The Bohr model for the electrons

Electronic structure – how the electrons are arranged inside the atom

Applying the quantum principle of energy

Two parameters: – Energy

– Position

Learning objectives

Describe the basic principles of the Bohr model

Distinguish between the “classical” view and the “quantum” view of matter

Define atomic orbitals

Distinguish between the Bohr orbit and atomic orbital

Apply quantum numbers and atomic orbitals to building atoms and the periodic table

Describe periodic trends in terms of electronic structure

Bohr’s theory of the atom: applying

photons to electronic structure Electrons occupy specific levels (orbits) and no others

Orbits have energy and size

Larger orbits are at higher energy

Electron excited to higher level by absorbing photon

Electron relaxes to lower level by emitting photon

Photon energy exactly equals gap between levels

Size of energy gap determines

photon energy

Small energy gap, low

frequency, long

wavelength (red shift)

High energy gap, high

frequency, short

wavelength (blue shift)

The full spectrum of lines for H

Each set of lines in the H spectrum comes from transitions from all the higher levels to a particular level.

The lines in the visible are transitions to the second level

The Bohr orbits

Bohr orbits have quantum numbers n

– n = 1 (capacity 2)

– n = 2 (capacity 8)

– n = 3 (capacity 8)

Bohr orbits and the periodic table

Elements in the same group have the same

number of electrons in outer Bohr orbit

Successes and shortcomings of Bohr

Couldn’t explain why orbits were allowed

Only successful agreement with experiment was with the H atom

Introduced connection between spectra and electron structure

Concept of allowed orbits is developed further with new knowledge

Nonetheless, an important contribution, worthy of the Nobel prize

Electrons are waves too!

Life at the electron level is very different

Key to unlocking the low door to the secret garden of the atom lay in accepting the wave properties of electrons

De Broglie wave-particle duality

All particles have a wavelength – wavelike nature. – Significant only for very small particles – like electrons or

photons

– As mass increases, wavelength decreases

Electrons have wavelengths about the size of an atom – Electrons are used for studying matter – electron microscopy

Electron microscopes can peer within –

waves interacting with matter

Heisenberg Uncertainty Principle:

the illusive electron We can predict the motion of a ball;

But not an electron: problems locating small objects

The Quantum Mechanics: waves of

uncertainty

System developed that incorporated these concepts and produced an orbital picture of the electrons

No longer think of electrons as particles with precise location, but as waves which have probability of being in some region of the atom – the orbital

Impossible with the classical mechanics of Newton

Orbits become orbitals

The Orbitron: a gallery

of atomic orbitals and

molecular orbitals

Orbitals are described by quantum

numbers

Each orbital has unique set

1s, 2p, 3d etc.

Number describes energy

Letter describes shape

– S zero dimensions

– P one dimension

– D two dimensions

– F three dimensions

Getting from the orbitals to the

elements

All elements have the same set

Atomic number dictates how many are

filled – how many electrons are added

Filling orbitals follows a fixed pattern:

lowest energy ones first

Orbital energy levels in H and other

elements

How many per orbital?

Electrons share orbitals (only two allowed)

A consequence of “spin”

How many electrons can be added

to the orbitals

1s, 2s, 3s etc. 2 electrons

2p, 3p, 4p etc. 6 electrons

3d, 4d etc. 10 electrons

4f, 5f etc. 14 electrons

Add electrons to the orbitals –

lowest first

2p

3d

3p

4p 4s

3s

2s

1s

H(z = 1)

2p

3d

3p

4p 4s

3s

2s

1s

He(z = 2)

Fill lowest orbital

2p

3d

3p

4p 4s

3s

2s

1s

Li(z = 3)

Begin next orbital

2p

3d

3p

4p 4s

3s

2s

1s

Be(z = 4)

Fill 2s

2p

3d

3p

4p 4s

3s

2s

1s

B(z = 5)

Begin filling 2p

2p

3d

3p

4p 4s

3s

2s

1s

C(z = 6)

Electrons don’t like to pair

2p

3d

3p

4p 4s

3s

2s

1s

O(z = 8)

2p

3d

3p

4p 4s

3s

2s

1s

F(z = 9)

2p

3d

3p

4p 4s

3s

2s

1s

Ne(z = 10)

Filled 2p – neon unreactive

Shape of the periodic table

explained by orbital picture

2

groups 10

groups

14

groups

6

groups

Connecting the table with orbitals: elements

per row matches capacity of orbitals

Simplifying with shells:

echoes of Bohr orbits The orbitals with the same Principal Quantum

number (1,2,3 etc) are grouped into shells

Filled shells have special significance

Unfilled

shell 2 Filled

shell 2

Filled

shell 1 Filled

shell 1

The periodic law and atomic size

Ionization energy and the periodic law

Ionization energy is energy required to remove electron from the neutral atom