• Try to find visual of uncertainty principle.

Chapter 6

Based his ideas on the relationship between matter and light (electromagnetic radiation)
We need background information on electromagnetic radiation before we can discuss
Bohr’s ideas

L I G H T
Properties of light:
• Part of the electromagnetic radiation spectrum
• Has a constant velocity (symbol is c) c = 3.00 X 10 8 m/sec
*watch the units!
on AP equation sheet!
There are 2 theories that can explain light: wave theory & particle theory.

light has a wave-like nature wavelength ( l ) - distance per cycle of wave frequency ( n ) - number of cycles per unit of time

on AP equation sheet!
For any wave, velocity = l • n
For any electromagnetic wave, c = l • n
3.00x10
8 m/sec= l • n l and n are inversely related for electromagnetic waves

Electromagnetic Radiation Spectrum - the type of radiation is related to its wavelength
-11 -9 -7 -5 -3 -1 1 3
10 10 10 10 10 10 10 10 ( l in cm)
| | | | | | | | | | | | | | |
| | | | | |
Gamma | X-rays | Ultra |V| Infrared | Micro | TV & radio radiation | | Violet |i| light | waves | signals
| | light |s| | | smaller wavelength larger frequency larger energy
VISIBLE LIGHT larger wavelength smaller frequency smaller energy

• Max Plank figured out that light can be released from and absorbed by atoms in “chunks” that have a minimum size.
• Light is made up of particles called photons or quanta- bundles of energy

There is a mathematical equation that ties these 2 theories together
Energy of photon = (Plank’s constant) (frequency)
E = h • n on AP equation sheet!
h = Plank’s constant = 6.63 X 10 -34 Joule-second (J-s)
E and n are directly related
E and l are inversely related

Electrons are found in “orbits” that surround the nucleus.
The concept of an orbit means that the electron travels around the nucleus at a fixed distance and at a fixed
speed that keeps it on that particular orbit.
Yes, Bohr was wrong about orbits….let’s pretend he meant energy levels
(because that’s what they really are!)

When an electron absorbs a certain amount of energy, called a “quanta,” it can jump to an energy
level farther from the nucleus.

When electrons absorb energy to jump up a level, they immediately
release the same amount of energy to get back to the ground state.

• Sometimes, the energy that is absorbed or released by excited electrons can be seen as different colors of light!
• The color depends on the wavelength or
energy of the light emitted!

• Examined the light released by atoms of certain elements
• Worked mostly with hydrogen since it has only 1 electron
• Bohr used a spectroscope to analyze the light released from a vacuum tube - the result is a bright line spectrum more commonly called an
EMISSION SPECTRUM (black background with bright colored lines)

Only certain wavelengths or photon energies of light are present in the emission spectrum
An element’s emission spectrum is unique to that element (like a fingerprint!)

E
N
E
R
G
Y
U V Visible
Energy level diagram
I R
7
6
5
4
A
H emission spectrum
B C
D
3
A B C D
2 violet red
4 x 10 -5 λ 7 x 10 -5 smaller λ larger λ larger E smaller E again, 21 possible emission lines again, also 21 possible absorption lines absorption lines (arrows) would point UP
1
Arrow D is the longest arrow in the visible region which means it has the highest energy and corresponds to the violet line in hydrogen’s emission spectrum.
Arrow A is the shortest arrow in the visible region which means it has the lowest energy and corresponds to the red line in hydrogen’s emission spectrum.

• To find the energy associated with each orbit in the hydrogen atom, Bohr used the equation:

2.178 10

18
E

2
Joules On AP Equation Sheet!
n
1.
Subtract the energy of the orbits between which an electron moves to find the energy released/absorbed.
2.
Use E=hν to convert energy to frequency
3.
Use c=λν to convert frequency to wavelength
4.
These correspond to the colors seen in the emission spectrum!

Bohr tried to apply his ideas about electron behavior of hydrogen atoms to atoms of other elements.
Bohr tried to explain the arrangement of elements on the periodic table by showing that elements
in a column of the periodic table have similar
chemical properties because they have
similar electron structures.

