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PART 2 – ATOMIC CHEMISTRY
OUTLINE
Lesson 5: The Nature of Light
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Electromagnetic Radiation (EMR)
Frequency and Wavelength
Frequency, Wavelength, and Energy
The Electromagnetic Spectrum
Line Spectra
Flame Tests
Applications of Line Spectra
Lesson 6: The Quantum Mechanical Model of the Atom
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The Model of the Atom…. So Far
Neil’s Bohr Explains Line Spectra
Quantum Theory
Orbitals
Lesson 7: Electron Configurations
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The Orbitals
The Electron Configuration
Aufbau Principle
Hund’s Rule
Zig-Zag Rule
Electron Configurations of Atoms
Valence Configurations
Electron Configurations and the Periodic Table
Lesson 8: Chemical Bonding: Electronegativity
 Defining Electronegativity
 Polar, Non-polar, and Ionic Bonds
 Predicting Bond Character
Lesson 9: Lewis Dot Diagrams
 Dot Structures for Atoms
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 Octet Rule
 Drawing Lewis Dot Structures
Lesson 10: Periodic Trends
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Force on an Electron
Atomic Radii
Trends in Atomic Radii
Ionization Energy
Periodic Trends in Ionization Energy
Lesson 5: The Nature of Light
The chemical properties of atoms are determined by their structure -- more specifically,
the arrangement of their electrons. In previous science courses you have learned about the
structure of the atom. These models do not sufficiently describe the observed properties
of atoms. In this unit we will further examine the structure of atoms and how they
combine to form molecules and compounds. We will relate their structure to their
chemical properties.
Before we study the atom we must first understand the nature of light. This may seem a
little strange, but in the early 1900’s scientists observed that some substances emitted
light when heated. This emitted light revealed information about the nature of the
substance.
When you have completed this lesson, you will be able to:
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Describe light in terms of electromagnetic energy.
Describe the electromagnetic spectrum.
Describe the relationship between frequency, wavelength and energy of light.
Identify an element based on its flame test.
Describe line and continuous spectra
Electromagnetic Radiation (EMR)
In 1678 Christian Huygens (1629-1695) proposed that light was in the form of waves.
But in 1704 Sir Isaac Netwon (1642-1727) proposed that light had a particulate nature
because this helped explain his observations from his experiments in optics. Netwon’s
theory of light was favoured over Huygens’ particle theory of light for over 100 years.
In 1807, Thomas Young (1773-1829) and Augustin Fresnel (1788-1827) performed
experiments where light was shone through slits. They observed an interference pattern
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that could only result if light had a wave-like nature. In 1860, James Clerk Maxwell
mathematically demonstrated that light was composed of both waves of electrical and
magnetic energy or electromagnetic radiation.
Frequency and Wavelength
Whether the waves are radio waves, waves of red light or X-rays, they all travel in a
wave pattern as shown on the following page:
The wavelength (symbol: λ Greek letter “lambda”) is the distance from one crest to the
next crest or trough to trough, usually measured in metres, m.
Another characteristic of waves is the frequency. Frequency (symbol: ν, Greek letter
“nu”) is the number of wavelengths, or wave cycles, that pass a point per unit time.
Frequency is measured in cycles per second, or the SI unit hertz (Hz). The frequency can
also be represented by the unit s-1, or the reciprocal of seconds.
Examine the diagrams below:
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What relationship do you notice between wavelength and frequency?
In all of the examples above, wavelength and frequency are inversely related, that is, as
wavelength increases the frequency decreases. This makes sense, since the longer the
wavelength the more time it will take to move past a point.
Frequency, Wavelength and Energy
In1900 Max Planck (1858-1947), a German physicist, heated objects until they glowed,
then studied the light given off by the glowing objects. He discovered that the frequency
of the light given off was directly related to the amount of energy released.
Since wavelength is inversely related to frequency, wavelength will also be inversely
related to the energy of the light. That is, the greater the wavelength, the lower the
energy. If you recall from the previous page, radio waves have a wavelength of about 2 m
(frequency about 1.5 x 108 Hz) and X-rays have a wavelength of about 1.25 x 10–10 m
(frequency about 2.4 x 1018 Hz). X-rays have a greater amount of energy than radio
waves. Radio waves are all around us (you can get a radio signal almost anywhere) and
are not harmful to us, where exposure to X-rays is limited because their high energy is
damaging. Gamma rays can result from nuclear reactions taking place in
objects such as pulsars, quasars, and black holes.
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The Electromagnetic Spectrum
If we shine sunlight through a prism we get a rainbow of colours, known as a spectrum.
Each colour in the rainbow represents light of a different frequency or wavelength.
This is known as a continuous spectrum because there are no breaks between the
individual wavelengths; the colours blend together. The visible spectrum is actually a
very small portion of the entire electromagnetic spectrum.
Line Spectra
When an electric current is passed through hydrogen gas in a tube the gas glows. If the
light produced by the glowing gas is focused through a slit then passed through a prism, a
spectrum with distinct lines is produced.
