# Absorption Spectra _amp; Quantum Mechanics

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Absorption
Spectra &
Quantum
Mechanics
Light is both a wave and a
particle???
Characteristics of Waves

• Wavelength
– The distance between consecutive peaks or troughs

• Frequency
– How many waves pass a point per second
• Speed
– How fast a wave moves through space
Electromagnetic Spectrum

include:
gamma rays
X-rays
Ultraviolet
Visible
Infrared
microwaves
Electromagnetic Spectrum
Electromagnetic Spectrum
• Visible light falls in the 400 to
700 nm range
• In the order of decreasing
wavelength
– Microwave: 1 mm
– Visible light: 500 nm
– X-rays: 1 nm
– Gamma rays: 10-3 nm
Atomic Spectra

• When visible light passes through a
prism, its components separate into a
spectrum.
• White light, such as sun light or light
from a regular light bulb, gives a
continuous spectrum:
Wavelengths of Visible Light
Which Colour Has More
Energy?
• If we look at light given off by electrified
gases of elements and pass it through a
prism then a unique and characteristic line
spectra is seen for each.
Spectral Lines
• Bright spectrum lines can be seen when a chemical substance is
heated and vaporized (Kirchhoff, ~1850)
• There are three different kinds of spectrum: continuous spectrum,
emission-line spectrum, and absorption line spectrum
Emission-Line Spectra
The Bohr Atom
• In 1913 Niels Bohr proposed that
the energies of electrons inside
atoms are quantized. By this he
meant that the electrons could
have only particular allowed
energies.

• This goes against the notions of
classical physics as well as our
own experiences.
• Niels Bohr, a student of Rutherford, studied
the line spectra of the hydrogen atom to try
to understand the electrons in the nuclear
model of the atom.
• Bohr came up with a planetary model, in
which electrons orbit the nucleus in circular
pathways.
• His model was based on the idea that
electrons and their energies are quantized:
they can have only certain values.
The Bohr Atom

Small but discreet packets of
energy called quanta are
absorbed when an electron
jumps to a higher orbital and
emitted when an electron falls to
a lower orbital.
Electrons in an atom have their
energies restricted to certain
specific energy levels that
increase in energy as they
increase in distance from the
nucleus.
When you climb a stairs you
must be on a step to climb. You
can’t climb from between steps.
The higher you climb the greater
earth.
Likewise electrons cannot be
between energy levels.
• In order to examine the
difference between the
energies      of     various
energy levels the lowest-
energy shell (n = 1) is set
at 1.0 electron volt.
• We may determine the
amount       of       energy
between energy levels by
subtracting the values
from level to level.
• The diagram on the left indicates a hydrogen atom
in the “ground state,” which is to say that the
arrangement of electrons in the atom is of the
lowest total energy.
• The diagram on the right indicates a hydrogen
atom in an “excited state” - one of several
possible excited states.
• The ground state of an atom is the condition
when all the electrons of an atom occupy the
lowest possible energy levels (those closest to
the nucleus).
• If the atom absorbs energy, an electron can be
raised to an excited state. The excited state is
the condition at which at least one electron in an
atom is at an energy level above the ground
state.
• The atom must absorb the exact amount of
energy required to raise the electron to the
excited state. This may be done by heating the
atom or by passing an electric current through a
gas.
• This explains why heated metals or gases glow
a certain colour.
• An electron in an
excited state is
unstable. It falls back
to the ground state,
sometimes in one
quantum jump,
sometimes in two or
more. In doing this, it
releases a photon
(packet) of light with a
frequency proportional
to the energy
difference between
the two levels.
Bohr’s Atomic Model for Hydrogen
• The strongest hydrogen
spectral line from the
Sun, Hα line at 656 nm, is
caused by electron-
transition between n=3
orbit and n=1orbit

• Lyman series lines:
between n=1 orbit and
higher orbits (n=2, n=3,
n=4,…)

