Notes: Chapter 11
- The source of all light is
the atom
- Light is generated when
the electrons in an atom are excited to a higher energy level and then
relax. When the atoms relax, they emit photons as the electrons drop to
lower energy levels:
·
Therefore, the more we understand about
light, the more we can understand about the arrangement of electrons in the
atom.
·
This last point is important because the
electron is the most important subatomic particle in chemistry. The arrangement
of an atom’s electrons, its electron configuration, is what determines an
element’s chemical properties.
·
We will first discuss light as a wave
phenomenon. Later, we will describe light in terms of being a particle.
·
Light is a transverse wave. A transverse
wave is one in which the medium travels at right angles to the direction of the
wave energy that is traveling through it.
·
A wave is simply some disturbance in a
medium
·
A good example of a transverse wave is
an ocean wave. The medium in this case would be the ocean.
- You should understand
these terms used to describe waves:
- Wavelength
- Symbol: l
(pronounced and written in English “lambda”)
- Measured in meters
- Equals the distance between
two crests (top of wave) or troughs (the bottom of the wave)
- Frequency
- Symbol: n
(pronounced and written in English “nu”)
- Measured in Hertz (Hz),
or s-1, or 1/s. The latter two are called “reciprocal seconds”
- Equals the number of wave
cycles that pass a given point per given amount of time.
- Speed
- Symbol is c in the case
of light, which is the only case we are concerned with
- Is typically measured in
meters per second (m/s)
- The speed of light is
ALWAYS equal to 300,000,000 m/s in a vacuum.
- c =3.00 x 108 m/s
- Because this is a speed
that never changes, the speed of light is called a constant.
- Speed = frequency X
wavelength (for any wave)
- For light,
c = ln
- From this equation, and
from the demonstrations in class, we can see that
- Frequency and wavelength
are inversely related
- Frequency and energy are directly
related
- We will see that this is
true of the different types of electromagnetic radiation (light) as well.
- Three examples of
speed/wavelength/frequency problems
- What is the frequency of
light that has a wavelength of 250 m?
- Three examples of
speed/wavelength/frequency problems
- What is the frequency of
light that has a wavelength of 250 m?
- All light is called
electromagnetic energy. These terms (light and E.R.) are synonymous for
our purposes.
Electromagnetic Spectrum
The electromagnetic
spectrum includes radio waves, microwaves, infrared light, visible light,
ultraviolet light, x rays, and gamma rays. Visible light, which makes up only a
tiny fraction of the electromagnetic spectrum, is the only electromagnetic
radiation that humans can perceive with their eyes.
Microsoft
® Encarta ® 2006. © 1993-2005 Microsoft Corporation. All rights reserved.
- Only a small portion
of the electromagnetic spectrum is visible light. That is, of all of the
light that exists, only part of it can be seen by humans.
- Visible light spectrum (p.
299)
The
visible light spectrum. From left to right, in order of decreasing energy,
decreasing frequency, and increasing wavelength.
- The electromagnetic
spectrum, in order of increasing frequency, increasing energy, and
decreasing wavelength:
- Radio waves
- Microwaves
- Infrared (IR) radiation
- Visible light
- Ultraviolet light
- X-rays
- Gamma rays
- The visible light
spectrum, in order of increasing energy, increasing frequency, and
decreasing wavelength:
- Red
- Orange
- Yellow
- Green
- Blue
- Indigo
- Violet
- Practice problems:
- What is the wavelength of
light that has a frequency of 1.2 x 1015 Hz?
- What is the frequency of
light that has a wavelength of 650 nm (109 nanometers = 1m)?
According to the diagram on p. 299, is this light visible? How did you
know?
- What is the wavelength of light that has a
frequency of 1.2 x 1015 Hz?
- What is the frequency of
light that has a wavelength of 650 nm? (109 nanometers = 1m)
According to the diagram on p. 299, is this light visible? How did you
know?
- Light can also be thought
of as a particle.
- The quantum theory was
first suggested by Max Planck in 1900. He suggested that energy can only
be absorbed or emitted in small packets called quanta (singular: quantum).
- Money can only be spent in
tiny amounts called cents. Every bit of money that you have ever spent or
earned in the U.S. has been some multiple of 1 cent.
