Atomic Structure & the Atom

1. Atomic Structure & the Atom Chapter 3 (p.65-80) Chapter 22 Section 2 (p.704-707)  Major Atomic Theories  Subatomic Particles  Symbols: Atomic # & Mass #  Isotopes  Radioactivity

2. I. Major Atomic Theories p.65-74+ outside sources

3. What do we know about atoms?

4. A. Atomic Theory Timeline • Democritus (400 B.C.) particle theory of matter atom is indivisible

5. A. Atomic Theory Timeline • John Dalton (1803) atoms are indivisible “Billiard Ball” model

6. A. Atomic Theory Timeline • J.J. Thomson (1897) Cathode Ray Tube (CRT) experiment particles in a cathode ray are negatively charged (ELECTRONS) “Plum Pudding” model

7. Voltage source A. Atomic Theory Timeline • J.J. Thomson - +

8. + - A. Atomic Theory Timeline • J.J. Thomson Voltage source - +

9. A. Atomic Theory Timeline • Max Planck (1900) founder of quantum theory used lots of MATH • Planck’s constant (use this next unit!) • h = 6.626 x 10-34 J•s

10. A. Atomic Theory Timeline • Ernest Rutherford (1911) • Gold foil experiment • positive charged alpha particles through goldfoil • something small in the middle must have a positive charge (NUCLEUS) “Nuclear Model”

11. A. Atomic Theory Timeline • What Rutherford EXPECTED

12. A. Atomic Theory Timeline • What Rutherford GOT

13. + A. Atomic Theory Timeline

14. A. Atomic Theory Timeline • Niels Bohr (1922) • electrons move around in orbits • bright-line spectrum experiment “Planetary Model”

15. A. Atomic Theory Timeline • Werner Heisenberg (1927) • Heisenberg Uncertainty Principle “Electron Cloud” Model

16. A. Atomic Theory Timeline • Erwin Shrödinger (1930) • quantum mechanics (math …) “Quantum Mechanic” model = orbitals

17. A. Atomic Theory Timeline • James Chadwick (1932) discovery of the NEUTRON

18. A. Atomic Theory Timeline

19. B. Dalton’s Atomic Theory • John Dalton Dalton’s Atomic Theory: 1. All matter is composed of atoms. 2. Atoms of a given element are the same, atoms of different elements are different. 3. Atoms cannot be subdivided, created, or destroyed. 4. Atoms of different elements combine in whole-number ratios to form chemical compounds. 5. Chemical reactions occur when atoms are combined, separated, or rearranged.

20. II. Subatomic Particles p.73-74

21. Subatomic Particles

22. Subatomic Particles

23. equal in a neutral atom Most of the atom’s mass. Subatomic Particles ATOM NUCLEUS ELECTRONS NEUTRONS PROTONS NEGATIVE CHARGE POSITIVE CHARGE NEUTRAL CHARGE

24. III. Symbols: Atomic # & Mass # p.75-76

25. 12 16 23 51 40 6C, 8O, 11Na, 23V, 18Ar A. Atomic Number • atomic number = # protons in the nucleus • whole number on the periodic table! • Tell us IDENTITY of the element • written at the lower-left corner (subscript) of the symbol ex: 6C, 8O, 11Na, 23V, 18Ar • symbol generally has atomic AND mass number mass number (P + N) atomic number (P)

26. B. Mass Number • mass number = # protons + # neutrons • Take molar mass & ROUND to nearest whole # • written in the upper-left corner (superscript) of the symbol ex:12C, 16O, 23Na, 51V, 40Ar • What is the mass number of an atom with: - 15 P, 15 e-, 16 N? 31 (phosphorus) - 86 P, 86 e-, 136 N? 222 (radon)

27. C. Hyphen Notation, Symbols, and the Periodic Table beryllium-994Be magnesium-262612Mg calcium-414120Ca

28. Example • carbon-12 126C atomic #: 6 mass #: 12

29. Example • Chlorine-37 • atomic #: • mass #: • # of protons: • # of electrons: • # of neutrons: 17 37 17 17 20

30. IV. Isotopes p.77-80

31. mass # atomic # A. Definition of ISOTOPE • atoms of the same element with diff mass numbers (due to a diff. # of NEUTRONS!) • symbol: • hyphen notation: carbon-12

32. A. Definition of ISOTOPE • same # P, diff # N • almost identical properties, except for radioactivity ex: 11H hydrogen has 1 proton, 0 neutrons (hydrogen-1) 21H deuterium has 1 proton, 1 neutron (hydrogen-2) 31H tritium has 1 proton, 2 neutrons (hydrogen-3)

33. A. Definition of ISOTOPE

34. B. Isotopes in Everyday Life • Medicine • Agriculture/Food • Pest Control • Smoke Detectors • Archaeology • a good EXTRA CREDIT activity - find a more specific everyday use of an isotope!

35. C. Relative Atomic Mass • 12C atom = 1.992 × 10-23 g • atomic mass unit (amu) • 1 amu = 1/12 the mass of a 12C atom 1 p = 1.007276 amu1 n = 1.008665 amu1 e- = 0.0005486 amu

36. Avg. Atomic Mass D. Average Atomic Mass • weighted average of ALL isotopes of a given element • average atomic mass is the decimal number on the periodic table • if not given, assume mass number is the “mass” • % is the percent abundance of that isotope • round to 2 decimal places Avg. Atomic Mass  (%) 100 (mass) =

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

38. 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.

39. D. Average Atomic Mass • ex: Copper is used in electrical cables and pennies, among other things. The atomic masses of its two stable isotopes, 29Cu (69.09%) and 29Cu (30.91%), are 62.93 amu and 64.9278 amu, respectively. Calculate the average atomic weight of copper. 63 65

40. V. Radioactivity p.704-707

41. What do you think of when you hear the word “Radioactive”?

42. A. Discovery of Radioactivity • 1896: French physicist Henri Bequerel • Nobel prize in 1903 with the Curies

43. B. Three Types of Radioactivity • alpha () rays/particles He nuclei, each with 2 protons and 2 neutrons, with an atomic number of 2 and a charge of +2 • beta () rays/particleselectrons, each with a charge of –1 • gamma () rays/particlesjust a form of energy (no charge)

44. B. Five Types of Radioactivity • Neutrons (n) • Mass of 1 (1 neutron) • No charge 10n • Positrons • No mass • Charge of +1 0+1e • gamma rays- pure energy 00 • No charge • No mass • alpha () rays/particle • Mass of 4 • