1.  Dismissed idea of the atom. Early Greeks Two schools of thought:  Matter is made of indestructible particles called “atomos” Plato (428-348 BC) Democritus (400 BC) Both theories lacked Scientific Evidence so idea of atom was lost for 2000 years

2. 2000 years of ALCHEMY  Alchemy began as an Arab mixture of Egyptian arts of dyeing, painting, glass making, pyrotechnics, medical drugs, mining, and metallurgy with the theories of the Greeks (mostly Aristotle's) to explain changes in color and appearance of materials Alchemy  Main goal was:  Philosopher’s stone Philosopher’s stone  Transmutation

3. 1700s  Experiments with air result in 1 st balloon flights and further investigations of gases.  Joseph Priestley: Discovers oxygen and is able to isolate it from the air.  Antoine Lavoisier: Identifies the role oxygen plays in combustion. Also discovers a natural law called: Law of Conservation of Matter.

4. Early 1800s  Joseph Proust: Believed substances always combine in a definite way and in the same proportions. Law of Definite Proportions  John Dalton: Showed that different substances with the same elements combined in ratios that were whole numbers. Law of Multiple Proportions next This indicated that matter exists as “ATOMS”. Atomic Theory

5. Dalton’s Theory  Experimental evidence  Scientific Laws  Law of Conservation on Matter You can’t create or destroy atoms.  Law of Definite Proportions Same compounds are same ratio by mass.  Law of Multiple Proportions Different compounds with the same elements are whole number multiples of the atoms.  Experiments  Scientific Laws  Atomic Theory (observations)(Patterns)(Explanations)  John Dalton realized that there must be an atom as Democritus first proposed.

6. Summary for Dalton’s Atomic Theory  Atoms are Tiny.  Atoms of the same element are the same.  Atoms of different elements are different.  Atoms can’t be divided, created or destroyed.  Atoms of different elements combine in simple whole-number ratios to form compounds.

7. Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle.J.J. Thomson Cathode ray tubes pass electricity through a gas that is contained at a very low pressure. Crookes Tube

8. Electrical Nature of Matter  Opposite charges attract each other. (+) (-)  Like charges repel each other. (+) (+) (-) (-)

10. Cathode Rays (electrons) Observations:  Cathode Ray Tube produces rays with constant charge to mass ratio. Conclusions and Hypotheses:  Cathode rays (electrons) were found in all substances tested.  Cathode rays (electrons) were attracted to the positive plate every time.  Electrons are negatively charged.  Neutral atoms are made up of equal amounts of (+) and (-) particles.  All substances have atoms that contain tiny particles called electrons.  The electron has a specific size and charge.

11. Thomson’s Atomic Model (1897) Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

12. Rutherford’s “Gold Foil Experiment”  Alpha particles are helium nuclei, He 2+  Particles were fired at a thin sheet of gold foil  Particle hits on the detecting screen (film) are recorded Radioactive source (+)

13. Rutherford’s Findings (1911) The atom is mostly empty space. The nucleus is dense. The nucleus is positively charged Electrons, e -, are moving large distances outside the nucleus. Observations:  Most of the alpha particles passed right through  Some alpha particles were deflected slightly  VERY FEW were greatly deflected “Like howitzer shells bouncing off of tissue paper!” Conclusions:

14. Rutherford’s Conclusion (1911)…  Small, dense, positive nucleus.  Equal amounts of (-) electrons at large distances outside the nucleus.  The neutron is not discovered until 1932 by James Chadwick (a student of Rutherford’s).

15. Neils Bohr’s Atomic model (1913)  Small, dense, positive nucleus.  Equal amounts of (-) electrons at specific orbits around the nucleus. This incorrect version of the atom is often used to represented atoms because it shows energy levels for electrons.

16. Current Atomic model  Small, dense, positive nucleus.  Equal amounts of (-) electrons occupy regions of space called “orbitals” in the electron cloud.orbitals Atom song

17. Explain the difference between mass number and average atomic mass. Mass number refers to the total protons and neutrons in a specific isotope. Atomic mass refers to all the isotopes masses averaged based on their relative abundance. Also called “Average Atomic Mass” or “Atomic Weight”. On the test you will have to:

18. Modern Atomic Theory  Atoms of the same element are chemically alike with a characteristic average mass which is unique to that element.  Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!  All matter is composed of atoms.  Atoms of any one element differ in properties from atoms of another element  The exact path of electrons are unknown and e - ’s are found in the electron cloud based on probability.

19. Atomic Scale  Most of the mass of the atom is in the nucleus (protons and neutrons)  Electrons are found outside of the nucleus (the electron cloud)  Most of the volume of the atom is empty space “q” is a particle called a “quark” Quarks:

20. About Quarks About Quarks …(particles that make up protons and neutrons)Quarks Protons and neutrons are NOT fundamental particles. Protons are made of two “up” quarks and one “down” quark. Neutrons are made of one “up” quark and two “down” quarks. Quarks are held together by “gluons”

21. IsotopeSymbolAtomic number Mass numberProtonsNeutronsElectrons Cobalt-60 Iron-56 6630 5228 5066 60 Co Cobalt 27 58.93 Atomic # Atomic mass (average) Fe Iron 26 55.85 Zn Zinc 30 65.39 27 Co 27 42 Ca 20 56 Fe 26 2727- 60-27= 33 60 26 26- 56 56-26= 30 42-20= 22 20 42 20- Calcium-42 Zinc-66 Chromium-52 30 30- 24 24- 66 Zn 30 52 Cr 24 66-30= 36 Cr Chromium 24 52.00 Sn Tin 50 118.71 Tin-116 116 Sn 50 50- 50+66= 116 52-28= 24

22. IsotopeSymbolAtomic number Mass numberProtonsNeutronsElectrons 14-12 130 -54 1123 11 26 Mg Magnesium 12 24.31 Atomic # Atomic mass (average) As Arsenic 33 74.92 Xe Xenon 54 131.29 12 Mg 12 130 Xe 54 75 As 33 12 12+14=26 33 33- 75 75-33= 42 130-54= 76 54 Xenon-130 Sodium-23 11- 11 11- 23 Na 11 23-11= 12 Na Sodium 11 22.99 Magnesium-26 Arsenic-75 11 11+11=22 Sodium-22 22 Na 11