October 16, 2014
Chapter 5: Electrons in Atoms
Honors Chemistry
Bohr Model
Niels Bohr, a young Danish physicist and a student of
Rutherford improved Rutherford's model.
Bohr proposed that an electron is found only in specific
circular paths, or orbits, around the nucleus. Each
electron orbit has a fixed energy.
Energy levels: the fixed energies of an electron.
Quantum: The amount of energy required to move an
electron from one energy level to another level.
Electrons can jump from one energy level to another.
The term quantum leap can be used to describe an
abrupt change in energy.
Light
The modern quantum model grew out of the study of light.
1. Light as a wave
2. Light as a particle
-Newton was the first to try to explain the behavior of light by
assuming that light consists of particles.
October 16, 2014
Light as a Wave
1. Wavelength-shortest distance between equivalent points
on a continuous wave.
symbol: λ (Greek letter lambda)
2. Frequency- the number of waves that pass a point per
second.
symbol: ν (Greek letter nu)
3. Amplitude-wave's height from the origin to the crest or
trough.
4. Speed-all EM waves travel at the speed of light.
symbol: c=3.00 x 108 m/s in a vacuum
The speed of light
c=3.00 x 108 m/s
Calculate the following:
1. The wavelength of radiation with a frequency of 1.50 x 1013
Hz. Does this radiation have a longer or shorter wavelength than
red light? (2.00 x 10-5 m; longer than red)
c=λν
The wavelength and frequency are inversely proportional
to each other.
2. The frequency of radiation with a wavelength of 5.00 x 10-8 m.
In which region of the electromagnetic spectrum is this radiation?
(6.00 x 1015 Hz; ultraviolet)
Electromagnetic Radiation
light consists of electromagnetic waves. This radiation includes
radio waves, microwaves, infrared waves, visible light,
ultraviolet waves, X-rays, and gamma rays.
http://www.youtube.com/watch?v=cfXzwh3KadE
Light as a Particle
• Einstein
> Light as a photon
> Photoelectric Effect
Photoelectric Effect: refers to the emission of electrons from
a metal when light shines on the metal.The photoelectric
effect causes electrons to be ejected from the surface of a
metal when light is of high enough frequency to hit the
metal's surface.
October 16, 2014
Light as a Particle
• Wave behavior of light cannot explain why heated objects
give off distinct colors (specific frequencies) of light.
• Max Planck (1858-1947)
> studied the different wavelengths of light emitted by
heated objects
> conclusion: matter can gain or lose energy only in
specific amounts.
• quantum-minimum energy that can be gained or lost by
an atom
> hot objects emit light in quantized amounts
E=hv
Planck proposed the previous relationship between a quantum of
energy and the frequency of radiation.
E= energy, in joules
v= frequency in s-1
h= Planck's constant 6.626 x 10-34 J·s
E=hc
λ
http://www.youtube.com/watch?v=Xmq_FJd1oUQ
Practice
The yellow vapor from a sodium lamp emits 3.37 x 10
What is the wavelength of this light?
E=hv
E=hc
λ
3.37x 10-19 J=(6.626 x 10-34 Js)(3.00x 108 m/s )
2
λ
λ= 5.89 x 10-7 m or 589 nm
-19
J.
Physics and the Quantum Mechanical Model
Wave-Particle Duality
October 16, 2014
Atomic Spectra
When atoms absorb energy, electrons move into higher
energy levels. These electrons then lose energy by
emitting light when they return to lower energy levels.
Atomic emission spectrum:
the frequencies of light that are emitted by an element into
separate discrete lines.
The fact that hydrogen atoms emit only specific frequencies of
light indicated that the energy differences between the atoms'
energy states were fixed.
The following figure is the line-emission spectrum for hydrogen.
Explanation of the Atomic Emission Spectra
Bohr's model not only explained why the emission
spectrum of hydrogen consists of specific frequencies of
light, but it also predicted specific values of these
frequencies.
Ground State:
When an electron has its lowest possible energy. The
principal quantum number (n) is 1.
*If the electron is then excited to a higher energy level, the
dropping of the electron to a lower energy level creates
the light emitted.
October 16, 2014
Quantum Mechanical Model
Quantum Mechanical Model
de Broglie's Hypothesis
• matter has wave-like properties
• consequence: whole number of wavelengths must fit
within the circumference of the orbit.
• The energy level number, n, is equal to the number of
waves.
Heisenberg uncertainty principle:
It is impossible to know exactly both the velocity and the
position of a particle at the same time.
We define electron energy exactly but accept that we do
not know the electrons definite position.
Impossible to take any measurement of an object
without disturbing it!
Quantum Mechanical Model
http://www.youtube.com/watch?v=7SjFJImg2Z8&feature=player_embedded
http://www.youtube.com/watch?v=uWMTOrux0LM
Erwin Schrödinger used results to devise and solve a
mathematical equation describing the behavior of the
electron in a hydrogen atom.
*The quantum mechanical model determines the allowed
energies of an electron and how likely it is to find
the electron in various locations around the nucleus.
Probability of finding a hydrogen electron.
We can apply to other elements.
Quantum Mechanical Model
October 16, 2014
Orbit vs. Orbital
Orbit: defined path of an electron (Bohr)
Orbital: defined area of space for finding
an electron. (Schrodinger)
Types of Orbitals
• The most probable area to find these electrons take on a
shape.
• So far, we have 4 shapes.
> s, p, d, f.
• No more than 2 electrons assigned to an orbital.
