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SCH4U
Grade 12
University Chemistry
Version C
SCH4U – Chemistry
Introduction
Introduction
Welcome to the Grade 12 University Chemistry Course, SCH4U. This full-credit course
is part of the Ontario Secondary School curriculum. Prerequisite – Grade 11 University
Chemistry.
This course enables students to deepen their understanding of chemistry through the
study of organic chemistry, energy changes and rates of reaction, chemical systems
and equilibrium, electrochemistry, and atomic and molecular structure. Students will
further develop problem solving and laboratory skills as they investigate chemical
processes, at the same time refining their ability to communicate scientific information.
Emphasis will be placed on the importance of chemistry in daily life, and on evaluating
the impact of chemical technology on the environment.
How to Work Through This Course
Each of the units is made up of four lessons. Each lesson has a series of assignments
to be completed. In this course you must complete ALL assignments. Be sure to read
through all the material presented in each lesson before trying to complete the
assignments.
Important Symbols
Questions with this symbol are Key Questions. They give you an
opportunity to show your understanding of the course content. Ensure that
you complete these thoroughly as they will be evaluated.
Questions with this symbol are Support Questions. They do not need to
be submitted to the marker, but they will help you understand the course
material more fully. Answers for support questions are included at the end of
each unit. Refer to these for suggestions of how to properly structure the
answers to questions.
Remember, you must complete the KEY QUESTIONS successfully in order to
achieve the credit in this course. Remember to write the unit number, lesson
number and key question number on all assignments. Make sure that your
assignments are submitted in the proper order.
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SCH4U – Chemistry
Introduction
What You Must Do To Get a Credit
In order to be granted a credit in this course, you must
9 Successfully complete the Key Questions for each unit and submit them for
evaluation within the required time frame. This course is made up of 5 units.
9 Complete the mid-term exam after Unit 3.
9 Complete and pass a final examination.
After you submit lessons for evaluation, begin work on your next lesson(s) right
away! Do not wait until you receive your evaluated assignments from the marker.
Your Final Mark
•
Each Unit has 4 lessons each worth 2% (10% per Unit x 4 Units)
Midterm Test
40%
30%
•
Final Examination
30%
•
Term
Materials
This course is self-contained and does not require a textbook. You will require lined
paper, graph paper, a ruler, a scientific calculator and a writing utensil.
Expectations
The overall expectations you will cover in each unit are listed on the first page of each
unit.
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SCH4U – Chemistry
Introduction
Table of Contents
Unit 1 – Structure and Properties
Lesson 1
Lesson 2
Lesson 3
Lesson 4
Atomic Theories
Quantum Mechanics
Chemical Bonding
Intermolecular Forces
Unit 2 – Organic Chemistry
Lesson 5
Lesson 6
Lesson 7
Lesson 8
Hydrocarbons
Functional Groups
Types of Organic Reactions
Polymers
Unit 3 – Rates of Reactions
Lesson 9
Lesson 10
Lesson 11
Lesson 12
Thermochemistry
Enthalpies of Reactions
Energy Options
Chemical Kenetics
Unit 4 – Electrochemistry
Lesson 13
Lesson 14
Lesson 15
Lesson 16
Oxidation and Reduction Reactions
The Activity Series of Metals
Galvanic Cells
Electrolytic Cells
Unit 5 – Chemical Systems and Equilibrium
Lesson 17
Lesson 18
Lesson 19
Lesson 20
Introducing Equilibrium
The Equilibrium Constant
Acid and Bases Equilibrium
Solubility Equilibriums
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SCH4U
Grade 12
University Chemistry
Lesson 1 – Atomic Properties
SCH4U – Chemistry
Unit 1 – Lesson 1
Unit 1: Structure and Properties
Have you ever seen a picture of the first computer? Figure 1.1 below depicts what the
first computer looked like. It may seem odd and funny to look back at such pictures, but
most technologies are constant works in progress. Computers now come in tiny
devices such as phones and laptops.
The development of a modern day atomic model is similar in nature to that of computer
technologies. There have been many modifications along the way to current modern
atomic theory. In this unit you will learn about these models and modification. You will
also learn more about the molecular structures, and chemical bonding.
Figure 1.1: The First Computer
Overall Expectations
•
•
•
demonstrate an understanding of quantum mechanical theory, and explain how
types of chemical bonding account for the properties of ionic, molecular, covalent
network, and metallic substances;
investigate and compare the properties of solids and liquids, and use bonding
theory to predict the shape of simple molecules;
describe products and technologies whose development has depended on
understanding molecular structure, and technologies that have advanced the
knowledge of atomic and molecular theory.
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SCH4U – Chemistry
Unit 1 – Lesson 1
Lesson 1: Atomic Theories
The atom is the smallest unit of an element that retains the chemical properties of that
element. In this lesson, you will learn about the various theories that led up to the
current model of the atom.
What You Will Learn
After completing this lesson, you will;
•
•
•
explain the experimental observations and inferences made by Rutherford and Bohr
in developing the planetary model of the hydrogen atom;
describe some applications of principles relating to atomic and molecular structure in
analytical chemistry and medical diagnosis (e.g., infrared spectroscopy, X-ray
crystallography, nuclear medicine, medical applications of spectroscopy);
describe advances in Canadian research on atomic and molecular theory
The Development of the Atomic Model
Matter is anything that has mass and takes up space. This means everything around
you (including yourself!) is matter. Matter is made up of tiny particles called atoms.
The term atom was first coined by the Greek philosopher Democritus, who proposed
that the atom was the smallest particle that could not be subdivided.
As experimentation and the scientific method gained importance, the model of the atom
began to evolve. Let’s take a look at the scientists involved in the development of the
atomic model.
Making a Model: Key Scientists
JOHN DALTON (1809)
Dalton was an English schoolteacher came up with his atomic theory based on many
years of experimentation by many scientists.
Dalton’s Atomic Theory
1. All matter is composed of tiny particles called atoms
2. Atoms can be neither subdivided nor changed into one another
3. Atoms cannot be created or destroyed
4. All atoms of one element are the same in shape, size, mass and all other properties
5. All atoms of one element differ in these properties from atoms of all other elements
6. Chemical change is the union or separation of atoms
7. Atoms combine in small whole-number ratios such as 1:1. 1:2, 2:3, etc.
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SCH4U – Chemistry
Unit 1 – Lesson 1
Figure 1.1: Dalton’s “Billiard ball” model
Dalton’s model did not account for the subatomic particles found in an atom (protons,
neutrons, and electrons).
J.J. THOMPSON (1897)
• studied the deflection of cathode rays by electric and magnetic fields
• results suggested that the atom was not the smallest unit of matter; there were
subatomic particles within the atom.
• J.J. Thompson proposed that the atom is a sphere of uniform positive electricity in
which negative electrons were embedded like raisins in plum pudding (or chocolate
chips in a cookie)
Figure 1.2 – Thompson’s raisin bun model for the atom
RUTHERFORD (1909)
• Radium gives off three different types of radiation (alpha and beta particles, and
gamma rays)
• Alpha particles are Helium nuclei (2 protons, 2 neutrons)
• Using these principles, Rutherford performed his famous gold foil experiment
The Gold Foil Experiment
Rutherford hypothesized that If Thompson’s model is true, then the high speed
positively charged alpha particles should pass through the gold foil without being
deflected. Although most alpha particles passed through the gold foil, some were
deflected, and some even reflected back towards the source. Since opposite charges
attract, this means that there must be a positive charge present in the centre or nucleus
of the atom. From these observations, Rutherford formulated his own nuclear model of
the atom.