This type of spectrum is known as an emission spectrum, since it is the separate
wavelengths of light emitted by the gas. This emission spectrum is also known as a line
spectrum because the light separates into discrete wavelengths of light that appear as
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lines of colour on a screen or photographic plate. Unlike a continuous spectrum, the
colours in a line spectrum do not blend into each other. The visible region of the emission
spectrum of hydrogen is shown in the figure below. Each coloured line corresponds to an
exact wavelength or frequency of light.
Spectroscopy, spectrophotometry and spectrometry are several names given to the
techniques used to determine a substance’s emission spectrum
To see the emission or line spectra of several other elements, visit this website:
http://astro.u-strasbg.fr/~koppen/discharge/
If you examine these spectra you will notice they are all different. The line spectrum of
each element is like the fingerprint of that element.
In 1870 the bright yellow line in helium’s spectrum was first seen when examining the
light emitted by the Sun. The discoverers determined this was a new element and named
the element helium, after the Greek Sun god “helios”. Twenty five years later helium was
discovered on Earth.
Flame Tests
When some elements are burned, they give off a distinctive colour of light. This is due to
the emission of light predominantly in that wavelength in their line spectrum. Preliminary
evidence of the presence of a metal is often the colour that results when it is placed in a
flame. This is known as a flame test.
Salts containing the metals, when placed into a flame will give the flame the colours as
shown below:
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Colour
Element
green
copper
yellow
sodium
red
strontium
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yellow-green
barium
orange-red
calcium
purple
potassium
purple-red
lithium
Some commercially made fireplace logs advertise that they will produce multi-coloured
flames, when burned. This is because the logs are made of pressed sawdust or smaller
chips of wood soaked in a solution containing several metal salts. You can also see
colourful flames if you burn the paper from newspapers or magazines with colour
pictures, since the coloured inks contain various metal salts.
Applications of Line Spectra
As mentioned earlier helium was first discovered on the Sun before it was found on Earth
form its distinctive line spectrum. Astronomers use line spectra to determine the elements
in various light sources, such as stars and nebulae. Astronomers don’t just study visible
light emissions. Many astronomers use radio telescopes to find distant objects.
The aurora borealis, or northern lights, are a result of ionized gases becoming excited.
This excitation leads to the release of light which you can see in the northern sky.
Another popular application of the spectra produces by elements is in the manufacturing
of fireworks. We saw the distinctive colours produced by metal ions in flame tests. These
metal salts are used in fireworks to produce the distinctive colours of light. Below is an
example of a fireworks shell that contains copper and strontium for a burst that will have
both red and green.
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Assignment 5
1. Put the following in order of increasing energy: green light, x-rays, radio, red
light, ultraviolet, microwaves, blue light, gamma rays.
2. Describe the relationship between frequency, wavelength and energy.
3. Why is light called electromagnetic radiation?
4. Compare and contrast line and continuous spectra.
5. Give the colour for the flame test of each of the following:
a. copper
b. strontium
c. lithium
d. potassium
e. barium
f. calcium
g. sodium
Lesson 6: The Quantum Mechanical Model of the Atom
Introduction
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Why do line spectra for elements exist? Why is the light emitted in discrete wavelengths?
What does this reveal about the structure of the atom?
These are questions posed by scientists in the early part of the 20th century. We will
answer these questions here and discuss the improvements on the Bohr model of the
atom.
When you have completed this lesson, you will be able to:
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Explain the development of the Quantum Mechanical Model of the atom.
Explain the formation of line spectra.
The Model of the Atom…So Far
In previous science courses you have learned the following historical development of the
model of the atom.
The early theories of matter were not based upon experiments, but by observations of
physical phenomena. In the Greek era, about 440 B.C., Democritus stated that we took
matter and broke it down into smaller pieces there would eventually be a point to which
we could not divide matter down further. He said the smallest parts of matter were too
small to be seen and he called them “atomos”, Greek for uncuttable or indivisible.
About a hundred years later Aristotle didn’t believe in the atomic model of matter. He
influenced scientific thought for several hundred years later.
In the 1600’s and 1700’s several scientists including Sir Isaac Newton and Robert
Boyle revived the atomic theory of matter.
In 1803, John Dalton (1766-1844) published his atomic theory of matter. Dalton’s model
assumed the atom was a solid, indivisible sphere. He based his theory on mass
experiments.
In 1897, Joseph John (J.J.) Thomson (1856-1940) used
cathode ray tube studies to propose what has been called the
“plum pudding model” or the “raisin bun model”. He believed
that the atom was a sphere of “doughy” positive charge with
small negatively charged electrons (raisins) that could be easily
removed. These electrons were a very small portion of the
atom’s mass.
In 1911, Ernest Rutherford (1871-1937) interpreted the results of the “gold foil
experiment”. He fired helium nuclei (alpha particles) at thin gold foil. If the Thomson
model of the atom was true, the particles should pass through the foil. He found that some
of the particles were deflected.
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He interpreted this as the atom having a very small nucleus, having most of the mass of
the atom, and electrons move around the nucleus. Rutherford discovered the nucleus.
Neils Bohr Explains Line Spectra
In 1922, Neils Bohr (1885-1962) developed a model that explained hydrogen’s line
spectra. His model dealt specifically with the movement of electrons around the atom.
