Download The Bohr model for the electrons

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.
2p, 3p, 4p etc.
3d, 4d etc.
4f, 5f etc.
2 electrons
6 electrons
10 electrons
14 electrons
Add electrons to the orbitals –
lowest first
4s
4p
3d
3s
3p
2s
2p
1s
H(z = 1)
Fill lowest orbital
4s
4p
3d
3s
3p
2s
2p
1s
He(z = 2)
Begin next orbital
4s
4p
3d
3s
3p
2s
2p
1s
Li(z = 3)
Fill 2s
4s
4p
3d
3s
3p
2s
2p
1s
Be(z = 4)
Begin filling 2p
4s
4p
3d
3s
3p
2s
2p
1s
B(z = 5)
Electrons don’t like to pair
4s
4p
3d
3s
3p
2s
2p
1s
C(z = 6)
4s
4p
3d
3s
3p
2s
2p
1s
O(z = 8)
4s
4p
3d
3s
3p
2s
2p
1s
F(z = 9)
Filled 2p – neon unreactive
4s
4p
3d
3s
3p
2s
2p
Ne
(z = 10)
1s
Shape of the periodic table
explained by orbital picture
2
groups
6
groups
10
groups
14
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
Filled
shell 1
Unfilled
shell 2
Filled
shell 1
Filled
shell 2
The periodic law and atomic size
Ionization energy and the periodic law
Ionization energy is energy required to remove electron
from the neutral atom