Chemistry 20
CHEMICAL BONDING
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Name: _______________________
Date: ___________________
Bonding Theory and Lewis Forces
A few things we need to keep in mind before we look at this unit!
Under standard Conditions.
- Metals (other than mercury) and all ionic compounds are solid (this is not totally true but in high
school you can assume this is true)
- Non-metals can combine to form gases, liquids or solids.
- Noble gases do not combine with other elements (again there are exceptions but we will not deal
with them in this course).
The Rule of Eight (Octet Rule)
Elements are most stable when their outer energy level (valence shell) is full like that of noble gases,
which are unreactive. Elements therefore usually react in ways to fill their outer energy level. Excluding
helium all noble gases have eight electrons in their valance electron level. We therefore generalize that
elements tend to react, bond in such a way to attain eight electrons in their valence shells. This is called
the octet rule.
Ionic bond – electrostatic attractions between two oppositely charged ions.
Covalent bond – electrostatic attraction between the nuclei of two adjacent atoms and a pair of shared
valence electrons.
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Valence Electrons - # of electrons in the highest energy level of the atom.
The general rule for determining the # of valence electrons a main group element goes as follows:
Group
Valence
Electrons
1
2
13
14
15
16
17
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1. Determine the # of valence electrons in each of the following atoms.
a) Oxygen __________________
e) Hydrogen ___________________
b) Nitrogen __________________
f) Chlorine _____________________
c) Carbon ___________________
g) Helium ______________________
d) Aluminum ________________
h) Magnesium____________________
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Bonding Theory
When considering bonding we are only concerned with the atom’s valence electrons.
In valence orbitals:
- Two electrons occupying the same orbital are referred to as a lone pair and are considered to be nonbonding.
- A half-filled orbital (orbital with one electron) is referred to as a bonding orbital.
We use electron dot diagrams (or lewis structures) to help us visualize valence electrons in atoms as they
interact (bond) with other atoms.
Drawing Electron Dot Diagrams
1. Write the symbol to represent the element.
2. Use a dot to represent each valence electron.
3. Place a single valence in each of the four valence orbitals before pairing electrons (in He and H there
will be only one valence orbital, boron only has three orbitals).
Element
Electron Dot Diagrams
Number of unpaired
electrons.
1. Sulfur
2. Aluminum
3. Carbon
4.
Chloride ion
(Cl-)
5.
Sodium ion
(Na+)
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Number of
electron
pairs (lone
pairs).
Unpaired electrons are likely to participate in bond formation so are sometimes referred to as bonding
electrons.
Electron pairs are less likely to participate in bond formation. For now we are going to consider them to be
non-bonding electrons although later we will see an exception to this as is the case with coordinate covalent
bonding.
Electron Dot Diagrams for Ionic Compounds
– Ionic bonds are the result of electrostatic forces between a positive and negative ion.
- Although we always express the ratios of atoms in an ionic compound in their reduced form, millions of atoms can exist
in a lattice structure (a regular repeating 3-D structure).
For main group elements (binary compounds)
1. Draw electron dot diagrams for atoms.
2. Show electron transfer, using arrows to show formation of positive and negative ions (usually from metal to
non-metal, always from less to more electronegative species).
3. All ions should be drawn in square brackets, showing all valence electrons
4. Charges should be shown as a superscript outside of brackets and ratios should be expressed as subscripts.
Example 1 – Draw the electron dot diagrams for the following ionic compounds.
a) NaCl
b) MgS
c) Al2O3
d) LiF
e) Li2O
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Electron Dot Diagrams for Molecular Compounds (Lewis Structures)
- Covalent bonds involve the sharing of electrons between nuclei.
- When drawing electron dot diagrams of molecular compounds we do not use arrows as there is no
transfer of electrons.
- We draw circles around electrons being shared or just draw the electron pairs between the nuclei.
- When drawing covalent compounds as in ionic compounds it is important to consider the octet rule
(except for Hydrogen and Boron).
- For the simple molecules we are dealing with now only unpaired electrons will pair with electrons from
another atom (later we will look at an exception to this rule called coordinate covalent bonding).
Example 3 – Drawing Simple Molecular Compounds (containing only single bonds).
a) HCl
b) Cl2
c) CH4
d) CCl4
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Sometimes atoms will form more than one bond with each other to fill their octet.
