1. Electromagnetic spectrum
Electromagnetic radiation: energy that exhibits wave-like behavior as it travels through space

2. How are wavelength and frequency related?
Wavelength (λ) : the distance between corresponding points on adjacent waves (measured in meters)
Frequency (ν) : the number of waves that pass a given point in a specific time, usually one second (measured in 1/s, s-1, Hz)
Speed of light (c) = 3.0 x 108 m/s
c = λ ν

3. Photoelectric Effect
Photoelectric effect: refers to the emission of electrons from a metal when light shines on the metal Ex. heating a piece of iron grey at room temp. → red → blue However, energy (light) is only observed when a certain frequency of light was used

4. The Particle Description of Light
German physicist, Max Planck suggested that the objects emits energy in quanta of energy.
quantum: the minimum quantity of energy
that can be lost or gained by an atom.

5. The Particle Description of Light (cont.)
Max Planck proposed the following relationship between a quantum of energy and the frequency of radiation:
E = hν
E = the energy, in joules, of a quantum of radiation,
ν = the frequency, in s−1, of the radiation emitted
h = a fundamental physical constant now known as Planck’s constant; h = × 10−34 J• s.

6. The Particle Description of Light (cont.)
A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy.
Ephoton = hν

7. How is light produced in a neon tube?
When a current is passed through a gas, the atoms gain energy and electrons “jump” from their ground state to a higher excited state.
The lowest energy state of an atom is its ground state.
A state in which an atom has a higher potential energy than it has in its ground state is an excited state.

8. The Hydrogen Atom
The four bands of light were part of what is known as hydrogen’s atomic (line) emission spectrum.
Every element has a unique atomic emission spectrum (a “fingerprint”)

9. Hydrogen’s Atomic Emission Spectrum

10. Absorption and Emission Spectra

11. Bohr Model of the Atom
An electron can circle the nucleus only in allowed
paths, or orbits
Energy must be added in order to move an electron from a lower energy level to a higher energy level. This process is called absorption.
When an electron falls to a lower energy level, a photon is emitted, and the process is called emission.

12. The Heisenberg Uncertainty Principle
The Heisenberg uncertainty principle states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.

13. Quantum Mechanical Model of an Atom
In 1926, Austrian physicist Erwin Schrödinger developed an equation that treated electrons in atoms as waves
Electrons do not travel around the nucleus in neat orbits, but they exist in certain regions called orbitals
An orbital is a 3D region around the nucleus that indicates the probable location of an electron (“fuzzy cloud”)

14. electron configuration: the arrangement of electrons in an atom
Noble Gas Notation:
Uses the symbol of a noble gas followed by
the electron configuration of the outer electrons

15. Periodic Table

16. Rules Governing Electron Configurations:
Aufbau Principle: electrons are added one at a time to the lowest energy orbitals available (aufbau = “building up”)
Hund’s Rule: electrons occupy equal energy orbitals one at a time before pairing up
Pauli Exclusion Principle: an orbital can hold a maximum of two electrons, each with an opposite spin direction

17. Orbital Diagrams An unoccupied orbital is represented by a line
(or a box), with the orbital’s name written
underneath the line.
An orbital containing one electron is
represented as:
An orbital containing two electrons is represented as:

18. Element symbol surrounded by dots representing the
Valence electrons: outermost electron in an atom
Electron dot diagrams (structures):
Element symbol surrounded by dots representing the
atoms valence electrons
Ex.