chemical nomenclature.
This type of bonding is called ionic bonding (see figure 6-17, page 179)
The number of bonds made by an atom can be predicted, based on the position of the atom on the Periodic Table:
Group
Valence
Electrons
Covalent Bonds
Ionic
Bonds
1
1
1
1 (1+)
2
2
1
2 (2+)
13
3
1
3 (3+)
14
4
1
-
15
5
16
6
17
7
18
8
1
electron promotion
2
3
4
2
3 (5 )
2
2 (6 )
2
1 (7 )
2
0 (8 )
2
valence level expansion
10
3 (3-)
2 (2-)
1 (1-)
0
Electronegativity

Electronegativity is a measure of how strongly an atom is holding on to its valence electrons.
 It is related to ionization energy (chapter 5).
 If an atom loses an electron fairly easily it has a low electronegativity (and tends to be a cation).
 If an atom tends not to lose electrons, but tends to steal them from other atoms (and become an
anion) it has a high electronegativity.

To



determine what type of bond exists between two atoms you subtract their respective electronegativities:
if the electronegativity difference is 0.2 or less, the bond is covalent
if the electronegativity difference is 1.7 or greater the bond is ionic.
if the electronegativity difference is greater than 0.2 but less than 1.7 the bond is called a polar
covalent bond.
- When the electronegativity difference is in this range the atom with the greater electronegativity is
strong enough to pull the bonding electrons so that they spend more time around it than the other
atom, but is not strong enough to pull the bonding electrons away completely and form ions.
- These bonding arrangements form electric dipoles where one end has a slightly positive charge
and the other end has a slightly negative charge. The greater the electronegativity difference, the
greater the dipole:
Electronegativity
Element bonded to F
Fluorine
Electronegativity
difference
Bond type
Compound
Bond
FF
FF
F
4.0
4.0
0.0
F2O
OF
O
3.5
4.0
0.5
NF3
NF
N
3.0
4.0
1.0
CF4
CF
C
2.5
4.0
1.5
BF3
BF
B
2.0
4.0
2.0
BeF2
BeF
Be
1.5
4.0
2.5
LiF
LiF
Li
1.0
4.0
3.0
COVALENT
│
↑
increasingly ionic │
│
│


│
│


│
│


│
│
│ increasingly covalent
↓
│
IONIC
Assignment:
Determine the electronegativity difference for each chemical bond. If the bond is polar covalent draw an arrow in the
direction of the dipole, from positive to negative:

