1. Atomic Timeline

2. 650 BC - Greek Philosophers Earth Water Fire Wind

3. 450 BC - Democritus Greek: atomus – means indivisible Coined the term “atom” Deep Thinker - No evidence was provided

4. The Dark Ages Religious and scientific persecution BUT only in EUROPE

5. Meanwhile… Trade routes between the Middle East and Asia kept Science going. Far Eastern alchemists designed modern glassware and apparatus

6. 1808 - Dalton ► First model based on experiments. 1.All matter is made up of tiny particles called atoms. They cannot be created, destroyed or divided into smaller particles. 2.The atoms of one element cannot be converted into the atoms of any other element. 3.All the atoms of one element have the same properties, such as mass and size. 4.Atoms of different elements combine in specific proportions to form compounds.

7. 1903 – J.J. Thomson ► Introduction of atomic charge. Sphere of positive mass with negative charges interspersed. Raisin Bun model

8. 1911 – Ernest Rutherford ► Gold Foil Experiment Anticipated Results Actual Results

9. 1911 - Rutherford ► 1. most of the atom is empty space ► 2. there must be a central component of the atom containing all of the positive charge. He called this the Nucleus

10. 1913 – Neils Bohr ► Explained the path of electrons (orbits) around the positive nucleus ► These orbits are specific distances from the nucleus ► Electron energy level model.

11. Modern Atomic Theory 1. All matter is made up of atoms. Each atom consists of subatomic particles: electrons, protons & neutrons (an atom is divisible, it is the smallest part of an element) 2. Atoms of on element cannot be converted into atoms of another element by a chemical reaction (nuclear reactions, alter the composition of the nucleus, so convert atoms of one element into another) 3. All atoms have the same properties such as size and mass. (exception is isotopes) 4. Atoms of different elements combine in fixed proportions to form compounds. (no changes yet)

12. Atomic Structure Subatomic particles Location Relative Mass Charge ProtonNucleus 1 amu Positive NeutronNucleus Neutral Electron Energy Levels/ Orbitals 0.00054 amu Negative * 1 amu (atomic mass unit) = 1.66 X10 -27 kg

13. Standard Atomic Notation Mass Number (A) – Each element is assigned a mass number which corresponds to the number of protons plus neutrons found in the nucleus of the atom of that element. A = #p + + #n Atomic Number (Z) – Each element is assigned an atomic number which corresponds to the number of protons found in the nucleus of the atom of that element. Z = #p + Standard Atomic Notation X A Z Mass # Atomic # Si 28 14 p = 14 e = 14 n = 14 Gold?

14. Isotopes Isotopes – atoms of an element that have the same number of protons but different numbers of neutrons. This change in neutrons is reflected in the Mass Number 12 C 6 13 C 6 14 C 6 Carbon -12 Carbon -13Carbon -14 Isotopes of the same type of atom exhibit the same chemical properties! The mass listed on the Periodic Table is an average atomic mass calculated from the masses of the naturally occurring isotopes of that element.

15. Applications – carbon dating, x-rays, cancer treatments… Radioisotopes Radioisotopes – (short for radioactive isotopes) have unstable nuclei. Radioisotopes decay into more stable atoms by giving off radiation. Radioactivity – process of releasing energy and/or particles from the nucleus of an atom as it decays.

16. Bohr – Rutherford Diagrams