Bohr proposed that the electrons in a hydrogen atom are arranged in stable orbits around
the nucleus depending on their energy. He said the orbits were like a ladder with
unequally spaced rungs. Or a stairway with unequal stairs.
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When radiation is absorbed by an atom, an electron jumps from the ground state (its
resting state) to a higher unstable energy level called the excited state.
The model then states that this electron eventually loses energy and changes to a lower
energy level by emitting energy in the form of light. Click the link below to view an
animation of this process.
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The individual lines of the hydrogen line spectra are due the hydrogen electron becoming
excited, moving to higher orbit, then falling to a lower orbit and releasing energy in the
form of light. The distinct lines are due the distinct orbits or energy levels the electron
can occupy.
Go to the site below for a tutorial on emission spectra and the Bohr model of the atom.
Both sites are links on my website.
http://www2.wwnorton.com/college/chemistry/gilbert/tutorials/interface.swf?chapter=cha
pter_03&folder=emission_absorption
http://www2.wwnorton.com/college/chemistry/gilbert/tutorials/interface.swf?chapter=cha
pter_03&folder=hydrogen_energies
Quantum Theory
Unfortunately for Bohr, this model only worked for atoms with a single electron, that is,
hydrogen. Still, physicists were both mystified and intrigued by Bohr’s Model for the
atom. They wondered why the energies of electrons were allowed only certain energies,
or quantized. Apparently, even Bohr was not able to provide a logical explanation.
To find a solution to this puzzle, we must go back a few years to Max Planck’s discovery
of the energy-frequency relationship. Planck determined that the light gained or emitted
by objects could only occur in discrete amounts he called quanta or quantum (singular).
Bohr said the energy released by the electron was quantized.
Albert Einstein (1879-1955) used Planck’s Quantum Theory, in 1905, to explain why
the photoelectric effect occurs. Scientists had long known about the photoelectric effect
but could not explain why it occurred. The photoelectric effect occurs when ultraviolet
light shines on a metal surface, resulting in an electron being emitted from the metal. The
electron can then be used to produce an electric current. This is the basis of solar
batteries. Einstein proposed that this occurs because electromagnetic radiation is a stream
of tiny bundles of energy called photons. These photons have no mass but carry a
quantum of energy. The energy of the photon is transferred to an electron giving the
electron sufficient energy to escape the metal. Using the Quantum Theory, Einstein was
able to use to calculate the energy of a single photon of light. In 1922, Einstein received
the Nobel Prize for his work on the photoelectric effect.
According to Planck and Einstein, as an electron “falls” from one Bohr orbit (or energy
level) to a lower Bohr orbit the electron releases a single photon of light energy with a
specific wavelength or frequency, determined by the energy of the photon.
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Source: http://intro.chem.okstate.edu/ChemSource/Atomic/concpt3.htm
Bohr’s orbits or energy levels become principal quantum numbers (n), also called
principal energy levels. Within each energy level are a set of orbitals or sublevels. In the
Bohr atom the “orbit” was a path the electron would follow, where orbitals represent
regions in space around the nucleus that the electron will probably be found. The
principal quantum number will indicate the size and energy of each orbital. The lowest
energy level is n = 1 and the highest is n = 7. As the value of n increases the energy levels
become larger and the electrons spend more time further from the nucleus.
Orbitals
Electron orbitals are described in terms of three variables: size, shape and orientation in
space. There are 4 types of orbitals with different regions of probability, given the letter
names: s, p, d and f. These regions of probability result in fuzzy electron clouds that have
different shapes.
The s-orbitals are spherical in shape. Each energy level has one s-orbital. The s-orbital in
the first energy level is given the designation 1s, the s-orbital in the second energy level
has the designation 2s, and so on. Below are representations of the s-orbitals for the first
3 energy levels. Notice how the size of the orbital increases as the principal quantum
number, or energy level, increases. The position of the nucleus is assumed to be the point
where the 3 axis intersect.
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The p-orbitals are dumbbell-shaped and each has two regions or lobes of high probability.
Since these are regions of probability, electrons spend equal amounts of time in both
lobes of the p-orbital.
The p-orbitals are only present in the second energy level and higher. There are 3 kinds
of p-orbitals, depending on their position in space:
The d-orbitals are only present in the third energy level and up. They have varied shapes
and occupy more complex regions in space than the p-orbitals. There are 5 types of dorbitals:
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The f-orbitals are only present in the 4th energy level and up. There are 7 different types
of f-orbitals and their shapes are considerably more complex than the d-ortbitals (some
containing as many as 8 lobes).
Lesson 7: Electron Configurations
The characteristics of elements are due to the arrangement of their electrons. If we
examine the arrangement of the electrons around the nucleus we can predict some
of the properties of that element.
Learning Outcomes
When you have completed this lesson, you will be able to:
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Write the electron configuration of atoms up to Krypton, Element 36
Write the valence electron configuration of an atom.
The Orbitals
From the last lesson, we learned that orbitals are regions in space, around the nucleus of
the atom, where there is a high probability of finding an electron. Each principle energy
level, n, has n2 orbitals or sublevels.
