1. Arrangement of electrons
Chapter 4 – Modern Chemistry

2. 4.1 Development of new model
Rutherford’s model did not address the arrangement of electrons around the nucleus
A model was developed after looking into the properties of light

3. Properties of Light Light behaves in a dual nature Wave
Light belongs to the electromagnetic spectrum which allows light to exhibit properties of waves as it moves through space
Parts
Wavelength (meters)
Frequency (Hertz)
Speed
Inversely proportional to wavelength and frequency (3.0 x m/s)

4. Particle Nature of Light
Photoelectric Effect
Natural phenomena in which a metal will emit an electron after absorbing light
From this experiment, light is thought to have a certain quanta of energy which is the minimum energy that can be gained or lost by an atom
Max Planck proposed this energy is equal to the frequency times Planck’s constant (6.626 x Jxs)

5. Photon
Einstein expanded Planck’s view introduced the dual wave- particle nature of light
He stated that all light has zero mass and carries a quantum of energy
This particle is known as the photon

6. Hydrogen Atom Emission Line Spectrum
When current is passed through a tube of hydrogen at low pressure, it causes the gas to emit a light
This is a result of ground state electrons becoming excited and changing energy levels
If this light is passed through a prism, it is separated into its own unique spectrum
Attempts to explain why only certain frequencies were present was explain in a new model of the atom

7. Bohr’s model
Niels Bohr was the first to come up with a mathematical representation of why hydrogen has its own unique spectrum (1913)
Planetary Model
Electrons essentially orbit around the nucleus in fixed energy levels

8. Bohr’s Model

9. 4.2 Quantum Model What changed from Bohr’s model to the Quantum model?
Louis De Broglie suggested that electrons would also have a dual nature
Therefore, they acted like waves as well as particles
Werner Heisenberg stated his uncertainty principle
Due to speed of electrons, it is impossible to know but the velocity and the location simultaneously
Erwin Schrodinger created his wave equation and quantum theory
This equation gave mathematical representation of areas of probability where electrons could be found

10. Quantum Numbers Principle Quantum numbers (n)
Main energy levels
Labeled 1 to 7
Angular Momentum number (l)
These are the sublevels (s, p, d, f)
Values are 0 for s, 1 for p, 2 for d, and 3 for f

11. Quantum Numbers Magnetic Quantum number (m)
Represent the orbitals in each sublevel
Number of orbitals is represented by the sublevel
S has 1
P has 3
D has 5
F has 7
Labeled from –l to l (that is L not one, angular momentum number)

12. Quantum Numbers Spin Quantum Number Either -1/2 or ½
All electrons in the same orbital must have opposite spins
No more than 2 electrons per orbital

13. 4.3 Electron Configuration
Aufbau’s Principle
Electrons fill in the lowest energy levels first
Follow the diagonal diagram
Pauli Exclusion Principle
No one electron can have the same four quantum numbers
Hund’s Rule
When there are multiple orbitals, each orbital must have one electron before they pair up. Those electrons will have the same spin

14. Diagonal Diagram

15. Types of Notations Orbital Notation Electron Configuration Notation
Illustrates orbitals and electron spin
Electron Configuration Notation
Simplifies the notation by placing number of electrons in the superscript of the sublevel
Noble Gas Notation
Replaces the last full s and p sublevel with corresponding noble gas