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Chapter 5 Chemistry :
EMR and Quantum Theory
How did scientist find out about energy levels and sublevels?
•When we previously found the electron configuration for elements, it was for electrons at ground state
(the lowest energy possible).
How did scientist find out about energy levels and sublevels?
•As energy is added to atoms, they absorb the energy by electrons going from ground state to an
excited state, where electrons are no longer in the lowest energy orbitals.
•Electrons can then only go back to ground state by releasing the energy, usually in the form of light in
discreet packets called photons.
•
If electrons could orbit anywhere, when they went from excited state to ground state they
would emit light of all wavelengths.
This doesn’t happen! Only certain wavelengths of light are emitted, which is different for each
atom.
How did scientist find out about energy levels and sublevels?
•
This is called an atomic emission spectra, and is different for every element!
•
A scientist named Max Planck studied the cooling of metal and how its color changes, and tied
the idea of frequency of light to energy.
•
The bands of light correspond to the specific energy levels.
•
White light (sunlight) is a blend of all colors (ROY G BIV)
combined together
•
The wavelength (λ) and frequency (υ) for each color are unique to
that color.
As light passes through a prism…
•the different wavelengths of the colors are separated.
•individual colors can be detected by the eye.
2…The Electromagnetic Spectrum
•
All substances (radioactive or not) emit electromagnetic radiation.
•
Only part of the spectrum that human eyes can detect is visible light
•
All other radiations have wavelengths that are either too long or too short for our eyes to
detect.
3…Wavelength vs. Frequency
•Wavelength (λ):
•Frequency (ν):
distance b/n crests of a wave.
# of wavelengths that pass a certain point in a given amount of time; SI is
Hertz (Hz)
4…Wave Calculations
•all waves on the EM spectrum travel at the speed of light (c).
•wavelength (λ) and frequency (υ) are inversely related.
•All waves move at the speed of light (c): 3 X 108
•C= λv = wavelength times frequency
Wavelength = 3 X 108 / Frequency
Frequency = 3 X 108 / Wavelength
Planks Assumptions:
Quanta or Photons are packets of light which are distinct bundles of energy.
Energy of the photon is directly proportional to the frequency of the light wave.
E = h(V) ; SI unit is in Joules
V = E/h
h = Plank’s constant = 6.626 X 10-34 J/Hz
Bohr Model and Energy Levels
•
In the Bohr model, electrons are in energy levels, or regions where they most probably
are orbiting around the nucleus.
•
The analogy is that energy levels are like
the rungs of a ladder—you
cannot be
between rungs, just like an electron
cannot be between energy levels.
•
A quantum of energy is the amount of
one energy
level to the next.
•
Quantum Mechanical Model
•
In 1926, Erwin Schrodinger used the new quantum theory to write and solve
mathematical equations to describe electron location.
•
The Quantum Mechanical Model, cont.
•
Today’s model comes from the solutions to Schrodinger’s equations.
•
Previous models were based on physical models of the motion of large objects.
•
This model does not predict the path of electrons, but estimates the probability of finding
an electron in a certain position.
•
There is no physical analogy for this model!
•
Evolution of Electron Models
•
•
The first model of the electron was given by J.J. Thompson—the electron’s discoverer.
His was the “plum pudding” model.
The Rutherford Model
•
•
energy it takes to move from
With Rutherford’s discovery of the nucleus of an atom, the atomic model changed.
The Bohr Model
•
Niels Bohr introduced his model, which answered why electrons do not fall into the
nucleus.
•
He introduced the concept of energy levels, where the electrons orbited similar to the
way the planets orbit the sun.
•
Bohr Model and Energy Levels
•
In the Bohr model, electrons are in energy levels, or regions where they most probably are
orbiting around the nucleus.
•
•
•
The analogy is that energy levels are like the rungs of a ladder—you cannot be
between rungs, just like an electron cannot be between energy levels.
•
A quantum of energy is the amount of energy it takes to move from one energy
level to the next.
Quantum Mechanical Model
•
In 1926, Erwin Schrodinger used the new quantum theory to write and solve
mathematical equations to describe electron location.
•
Today’s model comes from the solutions to Schrodinger’s equations.
•
Previous models were based on physical models of the motion of large objects.
•
This model does not predict the path of electrons, but estimates the probability of
finding an electron in a certain position.
•
There is no physical analogy for this model!
Where are the electrons?
•
In an atom, principal energy levels (n) can hold electrons. These principal energy levels
are assigned values in order of increasing energy (n=1,2,3,4...).
•
Within each principal energy level, electrons occupy energy sublevels. There are as
many sublevels as the number of the energy level (i.e., level 1 has 1 sublevel, level 2 has
2 sublevels, etc.)
•
There are four types of sublevels we will talk about—s,p,d and f. Inside the sublevel are
atomic orbitals that hold the electrons. Every atomic orbital can hold two electrons.
•
S has 1 orbital, P has 3, D has 5 and F has 7. How many electrons can each one hold?
Levels: 1, 2, 3, 4 ...
Sublevels: s, p, d, f
Orbitals: 1 for s, 3 for p, 5 for d, 7 for f.
Electrons: 2 for each orbital.
Orbital Shapes

S is shaped like a sphere

P is shaped like a peanut
•
D orbitals – fyi only
(daisy)
•
So how many electrons can each energy level hold?
–
Level 1 has an s sublevel=2 e-
–
Level 2 has an s and a p sublevel=8e-
–
Level 3 has an s, p and d sublevel=18e-
–
Level 4 has an s, p, d and f sublevel=32e-
Electron Configuration
•
In the atom, electrons and the nucleus interact to make the most stable arrangement possible.
•
The ways that electrons are arranged around the nucleus of an atom is called the electron
configuration.
Aufbau Principle
•
Electrons occupy orbitals of the lowest energy first.
Pauli Exclusion Principle
–
Each orbital can hold only TWO electrons with opposite spins.
–
“Sports Spectator Rule” (fill the lower stands first)
–
Electrons fill the lowest energy orbitals first.
Hund’s Rule “Empty Bus Seat Rule”
Within a sublevel, place one electron per orbital before placing a second electron.
Heisenberg Uncertainty Principle
•
Heisenberg concluded that it is impossible to make any measurement on an object with out
disturbing the object (at least a little).
•
The principle states:
“It is fundamentally impossible to know precisely both the velocity and the position of a particle at the
same time.”
•
B. Notation
»
•
Unabbreviated
•
C. Periodic Patterns
•
C. Periodic Patterns
•
Example - Hydrogen
Orbital Diagram
•
C. Periodic Patterns
•
Shorthand Configuration
–
Core e-: Go up one row and over to the Noble Gas.
–
Valence e-: On the next row, fill in the # of e- in each sublevel.
•
C. Periodic Patterns
•
Example - Germanium
•
D. Stability
•
Electron Configuration Exceptions
•
Valence Electrons
•
•
–
Electrons in the atom’s outermost orbitals.
–
These are the electrons that determine the atom’s chemical properties.
–
The fewer valence electrons an atom holds, the less stable it becomes and the more
likely it is to react.
Octets
–
When an atom has 8 electrons in its largest energy level it is said to have an octet.
–
This is the most stable and least reactive atom.
Lewis Dot Structures
–
An atom’s dot structure consists of the element’s symbol, (represents the atomic
nucleus and inner-level electrons), surrounded by dots representing the atom’s
valence electrons.
–
G. N. Lewis devised this method while teaching a college chemistry class in 1902.
–
Example:
Lithium
•
Atomic Number: 3
•
Electron Configuration: 1s22s2
•
Electron Dot Structure: Li·