1. Atomic Theories

2. Atomic Theories Atomic Theory – A Short History Fifth Century, BCE
Democritus
Believed matter was composed of very small, individual particles that were indestructible
He called them “atomos” (meaning uncuttable)
His ideas persisted for centuries even though there was no experimental proof

3. Atomic Theories JOHN DALTON - 1808 Revised early Greek ideas
into testable scientific theory
Based his Atomic Theory on
three important concepts:
1. Law of Conservation of Mass
2. Law of Multiple Proportions
3. Law of Definite Proportions

4. Atomic Theories Law of Conservation of Mass
States that mass cannot be created or destroyed
the mass of the reactants equals the mass of the products

5. Atomic Theories Law of Multiple Proportions (pg 77)
States that when two elements combine to form two or more compounds, the mass of one element that combines with a given mass of the other is in the ratio of small, whole numbers

6. Atomic Theories Law of Definite Proportions
States that a chemical compound always contains the same elements in exactly the same proportions by mass
Example: water

7. Atomic Theories Dalton’s Principles
All matter is composed of extremely small particles called atoms which cannot be subdivided, created or destroyed
Atoms of a given element are identical in their physical and chemical properties
Example: all water molecules freeze at 0 deg C and react with explosively with sodium

8. Atomic Theories Dalton’s Principles (continued)
3. All atoms of one element are different from those of any other element
4. Atoms combine in simple, whole-numbered ratios to form compounds
Based on the Laws of Definite and Multiple Proportions

9. Atomic Theories Dalton’s Principles (continued)
5. In chemical reactions, atoms are combined, separated or rearranged but NEVER created, destroyed or changed
Based on The Law of Conservation of Mass

10. Atomic Theories Dalton’s Principles (continued)
6.Atoms of one element are never changed into atoms of another element in a chemical reaction

11. (Protons, Neutrons, Electrons)
Atomic Theories
Dalton, however, did all this work in the early 1800’s without ever knowing about subatomic particles!
(Protons, Neutrons, Electrons)

12. MMMMM…..That plum pudding looks delicious!
Atoms
The Adventures of J.J. Thomson, Plum Pudding and the Electron!
MMMMM…..That plum pudding looks delicious!

13. Atoms Chapter One: The Discovery of the First Subatomic
Particle: The Electron
JJ Thomson was an English
Physicist who was studying
electricity. He decided to build
a glass tube which contained two
metal plates at either end. He
then pumped all the air out of the tube and attached a voltage
source to either end. A glowing beam came out of the cathode and struck the
anode and the walls of the glass tube. He called this a cathode ray.

14. Atoms JJ Thomson’s Cathode Ray Tube (CRT) Anode – attached to the
positive terminal of the
voltage source
Cathode – attached to the
Negative end of the voltage
source

15. Atoms
To test this, he placed a magnet near the tube and the beam was deflected away from it, proving it was negatively charged.
He also placed a small paddlewheel in the tube, which turned when hit by the beam. That meant the particles had mass.
The particles are called “electrons.”

16. Atoms Where did the beam come from?
Because most of the air had been removed from the tube, and the beam originated at the negatively charged cathode, Thompson reasoned the ray must be negatively charged.

17. CRT Tube with Paddle Wheel

18. Atoms
Atoms have no charge, yet they gave off negatively charged electrons. The scientists hypothesized that there must be positive charges also included in the atom.
JJ Thomson proposed the “Plum Pudding” model of the atom, which states that electrons were embedded in a
mass of positively charged matter.

19. Atoms
In 1909, one of his students, Ernest Rutherford, disproved the “Plum Pudding” model by doing is famous “Gold Foil” experiment.
BLING!

20. Atoms Atomic Nuclei History
1911: Ernest Rutherford does his famous experiment
“Gold Foil Experiment”
He shot radioactive alpha particles at an extremely thin gold foil
Most particles went right through to the other side
and were detected on the
screen behind it.
Only 1 in 8000 bounced
back.

21. Atoms

22. Atoms FOIL (sideway view) These particles pass through the foil to the
other side and hit the
detecting screen
Detecting
Screen
This particle hit
the dense nucleus
and bounced off of it

23. Atoms Rutherford’s experiment

24. Atoms Rutherford’s Conclusions: Most atoms are made of empty space
All of an atom’s positive charge and almost all of its mass is contained in an extremely small
area
He called this a “nucleus”
Eventually, other scientists later determined more details about the nuclei of atoms.

25. Atoms Rutherford’s Model of the Atom:
Electrons orbit the nucleus just as planets
orbit the sun. However, he could not explain
why the negatively charged electrons did not
crash into the positive charged nucleus.

26. Atoms
Two years later, Danish physicist, Niels Bohr, proposed the Bohr Model of the atom

27. Atoms Bohr’s Model has the following characteristics:
Electrons are located certain distances from the nucleus
Each distance is a certain quantity of energy that the electron can have
Electrons closest to the nucleus have the lowest energy, while the ones further away are in higher energy levels
The difference between two energy levels is called a quantum of energy.
Electrons can be only in an energy level, NOT between levels.
Electrons do not give off energy while they are in an energy level.
(Page 91)

28. Atoms Various Scientists: Created the quantum mechanical model:
Electrons are not located in specific, fixed, circular orbits
Electrons have certain allowed energies and one can determine how likely it is to find an electron with that particular amount of energy
Looks at probabilities of locating an electron