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SCH4U – Chemistry
Unit 1 – Lesson 1
Rutherford’s nuclear model of the atom (1911)
• the mass and the positive charge in the gold atoms is concentrated in a very small
region
• most of the atom is empty space
• this was dubbed the “beehive model” of the atom
Figure 1.3 Rutherford’s beehive model
BOHR (1913)
•
•
•
explained why the hydrogen atom does not collapse
explained the line spectrum of hydrogen
predicted undiscovered lines in the ultraviolet region of the Hydrogen spectrum
Figure 1.4 The Line Spectrum for Hydrogen (a line spectrum is created when
electricity is applied to elemental gas. Light is then passed through the sample and a
prism. The pattern revealed is called a line spectrum. Each element has its own
unique spectrum)
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SCH4U – Chemistry
Unit 1 – Lesson 1
Bohr’s theory states:
•
•
•
•
•
•
•
There are specific allowed energy levels (n) in which an electron can move.
The energy of an electron in each level is quantized.
The larger the n value, the more energy an electron possesses.
Each energy level corresponds to an orbit, a circular path in which the electron can
move around the nucleus.
An electron can travel in one of the allowed orbits without loss of energy
An electron can “jump” from one allowed orbit to another. The jump cannot be
gradual – it must occur all at once.
Only certain energies can be absorbed or emitted as the electron changes orbits.
Two principles of Quantum Mechanics and some of Bohr’s terminology can be used to
develop a straightforward view of the electron structure of the first 20 elements.
1. Electrons exist in energy levels in atoms. The number of the energy level, n, is
called the principle quantum number.
2. Each energy level can hold up to 2n2 electrons.
The 1st energy level can hold 2
The 2nd energy level can hold 8
The 3rd can hold 18
Energy Level Population = electrons in ground state
Na) 2e-)8e-)1e-
P)2e-)8e-)5e-
CHADWICK (1932)
• discovered the neutron
Subatomic Particles
Atoms can be broken down into smaller subatomic particles: The table below
summarizes some key information about the three subatomic particles (protons,
neutrons and electrons).
Table 1.1: Subatomic particles
Subatomic particle
Charge
Proton
Positive (+1)
Neutron
Neutral (0)
Electron
Negative (-1)
Location in atom
nucleus
nucleus
Orbits nucleus
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Relative mass
1
1
1/1800
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SCH4U – Chemistry
Unit 1 – Lesson 1
The number of protons in the nucleus is called the atomic number. This number
determines the identity of an atom. Atoms are electrically neutral; therefore the number
of protons in an atom must equal the number of electrons in an atom. For example,
oxygen has an atomic number of 8 and has 8 protons in its nucleus and 8 electrons
orbiting around the nucleus. Atoms also have a number called the mass number. The
mass number is the number of protons plus the number of neutrons an atom has. Let’s
look at the element oxygen for an example. Oxygen has a mass number of 16. The
number of neutrons = mass number – atomic number or 16-8 =8. Thus oxygen has 8
protons, 8 electrons and 8 neutrons.
Elements are organized on a chart called the periodic table. You should have received
a periodic table at the start of this course. The elements are organized into vertical
columns called groups, and horizontal rows called periods.
Figure 1.5: The Periodic Table of Elements
Let’s examine how you can retrieve information about subatomic particles from the
periodic table.
Example 1: Determine the number of protons, neutrons, and electrons a potassium
element has.
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SCH4U – Chemistry
Unit 1 – Lesson 1
Solution 1:
Atomic number
19
K
Mass number
39.10
The atomic number of potassium is 19, and since the atomic number is equal to the
number of protons and electrons, a Potassium atom has 19 protons and 19 electrons.
The number of neutrons is 40-19 or 21 neutrons.
Support Questions
(Reminder: these questions are not to be submitted but reinforce the material taught
and are strongly recommended – DO NOT write in this book).
1. Reproduce this chart in your notes and your periodic table to complete the missing
information:
Element
Atomic
Number
Mass
Number
Number of
Protons
Number of
Electrons
Number of
Neutrons
Hydrogen
2
7
4
5
Carbon
7
16
9
10
Sodium
12
27
14
15
Sulphur
17
40
19
20
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SCH4U – Chemistry
Unit 1 – Lesson 1
Structure of the Atom: Bohr-Rutherford Diagrams
Electrons move around the nucleus in circular paths called shells like planets around the
Sun.
Electrons are spinning so fast in their orbits that they seem to form a solid shell around
the nucleus. Electrons cannot exist between these orbits, but can move up or down
from one orbit to another. Electrons are more stable when they are at lower energy,
closer to the nucleus. Each orbit has a maximum number of electrons that it can hold.
The number of electrons found in the orbits of the first twenty elements:
1st orbit (K shell) – holds 2 electrons
2nd orbit (L shell) – holds 8 electrons
3rd orbit (M shell) – holds 8electrons
4th orbit (N shell) – holds 2electron
Drawing Bohr-Rutherford Diagrams
To draw Bohr-Rutherford diagrams, use the following steps:
1. Using the periodic table, determine the number of protons, neutrons, and electrons
for the element.
2. Draw a circle to represent the nucleus of the atom. The number of protons and
neutrons are written inside this circle.
3. Electrons are drawn in circular orbits around the nucleus. Remember that lower
orbits will fill up first!
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SCH4U – Chemistry
Unit 1 – Lesson 1
Example 2: Carbon
Draw the Bohr diagram for an atom of the element Carbon:
Solution 2:
Following the steps outlined above:
1. Since carbon has a mass number of 12 and an atomic number of 6, it has 6 protons,
6 electrons, and 6 neutrons.
2. Draw the nucleus and write the 6 protons (6p+) and 6 neutrons (6n0) inside.
6p
6n
+
0
3. Then add your electron shells. Remember the K shell holds 2 electrons, and the L
shell holds the remaining 4 electrons. The shells total 6 electrons.
6p
6n
+
0
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SCH4U – Chemistry
Unit 1 – Lesson 1
Example 3: Nitrogen
Draw a Bohr Diagram for an atom of the element nitrogen
Solution 3:
14
N
7
7p
7n
+
0
Often this can be written in a short-hand manner as follows:
N)2e-)5eThis method will be used more frequently in this course.
Support Questions
2. Draw the Bohr diagrams for the first twenty elements on the periodic table (i.e.
elements with atomic number 1-20). State any patterns you may observe based on
the locations of the elements on the periodic table.
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SCH4U – Chemistry
Unit 1 – Lesson 1
Key Question #1
1. Summarize, using labelled diagrams in chart form, the evolution of atomic theory
from Dalton to the Rutherford model. (10 marks)
2. Write a short biography for one of the following Canadians that made advances in
atomic and molecular theory. (15 marks)
a. Harriet Brooks
b. R.J. Leroy
c. Richard Bader
Include the following;
i)
ii)
iii)
A brief history of their life (birth death, family, where they lived, etc),
their research, and,
the significance/relevance of their research to the development of atomic and
molecular theory.
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SCH4U
Grade 12
University Chemistry
Lesson 2 – Quantum Mechanics
SCH4U – Chemistry
Unit 1 – Lesson 2
Lesson 2: Quantum Mechanics
The development of the Bohr model led to further study of the atomic model based of
wave properties of electrons, a branch of chemistry known as Quantum Mechanics. It
can be a very complex area of chemistry, so this lesson will cover the fundamental
concepts.