Example 4 – Drawing Simple Molecular Compounds (containing double and triple bonds).
a) CO2
b) N2
c) O2
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Drawing Lewis Structures for Complex Molecular Compounds
- If the molecule you are drawing has many atoms, it can be difficult to determine the central atom.
- In most cases you can assume that the atom with the highest number of unpaired electrons will be the
central atom.
- In some cases there is no central atom and you must use the information from the chemical formula to
determine the Lewis Structure.
See page 81 which lists 4 simple steps to drawing Lewis structures.
Example 4 – Drawing Complex Molecular Compounds.
a) CH3CH3
b) CH2Cl2
c) CH3COOH
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Coordinate Covalent Bonds
A special type of covalent bonding in which one atom donates both electrons to be shared between the nuclei of
two atoms. In IB this is referred to as dative bonding.
Example 5 – Drawing Lewis Structures with coordinate covalent bonds.
a) H3O+
b) NH4+
Example 6 – Drawing Lewis Structures for other ions.
a) NO3-
b) **CH3COO-
c) OH-
d) CN-
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Structural Formulas
- Another way to show bonding which is less tedious.
- In a structural diagram we do not show lone pairs of electrons and use a line to represent a bond (a pair
of shared electrons).
Example 7 – Drawing structural formulas.
a) H2O
b) NH3
c) H2CO
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Example 8 – Structural Diagrams and Lewis Dot Diagrams for Ionic and Molecular Compounds.
Compound
Lewis Structure/Electron Dot Diagram
Structural Diagram
a) H2S
b) CCl4
c) OF2
d) HCl
e) CH4
f) Na3N
N/A
g) CaCl2
N/A
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The Nature of Chemical Bonds
Electronegativity
- Relative measure of an atoms to attract the shared electrons in a chemical bond.
- relative electronegativities are listed on the periodic table.
- Fluorine has the highest at 4.0 while Cesium has the lowest at 0.8.
- Metals generally have low electronegativities – this is why they are often present as positive ions.
Trends in electronegativity
Decreasing Electronegativity doing downwards in each group
- As we go down groups in the periodic table the atomic size increase and electronegativity decreases. This can be
explained in looking at the force exerted by the nuclei on the valence electrons. As electrons become farther away
from the nucleus the force of the nuclei on the electrons decreases there decreasing electronegativity.
Image from: http://image.tutorvista.com/content/periodic-classification-elements/atomic-radii-in-group.jpeg
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Increasing electronegativity going from left to right in each period.
- As the magnitude of the charges increase (# of protons) the electronegativity also increases.
General Trend
Image from: http://voh.chem.ucla.edu/vohtar/winter02/20A/lecture5_files/image006.gif
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Bond Type and Electronegativity
If chemical bond is formed between atoms with significantly different electronegativity’s we can have polar covalent
bonds.
Non- polar covalent bonds – involves equal sharing of electrons between atoms.
Polar covalent bonds – involve sharing of electrons unequally between atoms. The more electronegative element holds
the electrons closer to itself than the less electronegative species. We use the symbols δ- to show that the species which is
more electronegative will be slightly negative and δ+ for the slightly positive, less electronegative atom.
Example – Determine if bonds between the following atoms would be polar or non-polar covalent. Use the greek letters
δ+ and δ- to identify any atoms that are slightly positive or negative.
a) O=O
b) H-H
c) H-Cl
d) O-H
e) N-H
f) H-F
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Ionic bonds – involves the transfer of electrons between atoms, the more electronegative element receives the electrons
and the least electronegative element donates the electron. The bond is then formed from the electrostatic attraction
between the anion and cation.
EXTENSION MATERIAL (another way to look at polarity)
Electronegativity difference (∆EN) between atoms involved in chemical bonds determines the character of the bond.
∆EN
Bond Character
<0.4
Non-polar covalent
0.4≤EN≤1.4
Polar colavent
>1.7
Ionic
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More Practice
1.i) Determine the type of bonding in each of the following compounds
ii) Determine the electronegativity difference between the atoms involved in the chemical bonds.
a) Methane (involves only carbon-hydrogen bonds)
b) silicone tetrafluoride
c) Magnesium oxide
d) iron-aluminum alloy
e) Aluminum Sulfide
f) dihydrogen monoxide
g) Molecular Oxygen
h) Carbon monoxide
2. Let’s classify the chemical bonds above (these do not include metallic bonds) using the following
classifications scheme.
