Moving from the Bohr Model to Quantum Mechanics
Figure 1:
Bohr’s model for hydrogen was based partly on
“___________” mechanics (the laws of motion
that apply to objects on the earth and in the solar
system). Bohr used the laws of classical
mechanics to calculate the possible orbits and
speed of the electron in the hydrogen atom. .
The major difficulty with the Bohr (orbital)
Model was that it could not predict the orbits
(energy levels) of atoms with multiple electrons.
Bohr’s Model needed to be modified.
“Quantum” mechanics are now used to explain that these orbits are __________ (of a
specific energy) and that the electron jumped (it didn’t change smoothly and continuously
as classical mechanic predicts) between orbits if that amount of energy is present. In the
modern quantum mechanical (____________) model, no orbit or path of the electron is
proposed. In this model, only the______________of finding an electron in a certain
region of the atom is given.
The region of highest probability for each energy level, is a plot of points. The average
position of these points is a spherical “shell” (cloud) centred around the nucleus. This
shell or energy level is the average of the points indicating the probability that the
electron will be there at any given time, not a path of movement of the electron.
Note:
Figure 2 Neils Bohr
Energy Sublevels
According to quantum mechanics, every atom has principal energy
levels and every principal energy level has one or more sublevels
within it. Each sublevel has a slightly different energy. Bohr’s model
did not include these sublevels.
The number of sublevels in any principal level is the same as “n”. That is, the first
principal energy level (n=1) has 1 sublevel (“s” sublevel). The second principal energy
level (n=2) has 2 sublevels (“s” and “p” sublevels). The third level (n=3) has 3 sublevels
(“s”, “p” and “d” sublevels) and so on.
The lowest sublevel in any principal level is the “s” sublevel
Ex:
1s (1st principal energy level) or______
2s (2nd “
“
“ ) or ______
3s (3rd “
“
“ ) or _______
The next higher sublevel is the “p” sublevel. Remember there is no “p” sublevel
when n=1. For n=2 and higher (n=3, n=4 etc) “p” sublevels exist in addition to the “s”
sublevels
Figure 3: First Four Principal Energy Levels:
Ex:
1s
2s
3s
4s
2p
3p
4p
n=1
n=2
n=3
n=4
n=4
When n=3 the 3rd sublevel appears, there
is a 3rd sublevel called the “d” sublevel.
The “d” sublevel only appears in the
energy levels of n=3 or higher
n=3
n=2
Ex.
1s
2s
3s
4s
2p
3p
4p
n=1
3d
4d
4f
When n=4 there is a fourth sublevel, “f”. Additional sublevels, (g,h, etc) are available in
higher energy levels (but won’t be discussed)
Note:
Orbitals
Calculations from quantum mechanics have given use to better probability plots of
energy levels. These plots have not only identified energy sublevels, but regions within
the sublevels where electrons are likely to be found, called Orbitals (NOT ORBITS). The
total number of orbitals available in an atom is predicted by:
____________
The laws of quantum mechanics allow for at most 2 electrons in each orbital. So the
three “p” orbitals can have 3x2 = 6 electrons occupying them, likewise, “d” orbitals can
have 5x2=10 electrons.
The maximum number or electrons can be predicted by:
_________
Ex.
n=1 can have a total of _______________
n=1 can have a total of _______________
n=1 can have a total of _______________
n=1 can have a total of _______________
Review:
Atoms contain energy levels, which contain sublevels, which themselves contain orbitals.
Energy Levels =
Orbitals =
Total electrons =
_______________
_____________
_______________
Shapes of Orbitals
Figure 3:
Plots of the charge clouds reveal that
different shapes exist for each type of
orbital. All “s” orbitals, regardless of
which principal energy level they are
part of (1s, 2s, 3s...), are ___________
in shape. The differences between them
are that those in higher energy levels
have ________________ than the lower
ones.
In a “p” orbital an electron has a
“__________l” or “_________” shape.
Evidence shows that there are three p
orbitals each having the same amount of
energy and the same shape and size
(degenerate). How they differ is in how
the orbital (dumbbell shape) is
orientated – along the x-axis, y-axis or z-axis. These orbitals are given the names, px, py,
pz respectively.
The shapes of the “d” and “f” orbitals are more complicated and won’t be considered in
chemistry 20.
Questions:
1. What was the major problem with Bohr’s atomic model?
2. What information about the location of an electron is given by its principal
quantum number?
3. What symbol represents the principal quantum number?
4. In the quantum mechanical model of the atom, what “new parts” were added to
the principal energy levels of Bohr’s model?
5.
a. How many sublevels can be present in:
i. Energy level 3
ii. Energy level 6
b. How many orbitals are present in i & ii above?
c. What is the maximum number of electrons that can be present in i & ii
above?
6. What is an orbital? How many electrons can occupy an orbital?
Taking care of those electrons - Energy Level Diagrams
To help us understand what energy levels, sublevels and orbitals electrons occupy in
atoms that have more than one electron (He and on) we use the energy level diagram:
But first, we require some tools to use it.
Note:
1. The energy of the electrons increase as we move add electrons (we move up the
chart)
2. The energy levels (what n is equal to)
3. Te energy sublevels (s, p, d, f…)
4. The orbitals as o
When filling the energy level diagrams, we use the following notations:
1 Electron
2 Electrons
Remember an orbital can have TWO electrons MAX!!
How to fill the chart:
3 Principles to follow:
1. Fill orbitals with the lowest energy first, starting at n=1 (Ground state), then move
up
2. Maximum 2 electrons can occupy an orbital
3. When electrons fill orbitals of the same energy and type (i.e. 2px, 2py, 2pz) each
orbital must be occupied by a single electron before they start “doubling up”
We will look at four ways to write electron configurations, the first two are:
1. Orbital Notation:
2. Electron Configuration
Examples:
The next notation is:
3. Shortcuts:
When writing out electron configurations and orbital notations for elements with large
numbers of electrons, we can use a shortcut!
Steps:
1. Find the last nobel gas that the element has gone past
2. Write that element in square brackets
3. Start the orbital notation/electron configuration with the next principal level
and continue as normal
Ex:
A quick review:
Element
Li
C
Ne
K
Si
# Electrons
Electron
Arrangement/Configuration
Number of electrons
per level
Orbital Notation and Electron Configuration of Ions
Recall: An ion is an atom that has a positive of negative charge and so has either lost or
gained electrons.
Metals (elements on the left side of the periodic table) tend to lose electrons.
Nonmentals (elements on the right side of the periodic table) tend to gain electrons.
This is because elements like to have the same number of electrons (become
isoelectronic) as the closest nobel gas. They remain as the same element (the number of
protons doesn’t change, only the number of electrons). As we will see later in this
course, this is an important concept when it comes to bonding and forming compounds.
Becoming isoelectronic with nobel gases is energetically favourable in elements, these
elements are very stable.
We can show orbital notation and electronic configuration for ions as well.
Ex:
The forth notation is:
4. Lewis Diagrams aka Electron dot notation
In Lewis diagrams we look only at the valence electrons.
Valence electrons are the outside shell s and p orbitals only, these are the most important
electrons and used in bonding!
Group #
I
II
III
IV
V
VI
VII
VII
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
Element
Note: Valence electrons are the
same as the group number!!
Time for Practice
Complete the following:
Element Electrons Electron Configuration
# of
Electron Dot (Lewis
Valence Formula)
Electrons
Be
N
Cl
K
Ne
He
I
Want to learn more about quantum mechanics? Check out this website:
http://www.howe.k12.ok.us/~jimaskew/cquantum.htm