Atomic Theory I
Chap 5
Wave Nature of Light
Atomic Theory
Learning Targets
5.5 Know the history of the
development of the atomic
theory
5.12 Apply quantum mechanics
Light and Elements
• Since ancient times that
elements emitted their own
specific colors when vaporized
• That light can be separated into a line
spectra by the use of a prism.
What is light?
• Electromagnetic radiation (a.k.a. light) is
a form of energy that behaves like a
wave.
• Waves have 3 properties:
– Wavelength = Distance between consecutive
peaks or troughs
– Frequency = How many peaks pass a given
point per time
– Speed = How fast a given peak moves through
space (3x108 m/s)
How are the wavelength & energy related?
• Shorter wavelength is higher frequency
which means greater energy
Let’s look at visible light
• Purple light is of slightly higher energy
than red light (far left versus far right of the
spectrum)
• Purple light has a higher frequency and a
shorter wavelength than red light.
Light behaves like a wave…
…but it also showed particle properties
• Debate on light properties raged until
the 1900’s because light:
– behaves like a wave AND
– behaves like a particle
Photoelectric Effect
• Light shined on metal could eject
electrons
– Only certain wavelengths/frequencies
ejected electrons
• Demonstrated Particle Nature of Light!
OK so why are Light Emission Spectra Unique for each Atom
Quantum Theory
• Planck suggested since only certain
wavelengths of light are given off that the
energy of an element is restricted to
“pieces” of particular sizes called a
quantum (plural is quanta)
Max Planck & Neils Bohr
Einstein said that light are packets
(particles) of a specific amount of energy
called photons.
How can you explain the spectrum
of an element such as hydrogen?
• How can something with a positive
nucleus and negative electrons moving
around it only give off energy in small
discrete areas
• We need some help
……along came Niels Bohr
Bohr’s Contribution
• Bohr proposed a planetary model of the
atom: electrons orbit the nucleus in
distinct energy levels (n = 1, 2, 3, ...).
• The lowest energy level is called the
ground state (n = 1)
• When an electron absorbs
energy it is promoted to an
excited state (n = 2,3,4…)
Contributions restated
• Energy increases as shell (level) number
increases.
• Light can be absorbed to make an electron
jump to a higher level (excited).
• Light is emitted when an electron falls to a
lower energy level
• Bohr predicted all the lines seen for the
spectrum of Hydrogen.
• Bohr’s model worked for hydrogen,
but he was unable to explain the
spectra of atoms with more than
one electron.
• However, his ideas led to the
current model of the atom we use
today.
A New Approach to the Atom
• Quantum-Mechanical Model:
– Energy of electrons is quantized (Planck)
– De Broglie reasoned: if light had a dual nature
matter could also (wave & particle properties).
– Heisenberg Uncertainty Principle
• It is impossible to know the exact position as
well as the momentum of an electron at any
given instant
Schrodinger
• developed equation for location of electron
Atomic orbital
• Exists around the nucleus
• Location where you have a chance of finding the
electron
• Orbitals have different shapes
Quantum Numbers
• Quantum numbers describe the position of
an electron.
1. Principle quantum number (n):
• Positive numbers starting with 1 (n = 1,2,3,….)
• n=1 is lowest energy level (ground state)
• Designates the energy level (higher n, farther from
nucleus, higher energy)
Quantum Numbers
2. Sublevel tell shape of orbital
Principle energy level divided into sublevels
Designated by letters (s,p,d,f)
Letter
s
p
d
f
Shape
spherical
dumbbell
complex
complex
# in a level
1
3
5
7
Increasing energy
•
•
f-orbitals
How do levels and sublevels fit?
• # of sublevels in a main level equals
the principle quantum number
– level 1 = has 1 sublevel = s
– level 2 = has 2 sublevels = s & p
– level 3 = has 3 sublevels = s, p, & d
– level 4 = has 4 sublevels = s, p, d, & f
– level 5 = has 5 sublevels = s, p, d, f & g
How do electrons fit into sublevels?
