2. More history…
• Niels Bohr - Developed a model based on
observation that energized electrons appear to
“jump” from one energy level to the next, and
release a discrete amount of energy (quanta) when
returning to lower levels.
3. • Werner Heisenberg - Uncertainty principle: The more
precisely the position of an electron is determined,
the less precisely it’s momentum is known and vice
versa.
4.
5. • Erwin Schrödinger – Contributed
his famous wave equation used
to predict the shapes of atomic
orbitals.
http://youtu.be/uWMTOrux0LM
6.
7. A New Model of the Atom
The Quantum Mechanical Model
• Quantum mechanics provides a mathematical description
of much of the dual particle-like and wave-like behavior
and interactions of energy and matter
• In the 1920’s Heisenberg and Schrodinger used quantum
mechanics to describe the allowed energies an electron
can have and how likely it is to find the electron in various
locations around the nucleus.
8. Atomic Orbitals and Energy Levels
• Atomic orbitals represent location
around the nucleus that an electron
has the greatest probability of
existing
Electron density distribution for H
• Orbital arrangement follows a predictable pattern
on the periodic table grouped by energy levels.
9. Atomic Orbitals and Energy Levels
• Energy level number (n) corresponds the the distance from the
nucleus
• n=1, n=2, n=3 … (n=1 is closest to nucleus and lowest energy)
• Within the energy levels there are sublevels containing orbitals
• Each orbital can hold up to two electrons
10. Atomic Orbitals and Energy Levels
• Definitions:
• Shell – The set of orbitals having the same n-value
• Ex. The third shell contains 3s, 3p & 3d orbitals
• Sublevel (subshell) – The set of orbitals of the same type
• Ex. The five 3d orbitals in the third shell are a sublevel
11. Types of Atomic Orbitals
• Lowest energy orbitals are the s orbitals
• Spherical shape around the nucleus
• Increase in size as n increases
• One s orbital found at each sublevel (2e-)
12. • Areas between orbitals are called nodes.
• Nodes are regions of space where the probability of
finding electrons is zero
13. Types of Atomic Orbitals
• Second lowest energy orbitals are the p orbitals
• Aligned along perpendicular axes
• Increase in size as n increases
• Three p orbitals found at each sublevel (6e-)
14. Types of Atomic Orbitals
• Third lowest energy orbitals are the d orbitals
• Complicated orbital shape
• Increase in size as n increases
• Five d orbitals found at each sublevel (10e-)
15. Types of Atomic Orbitals
• Highest energy orbitals are the f orbitals
• Complicated orbital shape
• Increase in size as n increases
• Seven f orbitals found at each sublevel (14e-)
16. Electron Configuration
Rows in the periodic table correspond to the different energy
levels and elements are grouped based on the type of
orbitals their valence electrons are stored in
17. Electron Configuration
• Three guiding rules:
1. Aufbau Principle “build up”
2. Hunds Rule “empty bus seat rule”
3. Pauli Exclusion Principle “opposite spins”
• Remember:
• Two electrons per orbital
• 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals
• For neutral atoms: Atomic number = protons = electrons
18. Electron Configuration
1. Aufbau Principle
• Build Up
• When filling atomic orbitals with electrons, the
lowest energy orbitals are filled first.
19. Electron Configuration Standard Notation
• Ex. Write the electron configuration of Nitrogen
• Nitrogen has 7 electrons: 2 electrons in the 1s orbital, 2
electrons in the 2s orbital, 3 electrons in the 2p orbitals
N 1s2 2s2 2p3
20. Electron Configuration
2. Hund’s Rule
• Empty bus seats fill first
• For same energy sublevels one electron is placed in
each orbital before doubling up
• Ex. Electron configuration for Carbon
21. Electron Configuration
3. Pauli Exclusion Principle
• No two electrons can occupy the same space
• Partnered electrons must have opposite spin
• Possible values for spin are indicated with up or
down arrows
• Ex.
24. Practice
• Write the electron configuration in standard
notation for the following atoms:
1. Helium
2. Beryllium
3. Oxygen
4. Sodium
5. Sulphur
6. Calcium
25. Practice
• Write the electron configuration in standard
notation for the following atoms:
1. Helium: 1s2
2. Beryllium: 1s2 2s2
3. Oxygen: 1s2 2s2 2p4
4. Sodium: 1s2 2s2 2p6 3s1
5. Sulphur: 1s2 2s2 2p6 3s2 3p5
6. Calcium: 1s2 2s2 2p6 3s2 3p6 4s2
26. Practice
• Show the electron configuration and spin states for
the following atoms:
1. Lithium
1. Boron
1. Fluorine
27. Practice
• Show the electron configuration and spin states for
the following atoms:
1. Lithium: 1s2 2s1 ____ ____
1. Boron: 1s2 2s2 2p1 ____ ____ ____ ____ ____
1. Fluorine: 1s2 2s2 2p5 ____ ____ ____ ____ ____
28. Core Notation
• Writing out the entire electron configuration can get
very long
• Standard notation can be shortened to core
notation by using the noble gas configuration
before the element as the core
• Ex. Write the standard and core notation for P
• Standard - P: 1s2 2s2 2p6 3s2 3p3
• Core - P: [Ne] 3s2 3p3
