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FAIR USE STATEMENT :

FAIR USE STATEMENT :

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FAIR USE STATEMENT :

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  1. FAIR USE STATEMENT: Please feel free to edit and use this presentation in your classroom. Please do not remove the credit line on the title page or republish the file in whole or in part as your own. Please do not distribute the file to individuals or at conferences or workshops. I am more than willing to share the presentation with anyone that contacts me at rhondaa@cox-internet.com. The images used in the presentation are not original and the presentation is distributed freely but only for classroom instruction. Rhonda Alexander

  2. Atomic Structure & Periodicity

  3. Chapter 3 - Atoms Democritus – a Greek philosopher 400 BC * ‘Atomos’ – atoms are indivisible. A – not Tomos - cutting

  4. Dalton’s Atomic Theory • John Dalton – English school teacher - 1808 • He proposed an explanation for several laws • All matter is composed of small particles called Atoms • Atoms of a given element are identical in size, mass, & • other properties • Atoms cannot be subdivided, created, or destroyed • Atoms combine in simple whole-number ratios • In a rxn, atoms are combined, separated, or rearranged

  5. Law of Conservation of Mass (Energy) • Matter is neither created nor destroyed. • It only changes form. • If 5 g of element A combines with 10 g of element B to form AB. How many grams of AB are formed?

  6. Law of Definite Proportion • A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or the source of the compound Table salt – NaCl 39.34 % Na & 60.66 % Cl

  7. Law of Multiple Proportions • Two or more different compounds are composed of the same two elements, the masses of the 2nd element combine with a certain mass of the 1st element can be expressed as ratios of small whole numbers. Water vs. Peroxide H2O H2O2

  8. Water vs. Hydrogen Peroxide

  9. Discovery of the Electron Cathode Ray tube • Different gases glow with different colors if a current is passed through the tube. • Glass directly opposite the cathode glows. • An object placed in between will cast a shadow. • A paddle wheel will roll along on its rails from cathode toward • the anode. • 5. Rays are deflected away from a negative electrode. • ***Sir Joseph John Thomson – 1897 *** Electrons are composed of negatively charged subatomic particles.

  10. Cathode Ray tube

  11. Not in handout

  12. Not in handout

  13. Development of Atomic Models plum pudding

  14. Robert A. Millikan 1909 • American physicist showed that the mass of the electron is 9.109 X 10 –31 kg Confirmed the electrons carry a negative charge and its mass. Brought about more questions about the atomic structure. • Atoms are neutral • Atoms have mass

  15. Not in handout Robert Millikan: Oil Droplet Experiment

  16. Not in handout Rutherford Experiment: Nuclear Atom

  17. Recording of Rutherford Not in handout

  18. Discovery of the Atomic Nucleus Ernest Rutherford – 1911 Gold Foil Experiment ‘As if you had fired a 15 – inch shell at a piece of tissue paper And it came back to hit you.’

  19. Not in handout

  20. Gold Foil Experiment IN NUCLEUS Protons - + charge p+ Nuetrons – neutral no Isotopes – atoms of the same element w/ different masses. In The early 20th century, Rutherford showed that most of an atom’s mass is concentrated in a small positively charged region called the nucleus. Electron cloud - 90% probability of finding the electron within this space Electron: - charge ( e-)

  21. Quantum Mechanical Bohr Model Modern Model 90 % probability of finding the electron within this space Bohr Model After Rutherford’s discovery, Bohr proposed that electrons travel in definite orbits around the nucleus

  22. Not in handout

  23. Structure of the Atom • Atom – is the smallest particle of an element that retains the chemicalproperties of the element Atom has 2 regions: Nucleus – protons & neutrons Electron Cloud - Electrons

  24. Weighing and Counting Atoms Atomic number (Z) – number of protons in a nucleus Mass Number – total number of protons & neutrons in a nucleus Atomic Mass – Weighted average mass of all the isotopes of the element Atomic Mass Units – amu 1/12 the mass of C-12 Charge = p+ - e- Mass # = p+ + no

  25. Examples of PEN Atomic number = number of protons If atom is neutral, then the number of protons must equal number of electrons. Carbon Mass Number – Atomic Number = # of Neutrons 12 6 C 12 - 6 = 6 neutrons • Hydrogen 2. Sodium 3. Oxygen • 4. Copper 5. Gold

  26. Isotopes • Isotopes of Hydrogen include • Hydrogen -1 • Hydrogen – 2, Duterium • Hydrogen – 3, tritium • They have the same number of protons but different numbers of neutrons and a different mass number

  27. Not in handout isotopes

  28. Average Atomic Mass – is the weighted average of the atomic masses of the naturally occurring isotopes of an element. • Weighted Average of Isotopes • Copper–63 69.17%, Copper-65 30.83%, Calculate the average atomic mass. (63 x .6917) + ( 65 x .3083) = 43.5771 + 20.0395 = 63.6166 Carbon-12 98.90% , Carbon-13 1.1% 12.011

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