Download Electrons in Atoms

Chapter 5
From Democritus to Rutherford, models of the
atom have changed due to new experiments.
 As technology develops, a more complete model
of the atom is developed.
 Rutherford’s model identified the nucleus
surrounded by electrons.
 His model DID NOT explain why some things
glow when heated.
 His model DID NOT explain the chemical
properties of elements.
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Niels Bohr (1885-1962) was one of
Rutherford’s students. He added to the
model.
Proposed that an electron is found only in
specific circular paths, or orbits, around the
nucleus.
Studied how energy of an atom changes
when it absorbs or emits light.
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Energy levels are fixed “paths” or energies an
electron can have when orbiting the nucleus.
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Close to nucleus = less energy
Further from nucleus = more energy
e- must reside on an energy level
Moving from one level to another is possible if the
right energy is lost or gained.
A quantum of energy is the amount of
energy required to move an electron from
one level to another.
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Energy gained or lost in an atom is
not always the same.
Energy levels are not evenly
spaced.
 Higher levels are closer together.
Therefore less energy needed to
move levels further from nucleus.
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Erwin Schrodinger (1887-1961) devised and
solved a mathematical equation describing
the behavior of the electron in a hydrogen
atom.
This model comes from his mathematical
solutions.
Determines the allowed energies an electron
can have and how likely it is to find the
electron in various locations around the
nucleus.
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Based on his equations, Schrodinger was also
able to explain atomic orbitals.
Orbitals explain the probability of finding an
electron at various locations around the
nucleus.
Often thought of as the region of space in
which there is the highest probability of
finding an electron.
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Used to describe the region of space with the
highest probability of finding an electron.
 Energy level (n)
▪ n = 1,2,3,4,5,6,7
 Sublevels correspond to an orbital of a different shape
▪ Each energy level has an equal number of sublevels
▪ Denoted by letters (s,p,d,f)
 Orbital
▪ Describes highest probability
▪ Contained within sublevels
▪ Shapes describe probability
Principal
Energy Level
Number of
sublevels
Type of
Sublevel
Orbitals
n=1
1
1s
1
n=2
2
2s
2p
1
3
n=3
3
3s
3p
3d
1
3
5
n=4
4
4s
4p
4d
4f
1
3
5
7
Each orbital is associated with a different shape. The shapes are
representations of mathematical probability rather than actual shapes of the
orbitals.
Click here to
see a model
of the shapes!
And here for a visual
representation of atomic
orbitals.
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Understanding Electron Arrangements (econfigurations) is the key to understanding
the chemical properties of elements.
As scientists began to determine these
configurations, they realized that energy and
stability play a huge part in how electrons are
configured.
We, in chemistry, need to gain a basic
understanding in order to proceed.
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Aufbau Principle: electrons occupy the orbitals
of lowest energy first. The boxes or the econfiguration chart will tell you the order.
Pauli Exclusion Principle: An atomic orbital
may describe at most two e- with opposite spin.
Hund’s Rule: When occupying a sublevel where
all orbitals have equal energy, orbitals must all
fill with one electron before any orbital contains
two electrons.
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Up to Vanadium (atomic # 23) the Aufbau
diagram will work.
Elements like chromium and copper are more
stable when they have half or completely
filled d sublevels.
Some e- configurations differ from those
assigned using Aufbau principle because halffilled sublevels are not as stable as filled
sublevels.