Each orbital is given the letter designation of s, p, d or f. Each of these orbitals has a
different shape. There are also a different number of each type of orbital possible in each
energy level, n.
The table below summarizes the types and number of orbitals available in each energy
level.
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Principle Energy Level
or Principle Quantum
Number (n)
Number of orbitals
(n2)
1
1
1 s-orbital
2
4
1 s-orbital + 3 p-orbitals
3
9
1 s-orbital + 3 p-orbitals + 5 d-orbitals
4
16
1 s-orbital + 3 p-orbitals + 5 d-orbitals + 7 f-orbitals
Orbital Types
If the maximum number of electrons per orbital is two, then the number of electrons per
energy level will be 2×n2. The table from the previous page can now be expanded to
include the number of electrons in each energy level.
Principle Energy
Level or
Principle
Quantum
Number (n)
Number of
orbitals
(n2)
1
1
1×s
2
2
4
1×s+ 3×p
8
3
9
1×s + 3×p + 5×d
18
4
16
1×s + 3×p + 5×d + 7×f
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Orbital Types
Number of electrons
(2n2)
The Electron Configuration
The electron configuration of an atom assigns a number to each electron. The number
includes the principal energy level of the electron, the name of the orbital and the number
of electrons in that orbital. The electron configuration of oxygen is shown below:
The electron configuration shows that there are 2 electrons in the s-orbital of the first
energy level (1s2), 2 electrons in the s-orbital of the second energy level (2s2) and 4
electrons in the p-orbital of the second energy level (2p4).
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The orbitals are always written as lower case letters and the number of electrons is
always written as superscripts.
Note: When submitting assignments using a word processor use the superscript function
for writing electron configurations. For example 1s22s22p4.
Aufbau Principle
Aufbau is German for “build-up”. The aufbau principle was proposed by Neils Bohr.
The aufbau principle states that as electrons are added to an atom, electrons fill the lowest
energy state before adding electrons to a higher energy level. According to the Pauli
Exclusion Principle, each orbital gets two electrons so the electrons will fill the energy
levels according to the diagram below.
In each energy level, the s orbital has the lowest energy so it is always the first orbital
filled in that energy level. The range of energies in every principal energy level allows for
an overlap between energy levels. Notice that the 4s orbital has a lower energy than the
3d orbitals. There is an overlap between the energy of the 4s orbital and the 3d orbitals.
The reason for this overlap is beyond the scope of this course.
According to the aufbau principle, electrons will fill the 4s orbital before filling the 3d
orbitals. When the 3d orbitals are full, electrons begin filling the 4p orbitals.
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Hund’s Rule
Friedich Hund (1896-1997) proposed that when electrons fill orbitals, they obey the
aufbau principle and fill such that the number of unpaired electrons is maximized. That
is, before filling the first p-orbital with two electrons, an electron is placed into the pxorbital, then an electron into the py-orbital then the pz-orbital before filling the px-orbital.
This is known as Hund's Rule.
Below is a table that includes both the orbital diagrams and the electron configurations
for the first ten elements. An orbital diagram represents the spin of an electron in terms of
up (↑) and down (↓) arrows. These spins can be thought of “up” and “down” or
“clockwise” and “counterclockwise”.
According to the Pauli Exclusion Principle, each orbital can hold two electrons, providing
they have opposite spins. In the hydrogen atom there is only one electron, which we
arbitrarily choose to represent with an upward arrow. The direction we begin with is not
important. The helium atom has 2 electrons to fill the s-orbital in the first principal
energy level. The electron must fill the orbital with an opposite spin, represented by an
upward and downward arrow. This trend continues until we reach the carbon atom.
According to the aubau principle, electrons fill lower energy levels before filling higher
energy levels. Once the 1s orbital is full, electrons begin filling the 2s orbital, then the 2p,
the 3s, etc. Carbon has 6 electrons. Carbon’s first four electrons fill the 1s and 2s orbitals.
According to Hund’s Rule, carbon’s next two electrons fill the px and py orbitals.
Nitrogen’s sixth electron fills the third p-orbital. Oxygen’s last electron adds a second
electron to the px-orbital with an opposite spin.
Zig-Zag Rule
The overlap between d-orbitals, f-orbitals and s-orbitals makes remembering electron
configurations difficult. There is a quick tool we can use to remember the filling order for
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electron configurations. This tool is called the zig-zag rule. To make this tool start by
writing the orbitals out in a list downwards:
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
We can draw a diagonal line through the orbitals.
To write the filling order we just follow the diagonal lines. When we reach the tip (end)of
an arrow, we start at the tail (start) of the next arrow. The filling order is:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s etc.
You do not have to use the zig-zag rule. Some students find it useful rather than
memorizing the filling order.
Electron Configurations of Atoms
To write the electron configuration of an atom:
1. Determine the number of electrons (atomic number).
2. Use zig-zag, if you wish, to add electrons until the number of electrons is reached.
Example 1. Write the complete electron configuration for magnesium.
Solution.
Step 1. Number of electrons.
