Post on 06-Nov-2020
transcript
Neils Bohr
• Niels Bohr (1913) – developed the “planetary”
model of the atom based upon the following:
– Rutherford’s Gold Foil Experiment
– E = mc2 – Albert Einstein (1905)
– Quantum Theory – Max Planck (1910)
• He postulated that the electrons were in
specific orbits about the nucleus.
• That the electrons were spinning so that
they would not crash into the nucleus.
• And he knew the model was very limited
and that it was going to be modified as
soon as he wrote it down!
• Bohr stated that the light must be from
energy given off from the element
• Different colors of light must be different
energy level transitions
• This means an element has specific
energy level transitions that it can give off
light
• Light can have only discrete amounts of
energy
– Energy is quantized (fixed levels like the steps
of a ladder or shelves of a bookshelf)
• Electrons can have only these values and
no others
• Similar to books on a shelf
– Can be on the first shelf or the second shelf,
but not in between
• Electrons “prefer” to be in the lowest
energy level
– levels closest to the nucleus
– Ground state
• Excited state
– electron goes from the lowest energy level to
a higher energy level when it absorbs energy
Ground State
Excited State
• Electrons cannot just jump to a higher state for no reason
• Something has to make them do it, otherwise they’d stay at the ground state
• If energy is put into the atom, the electron can take that energy and jump to another level
• This “taking in” of energy causes the absorption spectra, the releasing of energy causes emission spectra
• Bohr’s idea of the atom worked
well… for hydrogen
• Any other gas this was attempted
with, the spectra didn’t look like
they should have
• Needed something better
Neils Bohr
I pictured electrons orbiting the nucleus much like planets orbiting the sun.
But I was wrong! They’re more like bees around a hive.
WRONG!!!
• Rutherford said very little about them
• Neils Bohr said a lot!
• But we need to cover more before we
get to the Bohr Atom!
• So…. Back to Physics!
Equation for probability of an electron being found within a region of space
Erwin Schrodinger
E= H
• Schrödinger’s model:
probability of finding
electron in a given volume
– Orbitals
– Electron clouds
• Different shapes for
different types of orbitals
Orbital shapes are defined as the volume that contains 90% of the total electron probability. There are 4 Types of Orbitals, named s, p, d & f
An orbital is a region within an atom where there is a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level…
The s orbital has a spherical shape centered around
the origin of the three axes in space.
s orbital shape
There are three dumbbell-shaped p orbitals in
each energy level above the first, each assigned to its own axis (x, y and z) in space.
Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with the third energy level. To remember the shapes, think of “double dumbbells”
…and a “dumbbell
with a donut”! d orbital shapes
• We know where we might find the
electron, but…..
• Once we find it, it moves!
• Ok – anything else?
• What really matters to the Chemist?
• As it happens we are interested in the
Energy of the electron, not where it is.
You can find out where the electron is, but not where it is going.
OR…
You can find out where the electron is going, but not where it is!
“One cannot simultaneously determine both the position and momentum of an electron.”
Werner
Heisenberg
• Since Heisenberg demonstrated that you
cannot know both the energy and the
position of the electron,
• Chemists concentrate on the energy of
the electron – and according to Bohr
• That means we need to know the energy
level the electron occupies.
• This gives rise to:
• Electron Configurations
or
• Orbital Notation
• Aufbau Principle - The electron that
distinguishes an element from the
previous element enters the lowest
energy atomic orbital available.
• Or: electrons fill up the orbitals from
the bottom up… lowest energy to
highest energy
• Orbital Notation for carbon
• 1s 2s 2p
• Electron configuration for carbon
• element #6
• C - 1s2 2s2 2p2
1s 2s 2p
• Electrons fill sublevels of an orbital singly before
they spin pair.
1s 2s 2p
Nitrogen
• An Orbital can hold a maximum of 2 electrons –
but those electrons must have opposite spins.
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
• Get out your Periodic Table!
• Determine the energy levels used
• Determine the orbital type
• Determine the number of electrons in each orbital
• Continue to fill each higher level until all electrons are accounted for
The Orbitals Being Filled for Elements in
Various Parts of the Periodic Table
Modern View of Atom
From past to present