Transcript L2x

Recap – Last Lecture
• To balance a nuclear equation check the sum of the
charge (subscript) and mass (superscript) of reactants
equals those of the products.
eg 1 n
0
+
235
92
U

141
56
Ba +
92
36
1
0
Kr + 3 n
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Structure of Atoms – History
Atomic Theory
http://www.humantouchofchemistry.com/john-dalton.htm
1808 J Dalton
All matter consists of atoms
that cannot be created or
destroyed.
Atoms of one element cannot
be converted into atoms of
another element.
Atoms of an element are
identical and are different from
atoms of any other element.
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Structure of Atoms - History
1897 J J Thomson Cathode Rays
• Negatively charged particles.
• All metals produced the same particles.
• ~ 1000 times lighter than a hydrogen atom. Atoms
are divisible!
• Cathode rays were later renamed electrons.
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Structure of Atoms – History
1909 E Rutherford
Nucleus of atom
• Atoms are mostly empty
space occupied by
electrons.
• All the positive charge
and essentially all the
mass lies in the nucleus.
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Structure of Atoms – History
N Bohr
Electrons in orbits
• Electrons in atoms can only
occupy certain energy levels
(orbits).
commons.wikimedia.org/wiki/File:Niels_Bohr.jpg
1909
• When an electron moves from
one energy level to another,
energy is absorbed or emitted.
• This energy corresponds to light
of a specific energy/frequency.
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Silberberg fig 7.3
Electromagnetic Radiation
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Electromagnetic Radiation
Frequency, , nu
• The number of wave
crests passing a given
point per unit time.
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Blackman fig 4.1
Wavelength, , lambda
The distance between
two adjacent identical
points of the wave.
Electromagnetic Radiation
• All light waves travel at exactly the same
speed (in a vacuum) – the speed of light,
c, is a constant.
C = 2.998  108 ms-1
• Wavelength and frequency are related to
the speed of light.
c = 
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Electromagnetic Radiation
• All radiation may have the same speed
but the energy can vary.
• The higher the frequency, the more
rapidly the wave is oscillating and the
higher the energy.
Energy = Planck’s constant  frequency
E = h
h = 6.626 x 10-34 Js
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Electromagnetic Radiation
Large wavelength
Low frequency
Low energy
Short wavelength
High frequency
High energy 10
Converting wavelength  energy
eg: A radio station transmits at a wavelength of 2.84 m.
Calculate the frequency.
 = c so
=c/
= 3.00 x 108 ms-1 / 2.84 m
= 1.056 x 108 s-1 or 106 MHz
(this is the radio station 2JJJ)
Calculate the energy associated with this radiation.
E = h
= 6.626 x 10-34 J s x 1.056 x 108 s-1
= 7.00 x 10-26 J
and for one mole of radiation
E = 7.00 x 10-26 J x 6.022 x 1023 mol-1
= 0.0421 J mol11-1
Cu
Pass the light
through a prism
to see the lines:
commons.wikimedia.org/wiki/File:Flametest--Na.swn.jpg
Co
commons.wikimedia.org/wiki/File:Flametest--Cu.swn.jpg
Our eye sees the
combination of
wavelengths:
commons.wikimedia.org/wiki/File:Flametest-Co-.swn.jpg
Atomic Emission Spectra
Na
Co
Cu
Na
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http://chemistry.bd.psu.edu/jircitano/periodic4.html
Atomic Emission Spectrum
• Only light of certain energies is emitted.
• The pattern of lines is unique to hydrogen.
• Suggests the process of emitting light from the atom
is quantised (comes in discrete amounts).
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The Bohr Model
Energy
n=6
n=5
n=4
n=3
n=2
Silberberg Fig 7.11
n=1
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Theory & Experiment agree for H
Energy of the hydrogen atom orbits is inversely
proportional to the square of the orbit number:
E = - ER (1 /n 2)Z 2
As
then
E
E
n=6→n=2 n=5→n=2
ER = 2.18 x 10-18 J
Z = atomic number
=
Efinal - Einitial
= - 2.18 x 10-18 J (1/n
n=4→n=2
2
final
- 1/n
2
initial)
Z2
n=3→n=2
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Theory & Experiment agree for H
eg: Calculate the wavelength of light emitted when an
electron moves from the n = 3 to the n = 2 orbit of a
hydrogen atom.
E = - 2.18 x 10-18 J (1 / 22 - 1 / 32) (1)2
= - 3.03 x 10-19 J
(minus indicates light emitted)
Now E = h and E = hc / 
So  = hc / E
= (6.626 x 10-34 Js) (3.00 x 108 ms-1) / (3.03 x 10-19
J)
-7 m or 656 nm (red light)
=n=5→n=2
6.56 x 10
n=6→n=2
n=4→n=2
n=3→n=2
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Applications
• Na+ is the major ion in extracellular fluid. The body
requires 1-2 mmol/day of Na+ while a typical daily diet
contains 100 times this amount. The excess is excreted
by the kidneys.
• Commonly, atomic absorption spectroscopy is used to
determine sodium ion concentration in which the line in
the spectra used for this measurement is at 589 nm
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Learning Outcomes:
• By the end of this lecture, you should:
− Recognise the historical context of the Bohr
model of the atom.
− Be able convert between the wavelength,
frequency and energy of light.
− Be able to calculate the energy of a hydrogen
orbit.
− Be able to calculate the atomic emission
spectrum of a hydrogen atom.
− be able to complete the worksheet (if you
haven’t already done so…).
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Questions to complete for next lecture:
1. Calculate the energy of the light absorbed when an
electron in a hydrogen atom moves from the n=1 to n=3
orbit.
2. What is the wavelength of the light calculated in Q1?
3. To what region of the electromagnetic spectrum does the
light in Q1 belong?
Type of Light
Wavelength range
Ultra-violet
10 – 400 nm
Visible
400 – 700 nm
Infra-red
700 nm – 1 mm
4. Is the light in Q1 more or less energetic than the light
associated with the transition between n=2 to n=3?
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