1. Early Atomic Theories and the Origins of Quantum Theory

2. EARLY DEVELOPMENTS IN ATOMIC STRUCTURE
J.J Thomson (1900s):
Provided evidence for the existence of the electron, a negatively charged subatomic particle.
He applied high voltage to a partially evacuated tube with a metal electrode at each end called a cathode ray tube.
When a negative pole was brought towards the ray repelled from the pole indicating that there was a stream of negatively charged particles.
His model is sometimes referred to as the raisin bun method.

4. EARLY DEVELOPMENTS IN ATOMIC STRUCTURE
Rutherford (1911) :
He devised experiments in which positively charged alpha particles were fired at a thin sheet of gold foil.
He hypothesized that if Thomson’s model was correct, then, the massive alpha particles should break through the thin foil.
The results showed that most particles passed through however some were deflected at various angles while others were reflected back.
He concluded that there must be some sort of positive center which he later called the nucleus. He also concluded that since many particles did pass through that atoms were made up of mostly empty space.
Electrons move around the nucleus at a relatively far distance, similar to planets orbiting the sun.

6. THE NATURE OF MATTER AND ENERGY
Classical Theories of Light
Light Energy  Electromagnetic radiation.
Visible Light  The portion of this spectrum that can be seen by the human eye.

7. James Maxwell
He theorized that light could act on charged particles because it existed as an electromagnetic wave made of magnetic and electric fields.
His theory states that light is an electromagnetic wave composed of continuous wavelengths that form a spectrum.

8. Max Planck
He discovered that when a solid is heated to very high temperatures, it begins to glow, first red, then white, then blue.
The changes in colour do not depend on the composition of the solid.
Light is emitted in bursts of a distinct quantity of energy rather than a continuous stream. Quantum  A unit or packet of energy.

9. NEW DEVELOPMENTS IN ATOMIC STRUCTURE
Limits of the Rutherford Model
It became apparent that there was a problem with Rutherford’s idea.
A body that is moving in an orbit is constantly changing directions and a body that is changing direction or speed is accelerating.
An electron travelling in an orbit emits energy as photons and therefore, loses energy. If an electron loses energy as it orbits, it should spiral in toward the positively changed nucleus. Since the electron is negatively charged and opposite charges attract, the atom would eventually collapse. Generally most atoms are stable so Rutherford’s theory was unable to predict the stability of atoms.

10. Atomic Spectra
Spectroscopy  is the scientific study of spectra (plural of spectrum) in order to determine properties of the course of the spectra.
Light first passes through a sample and then is dispersed by a prism to form a spectrum of colours.
Continuous Spectrum  An emission spectrum that contains all the wavelengths in a specific region of the electromagnetic spectrum
Line Spectrum  An emission spectrum that contains only those wavelengths characteristic of the element being studied.

12. Bohr Model of the Atom
Bohr used the emission spectrum of the hydrogen atom to develop a quantum model for the hydrogen atom.
Bohr proposed that electrons could only move in specific orbits around the nucleus.
He assigned each orbit a specific energy level and that the energy level of an orbit increased with its distance from the nucleus.
When an electron gained more energy it could move into an orbit farther from the nucleus.
Transition state  the movement of an electron from one energy level to the next Ground state  the lowest energy state for an atom.

14. THE QUANTUM MECHANICAL MODEL OF THE ATOM
It became apparent by the mid 1920s that Bohr’s model could not explain and make predictions about multi-electron atoms. A new approached was introduced called Quantum Mechanics.

15. Orbitals
Orbital  is a region around the nucleus where an electron has a high probability of being found.

16. Heisenberg’s uncertainty principle  It is impossible to know the exact position and speed of an electron at a given time. Therefore the best we can do is describe the probability of finding an electron in a specific location.
Example students moving from classroom to classroom during a break. The students are like the electrons, the school is like the atom and the classrooms are like the orbitals. Someone who does not know the student’s exact schedule may be able to determine the probability of that student being in a particular classroom at a particular time, but it is not certain.

17. Wave Function  is a mathematical description of an orbital in an atom where an electron of a certain energy is likely to be found.
Note: An orbital is not a Bohr orbit, the electron is not moving around the nucleus in a circle. It is a mystery what electrons do in an atom

18. Electron Probability Density  The probability of finding an electron at a given location and it is used to determine the shapes of orbitals.
 The figure to the left illustrates different electron orbitals. The electrons can jump to any of these orbitals if it gets enough energy. These orbitals also overlap, rather than being distinct levels like Bohr suggested.