K Ca does not work for #21 scandium

• Loius De Broglie…if light can behave like matter
(photons)…can matter behave like a wave??
YEP!
l  h mv on AP equation sheet!
Λ = wavelength (m) h= Plank’s constant (J-sec) m= mass (kg)
V = velocity (m/sec)

• Werner Heisenberg….because of the dual nature of matter both the path and velocity of a particle with the size and mass of the
electron can not be determined at the same
instant in time - in other words, electrons cannot be in orbits!
• *The uncertainty in the calculation of the position of the electron is larger than the entire atom because the electron is so small!

1.
Does not explain the behavior of atoms with more than 20 electrons
2. Violates the Heisenberg Uncertainty Principle

• He developed the idea of wave mechanics - treating the electron as if it had wave-like properties.
• He mathematically considered the region of space in the atom where an electron of a given energy would MOST PROBABLY BE LOCATED.
• This is a radical departure from earlier work.
Schrödinger says to forget exactly where the electron is (Heisenberg says we can’t find out anyway) and concentrate on where the electron
MIGHT be.

• The math involved in wave equations is very complicated and beyond the scope of this course. Fortunately, the results of this theory are easier to understand than the mathematics.

ORBITAL
- The region of space in an atom where it is highly probable that an electron of a given energy will be located
- About 90 % of the solution set to the wave equation
- Max of 2 electrons/orbital

Each box = electron
Each column = orbital
Each group = sublevel
Big # out front = Energy level

• Atoms fill the lowest energy positions first
• 2 methods to know the order:
▫ Periodic table
▫ Wavy line diagram

n
4
5
6
1
2
3
7 s
1
p d
(n-1) f
(n-2)

• Indicates a specific energy LEVEL (energy levels values found by Bohr)
• Indicates the relative size of the energy level - larger value means larger size
• Indicate the relative energy of the electron
– larger value means higher energy.
• Possible values are n = 1, 2, 3, 4, 5, 6, 7, ...


• Indicates a specific energy
SUBLEVEL
• The number of energy sublevels per energy level is equal to the value of the energy level (example: 3 sublevels in the 3rd energy level)
• Possible values are

= 0, 1, 2, 3, ..., max. value of n-1
• more commonly

= s, p, d, f, ...
(alphabetically after f)

Indicates the shape of the orbitals in that sublevel
- s orbital has a spherical shape
- p orbital has a dumb-bell shape (technically, a pair of ellipsoids)

d orbitals - we won’t worry about their shape
- f orbitals - we won’t worry about their shape


• Indicates a specific ORBITAL of an energy sublevel
• Indicates the spatial orientation of the orbital on an x, y, z axis system in a magnetic field
• Possible values are m = -
 to +

(integer values)

s
• Indicates a specific ELECTRON of an orbital
• Indicates the electron spin direction - clockwise or counter-clockwise
• Possible values are s = -1/2 or +1/2

• No two different electrons in the same atom can have the EXACT same set of quantum numbers .

A Complete Set of Quantum Numbers
( n,  , m, s ) orbital electron sublevel level

- A representation of an atom’s electron structure using energy levels and sublevels
- All electrons are shown
Example: for
6
C
1s 2 2s 2 2p 2

-A representation of an atom’s electron structure using arrows for the electrons and lines/boxes for the orbitals
- Hund’s Rule: Within the same sublevel, the electrons go in separate orbitals until they are forced to pair up
Example: for
6
C
↑↓ ↑↓ ↑ ↑ __
1s 2s 2p

• Paramagnetic: an atom with one or more unpaired electrons (electron alone in an orbital)
• Diamagnetic: an atom with only paired electrons
(two electrons in each filled orbital)

• We predict ground state electron configurations.
• Sometimes when an atom gains energy an electron jumps to a higher energy level.
Ground state of iron: [Ar]4s 2 3d 6
But if it gains some energy one of the 4s electrons might jump into the 3d sublevel to make the electron configuration:
[Ar]4s 1 3d 7

• Half filled sublevels have an extra level of stability compared to other configurations – because of this some ground state configurations are not what we predict!
Examples:
• Cr & Mo – s 1 d 5 (not s 2 d 4 )
• Cu & Ag – s 1 d 10 (not s 2 d 9 )

• 6(6, 13, 15, 17, 28, 37, 45, 50a, 50b ,53, 55, 69, 85)