Transcript Atomic spectra and the Bohr atom

Electronic structure in atoms
AH Chemistry, Unit 1(a)
• Bohr’s theory cannot explain spectra for atoms
more complex than hydrogen, which show sublines.
• In 1926 Erwin Schrodinger devised a theory to
describe more complicated atoms: quantum
mechanics.
• This classifies regions of atoms into SHELLS, SUBSHELLS and ORBITALS.
Quantum numbers
• Each electron in an atom is described by 4
different quantum numbers:
n, l, ml, ms
• The first three (n, l and ml) describe an atomic
orbital – where an electron is found
– 3 numbers because three dimensions
• A fourth quantum number describes the spin of
an electron
Quantum Numbers 1 and 2
SHELLS
1. PRINCIPAL QUANTUM NUMBER (n): one on which
energy of an electron principally depends;
orbitals with same n are in same shell; also
determines size of an orbital; values = 1, 2, 3,…
Quantum Numbers 1 and 2
SUB-SHELLS
2. ANGULAR MOMENTUM QUANTUM NUMBER (l): energy of an
atom also depends to a small extent on l; distinguishes orbitals
of same n by giving them different shapes; any integer value
from 0 to n-1; orbitals of same n but different l are in different
sub-shells:
s p d f g
0 1 2 3 4
Example: 2p indicates shell 2 (energy), sub-level p (shape)
• Werner Heisenberg stated in 1927 that it is impossible to
know with precision both the position and the momentum
of an electron.
• The observer affects the observed.
• Only noticeable on the sub-atomic scale (e.g. an
electron, not a baseball, nor dust).
• It is not possible to define a point in space where an
electron will be found.
• However, we can obtain the probability of finding an
electron at a certain point.
An area where there is a greater than
90% chance of finding an electron =
atomic orbital
Quantum Number 3
ORBITALS
3. MAGNETIC QUANTUM NUMBER (ml):
distinguishes orbitals in same shell and subshell
(i.e. energy and shape) by giving them a
different orientation in space; values = -l to +l
s orbitals
p orbitals
d orbitals
n=3
l=2
-2
-1
0
l=1
-1
0
+1
3p
l=0
0
l=1
-1
+1
2p
l=0
0
+1
+2 3d
3s
0
n=2
2s
SHELL
n=1
l=0
0
1s
SUB-SHELL
ORBITAL
Pauli Exclusion Principle
• “No two electrons in any one atom can have
the same four quantum numbers”
• An orbital can hold at most two electrons, and
then only if the two electrons in that orbital
have opposite spins
Quantum number 4
4. SPIN QUANTUM NUMBER (ms): each orbital can
hold 2 electrons, and each has a different
direction of spin, -½ and +½
Hund’s Rule
• “Electrons fill orbitals singly and with
parallel spins before pairing occurs”.
Electronic configurations
Orbital notation:
Spectroscopic notation:
1s2
2s2
2p1
[He] 2s2 2p1
Outer electrons, control
chemical properties
Practise
• Write electronic configurations (obitalbox and spectroscopic) for:
–
–
–
–
–
–
Helium
Beryllium
Oxygen
Aluminium
Calcium
Bromine
Aufbau principle
• The Aufbau principle is a building up
principle, which helps you to write
electronic configurations
• “When electrons are placed in orbitals,
the energy levels are filled up in order of
increasing energy”
Filling orbitals
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f
Energy depends on n and l
Different orbitals in the
same sub-shell have the
same energy
Orbitals with the same
energy are called
degenerate,
There is interaction among
different sub-shells in
higher energy levels, which
lowers their energy and
explains the order of filling.