2. Unit 3 – Atomic Structure Bravo – 15,000 kilotons

3. Democritus 400 BC Greek philosopher 1 st to come up with idea of atoms

4. John Dalton – 1800’s Major contributor of Atomic Theory 1)All matter made of atoms 2)All atoms of an element are alike 3)Atoms cannot be created or destroyed 4)Atoms combine in whole-number ratios to form compounds

5. JJ Thomson – late 1800’s Cathode Ray Experiment – discovery of electrons “Plum Pudding” model of atom

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

8. Conclusions from the Study of the Electron  Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.  Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons  Electrons have so little mass that atoms must contain other particles that account for most of the mass

9. Millikan - 1910  Oil Drop Experiment  Discovered charge and mass of electrons

10. Rutherford - 1910  Discovered nucleus and that it was positive  Gold Foil Experiment 1) Most of atom is empty space (majority of particles went straight through) 2) nucleus is small, dense and positively charged (some positive charges were greatly deflected)

11. Rutherford’s Gold Foil Experiment  Alpha particles are helium nuclei  Particles were fired at a thin sheet of gold foil  Particle hits on the detecting screen (film) are recorded

12. Rutherford’s Findings  The nucleus is small  The nucleus is dense  The nucleus is positively charged  Most of the particles passed right through  A few particles were deflected  VERY FEW were greatly deflected “Like howitzer shells bouncing off of tissue paper!” Conclusions:

13. Niels Bohr - 1915 Proposed early model of atom “Planetary Model” electrons orbit nucleus like planets orbit sun Lacks math of modern version Has some errors/violates current theory Radiation is emitted when electrons move from one orbit to another

14. Chadwick Discovered neutrons

15. Modern Atomic Theory (changes from Dalton)  Atoms of an element have a characteristic average mass which is unique to that element (isotopes)  Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!

16. Foundations of Atomic Theory Law of Conservation of Mass: mass is neither created or destroyed in ordinary chemical reactions Law of Definite Proportions: compounds contain same elements in same ratio by mass Example: NaCl is always 39.9% Na and 60.66% Cl by mass Law of Multiple Proportions: 2 or more different compounds composed of same two elements have ratios of small whole numbers Example: CO vs CO 2 ratio of oxygen to oxygen is 2 to 1

18. What is AMU? Stands for atomic mass unit – used when describing “relative” atomic masses This system is used because the actual masses of atoms are so small Carbon-12 is the standard to which all other elements are compared (i.e. hydrogen-1 has a mass that is 1/12 that of carbon-12 so it’s mass would be 1 amu)

19. Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Element# of protonsAtomic # (Z) Carbon66 Phosphorus15 Gold79

20. Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope. Mass # = p + + n 0 Nuclidep+p+ n0n0 e-e- Mass # Oxygen -10 -3342 - 3115 8818 Arsenic753375 Phosphorus153116

21. Atomic Masses IsotopeSymbolComposition of the nucleus % in nature Carbon-12 12 C6 protons 6 neutrons 98.89% Carbon-13 13 C6 protons 7 neutrons 1.11% Carbon-14 14 C6 protons 8 neutrons <0.01% Atomic mass is the average of all the naturally isotopes of that element. Carbon = 12.011

22. Isotopes Isotopes are atoms of the same element having different masses due to varying numbers of neutrons. IsotopeProtonsElectronsNeutronsNucleus Hydrogen–1 (protium) 110 Hydrogen-2 (deuterium) 111 Hydrogen-3 (tritium) 112

23. ISOTOPES Example: Carbon, C, exists in 3 isotopes: IsotopeProtonsNeutronsMass Carbon-126612 Carbon-136713 Carbon-146814

24. Isotope symbols Hyphen notationNuclear notation Carbon – 12 12 C 6p+ and 6n o Carbon – 13 13 C 6p+ and 7n o Carbon – 14 14 C 6p+ and 8n o 6 6 6

25. How to Calculate the Average Mass What is the average atomic mass of sample of Cesium with 3 isotopes: 75% 133 Cs, 20% 132 Cs, and 5% 134 Cs. 0.75 x 133=99.75 0.20 x 132 = 26.40 0.05 x 134 = 6.70 Total = 132.85 avg. atomic mass

26. The Mole 1 dozen = 1 gross = 1 ream = 1 mole = 12 144 500 6.02 x 10 23 There are exactly 12 grams of carbon-12 in one mole of carbon-12.

27. Avogadro’s Number 6.02 x 10 23 is called “Avogadro’s Number” in honor of the Italian chemist Amadeo Avogadro (1776-1855). Amadeo Avogadro I didn’t discover it. Its just named after me!

28. Calculations with Moles: Converting moles to grams How many grams of lithium are in 3.50 moles of lithium? 3.50 mol Li = g Li 1 mol Li 6.94 g Li 45.1

29. Calculations with Moles: Converting grams to moles How many moles of lithium are in 18.2 grams of lithium? 18.2 g Li = mol Li 6.94 g Li 1 mol Li 2.62

30. Calculations with Moles: Using Avogadro’s Number How many atoms of lithium are in 3.50 moles of lithium? 3.50 mol Li = atoms Li 1 mol Li 6.022 x 10 23 atoms Li 2.11 x 10 24

31. Calculations with Moles: Using Avogadro’s Number How many atoms of lithium are in 18.2 g of lithium? 18.2 g Li = atoms Li 1 mol Li6.022 x 10 23 atoms Li 1.58 x 10 24 6.94 g Li1 mol Li (18.2)(6.022 x 10 23 )/6.94