C-H
N-H
B-F
S-O
P-H
Si - Cl
Cu - Br
N-I
Br - Cl
O-H
C - Cl
C-O
| 2.5 - 2.1 | = 0.4
polar covalent
11
Making diagrams of covalently-bonded molecules
Any covalently-bonded molecule is held together by shared valence electrons. The types of molecules this method
covers includes molecular compounds (non-metal bonded to non-metal) and organic compounds (containing C and
H). We will also include some group 1 and 2 elements.
Both orbital diagrams and electron-dot diagrams may be used to show how atoms bond to form molecules.
It is important to remember that only the valence electrons may be used for bonding, and of those, only the
bonding electrons may be shared. Thus an element like chlorine with 3 lone pairs and 1 bonding electron may
form only 1 bond.
The key for making diagrams of these molecules is to remember some simple facts:
i)
Atoms form covalent bonds in order get close to a complete valence set of electrons; an atom will
continue to form bonds until it is surrounded by a complete set of valence electrons (octet rule).
ii)
When you are given a chemical formula, the assumption is that each atom in the molecule has a
complete set of valence electrons surrounding it; you should be able to draw a structure which satisfies
the bonding capacities of each atom in the molecule.
iii)
When given a molecule for diagramming which has more than 2 atoms, look for bonding capacity. The
atom (or atoms) with the greatest bonding capacity will be the central atom; the atom at the middle of
the molecule which has all the other atoms attached to it.
There are two types of diagrams you should be able to draw; lewis diagrams and structural diagrams.
One Method for Drawing Lewis Diagrams
1. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.
 If it is an anion, add one electron for each negative charge.
 If it is a cation, subtract one electron for each positive charge.
2. The central atom is the least electronegative element that isn’t hydrogen. Connect the
outer atoms to it by single bonds.
3. Fill the octets of the outer atoms.
12
4. Fill the octet of the central atom.
5. If you run out of electrons before the central atom has an octet…
…form multiple bonds until it does.
Assignment:
Go back to the molecular model lab. You already have 3-D structural diagrams for several molecules. Make Lewis
diagrams and flat structural diagrams for the molecules in questions 1, 2 and 3. You will do several in-class as examples.
Coordinate Covalent Bond
Explains structures where conventional covalent bonding does not work. The best example is the ammonium ion, NH 41+.
It is a stable, covalently bonded polyatomic ion. It is the preferred form of ammonia (over NH 3). The question is, where
does the extra hydrogen go? A coordinate covalent bond is one where one atom provides both the bonding electrons; it
uses a lone pair to bond with an atom that has no bonding electrons.
Here is an example. Write down the drawings in class.
13
Resonance
Resonance refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure.
Sometimes there is a structure that includes both a double and a single bond. If the two bonds could reverse positions
without changing any other bonds in the molecule, they can and do reverse positions. Write down the examples in class of
ozone and benzene. The main result of resonance is in bond energy. Bond energy refers to the amount of energy that
must be applied to a chemical bond to break it. Single bonds take less energy to break than double bonds. Thus a
molecule held together with single bonds is easier to take apart than one held together with double bonds. When
resonance occurs the bonding electrons responsible for the double bond change position so often that the bond energy of a
resonant structure is approximately midway between that of a single bond and a double bond. This can lead to these
structures being very stable, like benzene. Sometimes resonance does not result in a stable molecule; the example is
ozone.
Here is a Lewis diagram of ozone:
You will notice that the second drawing is just as likely:
The reality is that the actual molecule is somewhere in-between. See the other examples in class.
Molecular Shape
The shape of molecules is determined by the nature of the central atom. The central atom is the atom in a
molecule with the greatest bonding capacity and can be an atom with an ability to make 2 or more bonds. This
limits this discussion to elements from groups 2, 13, 14, 15 and 16.
Molecular shape is guided by one principle; the tendency of electrons to repel one another. This includes the
electrons from individual atoms in a chemical compound. The electron cloud from each atom in a compound will
arrange itself so as to put the maximum distance between it and all the other electron clouds from the other atoms
in the compound. Because the electrons in question are the valence electrons, the theory which describes
molecular shape is called Valence Shell Electron Pair Repulsion (VSEPR) Theory.
Picture an element like beryllium; it has two bonding electrons and no lone pair electrons. It thus can bond with
two other atoms to create a compound like BeF2. Each bond consists of two pairs of electrons, each forming its
own electron cloud. These electron clouds repel one another and try to get as far from each other as possible.
This compound forms the following structure about the central atom:
F  Be  F
This is called a linear structure, meaning the three atoms arrange themselves in a line, so that the electron clouds
from the two fluorines are the maximum distance apart.
Look at figure 6-5 on page 186 of your text. It summarizes the electron configuration about the central atoms
from groups 2, 13 to 16, in addition to two arrangements which result from valence level expansion. Also see the
VSEPR table you were given in class.
14
Two shapes; pyramidal and angular, result when the bonding electrons (and the atoms bonded to them) must also
deal with repulsion from the lone pair electrons on the central atom. Group 15 elements, like nitrogen, have 3
bonding electrons and 1 lone pair electron. The three bonds are forced out of the planar position by the lone pair,
resulting in a pyramidal structure. Group 16 elements, like oxygen, have 2 bonding electrons and 2 lone pairs.
The two bonds made are forced out of a linear shape by repulsion from the two lone pairs, resulting in an angular
shape.
Following is a summary of shape relationships:
Group 2 Geometries
Central Atom
Bonding Electrons
Lone Pair
Electrons
Bond Type
Shape
Magnesium
2
0
all single
linear
one double
linear
Group 13 Geometries
Central Atom
Bonding Electrons
Lone Pair
Electrons
Bond Type
Shape
Boron
3
0
all single
trigonal planar
one double
linear
one triple
linear
Group 14 Geometries
Central Atom
Bonding Electrons
Lone Pair
Electrons
Bond Type
Shape
Carbon
4
0
all single
tetrahedral
one double
trigonal planar
one triple, or
two double
linear
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Summary for Groups 2, 13 & 14

because there are no lone pair electrons:
 central atom bonded to 1 atom:
 central atom bonded to 2 atoms:
 central atom bonded to 3 atoms:
 central atom bonded to 4 atoms:
linear
linear
trigonal planar
tetrahedral
Group 15 Geometries
Central Atom
Bonding Electrons
Lone Pair
Electrons
Bond Type
Shape
Nitrogen
3
1
all single
trigonal pyramidal
one double
angular
one triple
linear
all single
trigonal bipyramidal
5
0
Group 16 Geometries
Central Atom
Bonding Electrons
Lone Pair
Electrons
Bond Type
Shape
Oxygen
2
2
all single
angular
one double
linear
all single
octahedral
6
0
Summary for Groups 15 & 16

because lone pair electrons are on the central atom:
 central atom bonded to 1 atom:
linear
 central atom bonded to 2 atoms:
angular
 central atom bonded to 3 atoms:
trigonal pyramidal
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Geometries for Groups 17 & 18

group 17
 normally makes 1 chemical bond (linear)
 with valence level expansion can make 7 bonds (ClF 7).