- Energy is either
absorbed or emitted (gained or lost) in multiples of h, Planck’s
constant. h = 6.626 x 10-34 Js
- A quantum of light energy
is called a photon. A photon is light a tiny packet of
light energy.
- To calculate the energy of
a photon of light, we use the formula
E = hn
- E = energy of a photon
- h = Planck’s constant = 6.626 x 10-34 Js
- n = frequency of light
·
Practice problems:
o
Calculate the energy of a photon
that has a frequency of 1.20 X 1015 Hz.
o
Calculate the frequency of a
photon with an energy of 3.90 X 10-19 J.
·
Do problems 1-6, ch. 11, in
class.
·
Do worksheet in class
General Chemistry
Mr. MacGillivray
Worksheet:
Light and Quantum Theory
1.
Arrange the seven types of
electromagnetic radiation that we discussed in class in order of DECREASING
energy:
- _______ (highest E)
- _______
- _______
- _______
- _______
- _______
- _______ (lowest E)
- In the list above, use
words and arrows to indicate how the wavelength and frequency are
changing.
- Repeat #1 and #2 with the
colors of the visible spectrum.
- _______ (highest E)
- _______
- _______
- _______
- _______
- _______
- _______ (lowest E)
- “If the wavelength of
light is very short, then the energy is very __________ and the frequency
is very _____________.”
- “If the wavelength of
light is very long, then the energy is very __________ and the frequency
is very _____________.”
- Wavelength and frequency
are __________ly related. Energy and frequency are ___________ly related.
- Energy is measured in
these units: ________.
- Wavelength is measured in
these units: ________.
- Frequency is measured in
these units: ________, also written as ______ or ______.
- Convert the following
wavelengths to nm:
- l = 513 m
- l = 8.03 x 10-6 m
- Convert the following
wavelengths to m:
- l = 755 nm
- l = 0.272 nm
- Find the energy of a
photon of light with a frequency of 5.22 x 1021 1/s.
- Find the energy of a
photon of light with a wavelength of 425 nm.
- Find the wavelength of
light with a frequency of 5.28 x 1015 s-1.
- Using p. 299, answer
these questions:
- Is the light in question
#12 visible?
- How did you know?
- Is it too high in energy
or too low in energy to be seen?
- What type of light is it
(what region of the electromagnetic spectrum)?
·
Begin electron configurations
·
Chemistry: it’s the study of
matter and the changes it undergoes
·
These changes are due to atoms
combining, separating, or rearranging.
·
These changes occur when
electrons are given, taken, or shared between atoms
·
The more we know about the
electronic structure of the atoms,
·
There are 3 levels of
organization to the electron cloud.
o
Energy level (“n”)
o
Energy sublevel (“l”)
o
Orbital (“ml”). An
orbital is a region in space where electrons are likely to be found. It can
contain at most two electrons.
·
Think of the atom as a hotel with
different floors (energy levels), rooms on each floor (energy sublevels), and
closet/bathroom/bedroom/etc. (orbital) within each hotel room.
·
Each energy level has a different
number of sublevels
o
The first energy level has 1
sublevel, “1s”
o
The second has energy level has 2
sublevels, 2s and 2p
o
The third energy level has three
energy levels, 3s, 3p, and 3d
o
The fourth energy level has four
energy levels, 4s, 4p, 4d, and 4f
o
The fifth has five, and so on.
o
How many sublevels does the 7th
energy level have?
·
Sublevels are in turn divided up
into orbitals, which are regions in space where electrons are likely to
be found. It can contain at most two electrons.
o
Any s sublevel has 1 orbital
§
1s has one orbital
§
2s has one orbital
§
3s has one orbital
§
10s has one orbital, etc.
§
An s sublevel is spherical in
shape
§
Any s sublevel has one orbital,
which can hold 2 electrons at most. Therefore, any s sublevel can hold at most
2 electrons.
o
Any p sublevel has 3 orbitals.
§
2p, 3p, 4p, etc. all have three
orbitals each.
§
p orbitals look like dumbbells,
and are right angles to one another in 3-dimensional space. They can be thought
of as being oriented on three axes (x, y, and z).
§
They are called the px,
py, and pz orbitals.
§
Because any orbital can contain
at most 2 electrons, and because any p sublevel has three orbitals, any p
sublevel can contain at most 6 electrons.
o
Any d sublevel has 5 orbitals.
§
The 3d sublevel has 5 orbitals.
