Download Ch05ElectronConfig - Journigan-wiki

Electrons in Atoms
and
The Quantum Theory
Unit III
Ch. 5
Warm Up-02/06/13
Define isotope.
2. How do you calculate the number of
neutrons in an atom?
3. What does an atom’s atomic mass tell
you?
4. What can you never change about an
atom?
1.
Warm Up-02/12/13
1.
2.
3.
What does Pauli’s Exclusion Principle tell you
about how electrons enter their orbitals?
What does Hund’s Rule tell you about how
electrons enter their orbitals?
What does the Aufbau Principle tell you about
how electrons enter their orbitals?
Warm Up
Please complete a K-W-L chart concerning
electron configuration.
1.
2.
3.
What do I know about electron
configurations?
What would I like to know about electron
configurations?
What did I learn about electron
configurations (after class)?
Warm Up-02/08/13
1.
2.
3.
4.
5.
How many electrons will fit in the “s” suborbital?
How many electrons will fit in the “p” suborbital?
How many electrons will fit in the “d” suborbital?
How many electrons will fit in the “f” suborbital?
Please draw electron configurations for the
following: He, Na, Mg, Zn and Br.
Homework: Chapter 3
Due Date: 02/13/13
Page 137: 18-22.
Page 146: 46-52, omit 51
Page 147: 77-80
Warm Up
1.
2.
3.
4.
5.
Construct the
iron.
Construct the
silicon.
Construct the
potassium.
Construct the
titanium.
Construct the
cobalt.
electron configuration for
electron configuration for
electron configuration for
electron configuration for
electron configuration for
Warm Up
1.
2.
3.
4.
What percent are carbon and oxygen in
CO2?
What percent are calcium, sulfur and
oxygen in CaSO4?
What percent are sodium, nitrogen and
oxygen in NaNO3?
What percent are calcium and clorine in
CaCl2?
Electromagnetic Energy
Frequency (); measured in units of hertz (sec-1)
Wavelength () - measured units of length: mm, nm
Amplitude - the height of the wave
Speed (velocity) of electromagnetic energy =
3.00 x 108 m/s
c=
Electromagnetic Spectrum
Wavelength ()
1-800 m
10-1 - 10-2 m
10-2 - 10-4 m
10-4 - 10-6 m
400-700 nm
10-8 - 10-10 m
10-10 - 10-13 m
10-13 - 10-15 m
Description
Radio waves
Radar
Microwaves
Infrared
Visible
Ultraviolet
X-rays
Gamma Rays
Frequency ()
104 - 109
109 - 1010
1010 - 1012
1012 - 1014
1014 - 1015
1015 - 1018
1018 - 1021
1021 - 1023
Spectroscopy
 The method of studying substances
exposed to some sort of exciting energy.
Spectrum
 Observed when white light is dispersed
into the colors of the rainbow by a prism or
diffraction grating.
Emission Spectra
Absorption Spectra
The Photoelectric Effect
The emission of electrons from a metal when light
shines upon it. If the light frequency was below a
certain level, no electrons were emitted. The wave
theory of light predicted that light of any frequency
could supply enough energy to eject an electron.
Scientists could not explain why a certain
minimum frequency was required.
Radiation
Caused by an unstable nucleus which will eject
either a particle or energy until it reaches a
more stable arrangement.
Emission
Description
alpha
helium nucleus
beta
high speed electron
gamma
v. high energy X-rays
Planck’s Hypothesis
In 1900 Max Planck was studying black body
radiation. He suggested that hot objects emit
energy in small specific amount called
quanta. A quantum is the minimum amount of
energy that can be gained or lost by an atom.
Energy is given off (emitted) in little packets (or
quanta) called photons.
The amount of energy emitted is proportional to
the frequency of the light emitted according to
the equation:
E = hn
E = energy of a quantum or radiation
n = the frequency of the radiation
h = 6.626 x 10-34 J · s (h = Planck’s constant)
Einstein (1905) introduced the concept that
electromagnetic radiation has a dual waveparticle nature.