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Atomic number = 12
Number of electrons = 12
Step 2. Use zig-zag.
Follow the zig-zag (use the aufbau principle), adding electrons until 12 electrons.
1s22s22p6 so far, 10 electrons. Need 2 more.
1s22s22p63s2
Example 2. Write the complete electron configuration for germanium.
Solution.
Step 1. Number of electrons.
Atomic number = 32
Number of electrons = 32
Step 2. Use zig-zag.
Follow the zig-zag (use the aufbau principle), adding electrons until 12 electrons.
1s22s22p6 so far, 10 electrons (neon’s electron configuration).
1s22s22p63s23p6 18 electrons (argon’s configuration)
1s22s22p63s23p64s2 Remember the 3d-orbital now fills due to overlapping.
1s22s22p63s23p64s23d104p2
This can also be written using what is called a kernel. The kernel is the highest noble gas
structure within the configuration. In this case it is argon. We then write the symbol for
argon in square brackets followed by the remaining configuration:
[Ar]4s23d104p2
This kernel configuration is helpful when writing the configurations of atoms containing
many electrons.
Valence Configurations
he valence electrons are the electrons found in the outer-most or highest energy level.
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Example 4. Write the valence electron configuration for fluorine.
Solution.
Step 1. Write the complete electron configuration.
F = 9 electrons
1s22s22p5
Step 2. Choose the electrons in the highest energy level.
The highest energy level is the second, so the electrons in the second energy level are the
valence electrons.
The valence configuration is 2s22p5
Example 5. Write the valence electron configuration for germanium.
Solution.
Step 1. Write the complete electron configuration.
From a previous example
1s22s22p63s23p64s23d104p2
Step 2. Choose the electrons in the highest energy level.
The highest energy level is the fourth, so the electrons in the fourth energy level are the
valence electrons.
Even though the 3d orbital fills after the 4s orbital, the 4s orbital is in the highest or
outer-most energy level.
The valence configuration is 4s24p2.
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Electron Configurations and the Periodic Table
When we examine the electron configurations of the elements in the periodic table, we
can see a pattern. The periodic table seems to be arranged according to the electron
configuration of the atoms:
The electron configuration of atoms in these 4 areas will have the indicated orbital filling
last. For example, transition metals fill the d-orbitals.
Another trend is the energy level of the valence electrons is the same as the period of the
atom. The first period, which contains hydrogen and helium, has been omitted in the
diagram above. The first element in the s block would be lithium.
You learned in previous science courses that elements in columns or groups have similar
properties. This is because of their valence electron configurations. An element’s
chemical properties and bonding characteristics are determined by its valence electrons.
In this case atoms are quite shallow; they are more interested in what’s on the outside
than what is inside.
Below is a link to a good tutorial on electron configurations and orbital diagrams:
http://www2.wwnorton.com/college/chemistry/gilbert/tutorials/interface.swf?chapter=cha
pter_03&folder=orbital_filling
See my website for a direct link to this.
Assignment 7
1. How many electrons in an atom can have the designation
a. 1s
b. 2p
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c.
d.
e.
f.
g.
h.
TCLC
3px
6f
3dxy
n=2
4p
n=5
2. Write complete electronic configurations for the following atoms:
a.
b.
c.
d.
e.
f.
g.
h.
i.
P
Ca
Cu
Zn
S
Ge
Br
Kr
K
3. How many unpaired electrons are there in each of the following:
a.
b.
c.
d.
e.
f.
Mn
As
Sr
Sn
Lu
Te
4. Write the electronic configurations for the valence electrons of each of the following.
a.
b.
c.
d.
e.
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Mg
P
Se
Br
Ni
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Lesson 8: Chemical Bonding: Electronegativity
When you have completed this lesson, you will be able to:


Use electronegativity values to predict the type of bond between two atoms.
Describe a bond in terms of polar covalent, non-polar covalent and ionic.
Defining Electronegativity
Recall that chemical bonds form when electron pairs are shared by atoms (covalent
bonds) or electrons are transferred from one atom to another and the atoms are joined by
electrostatic forces (ionic bond).
The ability of atoms to attract electrons in a bond is called electronegativity. Fluorine is
the most electronegative of all atoms with a value of 4.0 (no units). The least
electronegative elements are in the lower left corner of the periodic table (cesium = 0.7).
Generally, electronegativity tends to increase as you move across the periodic table from
left to right and up the periodic table.
When atoms with different electronegativities form a bond the atom with the higher
electronegativity will draw electrons to itself in that bond. If electrons are closer to one
atom than another in a bond, the charge of the electrons is not shared equally. As a result,
one atom is more negatively charged than the other.
Bonds formed between atoms of differing electronegativities are called polar bonds
because one end of the bond has a partial positive charge and the other has a partial
negative charge because the electrons are shared unequally. A good example of a polar
bond is hydrogen fluoride, HF.
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Fluorine is the most electronegative atom. When bonded to hydrogen, the fluorine draws
the electrons towards itself. The arrow between the H and F indicates the direction the
electrons are drawn towards resulting in a more positive end. As a result, the fluorine end
of the bond has a partial negative charge (δ−) and the hydrogen has a partial positive
charge (δ+).