What You Will Learn
After completing this lesson, you will;
•
•
•
describe the quantum mechanical model of the atom (e.g., orbitals, electron
probability density) and the contributions of individuals to this model (e.g., those of
Planck, de Broglie, Einstein, Heisenberg, and Schrödinger);
list characteristics of the s, p, d, and f blocks of elements, and explain the
relationship between position of elements in the periodic table, their properties, and
their electron configurations;
write electron configurations for elements in the periodic table, using the Pauli
exclusion principle and Hund’s rule;
Schrodinger (1924) postulated that sometimes electrons behave as particles, and
sometimes like waves. Because of this we cannot measure both the position and
velocity of an electron at the same time. This exclusion is referred to as the Pauli
Exclusion Principle. What this really means is that we cannot determine the
momentum (velocity) and position (electron address) of an electron at the same time.
The best we can do is to calculate the PROBABILITY of an electron being in a certain
place at a certain time. There are calculations that can be done to describe the region in
space where the electron is MOST LIKELY to be found, however we will not focus on
the mathematical in this course,
Regions where electrons are most likely to be found are called orbitals.
For every value of n, there are n types of orbitals and n2 actual orbitals
1st energy level has 1 type of orbital and (12) orbital.
2nd energy level has 2 types of orbital and (22) 4 orbitals
3rd energy level has 3 types of orbital and (32) 9 orbitals
Each electron has a set of four numbers, called quantum numbers that specify it
completely; no two electrons in the same atom can have the same four. As mentioned
above, quantum numbers are determined by solving a mathematical wave equation,
which we will not do in this course rather we will focus on what each quantum number
represents.
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SCH4U – Chemistry
Unit 1 – Lesson 2
The Principle Quantum Number, n
The letter n was used by Bohr to identify the orbits and energies. In other words, n tells
you which of the "main" energy levels the electrons are in.
•
•
identifies the energy of an orbital
values: 1,2,3… to infinity (4)
You can use the analogy of uneven steps on a staircase to describe the energy levels.
If an electron “falls” from a higher energy level such as n=2 to n=1, the difference
between the two is released as a photon of light.
Figure 2.1 Energy Levels (Source: Nelson Chemistry 12)
The Secondary Quantum Number, l
The secondary quantum number, l was introduced by Arnold Somerfield in 1915 to
explain the line splitting of the hydrogen spectrum. The secondary quantum number l
identifies the electron subshells or sublevels that part of a main energy level.
•
•
identifies the shape of the orbital
values: 0 to (n-1) i.e: 0, 1, 2, 3, …n-1
s p d f
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SCH4U – Chemistry
Unit 1 – Lesson 2
The number of sublevels equals the value of the principle quantum number. For
example if n=3, then there are three sublevels, l= 0, 1, 2
Figure 2.2 – Energy Sublevels (Source: Nelson Chemistry 12)
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SCH4U – Chemistry
Unit 1 – Lesson 2
Support Questions
3. What is the difference in the electron orbits proposed by Bohr and those of
Somerfield?
4. Recreate and complete this table in your own notes.
Primary energy
level
Principle quantum
number
Possible
secondary
quantum numbers
Number of
sublevels per
primary level
1
2
3
4
5
5. Write a general rule that can be used to predict all possible values from lowest to
highest, of the secondary quantum number for any value of the principal quantum
number.
The Magnetic Quantum Number, ml
Recall from lesson one that a line spectrum is created by the passing of electricity
through a gaseous element sample contained within a gas discharged tube. Light is
passed through the sample and then a prism to obtain the spectrum of the element.
If a gas discharge tube is placed near a strong magnet, some single lines split into new
lines that were not initially present. This occurrence was discovered by Pieter Zeeman
is 1897 and is called the Zeeman Effect. For example he observed a single line
transformed into three lines when the magnet was applied. This effect was explained
using another quantum number, the magnetic quantum number, ml. The magnetic
quantum number explains that orbits could exist at various angles.
•
•
identifies direction of orientation (x, y, z) to an external magnetic field
values: -l…0…+ l
i.e. l = 2 p orbital split into 3 planes (-1, 0, +1)
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SCH4U – Chemistry
Unit 1 – Lesson 2
Figure 2.3: Spectra lines can be split in the presence of a magnetic field
(Source: Nelson Chemistry 12)
Table 2.1: Values for Magnetic Quantum Number
Value of l 0 to n-1
Values of ml -l to +l
0
0
1
-1, 0, +1
2
-2, -1, 0, +1, +2
3
-3,-2,-1,0,+1,+2,+3
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SCH4U – Chemistry
Unit 1 – Lesson 2
The Spin Quantum Number ms
Atoms have their own magnetism which is unique to the magnetism of a group of
atoms. This magnetism is referred to as paramagnetism. It was suggested that each
electron spins on its own axis. It can spin in either a clockwise or counterclockwise
direction.
•
•
identifies the spin direction of an electron
only 2 values: +½ (8) and -½ (9)
Table 2.2 Summary of Quantum Numbers
Principle
Secondary
Magnetic
Quantum
Quantum
Quantum
Number, n
Number, l
Number, ml
1
0
0
2
0
-1,0,+1
1
3
0
-2,-1,0,+1,+2
1
2
Spin Quantum
Number, ms
+1/2, -1/2
+1/2, -1/2
+1/2, -1/2
Electron Orbitals: Modernizing the Atom
Early Bohr atomic model theory were based on the idea of an electron travelling in
some kind of path or orbit, however modern day theory is that of an electron orbital.
Recall that an orbital is a volume of space where an electron is most likely to be found.
The first two quantum numbers, n and l describe electrons that have different energies
under normal circumstances in multi-electron atoms. The other two quantum numbers
apply to abnormal conditions in which a magnetic field is applied.
We will focus on the energy shells (n), or the primary quantum number, and subshells
(l), or the second quantum number. We will now communicate the values of l with
letters to denote the orbitals rather than numbers
Table 2.3 Values and Letters for the Secondary Quantum Number
Value of l
0
1
2
3
Letter of
s
p
d
f
designation
Name of
sharp
principal
diffuse
fundamental
designation
It is common for chemists to use the number for the main energy level and letter for the
energy sublevel. For example: a 1s orbital, a 2p orbital and a 3d orbital.
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SCH4U – Chemistry
Unit 1 – Lesson 2
Table 2.4 Number of Orbitals for each Energy Sublevel
Value of l
Sublevel symbol
Number of orbitals
0
s
1
1
p
3
2
d
5
3
f
7
Energy Level Diagrams
Modern day atomic theory infers that electrons in an atom have different energies. The
atomic spectra indicate the energy sublevels, defined by a quantum number. An energy
level diagram shows the relative energies of the electrons in various orbitals.
Figure 2.4: Energy Level Diagram. The circles represent orbitals which contain two
electrons. (Source: Nelson Chemistry 12)
Each electron orbital is represented by a circle and can contain two electrons. The
energy of the electrons increases with increasing principle quantum number, n. For a
given value of n, the sublevels increase in energy, in order, s<p<d<f. The restrictions
on the quantum numbers require that there can only be one s orbital, three p orbitals.
Five d orbitals, and seven f orbitals. The energy diagram for an atom is directly related
to its chemical properties and position on the periodic table.