• Balmer series lines:
between n-2 orbit and
higher orbits (n=3, 4,
5,…)
Bohr’s theory depended on the
following assumptions:
• An electron can travel indefinitely within an
energy level without losing energy.
• The greater the distance between the nucleus
and the energy level, the greater the energy
required for an electron to travel in that energy
level.
• An electron cannot exist between orbits, but can
move to a higher unfilled orbit if it absorbs a
specific amount of energy, and to a lower
unfilled orbit if it loses energy.
• When an electron drops back to its original
energy level, it is said to be in its ground state.
Successes of the Bohr Model
•   Explained the stability of the atom.
•   Explained the atomic line spectrum of the
hydrogen atom.
•   Introduced the concept of stationary
energy levels.
•   Introduced the concept of quantized
energy.
•   Introduced a model of the atom that
could be easily visualized.
Failures of Bohr Model
• Did not explain the following:
– Line spectra for many electron atoms
– Electron configurations of many electron
atoms
– The difference in energies of electrons
occupying the same energy level.
– The shapes and characteristics of molecules.
spdf sublevels
• Bohr’s model of the hydrogen atom
predicted the spectrum of hydrogen
but failed for other elements because it
required further sophistication. It was
refined to include subshells.
• The principal energy levels are further
divided into sublevels, labeled s, p, d,
and f
• The number of sublevels equals the
number of the principal energy level.
Orbitals
• Erwin Schrodinger
created a model
describing electrons
as waves - known as
wave mechanics
• won the Nobel Prize
in Physics in 1933
Orbitals
• Although we cannot locate an electron
precisely within an atom we can describe a
region in space around the nucleus where
there is a high probability of finding a given
electron.
• This region is called an orbital.
• n is the Principal Quantum Number. It is a
positive whole number that specifies the
energy level of an atomic orbital and its
relative size. A higher n indicates an orbital
with higher energy and larger size.
• The greatest number of electrons that are
possible in any energy level is 2n2

n= 1   #electrons=2
n= 2   #electrons=8
n= 3   #electrons=18
n= 4   #electrons=32
Orbital Shape Quantum Number
• l= orbital shape quantum number. It
indicates the shape of the sub-level.
• Has a maximum value of l=(n-1)
• So if n=1 l=0
•       n=2 l=1 or 0
•       n=3 l=2 or 1 or 0
•       n=4 l=3 or 2 or 1 or 0
Sub-levels
•   When l=0   s-shape
•   When l=1   p-shape
•   When l=2   d-shape
•   When l=3   f=shape
spdf sublevels
• The first principle energy level has one
sublevel: (1s) (since n=1 has only 2e).
• The second level has two sublevels: (2s)
and (2p) (since n=2 has only 8 e).
S orbital and p orbitals
P orbitals
spdf sublevels
• The first principle energy level has one
sublevel: (1s).
• The second level has two sublevels: (2s)
and (2p).
• The third energy level has three sublevels:
(3s), (3p) and (3d)
spdf sublevels
• The first principle energy level has one
sublevel: (1s).
• The second level has two sublevels: (2s)
and (2p).
• The third energy level has three sublevels:
(3s), (3p) and (3d)
• The fourth energy level has four sublevels:
(4s), (4p), (4d), and (4f)
D orbitals
The Magnetic Quantum #
• ml= The magnetic quantum number. It ranges in values
from +l to –l with 0 included. It indicates the number of
sub-levels.
• They may contain the following maximum electrons:
• s = 0 ml= 0  1 option so 1 s sub-level
• p = 1 ml= +1,0,-1  3 options so 3 p sub-levels
• d = 2 ml= +2,+1,0,-1,-2  5 options so 5 d sub-levels
• f = 3 ml= +3,+2,+1,0,-1,-2,-3  7 options so 7 f sub-
levels

• Note that every sub-level has 2 electrons
Filling Orbitals
1. No more than two electrons can occupy one
sub-level.
2. Electrons occupy the lowest energy orbitals
available.
3. Each orbital on a sublevel is occupied by a
single electron before a second electron
enters.
4. Oh what fun this is!!! I know you think so too.
You are probably going to go home and tell
Permissible Quantum States
Magnetic Spin Quantum Number
ms = spin magnetic  electron spin
ms = ±½    (-½ = ) (+½ = )
Pauli exclusion principle states:
Each electron must have a unique set of 4 quantum numbers.
*What this means is that no more than two electrons
can occupy the same orbital, and that two electrons in
the same orbital must have opposite spins.

Electron spin is a purely quantum mechanical concept.
How do we use these quantum
numbers? We use them to write
ELECTRON CONFIGURATIONS

Electron configurations of the first 11 elements, in subshell notation.
Notice how configurations can be built by adding one electron at a time.

Z    ground state electronic configuration
atom
H      1    1s1
He     2    1s2
Li     3    1s2 2s1
Be     4    1s2 2s2
B      5    1s2 2s2 2p1
C      6    1s2 2s2 2p2
N      7    1s2 2s2 2p3
O      8    1s2 2s2 2p4
F      9    1s2 2s2 2p5
Ne     10   1s2 2s2 2p6
Na     11   1s2 2s2 2p6 3s1
Note: The energy levels do not go
in order. As a result you need the
Aufbau Principle to determine the
order.
Aufbau Principle
• States:
– the number of electrons in
an atom is equal to the
atomic number;
enter the orbitals in the
order of increasing
energy;
– an orbital cannot take
more than 2 electrons.
atom
Orbital Box Diagrams
B                    2p
1s       2s

C                    2p
Examples of ground state
1s       2s
electron configurations in
N                    2p
the orbital box notation that
1s       2s
shows electron spins  .