> One spins clockwise, one spins counterclockwise.
October 16, 2014
Atomic Orbitals and Quantum Numbers
Quantum numbers specify the properties of atomic orbitals and the
properties of electrons in orbitals. This first three quantum numbers
result from the Schrödinger equation.
Quantum Number
Symbol
Principal Quantum Number (n)
Indicates the main energy level occupied by the electron.
n= 1, 2, 3 and so on.
As n increases, the electron's energy and its average
distance from the nucleus increases.
*The total number of orbitals that exist in a given shell
Definition
n
principle quantum number
l
angular momentum quantum number
(shape of the orbital)
ml
magnetic quantum number
(orientation)
ms
spin quantum number
Angular Momentum Quantum Number (l)
Indicates the shape of the orbital.
The number of orbital shapes possible is equal to n.
l=n-1
L value
Letter/Shape
Ex:
0
s
1
p
If n=2; L has two sublevels,
the s and p orbitals
because
2
d
3
f
l=n-1; l=0 and 1
How many sublevels would
n=3 have? _________
October 16, 2014
Magnetic Quantum Number (ml)
Spin Quantum Number (ms)
Indicates the orientation of an orbital around the nucleus.
Values of m are whole numbers; 0 from L to -L
Only two possible values (+1/2 and -1/2) which indicate the
two fundamental spin states of an electron in an orbital.
m=-L...0...L
A single orbital can then hold how many electrons? ______
Because an s orbital is spherical and is centered around
the nucleus, it has only one possible orientation.
How many orientations are present in the following orbitals?
Tying all Quantum Numbers Together
We can combine all four quantum numbers together to give us
1 atomic orbital contains 2 electrons.
p orbitals
___________
d orbitals
___________
Depending on the subshell, there may be different number of
atomic orbitals present.
The combination of the quantum numbers will tell us whether
we are talking about an entire subshell or one atomic orbital.
*This may seem confusing until we do an example
f orbitals
___________
Tying all Quantum Numbers Together
n=1, L= 0, ml=0, ms= +1/2, -1/2
n
l
m
This is one atomic orbital
ms
2
If I asked you to write all of the allowed quantum number
combinations for n=3, this would be all quantum numbers for
the n=3 subshell.
Table of Allowed Quantum Numbers
October 16, 2014
Electron Configuration
Introduction
1. Aufbau Principle
Electrons occupy the orbitals of lowest energy first.
Definition
The ways in which electrons are arranged in various
orbitals around the nuclei of atoms.
Change proceeds toward the lowest possible energy.
3 Rules: These three rules tell you how to find the
electron configurations of atoms.
1. Aufbau principle
2. Pauli Exclusion Principle
3. Hund's Rule
2. Pauli Exclusion Principle
an atomic orbital may describe at most two electrons.
*To occupy the same orbital, two electrons must have opposite spins.
3. Hund's Rule
electrons occupy orbitals of the same energy in a way that makes the
number or electrons with the same spin direction as large as
possible.
October 16, 2014
Orbital Notation vs.
Electron-Configuration Notation
Orbital Notation:
an unoccupied orbital is represented by a line, ___, with
the orbital's name written underneath the line. An orbital
containing one electron is represented as ____. A orbital
containing two electrons is represented as ____.
What element is represented by the
following orbital diagram?
How many unpaired electrons are present?
Electron-Configuration Notation:
Eliminates the lines and arrows. Instead, the number of
electrons in a sublevel is shown by adding a superscript to
the sublevel.
1s2
2s2
2p6
3s1
Ex: Hydrogen is represented by 1s1
Example: H, He, Li, Be, B
Sample Problem
The electron configuration of boron is 1s22s22p1. How many
electrons are resent in an atom of boron? What is the atomic
number for boron? Write the orbital notation for boron.
Steps for Writing Electron Configurations
1. Locate the element on the periodic table.
2. Start by filling in the electrons from the lowest possible
energy state first.
(Noble Gas Configuration ONLY after some practice!)
3. Continue filling higher in energy until all electrons are
used.
4. Do not forget to obey the three rules of Aufbau, Hund
and Pauli.
How many unpaired electrons are present?
Electron Configuration Practice
Noble Gas Shortcut
Examples (Cd)
1. Find the noble gas (group 18 on the periodic table)
with an atomic number less than the element...Kr
2. Put the noble gas in brackets...[Kr]
3. The noble gas filled the p sublevel with n=period
number...4.
4. Follow the periodic table/diagonal rule and
continue the notation from the next s-sublevel using
the remaining electrons...[Kr]5s24d10
H
He
Li
N
Fe
U
October 16, 2014
Nobel Gas Shorthand
When you are writing electron configurations for elements that
have large atomic numbers, you may use the nobel gas
configuration shortcut.
You use the most previous noble gas in brackets [He], and then
finish the electron configuration the normal way.
Ions
If you have an element that has a positive or negative
charge, you will add or take away electrons based on the
charge.
You MUST take electrons away from the highest energy
level (n) first!!!!!
Example:
Cl- : You must add an electron
W2+: Take 2 electrons away from the outermost energy
level!
Exceptional Electron Configurations
Write what you would expect the electron configuration for
copper (Cu) and chromium (Cr) to be.
Cu
Cr
*Some actual electron configurations differ from those
assigned using the aufbau principle because half-filled
sublevels are not as stable as filled sublevels.
3 Types of Notation
October 16, 2014
Practice Different Notations
Practice Different Notations