group 18
 normally makes no bonds
 with valence level expansion can make 8 bonds (ArF 8).
Assignment:
Go back to the molecular modeling assignment. Determine the shape about each central atom in each molecule
for questions 1, 2 and 3
Polarity
Polarity is concerned with whether a molecule has a positively charged end and a negatively charged end, or not.
If this is the case the molecule is polar and the chemical and physical properties of such molecules will be altered.
The importance of this will be discussed in the next unit.
Whether a molecule is polar or not depends on two things:
i)
ii)
the existence of bond dipoles
molecular symmetry
Bond dipoles exist when the electronegativity difference between the members of a bond is sufficiently high
(between 0.2 and 1.7). If a molecule contains bond dipoles it has the potential to be polar.
Molecular symmetry is concerned with how the atoms in a molecule are distributed about a central atom.
Molecules may be either symmetrical or asymmetrical. Of the groups on the periodic table we have studied, this
applies to groups 2, 13, 14, 15 and 16. Of these groups, 2, 13 and 14 form symmetrical molecules while groups 15
and 16 form asymmetrical molecules. Very simply, polarity and symmetry are related as follows:
IF NO BOND DIPOLES EXIST (i.e. all bonds are covalent):
molecule is NON-POLAR, no matter what the shape.
IF BOND DIPOLES EXIST:
a)
If the molecule is symmetrical the forces will cancel:
e.g. CCl4, BCl3 and BeCl2
molecule is non-polar
b)
If the molecule is non-symmetrical molecule the forces will reinforce:
eg. NH3, H2O
molecule is polar
IF THE CENTRAL ATOM IS BONDED TO MORE THAN ONE KIND OF ATOM:
Look for the direction of the bond dipoles in the molecule and decide on a case-by-case basis
e.g. CHCl3.
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Assignment:
Going back to the molecular modeling activity:
a)
Use electronegativity difference to determine if bond dipoles exist. Remember to determine the EN difference for
each different bond in each molecule. Show your work; state the bond type in each case. If the bonds are polar,
indicate the direction on the 3-D diagram.
b)
Combine molecular shape with the EN difference to determine whether each molecule is polar, or non-polar. Give
reasons for your answer. If the molecule is polar, indicate the overall direction on the 3-D diagram.
Properties of Molecules
Once we have determined whether a molecule is polar or not we can make certain predictions about properties.
The properties of a substance are determined by how the molecules interact with each other. How they interact is
in turn determined by the forces of attraction between molecules. These are called intermolecular forces; the
forces which act between molecules, to draw them together, forming the various phases of matter.
To facilitate this discussion an understanding of energy and its relationship to matter is necessary. Kinetic energy is the
energy of motion. All matter is in motion; the nature of the motion depends upon the phase of the substance, whether it
is solid, liquid or gas.
The motion of the particles of matter is connected to temperature. The higher the temperature, the faster the motion.
The lower the temperature, the slower the motion. At a temperature of 0 Kelvin (-273°C), all motion stops. This is
called absolute zero, and is the point where matter has no kinetic energy. As energy is put into matter, temperature
goes up. Thus temperature is a rough measure of the kinetic energy in a substance.
A solid has its atoms held relatively rigidly in place; the only motion allowed the atoms is for them to vibrate in place. A
liquid allows its molecules to slide by each other (which gives a liquid its properties), but the molecules are still very close
to one another and interact strongly. A gas gives the molecules free rein; they move about very rapidly and interact with
each other only weakly.
Whether a particular substance is a particular phase depends both upon the kinetic energy of the substance and the
forces which bind the molecules together.
 the weaker the intermolecular forces, the less energy which is required for the substance to go from solid to
liquid and from liquid to gas.
 Thus the physical properties of melting point and boiling point gives a rough measure of the intermolecular
forces which bind the molecules to each other.
 The higher the melting and boiling points, the greater the strength of the intermolecular forces which bind the
molecules together.
Following is a list of the intermolecular forces which we are concerned with:
1.
Van der Waals forces
a)
London dispersion forces
b)
Dipole-dipole attraction
c)
Hydrogen bonding
2.
Ionic bonding
3.
Metallic bonding
4.
Network covalent bonding
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Van der Waals forces are weak forces which bind all matter together (in many substances, this is the only attractive
force).
London Dispersion Forces are the dominant forces between covalently bonded, non-polar molecules
(ex. N2, O2, CH4). They result from the formation of "instantaneous dipoles" in electrically neutral matter. They are weak