So does the 4d and so on.
§
That means that any d sublevel
can hold a maximum of 10 electrons.
§
The shapes and orientations of
the d orbitals are complex.
o
Any f sublevel has 7 orbitals.
§
The 4d sublevel has 7 orbitals.
So does the 5d and so on.
§
That means that any f sublevel
can hold a maximum of 14 electrons.
§
The shapes and orientations of the
f orbitals are complex.
Sublevel type
|
# of orbitals
|
Max # of electrons in sublevel
|
s
|
1
|
2
|
p
|
3
|
6
|
d
|
5
|
10
|
f
|
7
|
14
|
g, but not important to us
|
9, but not important to us
|
18, for the sake of demonstration
|
Who cares
|
11, but who cares
|
22, but don’t lose sleep
|
Energy level
(=n)
|
Sublevels present |
# of orbitals
(=n2)
|
Max # of electrons
(=2n2)
|
1
|
1 (1s) |
1
|
2
|
2
|
2 (2s, 2p)
|
1+3 = 4 |
2 + 6 = 8
|
3
|
3 (3s, 3p, 3d)
|
1+3+5 = 9 |
2 + 6 + 10 = 18
|
4
|
4 (4s, 4p, 4d, 4f)
|
1+3+5+7 = 16 |
2 + 6 + 10 + 14 = 32
|
5
|
5 (5s, 5p, 5d, 5f, 5g)
|
1+3+5+7+9= 25 |
2 + 6 + 10 + 14 + 18 = 50
|
6
|
6, etc.
|
1+3+5+7+9+11= 36 |
2 + 6 + 10 + 14 + 18 +22= 72
|
7
|
7, etc.
|
1+3+5+7+9+11+13= 49 |
2 + 6 + 10 + 14 + 18 +22 + 26 = 98
|
- Don’t memorize any of
this info, except possibly the n2
and 2n2 shortcuts, because this info is all
summarized on the periodic table!
- Demonstration: determining
the available sublevels in the first 4 energy levels by looking at the
periodic table.
- Demonstration: determining
the number of electrons per each orbital type by looking at the periodic
table.
- Demonstration: determining
the number of orbitals per each sub level type by looking at the periodic
table.
- Time for a definition of
“orbital”
- An orbital is a region in
space where an electron is likely to be found
- Orbitals don’t need to
have any electrons, but can contain at most 2 electrons. That is, an
orbital can have 0, 1, or 2 electrons
- To repeat and
emphasize, ANY orbital - an s orbital, a p orbital, a d orbital
,etc. - can contain at most 2 electrons.
- If an orbital has two
electrons, the spins must be opposite
- “Spin” is a property of
electrons
- In reality, the
electrons do not spin, but it is convenient to describe this property as
spin.
- The values of ms
can be either + 1/2 or -1/2.
- We will call these “spin
up” () and “spin down” (¯).
- Now let’s look at a
simplified energy diagram of the atomic orbitals, arranging them from the
first to be filled to the last to be filled.
- Note: the orbitals within
a given energy level all have the same energy. However, for reasons we
will not cover in this course, the orbitals are filled in the order shown
below.
- Note: the 4s and 3d
sublevels “overlap”; that is, the 4th energy level (4s) starts
filling before the 3rd energy level is full of electrons.
- Note: no atom is actually
formed in this way. However, in order to understand why the electrons are
arranged as they are in an atom, it is often helpful to pretend that the
electrons are added to the electron cloud one at a time.
- This diagram is called an
Aufbau diagram.
- Let’s start with the
easiest atom: hydrogen. Hydrogen is atomic #1, thus a neutral atom of H
has one electron.
- Because H has one electron
in the 1s sublevel, its electron configuration is 1s1. The
Aufbau diagram for an element is a picture (shown above). The electron
configuration is a string of numbers and letters.
- The rules for filling the
Aufbau diagram are as follows:
- The Aufbau principle:
- Electrons enter orbitals
of lowest energy first.
- Pauli’s exclusion
principle:
- Orbitals may contain at
most two electrons, and if there are two electrons, they must have
opposite spins. (“No two electrons can have the same four quantum
numbers.”)
- Hund’s Rule
- When electrons enter
orbitals of equal energy, orbitals must first be singly occupied with
spins parallel before they are doubly occupied.