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SCH4U – Chemistry
Unit 1 – Lesson 2
Rules for Completing an Energy Level Diagram
¾ electron orbitals are filled from lowest energy up (Aufbau Principle)
¾ electrons go into the orbital corresponding to the lowest energy level available.
¾ an orbital can hold up to 2 electrons
¾ electrons in the same sublevel will not pair until all the orbitals in the sublevel
have at least 1 electron (Hund’s Rule)
Figure 2.5: Energy-level diagrams for lithium, carbon and fluorine. (Source:
Nelson Chemistry 12)
Aufbau Diagram for Filling Orbitals
Figure 2. Memory Tip for Filling Orbitals (Source
Nelson Chemistry 12)
To use this diagram to help with drawing energy level
diagrams, start at the bottom left side and add electrons
in the order shown by the diagonal arrows. Work your
way up to the upper right-hand corner.
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SCH4U – Chemistry
Unit 1 – Lesson 2
Example 1: Draw the energy level diagram for an oxygen atom
Solution 1: Oxygen has an atomic number of 8, so we must place these 8 electrons in
energy levels.
The first two electrons are placed in the 1s energy level
1s
The next two are placed in the 2s energy level
2s
The next three electrons are placed singly in each of the 2p orbitals
2p
Finally the last electron is paired up to fill one of 2p energy sublevels
2p
Now draw your energy diagram.
1s
2s
2p
Oxygen (O)
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SCH4U – Chemistry
Unit 1 – Lesson 2
Example 2: Draw the energy level diagram for nitrogen anion.
Solution 2:
Nitrogen (N) 1s22s22p3
Nitrogen anion (N-3) 1s22s22p6
Use the same procedure except remember that a nitrogen anion will have 7 + 3
additional electrons.
Electron Configurations
Electron Configurations provide the same information as energy level diagrams, but in a
more concise format. An electron configuration depicts the type of electron in
increasing energy level. For example the electron configuration for an oxygen atom is
1s22s22p2.
Example 3: Write the electron configuration for the following elements:
a) sodium atom
b) fluorine ion
Solution 3:
Sodium Atom
1s22s22p63s1
Fluorine ion
1s22s22p6
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SCH4U – Chemistry
Unit 1 – Lesson 2
Shorthand Form of Electron Configuration
The shorthand form abbreviates writing the electron configuration by using the nearest
Nobel Gas in its written format. For example:
Cl: 1s22s22p63s23p5
Sn:1s22s22p63s24s23d104p65s24d105p2
becomes [Ne]3s23p5
becomes [Kr]5s24d105p2
Support Questions
6. Draw an energy level diagram for the following
a) Carbon
b) Cl-1
7. Write the complete ground state electron configurations for the following:
a) lithium
b) oxygen
c) calcium
d) titanium
e) rubidium
f) lead
g) erbium
8. Write the abbreviated (shorthand method) ground state electron configurations for
the following:
a) helium
b) nitrogen
c) chlorine
d) iron
e) zinc
f) barium
g) polonium
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SCH4U – Chemistry
Unit 1 – Lesson 2
Key Question #2
1. Calculate the maximum number of electrons with principal quantum number;
(8 marks)
a) 1
b) 2
c) 3
d) 4
2. Show the permissible values of l and m for; (8 marks)
a) n=1
b) n=2
c) n=3
3. Draw energy level diagrams for beryllium, magnesium and calcium ions. What is the
similarity in these diagrams? (10 marks – 3 each for diagram – 1 for statement)
4. The sodium ion and the neon atom are isoelectronic (have the same electron
configuration). (8 marks – 4 marks each)
a) Write the electron configuration for the sodium ion and neon atom.
b) Describe and explain the similarities and differences in the properties of these
two chemical entities.
5. One important application of Quantum Mechanics is laser technology. Construct an
information pamphlet including;
•
•
•
uses of laser technology
how it applies quantum mechanics
advantage/disadvantages of laser technology
NOTE: Don’t forget to include the references for your information (Wikipedia or
search engine references are not acceptable – look for professional sites or journal
articles) (25 marks – 20 for content, references and design, 5 for grammar/spelling)
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SCH4U
Grade 12
University Chemistry
Lesson 3 – Chemical Bonding
SCH4U – Chemistry
Unit 1 – Lesson 3
Lesson 3: Chemical Bonding
Chemical compounds are formed by the joining of two or more atoms. You will review
some of the basics of chemical bonding that you learned in Grade 11 Chemistry;
however you will also extend your knowledge of chemical bonding using Quantum
Mechanics.
What You Will Learn
After completing this lesson, you will;
•
•
•
predict molecular shape for simple molecules and ions, using the VSEPR model;
use appropriate scientific vocabulary to communicate ideas related to structure and
bonding (e.g., orbital, absorption spectrum, quantum, photon, dipole);
predict the polarity of various substances, using molecular shape and the
electronegativity values of the elements of the substances
Ionic and Covalent Bonding: The Octet Rule
Atoms form bonds to become more chemically stable. The most chemically stable
elements on the periodic table are the noble gases. We know this because they are
extremely unreactive and tend not to form compounds.
According to the octet rule, atoms bond in order to achieve the same electron
configuration as a noble gas. This rule is called the octet rule because all the noble
gases (except helium) have eight valence electrons.
Generally, when an atom tends to gain or lose electrons, an atom will have the same
electron configuration (arrangement of electrons) as a noble gas. Atoms that have
identical electron configurations are said to be isoelectronic.
An atom that has lost an electron to become stable is referred to as an ion. If an atom
loses electrons, it becomes positively charged and is referred to as a cation.
Copyright © 2008, Durham Continuing Education
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SCH4U – Chemistry
Unit 1 – Lesson 3
Lewis Structures
Another way to draw atoms and ions is to use Lewis structures. Lewis structures depict
only the valence electrons an element has. The figure below shows the Lewis
structures for some common elements.
To draw Lewis structures, use the following steps:
1. Write the atomic symbol of the element. This will represent the atomic nucleus. It is
considered to have 4 “sides”.
2. Place the valence electrons, one per side, and then pair up if necessary.
Example 1: Chlorine - Draw the Lewis structure for a chlorine atom
Solution 1:
xx
x
Cl xx
xx
The table below shows the Lewis structures for the first twenty elements:
Figure 3.1: Lewis Structures for the first twenty elements
Notice that the group number (vertical column) indicates how many valence electrons
an atom has. Thus, lithium is in group one and has one valence electron, while
fluorine is in group seven and has seven valence electrons.
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SCH4U – Chemistry
Unit 1 – Lesson 3
Ions can also be represented by Lewis symbols. The Lewis symbol is enclosed by a
bracket as with the Bohr diagram:
Example 2: Lewis Chlorine Ion - Draw the Lewis structure for a chlorine ion
Solution 2:
-1
xx
x
x
Cl xx
xx
Lewis Structures and Quantum Mechanics
There is unity between the atomic and bonding theories you learned in Grade 11
chemistry. The octet rule comes from the maximum of two electrons in the s orbital and
the six electrons in the p orbital’s.
Figure 3.2 - Correlation between Atomic and Bonding Theories
(Source: Nelson Chemistry 12)
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Page 33 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
Support Questions
9. Recreate and complete the following table in your own notes;
Element
Atomic
Symbol
Electron
Configuration
Number of
valence
electrons
Valence
Oxygen
Chlorine
Sodium
Phosphorus
10. Identify the elements and write the Lewis structure for the following electron
configurations:
a) 1s22s22p4
c) [Ar]4s23d104p5
b) 1s22s22p63s23p3
d) [Kr]5s1
Electronegativity
Electronegativity is a measure of an atom’s ability to attract electrons in a chemical
bond. It is a periodic property.