Polar, Non-polar and Ionic Bonds
Generally, the greater the electronegativity difference the more polar the bond. If the
electronegativity difference is large enough the bonds are considered to be ionic. This is
why we use the general rule that metal and non-metal atoms make ionic bonds, since
metals usually have low electronegativities and non-metals have high electronegativities.
There is a Figure in your text you may refer to with electronegativity values for selected
elements. See page 372 of your text. You may have this to refer to on any test or exam.
Following is a table summarizing the differences in electronegativites that result in each
type of bond: Memorize the numbers in the first column and what type of bond they
produce.
Electronegativity
Difference
Character of Bond
Percent Ionic
Character
less than 0.4
non-polar covalent
0% − 5%
0.4−1.9
polar covalent
5% − 60%
greater than 1.9
ionic
>60%
Predicting Bond Character
Example 1. What type of bond forms between sodium and chlorine in NaCl?
Solution.
Find the electronegativity values for Na and Cl on the table.
Na = 0.9
Cl = 3.0
Find the EN difference
3.0 - 0.9 = 2.1
The difference is greater than 1.9 so the bond is ionic.
Example 2. What type of bond forms between sulphur and oxygen in SO3?
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Solution.
Find the electronegativity values for S and O on the table.
S = 2.5
0 = 3.5
Find the EN difference
3.5 - 2.5 = 1.0
The difference is between 0.4 and 1.9 so the bond is polar covalent.
Example 3. What type of bond forms between aluminum and chlorine in AlCl3?
Solution.
Find the electronegativity values for Al and Cl on the table.
Al = 1.5
Cl = 3.0
Find the EN difference
3.0 - 1.5 = 1.5
The difference is between 0.4 and 1.9 so the bond is polar covalent.
This is somewhat surprising because we used the general rule of metal and non-metal
makes ionic. After looking at electronegativities we find that aluminum chloride is more
covalent than ionic!
We will still use metal + non-metal to predict an ionic bond, but use electronegativity
values to confirm our prediction.
Assignment 8
Answer the following questions. Refer to the Electronegativity Table when needed.
1. Describe the trend of electronegativity on the periodic table.
2. Describe each of the following in terms of electronegativity.
a) non-polar covalent bond
b) polar covalent bond
c) ionic bond.
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3. Describe the following as non-polar covalent, polar covalent or ionic
a) N and H
b) F and F
c) Ca and O
d) Al and Br
e) H and I
f) K and Cl
4. List the following in increasing ionic character
a) Mg-F, Ca-I, Ca-Cl, Mg-Cl
b) Al-Cl, H-Cl, K-Cl, Cu-Cl
b) C-O, C-H, C-F, C-Br
c) S-Cl, H-C, H-F, H-H, H-Cl, H-O
Lesson 9: Lewis Dot Diagrams
When you have completed this lesson, you will be able to:

Draw dot diagrams for covalent compounds.
Dot Structures for Atoms
Since valence electrons are the only electrons involved in bonding, when drawing the dot
structures for atoms only valence electrons are drawn. We also take into account the
orbitals electrons occupy. Each dot represents an electron in the valence shell of the atom.
Shown below are the electron dot diagrams for the period 2 atoms.
Although it is not a hard and fast rule, we generally add electrons to atoms in the order
that matches the orbitals electrons occupy.
As we will see later, arranging th atoms in this pattern gives us some insights into the
bonding characteristics of the atom.
Octet Rule
In 1916 Gilbert Lewis (1875-1946) published a paper in which he described drawing the
valence electrons of atoms as dots at the vertices of a cube. Eventually, he would move
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away from the cube and focus on the electrons in a bond. Although he never actually
used the word “octet” in his published works, the “octet rule” is generally attributed to
Lewis.
In earlier science courses, you learned that atoms will tend to gain electrons, lose
electrons or share electrons to fill their valence shell. A full valence is shell is eight
electrons, or an octet. When drawing electron dot diagrams, we ensure that all atoms are
surrounded by eight electrons (represented as dots). This is known as the “octet rule”.
There are a few exceptions to the octet rule, but we will only examine one, hydrogen.
Since hydrogen only has a 1s orbital, a full valence shell only consists of two electrons.
In this course, the only atom that will not have a complete octet is the hydrogen atom.
Drawing Lewis Dot Structures
When drawing Lewis Dot Diagrams, follow the steps below.
1. Determine the total number of valence electrons for all atoms. Remember
to add electrons for negative ions and remove electrons for polyatomic
ions that are positively charged.
2. Position the least electronegative atom in the centre of the molecule or ion.
3. Write the other atoms around the central atom. Fluorine and hydrogen
atoms will usually be on the outside of the molecule or ion.
4. Place 8 dots/electrons around each atom (except hydrogen), ensuring they
all meet the requirements of the octet rule.
5. If there are too few electrons to give each atom a complete octet, add
double or triple bonds.
Example 1. Draw the Lewis Dot structure of HF
Step 1: Determine the total number of valence electrons.
Step 2: Position the least electronegative atom in the centre.
This molecule has only 2 atoms, so there is no central atom.