and act only over a short distance
Force increases with increasing number of electrons; the more electrons, the stronger the force. In very large molecules
it can be a very large force. (Analogy: Gravity)
Small molecules generally have a low melting point and boiling point (if molecules have a molar mass of less than 50,
they are generally gases at room temperature).
The force varies with shape of the molecule. A long molecule has a large number of electrons exposed; this increases
the force, so it has a relatively high boiling point. This same structure is very flexible and does not stack well, so it has a
low melting point. A more compact molecule exposes fewer electrons, so its boiling point is lower, but it stacks better, so
its melting point is higher:
Normal Pentane (C5H12)
Neopentane (C5H12)
m.p. -130C,
m.p. -20°C,
normal pentane
b.p. 36°C
b.p. 9°C
neopentane (dimethyl propane)
Dipole-Dipole Attraction is a force which acts between polar molecules (ex. H2S). It results from the attraction of
the opposite poles of the permanent molecular dipoles. These substances generally have higher melting and boiling
points than non-polar molecules with similar molecular weights.
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Hydrogen Bonding is a specialized form of dipole-dipole attraction. It occurs as when O, N, and F are bonded to H,
owing to the large electronegativity difference between these elements and hydrogen: This is a stronger force than
standard dipole-dipole attraction. Molecules with hydrogen bonding will have boiling points and melting points quite a
bit higher than molecules that have only dipole-dipole or London dispersion forces. Hydrogen bonding is responsible for
many of the unusual properties of water.
Electronegativity difference of H with O, N and F:
O - H
N - H
F - H
3.5 - 2.1 = 1.4
3.1 - 2.1 = 1.0
4.1 - 2.1 = 2.0
Activity:
All the molecules from the model building activity are covalently bonded. Based on the polarity of the molecules
(determined earlier) determine which of the three Van der Waals forces will be dominant for each substance. Do
this for the molecules from questions 1, 2, and 3.
Ionic Bonding generally occurs in compounds of metals and non-metals (salts). It is the result of the attraction of
oppositely charged ions. They come together in order to neutralize charge, and the attraction is relatively strong. The
structures formed are very orderly and are given the name crystal lattice. Ionic solids are called crystals. No sharing of
electrons occurs between the ions in the crystal lattice. As a result, ionic solids are brittle. Ionic solids conduct electricity
only in the molten state, and not very well. Ionic solids are characterized by very high melting and boiling points.
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Metallic Bonding is the bonding which occurs between metals in the Periodic Table. It is characterized by close
packing of the atoms, with the electrons delocalized; that is, they are free to jump from atom to atom, filling
unoccupied orbitals. This free sharing of electrons allows metals to conduct electricity freely (copper conducts
electricity 100 000 times better than molten NaCl). The free electrons also act as a lubricant, allowing metal atoms
to slide over one another without affecting the integrity of the material. Thus metals are malleable and ductile.
This bond is strong, giving most metals high melting and boiling points.
Network Covalent Bonding This is the traditional covalent bond, expanded to 2 or 3 dimensions in a network
which is theoretically infinite (much like an ionic crystal lattice). Network solids include diamonds, graphite, quartz,
and most rocks. Because the covalent bond is stronger than any other bond, the network solid is very hard
(diamonds are the hardest substance known). Because electrons are held tightly in their bonds, network solids are
brittle, and they do not conduct electricity. Because the orientation of the atoms is very specific to the bonding
orientation (tetrahedral, planar trigonal), network solids form distinct crystals. Because of the strength of the
bonds, network solids have very high melting and boiling points.
You will be given a fairly extensive assignment. Apply what you know about shape, electronegativity and polarity
to determine the intermolecular forces involved in the given compounds and make comparisons between them.
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Intermolecular Forces Affect Many Physical Properties
The strength of the attractions between particles can greatly affect the properties of a substance or solution.
Viscosity
•
•
•
Resistance of a liquid to flow is called viscosity.
It is related to the ease with which molecules can move past each other.
Viscosity increases with stronger intermolecular forces and decreases with higher temperature.
Surface Tension
•
Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.
Vapour Pressure
•
•
At any temperature, some molecules in a liquid have enough energy to escape.
As the temperature rises, the fraction of molecules that have enough energy to escape increases.
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