Predicting Bond Type Using Electronegativity
You can use the difference between electronegativities of two atoms to determine if the
bond formed between the two atoms is ionic or covalent, or polar covalent. The
symbol ΔEN stands for the difference in electronegativity between two values
3.3
1.7
MOSTLY IONIC
0.5
POLAR COVALENT
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0
COVALENT
Page 34 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
Table 3.1– Electronegativities of Various Elements
Element
H
Li
Metals
Be
Na
Mg
K
Ca
C
Nonmetals
N
O
F
P
S
Cl
Electronegativity
2.1
1.0
1.5
0.9
1.2
0.8
1.0
2.5
3.0
3.5
4.0
2.1
2.5
3.0
Example 3: Determine if the elements below would form ionic, covalent or polar
covalent bonds:
Solution 3:
Substance
KF
O2
HCl
EN Element 1
EN Element 2
ΔEN
0.8
4.0
3.2
Ionic or
Covalent?
ionic
3.5
3.5
0
covalent
2.1
3.0
0.9
Polar covalent
Bonds within Molecules: Intramolecular Bonds
When an ion loses or gains an electron, it is forming an ion. An atom always loses or
gains an electron in conjunction with another atom forming a chemical bond. There are
three main types of intramolecular bonds we will explore in this lesson, ionic, covalent
and polar covalent.
Ionic Bonding
•
•
electrons are transferred from one atom to another
usually occurs between metals and non-metals.
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Page 35 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
Example 4: Potassium and Fluoride - Using Lewis structures draw the formation of a
bond between potassium and fluorine.
Solution 4:
Step 1: Write out the correct number of valence electrons for each atom.
xx
K
+
x
F
x
x
xx
Step 2: Analyze the valence electrons. Potassium will lose its one valence electron to
obtain a full octet and fluorine will gain one.
xx
K
F
x
x
x
xx
Potassium loses its one valence electron, becoming a cation with a charge of +1, and
fluorine gains one valence electron, becoming an anion with a charge of -1.
+1
-1
xx
K
x
F
x
x
xx
Notice how the ions are written with square brackets and the overall charge is indicated
in the upper right hand corner.
When naming ionic compound, we name the cation first, then the anion. The ending of
the anion is replaced with “ide”.
Thus, the name of this compound is potassium fluoride. The positive potassium
cation (+) is attracted to the negative fluorine anion (-). This is what forms the ionic
bond.
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Page 36 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
Example 5: Magnesium and Nitrogen - Using Lewis structures draw the formation of a
bond between magnesium and nitrogen.
Solution 5:
Again, first draw the Lewis structures
x
Mg
+
x
N
x
x
x
In this case, magnesium will lose two electrons and nitrogen will gain. However, since
this does not balance we need three magnesium atoms and two nitrogen atoms to
make this balance. The transfer of the electrons from magnesium to nitrogen is shown
below.
Mg
x
x
N
x
x
x
Mg
x
x
N
x
x
x
Mg
We can then summarize this with brackets. The upper right hand corner indicates the
charge of the ion and the lower right hand corner indicates the number of ions.
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SCH4U – Chemistry
Unit 1 – Lesson 3
This compound is named magnesium nitride.
Support Questions
11. Recreate and complete the following table. Use Lewis structures to depict atoms.
Bond
formation
Name of
compound
Chemical
formula
Anion
Cation
a) lithium
and fluorine
b) calcium
and
phosphorus
Covalent Bonding
Covalent bonds form when two or more non-metals share one or more pairs of
electrons. As a result of forming covalent bonds through sharing electrons, the atoms
end up with a stable electron arrangement in their outer orbit similar to that of a noble
gas.
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Page 38 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
Example 6: Using Bohr diagrams, draw the formation of Chlorine gas (Cl2)
Solution 6:
Chlorine gas is a molecule that consists of two chlorine atoms held together with a
covalent bond.
→ Each chlorine atom has 7 electrons in its outer orbit and needs to gain 1 electron to
become stable.
→ Two chlorine atoms share a pair of electrons to form a covalent bond. Each chlorine
atom now has 8 electrons in its outer orbit (which forms a stable octet).
Example 7: Draw the formation of chlorine gas using Lewis structures
Solution 7:
Other examples of covalently bonded molecules include:
Methane (CH4)
Water (H2O)
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Ammonia (NH3)
Page 39 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
Drawing Lewis Structures for Polyatomic Ions
When an ion is composed of more that one atom, it is termed a polyatomic ion.
Polyatomic ions are groups of atoms that tend to stay together and carry an overall
ionic charge.
Table 3.2– Common Polyatomic Ions and Their Ionic Charges
Name of polyatomic ion
Ion formula
Ionic charge
nitrate
NO3–
1–
–
hydroxide
OH
1–
bicarbonate
HCO3–
1–
–
chlorate
ClO3
1–
carbonate
CO32–
2–
2–
sulfate
SO4
2–
3–
phosphate
PO4
3–
Drawing Lewis Structures for polyatomic Ions can be tricky. There are some general rules
for drawing these structures. We will use the nitrate ion (NO3-1) as our example.
Example 8: Draw the Lewis structure for nitrate, NO3-1
Solution 8:
Step 1: Arrange atoms symmetrically around the central atom (usually listed first, and
singular although not usually hydrogen or oxygen). In this case, nitrogen is our central
atom.
Step 2: Total valence electrons for all atoms. Add one electron for each -1 charge, subtract
for each +1 charge. For nitrate there are 5 valence electrons for nitrogen + 1 electron for -1
charge and 18 electrons for the oxygen atoms. This totals 24 electrons.
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SCH4U – Chemistry
Unit 1 – Lesson 3
Step 3: Place a bonding pair of electrons between the central atom and each of
The surrounding atoms
Step 4: Complete the octets of the surrounding atoms
Step 5: If the central atom does not have an octet, move lone pairs from the surrounding
atoms to form double or triple bonds.
Step 6: Draw the Lewis structure and enclose polyatomic ions within square brackets
showing the ion charge.
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Page 41 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
Support Questions
12. Draw Lewis Structures for each of the following molecules or polyatomic ions.
a) ClO4-1
b) CN-1
c)HCO3-2
Valence Bond Theory
The valence bond theory explains why and how electrons are shared between atoms.
The VB theory postulates how individual atoms, each with its own orbitals and electrons
come together and form covalent bonds in a molecule.
A bond between two atoms is formed when a pair of electrons is shared by two
overlapping orbitals,
For example, in a hydrogen molecule, the two 1s orbitals from each H atoms overlap
and share electrons.
H
H
Æ
H2(g)
1s1
1s1
1s2
In water, there is overlap of the 1s orbital of hydrogen and the 2 p orbital of oxygen.
Figure 3.3: Atoms overlapping orbitals (Source: Nelson Chemistry 12)
Copyright © 2008, Durham Continuing Education
Page 42 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
According to Valence Bond Theory:
• A half-filled orbital in one atom can overlap with another half-filled orbital of a
second atom to form a new, bonding orbital
• The new bonding orbital contains two electrons of opposite spin
• When atom bond, they arrange in space for maximum overlap
Hybrid Orbitals
You may have noticed in the figures above that when the orbital bond and overlap they
change shape. They form new orbitals called hybrid orbitals. Hybridization is the
process of combining two or more atomic orbitals to create new orbitals, called hybrids
that will fulfill the geometric demands of the system.
