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