O                    2p
1s       2s

F                    2p
1s       2s

Cl                   2p
1s       2s          3s          3p

Mn                   2p
1s       2s          3s          3p

…
3d     4s
Hund's Rule
• every orbital in a subshell is singly
occupied with one electron before any one
orbital is doubly occupied
• and all electrons in singly occupied orbitals
must have the same spin.
atom
Orbital Box Diagrams
B                    2p
1s       2s

C                    2p
1s       2s
Look at the p subshells and
N                    2p
1s       2s

O                    2p
1s       2s

F                    2p
1s       2s

Cl                   2p     3s
1s       2s                    3p

Mn                   2p     3s
1s       2s                    3p

…               4s
3d
Practice
• http://www.chempractice.com/drills/java_A
O.php
• Now do the electron configurations and
orbital box diagrams for the first 20
elements
• Homework: page 136 #1-5, page 138 #1-
2,5-6, page 145-146 #6
Short Forms
• For atoms with many electrons, this notation can
become lengthy.
• It is often abbreviated by noting that the first few
subshells are identical to those of one or another
noble gas.
• Phosphorus, for instance, differs from neon (1s2
2s2 2p6) only by the presence of a third shell.
Thus, the electron configuration of neon is pulled
out, and phosphorus is written as follows:
[Ne]3s2 3p3.
Look at how the periodic table
takes electron configurations into
account
Exceptions
• Look at copper and chromium
• They do not follow the Aufbau principle
• Many of the transition elements like a half
filled s sub-level.
• You need to know these exceptions
Element    Z   Electron configuration               Short electron conf.
2   2   6   2   6   2     1              2     1
Scandium   21 1s 2s 2p 3s 3p 4s 3d                  [Ar] 4s 3d
2   2   6   2   6   2     2              2     2
Titanium   22 1s 2s 2p 3s 3p 4s 3d                  [Ar] 4s 3d
2   2   6   2   6   2     3              2     3
Vanadium   23 1s 2s 2p 3s 3p 4s 3d                  [Ar] 4s 3d
2   2   6   2   6   1     5              1     5
Chromium   24 1s 2s 2p 3s 3p 4s 3d                  [Ar] 4s 3d
2   2   6   2   6   2     5              2     5
Manganese 25 1s 2s 2p 3s 3p 4s 3d                   [Ar] 4s 3d
2   2   6   2   6   2     6              2     6
Iron       26 1s 2s 2p 3s 3p 4s 3d                  [Ar] 4s 3d
2   2   6   2   6   2     7              2     7
Cobalt     27 1s 2s 2p 3s 3p 4s 3d                  [Ar] 4s 3d
2   2   6   2   6   2     8              2     8
Nickel     28 1s 2s 2p 3s 3p 4s 3d                  [Ar] 4s 3d
2   2   6   2   6   1     10             1     10
Copper     29 1s 2s 2p 3s 3p 4s 3d                  [Ar] 4s 3d
2   2   6   2   6   2     10             2     10
Zinc       30 1s 2s 2p 3s 3p 4s 3d                  [Ar] 4s 3d
2   2   6   2   6   10    2    1         10        2   1
Gallium    31 1s 2s 2p 3s 3p 3d           4s 4p     [Ar] 3d    4s 4p
Electron Configurations for Ions
• Because scandium is a metal in group 3
on the periodic table it can lose three
electrons and form +3 cation with the
stable 3s23p6 configuration of argon.
•      Sc                    Sc3+ + 3e-

• [Ar] 3d1 4s2      [Ar] or [Ne] 3s2 3p6
electrons in Neon
1s2 2s2 2p6
s (l=0)   p (l=1)   d (l=2)   f (l=3)

n=1

n=2

n=3

n=4

n=5

n=6

n=7
• Homework:
• page 150 #10-13
• page 165 #1-4
Variations in Ionization Energy
What Happens?
• Draw the orbital box diagrams for Be and
B, then N and O.
1s 22s2        2 2s2     1
1s        2p

Beryllium     Boron
1s2 2s2    2p3   1s2 2s2    2p4

Nitrogen           Oxygen
Chapter 4
• NOW YOU MUST READ SECTION 4.1!!!!
• If you do not, you will suffer!!!!
• How can I make you suffer? Well on Friday you
are going to be shown 5 substances and asked
to figure out what kind of substances they are.
There is NO LAB PROCEDURE!!! He
heheehehe, you must design it yourself and then
complete the lab on Monday March 26.
HAHAHAHAHAHA.

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