Atomic Theories

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

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

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

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

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

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

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

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

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

(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)

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

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.

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

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.”

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.

CRT Tube with Paddle Wheel

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.

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

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.

Atoms

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

Atoms Rutherford’s experiment http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/ruther14.swf

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.

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.

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

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)

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