When an atom hybridizes, it will restructure its original set of s and p atomic orbitals into
a new set of hybrid orbitals. The process of hybridization is driven by the needs of the
atoms to produce specific geometric patterns.
Example 9: What are the bonding orbitals of the BF3 molecule?
Solution 9:
Boron has a valence of +3 and Fluorine is -1.
Write the electron configuration of boron. Focus on the valence electron orbitals. In this
case the 2s has two electrons, and p orbital has one.
B: 1s22s22p1
2s2
2px1
2py
2pz
In this case, one of the 2s2 electrons can be promoted to an empty p orbital
2s1
2px1
2py
2pz
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SCH4U – Chemistry
Unit 1 – Lesson 3
By promoting an s electron to an empty p orbital, there are one s orbitals and two p
orbitals available for hybridization. A total of three identical sp2 (i.e one s, two p
orbitals)
Therefore the hybridization is sp2
More forms of hybridization are summarized in the table following.
Table 3.3: Forms of hybridization
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Page 44 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
Double and Triple Bonds
Two kinds of orbital overlap are possible.
a) The end-to-end overlap of s orbitals, p orbitals, hybrid orbitals, or some pair of
these orbitals. These form sigma bonds (σ)
b) Side by side overlap forms pi (π) bond
A double bond contains a sigma and a pi bond
Consider the molecule C2H4
In this case we only have partial hybridization. We have three “new” sp2 orbital and one
“normal” 2pz orbital.
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Page 45 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
This is shown in figure 3.4 following:
Figure 3.4: Partial hybridization (Source: Nelson Chemistry 12)
Partial hybridization can also be used to explain how triple bonds form
Consider ethyne gas C2H2
In this case, two sp hybrid orbitals are formed for each carbon, and two unhybridized
orbitals are formed.
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Page 46 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
Figure 3.5 Formation of a Triple Bond (Source: Nelson Chemistry 12)
Support Questions
13. What atomic orbital or orbitals are available for bonding for each of the following
atoms?
a) H
b) F
c) S
d) Br
14. Provide ground state and promoted state electron configurations for each of the
following atoms and indicate the type of hybridization involved when each atom
forms a compound;
a) carbon in CH4
b) boron in BH3
c) beryllium in BeH2
15. When are π bonds formed?
16. Provide an explanation for bonding in each of the following;
a) C2Cl4
b) C2F2
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Page 47 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
The Valence Shell Electron Pair Repulsion (VSEPR) model:
One of the more important properties of any molecule is its shape. It is very important to
know the shape of a molecule if one is to understand its reactions. It is also desirable to
have a simple method to predict the geometries of compounds.
The underlying assumptions made by the VSEPR method are the following.
•
•
•
•
•
Atoms in a molecule are bound together by electron pairs. These are called
bonding pairs. More than one set of bonding pairs of electrons may bind any two
atoms together (multiple bonding).
Some atoms in a molecule may also possess pairs of electrons not involved in
bonding. These are called lone pairs.
The bonding pairs and lone pairs around any particular atom in a molecule adopt
positions in which their mutual interactions are minimized. The logic here is simple.
Electron pairs are negatively charged and will get as far apart from each other as
possible.
Lone pairs occupy more space than bonding electron pairs.
Double bonds occupy more space than single bonds.
Using VESPR Theory
Example 10: What is the shape of the BeH2?
Solution 10:
Step 1: Draw the Lewis Structure
H
Be H
If we consider the central atom, Be, it has 2 bonding pairs and no lone pairs. It has the
general formula AX2 and is considered linear.
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Page 48 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
Table 3.4 VESPR predicts Molecular Shapes
Example Predict the shape of a sulphate ion, SO42The central atom sulphur, has 6 valence electrons plus the two additional electrons
totalling eight. This gives 4 bonding pairs and 0 lone pairs. Thus the general formula is
AX4, and the molecule is tetrahedral.
Support Questions
17. Use Lewis Structures and VESPR theory to predict the shapes of the following
molecules:
a) CO2
b) HCN
c) BF3
d) SiCl4
e) CH4
f) OCl2
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Page 49 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
Polar Covalent Bonds
A polar covalent intramolecular bond is formed when there is unequal sharing between
valence electrons resulting in dipoles (ends) that are slightly positive or slightly negative
Example 11: Draw the polar bond formed in a molecule of carbon tetrachloride (CCl4)
between the elements carbon and chlorine
Solution 11:
Referring to table 3.3 above, chlorine has an electronegativity of 3.0 and carbon 2.5.
Thus the difference, ΔEn is 0.5, and the intramolecular bond is considered polar
covalent. Since the chlorine has a higher electronegativity, the electrons are pulled
closer to the chlorine’s atomic nucleus. This makes the chlorine ends or dipoles slightly
negative (δ-). Conversely, the electrons are farther away from carbons nucleus, making
it slightly positive (δ+).
Molecular Polarity
Covalent bonds can be polar or nonpolar. However the polarity of a molecule as a
whole is dependent on bond polarity and molecular shape.
¾ Symmetrical molecules produce nonpolar molecules (whether the bonds are
polar or not).
¾ Asymmetrical molecules produce nonpolar molecules if the bonds are all
nonpolar
Example 12: Predict the Polarity of the ammonia, NH3 molecule including your
reasoning.
Solution 12:
First, draw the Lewis structure
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Page 50 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
Next, check the electronegativities. Nitrogen has an electronegativity of 3 and hydrogen
of 2.1. This difference is 0.9, making the hydrogen end positive and the nitrogen end
negative.
Now we must also consider symmetry. Because ammonia is asymmetrical and
contains polar bonds, the molecule is polar.
Support Questions
18. Recreate and complete the table below. If the molecule is covalent, indicate if it is
polar covalent or not.
Substance
EN Element 1 EN Element 2 ΔEN
Ionic or
Covalent?
NaCl
Cl2
HF
19. Predict the polarity of following molecules.
a) BF3
b) OF2
c) CI2
d) PCl3
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Page 51 of 64
SCH4U – Chemistry
Unit 1 – Lesson 3
Key Question #3
1. Recreate and complete the following table; (24 marks)
Molecule or
Compound
ion
HCl
CH4
CH3Cl
CO2
H2O
NH3
Lewis
structure
Name of
Shape
Bond
Polarity
Molecular
Polarity
2. Identify the types of hybrid orbitals found in molecules of the following substances;
(8 marks – 2 marks each)
a) CCl4(l)
b) BH3(g)
c) BeI2(s)
d) SiH4(g)
3. The polarity of a molecule is determined by bond polarity and molecular shape.
a) Compare the polarity of the bonds N-Cl and C-Cl. (2 marks)
b) Predict whether the molecules, NH3(l) and CCl4(l) are polar or non-polar. Explain
your predictions. (4 marks)
4. Indicate the number of sigma and pi bonds in each of the following molecules:
(8 marks – 2 marks each)
a)
b)
c)
d)
H2O
C2H2
C2H4
C2H6
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SCH4U
Grade 12
University Chemistry
Lesson 4 – Intermolecular Forces
SCH4U – Chemistry
Unit 1 – Lesson 4
Lesson 4: Intermolecular Forces
In this lesson, we will examine the other types of bonding that have not yet been
covered in the course. We will first examine the forces that exist between molecules,
and how such forces can lead to the formation of large aggregates such as crystals.
