Transcript Electron Orbital

-The Bohr Model
-The Quantum
Mechanical Model
Warner SCH4U Chemistry
Dalton’s Atomic Model
Plum Pudding Model (Thomson)
Niels Bohr
(Born in Denmark 1885-1962)

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Student of Rutherford
Through his study of
light, Bohr came up
with a theory of how
electrons were
arranged within an
atom.
What is Light?
Light: Electromagnetic radiation that
travels through space or matter in wave
like oscillations.
 Sometimes we refer to it as a wave and
sometimes we refer to it as a particle
(photon)

Photon

Bundles (package) of light energy
Characteristics of Light –
frequency, wavelength, energy
Max Planck (1858-1947)
Planck proposed that photons of light
carry energy and devised E=hn
E=energy
n=frequency
h=Planck’s constant 6.7x10-34Js
Quantum – term first used by
Planck
A
quantum is the Latin word for
discrete amount of energy
Niels Bohr’s Model (1913)
 Based
on his
study of light he
stated that
electrons orbit
the nucleus in
circular paths of
fixed energy
(energy levels).
Niels Bohr’s Atom Cont’d
 Electrons
can jump from energy level
to energy level.
 Electrons absorb or emit light energy
when they jump from one energy level
to another.
 Quantum jump is amount of energy
required to move an electron from one
energy level to another.
Excited State and Ground State

Ground state: the lowest possible energy
level an electron be at.

Excited state: an energy level higher than
the ground state.
Colour of light
Energy

Each electron that jumps back emits one
photon of light
What colour is this light?
 Depends on how big the jump between
orbits was
 The bigger the jump, the higher the
energy.
Energy of Emitted Photon
Energy of the emitted photon =
Difference in energy between two states

Energy emitted by the electron as it falls
back from the higher to the lower energy
level is proportional to the frequency of the
light wave.
The energy levels are like the rungs
of a ladder but are not equally
spaced.
Excited states are unstable. Electrons quickly falls back to
the ground state, but not always in a single step. For
example, if the electron is initially promoted to the n=3
state, it can decay either to the ground state or to the n=2
state, which then decays to n=1.
Emission Spectrum

Light emitted produces a unique
emission spectrum.
Hydrogen Emission Spectrum
Violet
Blue
Red
Balmer
Series
Flame Test Colour

Colour seen is a result of different
wavelengths of light (colours) emitted when
the electrons go down the step(s) to their
ground state.

Each element will have its own set of steps,
therefore each will have its own colour.

Some colours are very similar so a more
exact method can be used to identify the
elements.
23
Bohr Model for Hydrogen
 The
Bohr model explained the
emission spectrum of the hydrogen
atom but did not always explain those
of other elements.
Quantum Mechanical Model
1920’s
 Werner Heisenberg (Uncertainty Principle)
 Louis de Broglie (electron has wave
properties)
 Erwin Schrodinger (mathematical equations
using probability, quantum numbers)

Werner Heisenberg: Uncertainty Principle

We can not know both
the position and
momentum of a
particle at a given time.
The Wave Model


Uncertainty Principle makes it impossible to plot an
orbit for an electron. Is this a problem? NO
The probable location of an electron is based on how
much energy the electron has.
Louis de Broglie, (France, 1892-1987)
Wave Properties of Matter (1923)
Since
light waves have a
particle behavior (as shown
by Einstein in the
Photoelectric Effect), then
particles could have a wave
behavior.
de Broglie wavelength
l= h
mv
Electron Motion Around Atom
Shown as a de Broglie Wave
What is an electron?

https://www.youtube.com/watch?v=O55XiriEaQI#t=11
1. The electron was treated as a wave
 2. Light was treated as a particle (photon)
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What is an electron? Is it a particle or a
wave? And what is light? A wave or a
photon?
The Electrons: wave or particle
Electrons display properties of both.
 To think of them as a particle is easy
because they have a small amount of mass.
 There is evidence of wave behavior though.
 In this sense they are neither particles nor
waves in the absolute sense, but only
exhibit wave or particle properties,
depending on the experiment being
performed.

Erwin Schrodinger, 1925
Quantum (wave) Mechanical Model
of the Atom

Four quantum
numbers are required
to describe the state of
the hydrogen atom.
FYI: Schrodinger’s Equations!!!

y is called the wave function and indicates
the probability of where an electron may
be found.

According to the modern atomic model,
an atom has a small positively charged
nucleus surrounded by a large region in
which there are enough electrons to make
an atom neutral.

We can plot the position of an electron
over and over again and gradually build a
3D map of the places that the electron is
likely to be.

These defined regions are called orbitals
Atomic Orbital:
A region in space in which there is high probability of
finding an electron.

Electrons whirl about the nucleus billions of times in one
second

They are moving around in random patterns.
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Location of electrons depends upon how much energy the
electron has.
Quantum-Mechanical Model
Electron Orbital:

Depending on their energy, electrons are locked into a
certain area in the cloud.

Electrons with the lowest energy are found in the
energy level closest to the nucleus

Electrons with the highest energy are found in the
outermost energy levels, farther from the nucleus.
The Electron Cloud
 The
electron cloud represents
positions where there is probability of
finding an electron.
The Electron Cloud
The higher the
electron density,
the higher the
probability that
an electron may
be found in that
region.
http://www.chemeng.uiuc.edu/~alkgrp/mo/gk12/quantum/H_S_orbital.jpg
Quantum Mechanical Model

Electrons are located in specific energy
levels.

There is no exact path around the nucleus.
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The model estimates the probability of
finding an electron in a certain position.
Quantum Numbers:
specify the properties of atomic orbitals
and their electrons.
Four Quantum Numbers
1.
2.
3.
4.
Principal Quantum Number
Orbital Quantum Number
Magnetic Quantum Number
Spin Quantum Number
Principal Quantum Number, n

Indicates main energy levels
n = 1, 2, 3, 4…

Each main energy level has sub-levels
The maximum number of electrons
in a principal energy level is given
by:
Max # electrons = 2n2
n= the principal quantum number
Orbital Quantum Number, ℓ
(Angular Momentum Quantum Number)

Indicates shape of orbital sublevels
 ℓ = n-1
ℓ
0
1
2
3
4
sublevel
s
p
d
f
g
Atomic Orbital s
2s
The 3 p orbitals
http://www.rmutphysics.com/CHARUD/scibook/crystal-structure/porbital.gif
 The
d orbitals
f orbitals
Magnetic Quantum Number, ml
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Indicates the orientation of the orbital in space.
Values of ml : integers -l to l
The number of values represents the number of
orbitals.
Example:
for l= 2, ml = -2, -1, 0, +1, +2
Which sublevel does this represent?
Answer: d
Electron Spin Quantum Number, (ms or s)
Indicates the spin of the electron
(clockwise or counterclockwise).
 Values of ms: +1/2, -1/2
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Example:
List the values of the four quantum
numbers for orbitals in the 3d sublevel.
 Answer:
n=3
l=2
ml = -2,-1, 0, +1, +2
ms = +1/2, -1/2 for each pair of electrons
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