Step 3: Place 8 dots/electrons around each atom.
We use a little x to represent the electron coming from the hydrogen atom.
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Step 4: Count the dots.
There are 8 dots. This equals the total number of valence electrons shared by the two
atoms.
This is the Lewis structure for HF. An alternate way to show this is to use a line to
represent the two electrons in the bond as shown below:
Example 2. Draw the Lewis structure for CO2.
Step 1: Determine the total number of valence electrons.
Step 2: Position the least electronegative atom in the centre.
In atoms containing carbon, carbon is usually the central atom
O
C
O
Step 3: Place 8 dots/electrons around each atom.
The x’s represent the electrons coming from the carbon atom.
Step 4: Count the dots.
There are 20 dots. We only have a total of 16 electrons, so we will take a pair off the
carbon and a pair off an oxygen and replace them with an extra shared pair. This will
create one double bond.
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Notice that the carbon and the oxygen on the right side still have 8 electrons, including
the 4 shared electrons.
Count the electrons again. How many electrons? That’s right, 18 electrons. Still too
many.
Let’s do the same with the oxygen on the left. Take away two of its electrons and two of
carbon’s electrons and make a double bond.
Count the electrons again. Now there are 16 electrons. This is the Lewis structure for
CO2. Another way to draw the structure, using lines for share pairs would look like the
structure below.
Example 3. Draw the Lewis structure for the sulphate ion, SO42-.
Step 1: Determine the total number of valence electrons.
We must also add another two electrons for the charge on the negative ion, making the
total 32 electrons.
Step 2: Position the least electronegative atom in the centre.
The least electronegative atom is sulphur. Often in molecules or ions with a single atom,
like sulphur, and multiple atoms, like oxygen, the single atom will be the central atom.
Step 3: Place 8 dots/electrons around each atom.
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For ions, we place square brackets around the ion and put the charge outside the brackets.
Step 4: Count the dots.
There are 32 dots. Just enough. This is the Lewis structure for the sulphate ion.
Assignment 9
Draw the Lewis structure for each of the following.
1. NH3
2. SO3
3. BrF
4. O2
5. PO43–
6. CH4
7. HCN
8. N2
9. Cl2O
10. ClO3–
11. CO32–
12. SiF4
Lesson 10: Periodic Trends
Introduction
In 1870 Dmitri Mendeleev (1834-1907) arranged the 65 elements known at that time
into a periodic table. He proposed the “periodic law”, which states that when the elements
are arranged according to their mass they have a regular periodic repeating of similar
properties. Using the periodic law he was able to predict the existence of several elements
and their properties that were as yet undiscovered for which he left spaces in his table.
In this lesson we will examine some periodic trends and explain them using principles of
atomic structure.
When you have completed this lesson, you will be able to:


Describe the factors that affect the force on an electron.
Explain the trends in atomic radius, ionic radius and ionization energy.
Force on an Electron
The positively charged nucleus of the atom applies a force on each negatively charged
electron, holding the electrons around the atom. The force on an electron will affect its
position around the atom. The force on an electron is dependent upon two factors:
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1. The nuclear charge.
The more protons that are in the nucleus, the greater will be the nuclear charge. The
greater the nuclear charge, the more force there will be on each electron in the atom.
2. Distance of the electron from the nucleus.
According to Coulomb’s Law, the further two charges are from each other, the lower
their force of attraction on each other. This is very similar to Gravitational Law in that the
further a body is from the surface of the Earth, the lower the force of gravity.
3. Shielding Effect
Inner electrons tend to screen or shield the force of the nucleus from outer electrons.
The inner electrons almost cast a shadow on the outer electrons. As a result, the force of
the nucleus on outer electrons is reduced. The effective nuclear charge on outer electrons
is lower as more orbitals are added to an atom.
Atomic Radii
As you learned in earlier lessons, the electron orbitals are based upon regions of
probability. As such, atoms do not have clearly defined borders. Since the border of the
atom is fuzzy, chemists developed a method for measuring the size of an atom. The
atomic radius is the distance between two nuclei when two like atoms are bonded
together.
In the figure above, the distance between the two nuclei in the molecule is 50 pm (1 pm =
10−12m). The atomic radius of the atom is then 25 pm.
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Trends in Atomic Radii
As you move to the right, a proton is added to each atom. For each proton added, an
electron isadded to the same principal energy level. As a result the amount of shielding
occurring on each electron remains constant. If the nuclear charge increases, while the
shielding remains constant, the force one each electron increases. The increased force on
each electron draws the outermost electrons closer to the nucleus and the size of the atom
decreases.
As you move down a group, like group 1, the size of the atom increases.
Atomic Radii (pm) of Group 1 Elements
Li
Na
K
Rb
Cs
152
112
85
77
75
As you move down a group, a new energy level is added for every atom in that group.
With each added energy level, the electrons are situated further from the nucleus,
reducing the force of the nucleus on the outermost electrons. The addition of a new
energy level increases the amount of shielding on the outermost electrons, resulting in a
lower effective nuclear force. The lower force on each outermost electron allows the
electrons to move further from the nucleus, producing a larger atomic radius.