What you will learn
After completing this lesson, you will;
•
•
•
•
explain how the properties of a solid or liquid (e.g., hardness, electrical conductivity,
surface tension) depend on the nature of the particles present and the types of
forces between them (e.g., covalent bonds, Van der Waals forces, dipole forces, and
metallic bonds)
predict the type of solid (ionic, molecular, covalent network, or metallic) formed by a
substance, and describe its properties;
conduct experiments to observe and analyse the physical properties of different
substances, and to determine the type of bonding present
describe some specialized new materials that have been created on the basis of the
findings of research on the structure of matter, chemical bonding, and other
properties of matter (e.g., bulletproof fabric, superconductors, superglue);
Bonds between Molecules: Intermolecular Bonds
Intermolecular bonds are the chemical bonds between molecules. These bonds
determine the physical state of molecular substances. These bonds are broken as a
substance undergoes a change of state. Intermolecular forces are much weaker than
covalent bonds.
There are generally three types of intermolecular bonds: London forces, dipole-dipole
forces, and hydrogen bonds. These intermolecular forces are collectively called Van
der Waals forces. These are summarized in table 4.1 below.
Force
London forces
Dipole-Dipole
Hydrogen
bonding
Table 4.1- Vander Waals forces
Description
Hold covalent molecules together. Very weak
forces of attraction. Momentary dipoles are
created by the electrons contained within the
compound, which are constantly in motion
Hold polar covalent molecules together. These
forces are stronger than London forces.
Formed between the electropositive Hydrogen
dipole and an electronegative dipole of
Oxygen, Chlorine, or Fluorine.
Copyright © 2008, Durham Continuing Education
Example
All molecules
Methane gas
(CH4)
Hydrogen
Chloride(HCl)
Pure distilled
water
Page 54 of 64
SCH4U – Chemistry
Unit 1 – Lesson 4
Using Vander Waals forces to predict Boiling Points
As the number of electrons in the molecule increases, the boiling point increases as
well. There are some general guidelines for predicting boiling points;
¾ The more polar a molecule is, the higher the boiling point
¾ The greater the number of electrons, the stronger the London force and
therefore, higher boiling point
¾ Isoelectronic molecules have the same London force
Support Questions
20. Determine the type of Vanderwaals forces that would occur between the following
molecules;
a) water, H2O
b) butane C4H10
c) hydrogen chloride, HCl
21. Which of the following pure substances has a stronger dipole-dipole force than the
other? Discuss your reasons for your conclusion.
a)
b)
c)
d)
hydrogen chloride or hydrogen fluoride
chloromethane or iodomethane
nitrogen tribromide or ammonia
water or hydrogen sulfide
The Structure and Properties of Solids
Although solids have many commonalities, the physical properties of solids vary greatly
in such physical properties such as hardness, melting point, mechanical characteristics
and conductivity.
The structure and properties of solids are related to the forces between the particles.
There are four different categories of solids, summarized in Table 4.2 following.
Table 4.2: Categories of Solids
Class of Substance
Elements combined
Ionic
Metal + nonmetal
Metallic
Metal (s)
Molecular
Non-metal(s)
Covalent network
Metalloids/carbon
Copyright © 2008, Durham Continuing Education
Examples
NaCl(s), CaCO3(s)
Cu(s), CuZn3(s)
I2(s), H2O(s), CO2(s)
C(s), SiC(s), SiO2(s)
Page 55 of 64
SCH4U – Chemistry
Unit 1 – Lesson 4
Ionic Crystals
Earlier in this unit, you learned that an ionic bond is formed when there is a transfer of
electrons from a metal atom to a non-metal atom. However, the transfer of electrons is
not what forms the bond. The ions that are formed as a result of electron transfer allow
the resulting ions to form arrangements of in a definite crystal pattern called a crystal
lattice.
The crystal lattice for sodium chloride (NaCl) is depicted below in figure 4.1
Figure 4.1 – Sodium Chloride Crystal Lattice
There is a huge variety of crystal shapes that we will
not discuss here. Ionic compounds are relatively hard
but brittle substances. They conduct electricity in
liquid but not in solid state. They also have high
melting points.
Metallic Crystals
You are likely aware of the properties of metals. They are shiny, silvery solids that are
good conductors of heat and electricity, their hardness ranges from hard to soft. The
melting points also range from high to low. Metals have a
continuous compact crystal structure.
The properties of metal are the result of bonding between
fixed positive nuclei and mobile “loose” valence electrons. In
more simplistic terms, the valence electrons float around the
network between positive centers. The motility of the
electrons explains why metals are such good conductors of
heat and electricity.
Figure 4.2 – Metallic Bonding – A sea of negative electrons floats freely amongst
positive nuclei.
Molecular Crystals
Molecular solids may be elements such as iodine and sulphur or compounds such as
ice or carbon dioxide. They are crystals that have relatively low melting points, they are
not hard, and are not conductors of electricity.
They are arranged by neutral molecules with weak intermolecular forces.
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Page 56 of 64
SCH4U – Chemistry
Unit 1 – Lesson 4
Covalent Crystals
Massive aggregates of atoms, as distinct from ions, can exist where neighbouring
atoms share a pair of electrons to form multiple, single, covalent bonds. This type of
crystal is usually found in the elements carbon and silicon.
Carbon and Silicon both have four valence electrons. Diamond and Graphite are both
made up of carbon. Diamond is a huge network of C-C linkages where all four of the
carbon bonds are equal in strength and the whole structure is enormously strong.
Graphite, on the other hand, is a carbon crystal in
which each carbon atom uses only three of its four
available valence electrons to bond with adjacent
carbons atoms. This results in layers of carbon
that are not bonded. The unused bonded
electrons can move through layers, explaining why
graphite can conduct electricity, whereas diamond,
does not. The layers of carbon in graphite are held
together by Vander Waals forces. The layers in
graphite can slip over one another easily, and
although graphite is strong in two dimensions, it is
weak in third. This makes graphite useful for
products such as pencils, golf clubs, and as a
lubricant. Diamond on the other hand is used to
cut glass, and is considered precious due its
hardness.
Figure 4.3 - Diamond
Covalent crystals tend to exceptionally hard,
with very high melting points. They are
invariably insoluble in water, since there are
no ions to attract the polar water molecule.
Figure 4.4 – Graphite
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Page 57 of 64
SCH4U – Chemistry
Unit 1 – Lesson 4
Key Question #4
1. Use the Internet to write a short report on X-ray crystallography, describing how Xrays can be used to give information on structures of solids at the atomic level.
(10 marks)
2. How does the melting point relate to the type of particle and forces present?
(2 marks)
3. Identify the main type of bonding and the type of solid for each of the following;
(6 marks)
a) SiO2
e) C
b) CH4
f) CaO
c) Cr
d) Na2S
4. Water beads on the surface of a freshly waxed car hood. Use your knowledge of
intermolecular forces to explain this observation. (5 marks)
5. All molecular compounds may have London, dipole-dipole, and hydrogen-bonding
intermolecular forces, affecting their physical and chemical properties. Indicate
which intermolecular forces contribute to the attraction between molecules in each of
the following classes of organic compounds: (8 marks)
a) pentane, C2H5
c) acetic acid, CH3CHOHCH
e) dimethylether, CH3OCH3
g) diamond, C
b) 2-propanol, CH3CHOH
d) ethybenzoate, C6H5COOCH2CH5
f) ethylamide, CH3CONH2
h) calcium carbonate, CaCO3
6. Use the theory of intermolecular bonding to explain the sequence of boiling points in
the following alkyl bromides: CH3Br(g) (4oC), C2H5Br(l) (38oC), and C3H7Br(l) (71oC).