- Now let’s look at the
Aufbau diagram and electron configuration for helium, He, Z=2.
- Here is the diagram of
lithium:
- Here is the diagram for
Be, Z=4:
- Here is the diagram for B
(boron), Z=5:
- Now we need to invoke
Hund’s rule to figure out where the next electron will go. Here is the
electron configuration and Aufbau diagram for carbon (Z=6):
- Here is vanadium (V,
Z=23).
- How to use the periodic
table to determine the electron configuration: the order of orbital
filling is summarized by the guide shown below:
- See if you can determine
the electron configuration of the elements from the above guide. There are
exceptions to this rule, but the exceptions are unrelated to the order of
filling of the orbitals.
- Noble gas abbreviation:
- The electron
configuration for an element can be abbreviated by putting the symbol of
the last noble gas before that element occurs in the periodic table in
brackets, and then adding the “remainder” of the configuration notation
after that.
- Use the “Last Noble Gas
That You Passed”
- Sodium is 1s22s22p63s1,
or simply [Ne]3s1.
- Calcium is 1s22s22p63s23p64s2,
or simply [Ar]4s2.
- Chlorine is 1s22s22p63s23p5,
or simply [Ne] 3s23p5.
- Radium is 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s2,
or simply [Rn]7s2.
- Valence Electrons:
- The elctrons that are
involved in bonding.
- Generally, they are the s
and p electrons in the highest occupied energy level.
- Exceptional
configurations:
- Some electron
configurations violate the Aufbau principle
- Many inner transition
metals (a.k.a. rare earth elements, a.k.a. f-block elements) do this.
- You will be responsible
for examples found in the transition elements (a.k.a., the “d” block).
- Common examples that you
should know: s2d4 and s2d9
configurations do not occur.
- They instead occur as
s2d4
and s2d9 s1d5 and s1d10.
- An s electron is promoted
to the nearby d orbital whenever the d orbital is one electron away from
being filled (10 electrons) or half-filled (5 electrons).
- Examples:
- Copper (Cu) does not end
in “4s23d9”. Instead, [Ar]4s13d10.
- Silver (Ag) does not end
in “5s24d9”. Instead, [Kr]5s14d10.
- Chromium (Cr) does not
end in “4s23d4”. Instead, [Ar]4s13d5.
- Molybdenum (Mo) does not
end in “5s24d4”. Instead, [Kr]5s14d5.
- Periodic table: table
showing all of the elements in increasing atomic number.
- Development of the
periodic table
- Our current periodic
table is not the first periodic table.
o
What does the word “periodic” mean? “periodical”? “periodicity”?
Dmitiri Mendeleev, 1834-1907
o
The first periodic table was devised by a Russian chemist, Dmitri
Mendeleev. Mendeleev arranged his periodic table by increasing atomic mass.
- Mendeleev was writing a
chemistry textbook, and wanted to organize the information that was known
about all of the known chemical elements. At the time, scientists had
isolated and identified about 70 elements.
- The story goes that he
wrote the symbol of each element on a card, and then wrote the chemical
and physical data about each chemical on the back of each card.
- As he arranged the cards
in order of increasing atomic mass, he started to notice certain
recurring patterns in the elements’ properties.
- These properties
re-occurred at regular intervals. For instance (this is a
simplification), when he arranged the elements in order of increasing
atomic mass, he noticed that every eight elements, and then every
eighteen elements, there was an explosive metal (Na, K, Rb, etc.)
- When you lined up a set
of such elements into a “family” according to shared characteristics,
other sets of elements with shared characteristics also fell into line.
- Mendeleev was able to see
this pattern and use these patterns to predict the existence and
properties of as-yet undiscovered elements.
- Mendeleev was able to win
respect for his bold predictions when these elements were discovered and
studied. Examples: eka-silicon (germanium), eka-aluminum (gallium), and
eka-boron (scandium).
- He was able to predict
the properties of these elements by interpolating the patterns of
the surrounding elements on his periodic table. He also stated that
several elements had incorrectly stated atomic masses, since they did not
fit into the correct spot on his table. He was mostly right on this
point, but not always (see below).
- Fun fact: he wrote his
doctoral thesis, essentially, about vodka ("On Composing Alcohol
with Water").
- There are problems with
arranging the periodic table by increasing atomic mass. Can you find examples on our periodic table
where the elements are not in order of lightest to heaviest?