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For negative ions, one or more electrons are added to the atom. The more electrons that
are added to an atom, the lower the force of the nucleus on each atom and the ion
becomes larger. In the diagram above, both the chloride ion and the fluoride ion are
larger than the chlorine and fluorine atoms.
For positive ions, one or more electrons are removed from the atom. As a result, there is
more than one proton for every electron in the ion. The force on each electron is
increased and the ion is much smaller than the atom.
Periodic Trends in Ionization Energy
From the periodic trends graphing activity we can identify two major trends.
Trend 1. As you move down a group the ionization energy decreases.
As mentioned while discussing atomic radii, as you move down a group an energy level
is added for each period. The added energy level means more electrons are between the
outer electrons and the nucleus, increasing the shielding effect. The added energy levels
also increase the size of the atoms.
The increased distance of the outer electrons from the nucleus and the increased
shielding, decreases nuclear force. The lower the force on the outer electrons, the easier
they are to remove resulting in a lower IE.
Trend 2. As you move across a period from left to right, the ionization energy
increases.
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As you move across a period the nuclear charge increases and shielding remains constant.
This causes the atomic radii to decrease. The decreased radius means the distance
between the nucleus and the outer electrons is lower resulting in a greater force on the
outer electrons. The greater force on the electrons means it is more difficult to remove the
outer electrons and a higher IE.
Ionization energy (IE) is defined as the amount of energy needed to remove an electron
from a gaseous atom.
A(g) + energy → A+(g) + 1e−
The amount of energy required depends upon the force on the electron. The greater the
force on the electron, the more energy needed to remove the electron, hence the greater
the value of the ionization energy.
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Trend 3: Ionization energy tends to increase with each successive electron removed.
For example, sulfur’s ionization energy increases from 1005 kJ/mol for IE1 to 2260
kJ/mol for IE2 to 3375 kJ/mol for IE3, etc.
Why does this occur?
Each electron removed increases the proton to electron ratio. The nuclear force on each
electron will increase as each electron is removed because the number of protons is larger
than the number of electrons. Since nuclear force increases, it is more difficult to remove
an electron and the ionization energy increases.
Trend 4: Removing an electron from a noble gas electron configuration causes the
ionization energy to increase considerably. The bold lines in the tables represent
noble gas structures.
The sodium atom has one valence electron. Its full
electron configuration is 1s22s22p63s1. Removing one
electron results in a configuration of 1s22s22p6, the
same configuration as neon. The ionization energy
increases almost ten times removing the first electron to
removing the second, which is removed from the neon
configuration. Look at the bar graph below to see the
significant change. After the second electron is
removed, the remaining changes are similar.
We see a similar pattern for magnesium. After two
electrons are removed, the magnesium atom’s configuration goes from 1s22s22p63s2 to
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the same electron configuration as neon. The value of
ionization energy jumps considerably from the second to
third ionization energy.
The results for aluminum and silicon are similar as well.
See graphs below:
Aluminum has 3 valence electrons and silicon has 4. Notice the large increase in
ionization energy after all valence electrons are removed and an electron is removed from
the noble gas structure.
What is the explanation for this trend?
Once noble gas structure is achieved, the valence electrons are removed − the electrons in
the outermost orbitals. When the outermost electrons are removed, the next electron is
removed from a lower energy level. The valence electrons for sodium, magnesium,
aluminum and silicon are all in the third energy level. Once noble gas structure is
achieved, the next electron to be removed is from the second energy level. The electrons
in the second energy level are closer to the nucleus and experience less shielding effect
than those in the third energy level, resulting in a much higher ionization energy. After
this, the ionization energies increase at a regular rate.
A Summary of the Periodic Trends
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Assignment 10
1. Briefly explain why barium has a lower first ionization energy than calcium.
2. Given the following elements and their electron configurations, which element
will have the lowest first ionization energy? Why?
element A 1s22s22p63s23p64s2
element B 1s22s22p63s23p5
element C 1s22s22p63s23p64s1
element D 1s22s22p63s23p6
3. The first four ionization energies (IE) for the element aluminum are as follows:
IE1 = 577 kJ/mol
IE2 = 1 817 kJ/mol
IE3 = 2 745 kJ/mol
IE4 = 11 580 kJ/mol
How many valence electrons does aluminum have.
4. An atom has the following successive ionization energy levels.
IE1 = 737 kJ/mol
IE2 = 1 450 kJ/mol
IE3 = 7 731 kJ/mol
IE4 = 10 545 kJ/mol
IE5 = 13 627 kJ/mol
How many valence electrons does this element have? Explain.
5. Which of these elements would have the highest value for the second ionization
energy? Why?
a. K
b. Si
c. Ar
d. Br
6. Which of the following has the largest atomic radius and which has the smallest?
Explain.
a. nitrogen
b. antimony
c. arsenic
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7. For each of the following properties
i) atomic radius
ii) first ionization energy
iii) ionic radius
indicate which has the larger value fluorine or bromine.
8. Arrange the following from largest to smallest. Explain the order.
Ne, Mg2+, F–, Na+, O2–
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