2. Ch. 4 - Atomic Structure I. Subatomic Particles (p.113 - 114)

3. Subatomic Particle Properties ParticleSymbolLocationChargeRelative Mass (amu) Actual Mass (g) electron proton neutron e-e- p+p+ n0n0 Electron cloud nucleus – + 0 1/1840 1 1 9.11 x 10 -28 1.67 x 10 -24 approx 0

4. Symbols Elements are listed by their chemical symbols Symbols are usually either one capital letter like C for Carbon, or one capital and one lowercase letter like Ne for Neon

5. Periodic Table The periodic table gives much information we need to learn more about the atom of each element

6. Atomic Number Atomic number = # of protons in an atom Whole number shown on periodic table Periodic table is arranged by atomic number

7. Atomic Mass The average atomic mass is the number at the bottom of this square Found by averaging the natural abundances of its isotopes

8. Atom Math Protons Electrons ProtonsNeutrons # n 0 = Atomic mass – Atomic number

9. Subatomic Particles Most of the atom’s mass. NUCLEUS ELECTRONS PROTONS NEUTRONS NEGATIVE CHARGE POSITIVE CHARGE NEUTRAL CHARGE ATOM #n 0 = Atomic mass - Atomic # equal in a neutral atom Atomic Number equals the # of...

10. Ch. 4.3 - Atomic Structure II. How Atoms Differ (p. 114 - 121)  Mass Number  Isotopes  Relative Atomic Mass  Average Atomic Mass

11. A. Mass Number mass # = protons + neutrons always a whole number NOT on the Periodic Table! © Addison-Wesley Publishing Company, Inc.

12. B. Isotopes Atoms of the same element with different numbers of neutrons Mass # Atomic # Isotope notation: Isotope name: carbon-12 Element name Mass #

13. B. Isotopes Chlorine-37  atomic #:  mass #:  # of protons:  # of electrons:  # of neutrons: 17 37 17 20 Isotope notation:

14. Natural Abundances of Isotopes Most elements are found as mixtures of isotopes Relative abundance of each isotope is the same in each source

15. Ion: an atom or molecule in which the total number of electrons is not equal to the total number of protons, giving it a net positive or negative charge. Cation: has a positive charge, due to the loss of electrons Anion: has a negative charge, due to the gain of electrons

16. C. Relative Atomic Mass 12 C atom = 1.992 × 10 -23 g 1 p= 1.007276 amu 1 n = 1.008665 amu 1 e - = 0.0005486 amu © Addison-Wesley Publishing Company, Inc. atomic mass unit (amu) 1 amu= 1 / 12 the mass of a 12 C atom

17. D. Average Atomic Mass weighted average of all isotopes on the Periodic Table Avg. Atomic Mass

18. Avg. Atomic Mass D. Average Atomic Mass EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16 O, 0.04% 17 O, and 0.20% 18 O. 16.00 amu

19. Avg. Atomic Mass D. Average Atomic Mass EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine- 35 and 2 are chlorine-37. 35.40 amu

20. Atomic Theory Development of our understanding of the atom

21. Early Models Battle of Philosophers Aristotle vs. Democritus -Matter is infinite - Matter is composed of extremely small particles -4 Basic ‘elements’ - Called these particles ‘atoms’ - Earth (From ‘atmos’ meaning invisible) -Air -Fire *Eventually after many years -Water Democritus is proven right

22. Early Models Dalton’s model was the “Billiard Ball” Published an Atomic Theory 1. All matter is composed of atoms that are indivisible (did not know about protons, electrons, or neturons) 2. Atoms of a given element have same size, mass and chemical properties and are different from those of another element (no longer true, b/c of isotopes) 3. Different atoms combine in whole number ratios to form compounds and are separated, combined and rearranged in chemical reactions 4. In chemical reactions atoms can combine or separate, but are neither created or destroyed

23. Law of Definite Proportions n Each compound has a specific ratio of elements n It is a ratio by mass n Water is always 8 grams of oxygen for each gram of hydrogen

24. Law of Multiple Proportions n if two elements form more than one compound, the ratio of the second element that combines with 1 gram of the first element in each is a simple whole number.