Relationship between electromagnetic
energy and electrons
An electromagnetic wave of a certain frequency
has only one possible wavelength:  = c/n
It has only one possible amount of energy:
E = hn
c and h are constants. If frequency, wavelength
or energy is known, we can calculate the
other two. White light can be thought of as a
wave or as a stream of particles, which
Einstein called photons. A photon is a particle
of radiation having zero rest mass and carries
a quantum of energy.
Therefore,
Ephoton = hn
The Hydrogen Atom Line-Emission Spectrum
See p. 127 for an explanation of ground state and
excited state for an atom.
Ground state - The smallest orbit an electron can
occupy.
The energy of the photon, Ephoton , corresponds to the
energy difference between the different energy
levels in an atom [ E1 and E2 , for example].
Rutherford-Bohr Atom
This is referred to as the Planetary Atomic
Model because they proposed that the
negatively-charged electrons stay "in orbit"
around the positively-charged nucleus in the
same way that the planets stay in orbit
around the sun.
The Quantum Theory and the
Hydrogen Atom
Energy is given off in quanta. Bohr
pointed out that the absorption of light
by hydrogen at definite wavelengths
corresponds to definite changes in the
energy of the electron.
He concluded that the orbits must have
orbits of definite diameter.
Mechanics
Newtonian Mechanics - describes visible
objects at ordinary velocities.
Quantum Mechanics - describes extremely
small particles at velocities near that of light
Modern Atomic Structure
E = mc2
E = h
 mc2 = h
mv2 = h
mv2 = hv/
  = hv/mv2 = h/mv
momentum = mv = p
 = h/p
Heisenberg (1927)
Heisenberg Uncertainty Principle
It is not possible to know precisely the
position of an electron and its momentum
(velocity) at the same instant.
Schrödinger
Developed a mathematical equation to
describe the wave-like behavior of the
electron.
The equation is very complicated (using
second partial derivatives)
 is
the wave function. The equation
related the amplitude of the wave
function () to any point in space
around the nucleus.
Max Born showed that ||2 gives the
probability of finding the electron at
the point in space for which the
equation was solved.
Einstein
Proposed that electromagnetic radiation
can be viewed as a steam of “particles”
called photons.
The energy of each photon is given by:
E = h = h(c/)
E = mc2
“Energy has mass”
m=E/c2 = (hc/)/c2 = h/c
m=mass of a photon of light with a
wavelength 
Conclusions
Energy is quantized
Electromagnetic radiation shows some
characteristics of matter
Light as a wave: (sine wave)
Light as a stream of photons:
         
Wave Mechanical Model of the Atom
Bohr’s model was based on classical physics
and was shown to be inadequate.
Mid-1920’s: a new approach was taken by de
Bröglie, Heisenberg and Schrödinger.
De Bröglie proposed that the electron, which
had been considered a particle only, also
showed wave properties.
Schrödinger attacked the problem by putting
emphasis on the wave properties.
Atomic Orbitals & Quantum Numbers
The quantum theory describes the behavior of
electrons. There are four quantum numbers
which are needed to describe the electron in an
atom (n, l, m, s). Remember, no two electrons in
an atom can have the same four quantum
numbers.
n, The Principal Quantum Number represents the
main energy level, its "size".