- Non polar covalent – if EN difference is 0, the bond is covalent and bonding electrons are equally
shared.
- Polar covalent - If the EN difference is between 0.4 and 1.7 the bond is covalent and the bonding
electrons are not equally shared
- Ionic – Usually ionic if EN difference is 1.7 or higher although are exceptions.
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Three Dimensional Structures
Ionic Crystals
Thus far in the course we have been using the formula unit of ionic compounds to describe the ratio of ions in
the compound. Although the formula unit does give a clear indication of the ratio of ions we must understand
that many ionic compounds exist in crystal forms. Crystals are made up of many ions which arrange
themselves in what we call the crystal lattice, which often forms a distinct shape characteristic of that ionic
compound. The smallest pattern that is repeated over and over again of an ionic crystal is called the unit cell.
Image from: http://wikis.lib.ncsu.edu/images/9/9e/CH795C_Halite-2.png
The diagram above shows the unit cell of some different ionic compounds. The shape and structure of the ionic
compound is dependent on numerous factors including the radius of atoms involved, the charges and lone pairs.
Molecular Compounds
-
It is important to remember that when we find the formula for a molecular compound the numbers are
exact and cannot be reduced.
Molecular compounds contain covalent bonds between two atoms which share a pair of electrons.
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VSEPR and the Structures of Molecular Compounds
Valence-Shell Electron-Pair Repulsion theory is what we will be using to predict the three-dimensional shape of
molecules. VESPR theory is based on electrostatic repulsion of electrons from one another.
-
-
The theory states that electrons repulse one another and are therefore going to space themselves as far
apart as possible in a molecule. The repulsion between electrons decreases in the following order.
o Lone pairs –lone pairs> Lone pair-bonding pair>Bonding pair-bonding pair.
o The more repulsion between pairs the larger the bond angle.
In VSEPR theory we also treat any double or triple as if it were just a single bond.
We always look at the shape around the central atom, therefore if there is no central atom we must
specify the atom around which we are giving the geometry.
Linear Molecules (AX and AX2)
- Any compound that consists of only two atoms would be linear.
- Any compound where there are only two bonds (2 bonding pairs) around the central atom and no lone
pairs would be linear.
- Bond Angle (angle between bonds): 180o
Lewis
Formula
Bond
Pairs
Lone
Pairs
Total
Pairs
General
Formula
H-H
H-Cl
O=C=O
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Electron pair
arrangement
Stereochemical
Formula
Trigonal Planar (AX3)
- Consists of a central atom with three bonds around it and no lone pairs.
- Bond Angle: 120o (all pairs of electrons around the central atom are in the same plane 120o apart).
Lewis
Bond
Lone
Total
General
Electron pair
Stereochemical
Formula
Pairs
Pairs
Pairs
Formula
arrangement
Formula
BH3
BH2Cl
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Tetrahedral (AX4)
- Consists of a central atom with four bonds around it and no lone pairs.
- Bond Angle: 109.5o (all pairs of electrons arrange themselves in three dimensions).
Lewis
Bond
Lone
Total
General
Electron pair
Stereochemical
Formula
Pairs
Pairs
Pairs
Formula
arrangement
Formula
CH4
CH3Cl
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Trigonal Pyramidal (AX3E)
- Consists of a central atom with three bonds around it and one lone pair.
- Bond angle: 107.3o (similar to tetrahedral but lone pair exerts more repulsion so distorts the shape
slightly).
Lewis
Formula
Bond
Pairs
Lone
Pairs
Total
Pairs
General
Formula
NH3
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Electron pair
arrangement
Stereochemical
Formula
V-shaped (angular) (AX2E2)
- Consists of a central atom with two bonds around it and two lone pair.
- Bond angle: 104.5o (similar to tetrahedral but lone pairs exerts more repulsion so distorts the shape
slightly).
Lewis
Formula
Bond
Pairs
Lone
Pairs
Total
Pairs
General
Formula
Electron pair
arrangement
H2O
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Stereochemical
Formula
Bond Angle for least repulsion: The bond angles are 104.5o, which scientists believe is due to more repulsion
from lone pairs as compared to the tetrahedral and trigonal planar.