- It is obvious by looking
at iodine’s chemical properties that it belong in its group, Group 7A
(a.k.a. Group 17). For instance, it reacts with Na in a 1:1 ratio, making
NaI (just like NaCl, NaBr, NaF, etc.).
- However, if you put iodine
in the correct group, then it is out of order with respect to mass.
If you put Iodine in the “correct” spot according to its mass, then it
winds up in the wrong family.
- Henry Moseley helped to
develop the modern periodic table.
- Moseley was able to
determine the “atomic charge” of an element, what we would today call the
number of protons. This atomic charge is now called the atomic number.
- Henry Moseley is
responsible for today’s arrangement of the periodic table. Today the
periodic table is arranged by increasing atomic number.
- If the elements are
arranged by increasing atomic #, then there are no “mistakes” in how the
elements are organized into the proper groups.
- The periodic table
o
Groups, a.k.a. families – these are the vertical columns of
elements. Elements in the same families have similar chemical and physical
properties.
o
Periods – these are the horizontal rows of the periodic table.
Elements in the same period have very different properties, for the most part.
- The periodic table has
three main sections: The metals (left side of staircase), the nonmetals
(right side of staircase), and the metalloids (touching the staircase with
a whole side of their square).
- Exceptions: Hydrogen is a
nonmetal. Aluminum is a metal.
- Metals (p.103)
o
Good conductors of heat & electricity
o
Malleable
o
Ductile
o
Lustrous
o
Poor conductors (insulators)
o
Brittle
o
Dull
o
Properties that are somewhere in between those of metals and nonmetals
o
Example: silicon (Si). It is shiny but brittle. It is a semiconductor,
which is a material that is neither a good conductor nor good insulator. It can
control the flow of electrons (electricity).
·
Numbering system of the groups: we will use the easy, new system:
1 through 18.
·
Group names that you should know:
o
Alkali metals (group1)
o
Alkaline earth metals (group 2)
o
Chalcogens (group 16)
o
Halogens (group 17)
o
Noble gases, a.k.a. inert gases
(group 18)
o
Transition metals (d-block
elements)
o
Inner transition metals (f-block
elements)
- Elements found alone in
nature (uncombined):
o
Noble gases
§
Unreactive elements, always found
alone in nature
o
Gold, Silver, Platinum
§
Relatively unreactive metals,
often found chemically pure (or nearly so) in nature
o
Most other elements are found
combined with other elements as compounds.
o
Some elements combine with
themselves to form diatomic molecules:
Periodic
Trends
- The Periodic Law: there is a periodic repetition of the physical
and chemical properties of the elements when they are arranged in order of
increasing atomic number.
- There are a number of
atomic properties that vary in a regular way across the periods and down
the groups.
- Atomic radius
o
Atomic radius is the distance
from the center of the atom to the outer edge of the atom
o
It is a way of measuring the size
(in distance) of atoms
o
The periodic and group trends are
§
L to R, the atoms decrease in
size
§
T to B, the atoms increase in
size
o
Reasons for the trend:
§
As you go down a group, the
number of “shells” or energy levels of electrons between the positive nucleus
and the negative nucleus increases.
§
This decreases the attraction
between the nucleus and the outermost electrons.
§
The effect is that the electrons
are less tightly held, and therefore are able to get further away. This is
called the shielding effect.
§
As the atomic number across a
period increases (L to R), the number of electrons (in the outermost energy
level) and the number of protons (in the nucleus) both increase.
§
Without the shielding effect of
adding more “layers” of electrons between the nucleus and these additional
electrons, the net effect is to draw the outermost electrons closer to the
nucleus.
§
“more protons + more electrons =
stronger attraction”: . . . within a given period.
o
Ionization energy is the energy
required to remove an electron from an atom
o
Metals give up electrons more
easily than nonmetals; metals will attain a full octet by losing relatively few
electrons.
o
Nonmetals will attain a full
octet more easily by gaining electrons than by losing them.
o
Thus, the trends are
§
L to R ionization energy
increases
§
T to B ionization energy
decreases
o
Electronegativity is the tendency
of an atom to attract electrons toward itself when bonded to another atom
o
Metals have low
electronegativities
o
Nonmetals have high
electronegativities
o
Thus, the trends are
§
L to R electronegativity
increases
§
T to B electronegativity
decreases