Can have values of 1, 2, 3, 4, …
l, Angular Momentum Quantum Number,
describes the "shape" of the orbital. These
are multiple energy states that are grouped
very close together. Can have values from 0
to n-1. The number of sublevels for energy
level "n" = n
n = 1 » 1 sublevel
n = 2 » 2 sublevels
n = 3 » 3 sublevels
n = 4 » 4 sublevels
m, Magnetic Quantum Number, describes
the orientation (direction) of the orbital
in space; that is, the direction in which it
points. Can have values from -l to +l
Degenerate orbitals are those orbitals
with the same size (n) and shape (l)
which have the same energy.
the three 2p orbitals
the five 3d orbitals
s, Electron Spin Quantum Number. Electrons can
spin either clockwise or counterclockwise.
s can have one of two values depending on the
direction of the rotation: +1/2 or -1/2
Quantum Number Overview
n - principal quantum number - size of energy
level
 values: 1, 2, 3, 4,…
l - energy sublevel - shape of the orbital
 values: 0 to n-1
m - orbital Q.N. - orientation in space (direction)
 values: - l to + l
s - spin Q.N.
 values: +1/2, -1/2
Rules for Filling Orbitals
Aufbau Principle - Build up the electrons from
the bottom
Pauli Exclusion Principle - No two electrons can
have the same set of four (4) quantum
numbers.
Hund's Rule - Add one electron to each orbital
of degenerate orbitals until all orbitals have at
least one electron. Then start pairing up the
remaining electrons.
The Apparent Contradiction
Waves can act as particles, and
particles can act as waves
Bohr’s atomic model explained light in
terms of particle properties.
Electrons (like light) have properties of
both waves and particles
Wave-particle duality of nature applies
to all waves and all particles
Electron Configuration
When we write the electron configuration for a
specific atom, we must specify the energy
level (principal quantum number, 1,2,3,...),
the sublevel (angular momentum quantum
number, s, p, d, f) and the number of
electrons in each sublevel (indicated via a
superscript).
Electron Configuration
For example, the electron configuration for
magnesium (which has 12 electrons) is:
1s22s22p63s2
When you add up all the exponents, you should
get the total number of electron for that
particular atom (in this case, 2 + 2 + 6 + 2 = 12)
Modern Atomic Structure
1. The division between matter and energy is
becoming even less clear.
2. de Bröglie Hypothesis (1923) led the way to
the present theory of atomic structure.
Electron Dot Diagrams
>>>Rules <<<
The elemental symbol represents the
nucleus and all electrons not in the
outer shell
Write out the electron configuration
(1s22s2…) selecting those electrons in
the outer energy level only
Each side represents an orbital. Draw
dots to represent electrons in that orbital
Quantum Theory
The quantum theory describes the behavior of the
electrons.
There are four quantum numbers needed to
describe the electron in an electron (n, l, m, s)
No two electrons can have the same four quantum
numbers
Principal Quantum Number, n
Represents the “size” of the energy level
(orbital)
Energy Sublevels, l
Describes the “shape” of the orbital
These are multiple energy states that are
grouped very close together
The number of sublevels (for each level) = n
n=1 
1 sublevel
n=2 
2 sublevels
n=3 
3 sublevels
n=4 
4 sublevels
Orbital Quantum Number, m
Describes the orientation of the orbital in
space; the direction in which it points.
Degenerate orbitals: those orbitals with
the same size (n) and shape (l) which
have the same energy. e.g.,
 the three 2p orbitals
 the five 3d orbitals
Electron Spin Quantum Number, s
Electrons can spin either clockwise or
counterclockwise
“s” will have one of two values depending on
the direction of rotation: +1/2 or -1/2
Distribution of Electrons
How are electrons distributed among the energy
levels?
In a neutral atom:
 #e-’s = # protons = atomic no.
Electrons always fill the energy level and
sublevel to produce the lowest energy
arrangement
No two electrons can have the same 4 quantum
numbers
 (Pauli Exclusion Principle)
The max. no. of e in energy level “n” = 2n2
Aufbau Principle
Build up the electrons from the bottom
 “The Aufbau Hotel”
Hund’s Rule
Add one electron to each orbital of degenerate
orbitals until all orbitals have at least one
electron. Then start pairing up the remaining
electrons.
Pauli Exclusion Principle
No two electrons can have the same set of
four (4) quantum numbers
Heisenberg Uncertainty Principle
It is not possible to know precisely the position
of an electron and its momentum (velocity)
at the same instant.