(6 marks – 3 marks each)
7. Compare the particles and forces in the following pairs of solids; (4 marks)
a) metallic and covalent
b) molecular and ionic
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Page 58 of 64
SCH4U
Grade 12
University Chemistry
Support Question Answers
SCH4U – Chemistry
Unit 1 – Support Question Answers
Answer to Support Questions
1.
Atomic
Number
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
Element
Hydrogen
Helium
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
Sodium
Magnesium
Aluminum
Silcon
Phosphorus
Sulfur
Chlorine
Argon
Potassium
Calcium
2.
H)1e
-
IIA
-
-
Li)2 e )1 e
-
IIIA
-
-
-
-
-
-
Be)2 e )2 e
-
Na)2 e )8 e
)1 e
Mg)2 e )8 e
)2 e
-
Ca)2 e )8 e
)8 e )2 e
-
K)2 e )8 e
)8 e ) 1 e
Mass
Number
1
4
7
9
11
12
14
16
19
20
23
24
27
28
31
32
35
40
39
40
Number of
Protons
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
IVA
-
-
B)2 e )3 e
-
VA
-
-
-
-
C)2 e )4 e
-
Al)2 e )8 e )3
e
Si)2 e )8 e
)4 e
Number of
Electrons
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
VIA
-
-
-
-
N)2 e )5 e
P)2 e )8 e
)5 e
Number of
Neutrons
0
2
4
5
6
6
7
8
10
10
12
12
14
14
16
16
18
22
21
20
VIIA
-
-
-
-
O)2 e )6 e
S)2 e )8 e
)6 e
He)2 e
-
-
-
F)2 e )7 e
-
-
Cl)2 e )8 e
)7 e
-
Ar)2 e )8 e
)8 e
Trend: The group number indicates the number of valence electrons, the period
indicated the number of shells the atom has.
3. Bohr- orbits are 2D, fixed distance from nucleus, circular or elliptical path, 2n2
electrons per orbit, Sommerfield- 3D region in space, variable distance from nucleus,
no path, 2 electrons per orbital
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-
Ne)2 e )8 e
-
Page 60 of 64
-
SCH4U – Chemistry
Unit 1 – Support Question Answers
4.
Primary energy
level
Principle quantum
number
1
2
3
4
5
1
2
3
4
5
Possible
secondary
quantum numbers
0
0,1
0,1,2
0,1,2,3
0,1,2,3,4
Number of
sublevels per
primary level
1
2
3
4
5
5. l =n-1
6. Draw an energy level diagram for the following;
a) Carbon
b) Cl-1
1s
1s
2s
2s
2p
2p
3s
3p
7. Write the complete ground state electron configurations for the following;
a) 1s22s1
b) 1s22s22p4
c) 1s22s22p63s23p64s2
d) 1s22s22p63s23p64s23d2
e) 1s22s22p63s23p64s23d104p65s1
f) 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2
g) 1s22s22p63s23p64s23d104p65s24d105p66s24f12
Copyright © 2008, Durham Continuing Education
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SCH4U – Chemistry
Unit 1 – Support Question Answers
8. Write the abbreviated (shorthand method) ground state electron configurations for
the following;
b) [He]2s22p3
d) [Ar]4s23d6
f) [Xe]6s2
a) [He]
c) [Ne]3s23p5
e) [Ar]4s23d10
g) [Xe]6s24f145d106p4
9.
Element
Oxygen
Chlorine
Sodium
Phosphorus
Atomic
Symbol
Electron
Configuration
1s22s22p4
1s22s22p63s23p5
1s22s22p63s1
1s22s22p63s23p3
O
Cl
Na
P
#of valence
electrons
6
7
1
5
Valence
-2
-1
+1
-3
10. Identify the elements and write the Lewis structure for the following electron
configurations;
a) 1s22s22p4
b) 1s22s22p63s23p3
c) [Ar]4s23d104p5
d) [Kr]5s1
a) oxygen
b) phosphorus
c) bromine
d) rubidium
xx
x x Ox x
xx
xx
P
Br
Rb
11.
Bond formation
+1
a) lithium and
fluorine
-1
xx
Li
x
F
Name of
compound
Lithium
fluoride
Chemical
formula
LiF
Anion
Cation
fluorine
lithium
Calcium
phosphide
Ca3P2
phosphorus
calcium
x
x
xx
b) calcium
and
phosphorus
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SCH4U – Chemistry
Unit 1 – Support Question Answers
12. Draw Lewis Structures for each of the following molecules or polyatomic ions.
a) ClO4-1
c)HCO3-2
b) CN-1
-1
O
O
Cl
-1
O
O
N
O
HO
C
-2
C
O
13. What atomic orbital or orbitals are available for bonding for each of the following
atoms?
a) H -1s
b) F -2s, 2p
c) S- 3s,3p
d) Br – 4p
14. Provide ground state and promoted state electron configurations for each of the
following atoms and indicate the type of hybridization involved when each atom
forms a compound;
a) carbon in CH4, 1s22s22p2 promoted to1s22s12px12py12pz1 (sp3)
b) boron in BH3 1s22s22p1 promoted to 1s22s12px12py1 (sp2)
c) beryllium in BeH2 1s22s2 promoted to 1s22s12px1 (sp)
15. When are π bonds formed?
Pi bonds are created when there is overlap sideways of non hybridized orbitals,
usually p orbitals
16. a) C2Cl4 - sp2 –There will be a double bond between the carbon atoms, one pi bonds
and one sigma. The two half-filled p orbitals of the adjacent atom overlap sideways.
b) C2F2 - sp – The sigma bonds use sp orbital, the two pairs of half-filled p of the
adjacent carbon overlap sideways.
17. a) CO2(g) - AX2- Linear
b) HCN - AX2 –Linear
c) BF3 - AX3- Trigonal planar
d) SiCl4 - AX4 -tetrahedral
e) CH4 - AX4 - tetrahedral
f) OCl2 - AX2E2 – V-shaped or bent
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SCH4U – Chemistry
Unit 1 – Support Question Answers
18.
Substance
EN Element 1
EN Element 2
∆ EN
NaCl
Cl2
HF
Na = 0.9
Cl = 3.0
H = 2.1
Cl = 3.0
Cl = 3.0
F = 4.0
2.1
0
1.9
Ionic or
Covalent?
Ionic
Covalent
Polar covalent
19. Predict the polarity of following molecules.
a)
b)
c)
d)
BF3 -polar
OF2 – non-polar
CI2 – non-polar
PCl3 -polar
20. Determine the type of Vanderwaals forces that would occur between the following
molecules;
a) water, H2O – hydrogen bonding
b) butane C4H10 –London forces
c) hydrogen chloride, HCl –dipole-dipole
21. Which of the following pure substances has a stronger dipole-dipole force than the
other? Provide your reasoning.
a) hydrogen chloride or hydrogen fluoride: HF – greater electronegativity difference
between atoms
b) chloromethane or iodomethane: chloromethane - greater electronegativity
difference between atoms
c) nitrogen tribromide or ammonia: ammonia - greater electronegativity difference
between atoms
d) water or hydrogen sulphide: water - greater electronegativity difference between
atoms
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