25. What? n Water is 8 grams of oxygen per gram of hydrogen. n Hydrogen Peroxide is 16 grams of oxygen per gram of hydrogen. n 16 to 8 is a 2 to 1 ratio n True because you have to add a whole atom, you can’t add a piece of an atom.

26. Parts of Atoms n J. J. Thomson - English physicist. 1897 n Made a piece of equipment called a cathode ray tube. n It is a vacuum tube - all the air has been pumped out.

27. Thomson’s Experiment Voltage source +- Vacuum tube Metal Disks

28. Thomson’s Experiment Voltage source +-

29. Thomson’s Experiment Voltage source +-

30. Thomson’s Experiment Voltage source +-

31. n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

32. n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

33. n Passing an electric current makes a beam appear to move from the negative (cathode) to the positive end (anode) Thomson’s Experiment Voltage source +-

34. n Passing an electric current makes a beam appear to move from the negative (cathode) to the positive end (anode) Thomson’s Experiment Voltage source +-

35. Thomson’s Experiment n By adding an electric field

36. Voltage source Thomson’s Experiment n By adding an electric field + -

37. Voltage source Thomson’s Experiment n By adding an electric field + -

38. Voltage source Thomson’s Experiment n By adding an electric field + -

39. Voltage source Thomson’s Experiment n By adding an electric field + -

40. Voltage source Thomson’s Experiment n By adding an electric field + -

41. Voltage source Thomson’s Experiment n By adding an electric field he found n By adding an electric field he found the ratio of electrical charge to mass (e/m) for an electron n The e/m ratio is (negative)  1.76 x 10 8 coulombs per gram (or C/g in SI units). + -

42. Thomsom’s Model n Thomson always found the same value for the e/m ratio no matter what the tube materials or the gas inside. n Reinforced the notion that the electrons are a fundamental component of matter. n ‘Plum Pudding’ model: a thin positive fluid, which contains most of the mass, w/ negative electrons embedded to balance the charge

43. Millikan used oil drop experiment Would spray a fine mist of oil droplets above a pair of parallel plates. Some of the oil drops would pass through the hole in the top plate. He then used X-rays to knock electrons off of the air molecules in the barrel and some of those electrons attached themselves to the oil drops. The oil drops, which were now negative, could now be affected by the electrical field. He then could now measure the charge of the oil drops.

45. Millikan found that all the values he obtained were whole-number multiples of -1.60 x 10 -19 coulomb. This value must be the charge of an electron. The electron’s charge was -1.60 x 10 -19 coulombs Using two values and solving for m - 1.60 x 10 -19 coul = - 1.76 x 10 8 coul/g m m = 9.11 x 10 -28 grams (a negligible mass even in the smallest atom) Confirmed the negative charge of an electron Determined mass of the electron

46. Rutherford’s experiment n Ernest Rutherford English physicist. (1910) n Believed in the plum pudding model of the atom. n Wanted to see how big they are using radioactivity n Alpha particles - positively charged pieces given off by uranium n Shot them at gold foil which can be made a few atoms thick

47. Lead block Uranium Gold Foil Florescent Screen

48. He Expected n The alpha particles to pass through without changing direction very much n Because n The positive charges were spread out evenly. Alone they were not enough to stop the alpha particles

49. What he expected

51. Because, he thought the charge was evenly distributed in the atom

52. What he got

53. +

54. How he explained it + n Atom is mostly empty n Small dense,positive piece at center (nucleus) n Small dense,positive piece at center (nucleus) n Proposed planetary model… (not as refined as the solar sytem model) n Refined the concept of the nucleus & concluded it was composed of positively charged particles called protons n James Chadwick: discovered a neutral atomic particle with a mass close to a proton. Thus was discovered the neutron. n James Chadwick: discovered a neutral atomic particle with a mass close to a proton. Thus was discovered the neutron.

55. Moving Forward… Neils Bohr said electrons move in orbits Found in energy levels Explains bright-line spectrum Called “Solar System Model” where Electrons move in orbits around the nucleus

56. What we believe now Heisenberg/Schrodinger Heisenber Uncertainty Principle: You can know either the eˉ position or velocity but not both Schrodinger said the eˉ are located in orbitals, (regions of probability) around the nucleus… not orbits “Electron Cloud” model