To determine the shape of a molecule follow the steps below:
1. Draw Lewis formula including all electron pairs around central atom.
2. Count the total # of bonding pairs and lone pairs.
3. Use table 2.1 (p.55) to predict the shape later by the time we have finished this chapter this table must
be memorized and for IB the bond angles must be memorized.
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Polar Bonds and Polar Molecules
Polar Molecule a molecule in which the negative charge is not distributed equally among all atoms.
Nonpolar Molecule a molecule with symmetrical distribution of charge.
Electronegativity and Bond Polarity
- The greater an atoms electronegativity the more pull it has on electrons. If one atom has a far
greater pull on electrons than the other, electrons will spend more time closer to that atoms
nucleus.
-
Nonpolar bond- when electronegativities of two atoms involved in the bond are the same.
-
Polar bond – when electronegativities of two atoms involved in the bond are very different.
Bond polarity and Molecular Polarity
-
A molecule may have polar bonds but may not be polar itself.
Molecule polarity depends on shape.
To predict molecule polarity use the following steps.
Steps to Determining Polarity of Molecules
Step 1: Draw Lewis structure of molecule.
Step 2: Use the electron pairs and VSEPR rules to determine molecule shape and draw stereochemical
representation of molecule.
Step 3: Use electronegativities to determine polarity of each bond.
Step 4: Draw bond dipole vectors and add them together to determine whether final result will be polar or nonpolar.
Example 1- Is Water a polar molecule?
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Molecule Symmetry and Polarity
In the next table we are going to use “A” to represent the central atom. “X” and “Y” are atoms with different
polarities bonded to the central atom.
A molecule is polar if there is an asymmetrical distribution of charge. Asymmetrical molecules like trigonal
pyramidal or bent/v-shaped molecules are therefore always polar.
VSEPR Shape
Stereochemical
Representation
Molecule Polarity
Example
Linear
H-H
Linear
H-Cl
Linear
O=C=O
Bent/Angular/V-shaped
H2O
Trigonal Planar
BH3
Trigonal Planar
H2CO or BH2Cl
Tetrahedral
CH4
Tetrahedral
CH3Cl
Trigonal Pyramidal
NH3
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Intermolecular Forces
Thus far we have only discussed intramolecular bonding, the forces that exist between the atoms or ions in a
molecule or compound. In this section we are going to look at the forces that exist between molecules.
Types of Intermolecular Forces
London Forces: (dispersion forces)
- Attraction b/tw a momentary dipole in a molecule and in surrounding molecules.
- More electrons ↑ chance of momentary dipole.
- Boiling point and melting point ↑ as electron # ↑ therefore, London forces ↑ as molecule size ↑
Image from: http://studentweb.usq.edu.au/home/W0099066/images/Bonding/london%20disperson.png
Dipole-Dipole forces:
- The attraction between the slightly negative and slightly positive ends two molecules with dipoles (which
are polar covalent)
- In the diagram below the dotted line represent the dipole-dipole forces between H-Cl molecules.
Image from: http://3.bp.blogspot.com/_HbJwA7v1Jss/TQojpBdry_I/AAAAAAAAAA0/AUQSTs9hASM/s1600/Dipole-dipole-interaction-in-HCl-2D.png
Hydrogen bonding: N-H/O-H/F - H
Hydrogen Bonds:
- a type of intermolecular attraction in which a hydrogen atom covalently bonded to a Fluorine, Oxygen or
Nitrogen will be attracted to the lone pairs of another nitrogen, oxygen or fluorine in a second molecule.
- A special and relatively strong dipole-dipole force.
- In the diagram below the hashed line represents the hydrogen bond, which exists between the oxygen of
one molecule of water and the hydrogen of another.
Image from: http://upload.wikimedia.org/wikipedia/commons/thumb/b/b5/Hydrogen-bonding-in-water-2D.png/800px-Hydrogen-bonding-in-water-2D.png
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Boiling Points and Intermolecular Bonding
The stronger the intermolecular force the higher the boiling point (the more energy it takes to break
intermolecular bonds).
Isoelectronic molecules - molecules with same # of electrons.
Intermolecular Forces
Strongest –
Weakest
Hydrogen
Dipole – Dipole
London
These can also be referred to as
Van der Waals forces.