• Pauli Exclusion Principle: an individual
orbital can hold up to 2 electrons
Sublevel # of orbitals Maximum number of
electrons in sublevel
s
1
2
p
3
6
d
5
10
f
7
14
Electron Configuration
• A distribution of atom’s electrons among
levels, sublevels, and orbital's
•Aufbau Principle – electrons added to
lowest energy orbital's available until all
electrons of atom have been used
Electron Configuration
• Beryllium (4 electrons)
Orbital Type
# of Electrons
»1s2 2s2
Energy Level (n)
Electron Configurations
• Start in 1st energy level
• Fill electrons into orbitals and levels starting
with lowest and going to highest.
• Note: starting in period 3, 4s fills before 3d,
use diagonal diagram to get order correct
Examples
H
1s 1
Mg 1s2 2s2 2p6 3s2
V 1s2 2s2 2p6 3s2 3p64s2 3d3
Electron filling clues come from
the periodic table?
 If you look at the table there are 7 rows:
 Row (periods) = 7 levels
 there are 4 major areas (blocks):
 Blocks = 4 sublevels
 there are 2 columns in the first area:
 s sublevel(1 orbital can hold 2 electrons)
 There are 6 columns in the next main block;
 p sublevel has 3 orbitals w/ 2 e- per orbital
Electron filling clues come from
the periodic table?
 the transition area has 10 columns:
 d sublevel(5 orbitals w/ 2 e- per orbital)
 14 columns in the inner transition;
 f sublevel has 7 orbitals w/ 2 e- per orbital
d (n-1)
f (n-2)
back
• Example – Germanium (atomic # 32)
Row #
11
22
33
44
55
66
77
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2
Short Hand
• Valence electrons
– The electrons in the highest energy level
– they are involved in all reactions
• Electron configuration short hand only has
valence electrons written with core
electrons represented by the Nobel gas
from the row before.
Electron Configuration Shorthand
• Longhand Configuration
Cl
17e
1s2 2s2 2p6 3s2 3p5
Core Electrons
Valence Electrons
• Shorthand Configuration (noble gas
configuration)
2
5
Cl 17e [Ne] 3s 3p
Shorthand Configuration
• Example - Germanium
1
2
3
4
5
6
7
[Ar] 4s2 3d10 4p2
Try the noble gas shortcut:
53I
30Zn
[Kr] 5s24d105p5
[Ar] 4s23d10
Electron Configs of Ions
• Cations lose the valence
electrons first (outer most)
• Na 1s2 2s2 2p63s1
– loses 3s1 becomes Na+1
• Ca 1s2 2s2 2p63s2 3p6 4s2
– loses 4s2 to be Ca+2
Electron Configs of Ions (cont)
• Anions(-) gain electrons to fill
outer level
• F 1s2 2s2 2p5
– gains e-  F 1s2 2s2 2p6
-
• O 1s2 2s2 2p4
– gains
e-
 O 1s2 2s2 2p6
-2
Orbital Filling (Box) Diagrams
•
•
•
•
Show distribution of electrons in orbitals.
Electrons are represented by arrows
Orbitals are represented by boxes
Hund’s Rule
– Within a sublevel, place one e- per orbital
before pairing them. (“Empty Bus Seat Rule”)
WRONG
RIGHT
# of Orbitals
•
•
•
•
-
s sublevel = 2 e = 1 orbital (1 box)
p sublevel= 6 e = 3 orbitals (3 boxes)
d sublevel=10 e =5 orbitals (5 boxes)
f sublevel=14 e =7 orbitals (7 boxes)
Other Electron Notations
• Electron Configuration
2 2s2 2p4
O
1s
8
• Orbital Diagram
O
8e-
1s
2s
2p
What will be the correct filling?
Example
• Na 1s2 2s2
2p6
3s1
Example
47Ag = [Kr]
[Kr]
5s2
4d9
Example
47Ag = [Kr]
[Kr]
5s2
4d9
Electron Dot Structure
• An atom’s valence electrons are
represented by an electron dot structure.
Na = 1s22s22p63s1 (1valence electron)
Na
Mg = 1s22s22p63s2 (2valence electrons)
Mg
the number of valence electrons is the same as
the column number in periodic table (except
helium)
All the elements within a Group have the same number of valence
electrons, therefore they will have the same number of dots in their
Lewis Structures.