Transcript Modern Atomic Theory

History of Atomic Theory
Atomic models from
Dalton to Bohr
A ‘Model’…
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is not a real thing, but is used to explain, mimic
or simulate reality,
is used as a tool,
is used to predict what happens in the real
world,
is changed or modified until it best fits new
information,
may have some limitations or be valid only
under certain conditions.
Examples: globes, computer simulations,
product prototypes
Historical Models of the Atom
John Dalton
‘Billiard ball’ model (1803)
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All matter consists of atoms
Each element has its own atom type
Atoms of different elements have different properties
Atoms of two elements can combine to form compounds
Atoms are never created, destroyed or subdivided
Historical Models of the Atom
J. J. Thomson
‘Raisin bun’ model (1897)
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First to include sub-atomic particles
(electrons) that had been seen in cathode
ray tube experiments
Model is of a positively charged sphere with negatively
charged electrons embedded in it
Positive ‘dough’ and negative ‘raisins’ make up an
atom that is neutral over all
Historical Models of the Atom
Ernest Rutherford
Nuclear (‘beehive’) model (1911)
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Tested Thomson’s theory with the famous
“gold foil experiment”
His results suggested that the atom was mostly empty
space with a very dense positively-charged ‘nucleus;’
he later discovered that protons were the positivelycharged part
In this model, negatively-charged electrons existed
within the empty space
**Neutrons were discovered much later by
James Chadwick (in 1932); why so late?
Historical Models of the Atom
Niels Bohr
Refined nuclear model (1913)
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Bohr knew that a new model was needed, primarily
because of a major problem with the Rutherford model:
Bohr and his contemporaries knew that if a charged particle
accelerates, it must give off energy, likely in the form of light.
The electrons, which are definitely charged (-) and accelerating
(changing direction constantly) should therefore give off energy
and eventually spiral in towards the nucleus and cause the
atom to collapse. Problem? Yes: Atoms don’t collapse!
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Bohr was very interested in the newly-developed
quantum theory of light proposed by Einstein and
Planck, and thought it could be applied to the problem
with Rutherford’s model…so…
more about quantum theory next, then back to Bohr…
The Quantum Model
Bohr’s Inspiration for a Better
Model of the Atom…but still not
the best…
Light defined as a Wave
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Light travels through space as an electromagnetic wave.
Waves are characterized by their wavelength, λ, and
frequency, f, and amplitude, A.
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Light colour is related to the wavelength (and frequency)
of a wave. Red light has a longer wavelength and lower
frequency than blue light.
Relationship Between Wavelength
and Colour of Light
A spectrum containing all colours of visible light is called a
continuous spectrum. This is what we see if we pass white light
through a prism.
The Quantum Model
Quantum – a specific allowable value.
Quanta – a set of specific allowable values.
Example – A staircase is like a set of quanta.
Each stair is an allowable position. Other
than when travelling between steps, an
individual step is the only place you may
exist. Compare to a ramp.
Origins of Quantum Theory
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Max Planck first hypothesized that the energy of an
oscillating atom was not continuous (or wavelike) when
he studied blackbody radiation
Albert Einstein stated that if the energy of the vibrating
atoms was quantized, the light they emit must also be
quantized
Einstein earned a Nobel Prize when he used this new
‘Quantum Theory of Light’ to explain the photoelectric
effect; one quantum of light (called a ‘photon’) could
release one electron from a metal surface; higher-energy
photons were more likely to liberate electrons than lowenergy photons
Light Defined as a Particle
(quantum model)
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Light is also thought to propagate through space
as individual particles called photons.
Each photon has a specific amount of energy that
is related to the wave characteristics and to the
colour of light, that is, blue photons have more
energy than red photons
Atoms and Light
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An element in a gaseous state produces light
when it is heated to a certain temperature
By passing this light through a prism, we can
see its ‘bright-line’ or ‘emission’ spectrum
Bohr wanted his atomic model to explain
the bright-line spectra of the elements
 Since only certain distinct colours of light
could be absorbed or emitted by atoms,
Bohr reasoned that this related to distinct
‘energies’ of the electrons inside the atom
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Conclusion: Electrons have distinct
energies, and are therefore ‘quantized’
And now we can finish the story…
Niels Bohr
Refined nuclear model (1913)
a.k.a.“Bohr-Rutherford Model”
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Nucleus containing protons (+) ((and neutrons)),
Electrons (-) are organized into specific energy levels orbiting the
nucleus, and are thereby ‘quantized.’
Bohr’s model allows the electrons to exist in specific allowable energy
levels that are identified by the principle quantum number, n. The
allowable values of n are 1,2,3, …
Electrons follow “occupancy rules”
Although technically ‘historical,’ this model is very useful.
It is still used daily by students and scientists alike.
Bohr-Rutherford Model
Electron occupancy rules:
2n2
n = energy level
* Electrons will always occupy the lowest energy level available. *
What Bohr proposed…
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Electrons are arranged in fixed energy states.
When an element is heated, electrons are promoted to
higher energy states (excited states). When electrons
return to a lower energy state, energy is given off in the
form of light (a photon is emitted).
The movement of an electron between energy levels is
referred to as a transition.
The type (colour) of light emitted is related to the size of
the transition. Many transitions produce photons that are
not in the visible region.
It is important to note that Bohr primarily studied
hydrogen.
Electron transitions of Hydrogen
Usefulness of the Bohr-Rutherford Model
1.
Periodic Trends
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Valence electrons = group # (A groups)
Common ion charges (A groups)
Ionization energy
Stability of Noble Gases and trends with successive
ionization energies.
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Elements in group 1 have an unusually high 2nd I.E ; those in
group 2 have an unusually high 3rd I.E. etc.. The B-R model
explains this by suggesting that once an element has
achieved an octet, it is in a stable arrangement that matches
a noble gas.
Atomic radii
Reactivity
Usefulness of the Bohr-Rutherford Model
Stability of Noble Gases; trends with successive ionization energies
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Ionization Energy (kJ/mol)
Elemen
t
1st
2nd
3rd
4th
5th
6th
7th
8th
H
1 313
He
2 374
5 251
Li
521
7 297
11 814
Be
898
1 757
14 855
21 013
B
801
2 423
3 658
25 028
32 827
C
1 091
2 355
4 623
6 226
37 836
47 276
N
1 400
2 857
4 575
7 480
9 449
53 270
64 360
O
1 313
3 388
5 299
7 470
10 994
13 328
71 339
83 600
F
1 679
3 378
6 042
8 416
11 023
15 163
17 866
92 042
Ne
2 085
3 967
6 177
9 382
12 200
15 241
Na
492
4 565
6 920
9 546
13 378
16 640
20 115
Mg
734
1 448
7 731
10 550
13 629
18 040
21 747
25 675
Al
579
1 815
2 741
11 583
14 845
18 378
23 348
27 518
Si
781
1 573
3 223
4 353
16 090
19 796
23 783
29 333
P
1 062
1 902
2 915
4 961
6 274
21 273
25 414
29 854
S
1 004
2 259
3 378
4 565
6 998
8 494
27 122
31 736
Cl
1 255
2 297
3 851
5 164
6 544
9 334
11 032
33 618
Ar
1 525
2 664
3 948
5 772
7 391
8 812
11 969
13 811
9th
10th
105 754
130 834
Usefulness of the Bohr-Rutherford Model
2. Predictions For Compounds
 Bonding ratios (MgF2, CaF2, SrF2, etc.)
 Bond polarity (electronegativity)
 Bonding types – covalent vs. ionic
3. Physical Properties
 Solubility
 Melting and boiling points
 Viscosity