The strength of the intermolecular forces also determines the state of the compound.
Example 1 – Using what you know about intermolecular forces explain the following.
1) At room temperature chlorine (Cl2) is a gas but bromine (Br2) is a liquid.
2) Hydrogen fluoride has a higher boiling point than hydrogen chloride.
3) Ethane (C2H4) has a much lower boiling point propane (CH3CH2CH3)
4) CH3CH2OH has a higher boiling point than (CH3) 2O.
5) H2O has a higher boiling point than H2S.
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Chemistry 20 Polarity and Intermolecular Forces Assignment
Formula
Lewis Dot
Draw VSEPR
Shape and Label
Diploles
Shape
name
a) CH3Cl
b) H2O
c) NF3
d) CO2
e) HI
f) BH3
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Polarity
(YES/NO)
Inter- bonds
(L, D-D, H)
Intra-bonds
(C, I or M)
Formula
Lewis Dot
Draw VSEPR
Shape and Label
Diploles
Shape
name
g) NH4+
i) H3O+
i) C2H4
j) SO2
k) CO
l) NO3-
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Polarity
(YES/NO)
Inter- bonds
(L, D-D, H)
Intra-bonds
(C, I or M)
Name: _____________________________
Date: _____________
Practice Quiz
1. For each of the following compounds.
i) Draw the Lewis diagrams.
ii) Using VSEPR theory determine the shapes of each molecule.
iii) Determine the major type of intermolecular interaction in each molecule.
iv) Determine which molecule would have a higher boiling point and melting point. Provide reasons.
a) methyl chloride (CH3Cl)
or
b) carbon tetraiodide
or
c) methane
or
methyl bromide (CH3Br)
carbon tetrafluoride.
ammonia
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2. Draw the Lewis Diagrams for the following ions.
a) CN-
b) NO3-
c) Mg2+
d) F-
e) Na +
f) NH4+
3. Show the equation for the formation of the following ionic compounds.
a) CaCl2
b) NaCl
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Relating Structures and Properties
Metallic Bonding
-
The bonds holding metal atoms together.
Valence electrons are delocalized in metallic compounds; they are composed of a lattice of positive ions
electrostatically attracted to a “sea” of delocalized electrons.
Metallic compounds are both malleable and conductive.
Conductivity – metallic compounds are very conductive. The delocalized electrons are free to move
around the lattice of positive ions.
Malleability – metallic compounds are malleable. The ‘sea’ of electrons enable the metal ions to roll over
each other.
The properties of different compounds are very dependent on the bonding present.
Solubility in water
Water is a polar solvent and therefore dissolves polar solutes. Polar covalent compounds dissolve well in water
and so do some ionic compounds. The solubility of ionic compounds is a little more complicated but we will
not be
Electrical Conductivity
In order for a substance to be conductive electrons must be able to flow.
Metals – metals can conduct electricity in their solid or molten state.
Ionic Compounds – the crystalline structure of a solid ionic compound does not allow electrons to flow. When
dissolved in water however soluble ionic compounds can conduct electricity as do molten ionic compounds. In
solution and in the molten state electrons are able to flow to conduct the electrical current.
Molecular Compounds – molecular compounds do not conduct electricity in any state. Water as a molecular
compound does not conduct electricity, however with ionic compounds dissolved in it, it does conduct electrical
current.
Network Solids – some network solids conduct electricity. Graphite is able to conduct electricity because there
are some electrons that are free to move whereas in a diamond there are no electrons free to conduct electricity.
State at Room Temperature
Ionic bonding is stronger than any intermolecular force present between covalent molecules.
Ionic compounds are all solid at room temperature.
Molecular compounds can be gas, liquid or solid at room temperature. Compounds with weaker intermolecular
forces are likely to be gases and compounds with stronger intermolecular forces will be liquid or solid.
Metals with the exception of mercury are all solid at room temperature.
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Solubility
in Water
Electrical
Conductivity
State at room
Temperature
Mechanical Properties
Ionic
Compounds
Molecular
Compounds
Metallic
Compounds
Network
Solid
The state at room temperature just like boiling point is dependent on the strength of the intermolecular forces.
Which substance has the lowest electrical conductivity?
A.
Cu(s)
B.
Hg(l)
C.
H2(g)
D.
LiOH(aq)
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