Transcript Arrangement of Electrons in Atoms The Development of a New Atomic Model The Development of the Atom  Dalton’s model     Characteristics Dalton’s atomic theory Modifications to Dalton’s atomic.

Arrangement of Electrons in
Atoms
The Development of a New
Atomic Model
The Development of the Atom
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Dalton’s model
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Characteristics
Dalton’s atomic theory
Modifications to Dalton’s atomic theory
Thomson’s model
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Characteristics
Grapes in jello
Plum pudding model
The Development of the Atom
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Rutherford’s model
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Gold foil experiment
Characteristics
Problems
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What was the exact location of electrons?
Why did electrons not fall into the nucleus?
Properties of Light
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The study of light revealed the relationship
between it and an atom’s electrons.
Light is a form of electromagnetic
radiation(EMR).
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It is a form of energy.
Other forms of EMR are gamma rays, X-rays,
microwaves, ultraviolet rays, infrared rays, and
radio waves.
Speed is 3 x 108 m/s(c) in a vacuum.
It has wave-like features.
The Electromagnetic Spectrum
Wave Features
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Repetitive nature.
Wavelength – distance from crest to
crest(m, cm, nm)(l)
Frequency – how many crests pass a
point in one second(cps, Hz)(f or n)
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1cps = 1Hertz
Crest - peak
Trough – valley
Wavelength and Frequency
The Photoelectric Effect
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The photoelectric effect refers to the emission
of electrons from a metal when light shines
on the metal.
The effect depends on the frequency of the
light striking the metal. If the frequency is
too low, nothing happens regardless of time
of exposure or intensity. A certain minimum
frequency was required.
This was a problem because classical wave
theory physics predicted that any frequency
of light would cause the effect.
The Photoelectric Effect
The Particle Description
of Light
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Planck was studying the emission of light by
hot objects(black body radiators).
He proposed that such objects emitted
energy in small, specific packets of
energy(quanta).
His relationship was E = hn or hf, where E is
the energy in joules, h is Planck’s
constant(6.626x 10-34 Js) and n or f is
frequency.
The Particle Description
of Light
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Einstein expanded on Planck’s ideas and
introduced the idea that EMR had a
dual wave-particle nature.
According to him, light could travel in
both wave motion and particle motion.
He called particles of light photons.
Each photon carried a quantum of
energy.
Einstein explains the
photoelectric effect
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Einstein proposed that EMR is absorbed by
matter only in whole numbers of photons. In
order for an electron to be ejected from a metal
surface,the electron must be struck by a single
photon possessing the minimum energy required
to knock the electron loose. This energy
corresponds to a threshhold frequency. If the
frequency is too low, nothing happens. If the
frequency is too high, more energetic electrons
are emitted. If the intensity of the light is
increased at the threshhold frequency, more
electrons are emitted.
The Hydrogen Atom Line
Emission Spectrum
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When current is passed through a gas at low
pressure, the potential energy of some gas
atoms increases.
The lowest energy state of an electron is the
ground state.
A state in which an atom has a higher potential
energy than it has in its ground state is an
excited state.
When an electron returns to the ground state, it
emits absorbed energy in the form of light.
When this light was passed through a prism, a
bright line emission spectrum was received.
Hydrogen’s Line Emission
Spectrum(Balmer series lines)
In other areas of the
EMR spectrum, the
Lyman series lines(UV)
and Paschen series
lines(IR) are seen in
the spectrum.
Hydrogen Atom Line Emission
Spectrum
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Classical physics predicted that the hydrogen
atoms would be excited by any amount of
energy and that their electrons could be in any
energy state.
If that was true, then a continuous
spectrum(rainbow) would be seen.
As shown, only specific frequencies of light are
given off.
This was even the case in non-visible spectra
from the infrared and ultraviolet regions of the
EMR spectrum.
Bohr’s model of the
hydrogen atom
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In 1913, Bohr proposed a model of the
hydrogen atom linking the atom’s
electron with photon emission.
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The electron can circle the nucleus in
allowed orbits only. It has fixed energy.
It has lowest energy closest to the nucleus.
As it moves away from the nucleus, it gets
higher in energy.
There is a common analogy between
Bohr’s atom and the rungs of a ladder.
What really happens when energy
is absorbed and re-emitted?
The Quantum Model of the
Atom
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New developments – De Broglie
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Heisenberg Uncertainty Principle
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De Broglie wavelength l = h/mv
It is impossible to know both the location and the
velocity of an electron or any other particle.
Schrodinger’s model
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Quantum mechanical model
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Considers probability of finding electron in a position.
Orbitals are what electrons travel in.
Energy sublevels correspond to orbitals.
Describes mathematically the wave properties of
electrons and other very small particles.
Atomic Orbitals and Quantum
Numbers
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Energy levels designated by the principal
quantum number, n, where n may be 1, 2, 3,
4, ………
Energy sublevels correspond to different
cloud shapes or regions of high electron
probability. These are called atomic orbitals.
They are designated by the azimuthal (or
angular momentum) quantum number, l,
where l may be 0, 1, 2, 3,…. n-1.
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Example: If n = 3, then l may be 0, 1, 2.
Atomic Orbitals
The “s” orbital is
spherical. All
main levels have
an “s” orbital.
“p” orbitals first appear in the 2nd main
level. There are 3 of them, px, py, and
pz. They are perpendicular to one
another.
Atomic Orbitals
“d” orbitals first appear in the 3rd main level. There are 5 of
them, dxz, dyz, dxy, dx2-y2, and dz2.
Atomic Orbitals
“f” orbitals first
appear in the 4th
main level. There
are 7 of them. They
are much more
complex than the
previous three types
of orbitals.
Atomic Orbitals
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The orbital is designated by the magnetic
quantum number, ml, which describes the
orientation around the nucleus. It has values
from –l to 0 to +l.
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Example: If l = 2, then ml can have the values -2,
-1, 0, 1, 2.
There may be only 2 electrons in an orbital
and they must be spinning in opposite
directions according to Pauli Exclusion
Principle. Thus the spin quantum number, ms,
indicates this with values of –1/2 or +1/2.
Example question about
quantum numbers
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What is the set of quantum numbers
associated with a 3d electron?
n= 3
l=2
ml = -2, -1, 0, 1, 2
ms = -1/2, +1/2
Important Points about Atomic
Orbitals
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The lowest energy level(closest to the
nucleus) has only 1 sublevel and in that
sublevel(“s”) is only one orbital.
The second energy level has 2 sublevels(“s”
and “p”), the first of which has one orbital
and the second has 3 degenerate orbitals.
The third energy level has 3 sublevels(“s”,
“p”, and “d”) having 1, 3, and 5 orbitals
respectively.
Important Points about Atomic
Orbitals
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Remember that each atomic orbital can have
only 2 electrons.
The maximum number of electrons in any
main energy level is 2n2.
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Example: If n = 3, then the maximum number of
electrons that can be placed there is 2(32) = 18.
The maximum number of orbitals in a given
main energy level is n2.
Summary
Main
Sublevel
Max e
Orbitals
1
s
2
2
1
1
2
s
p
2
6
8
1
3
4
3
s
P
d
2
6
10
18
1
3
5
9
4
5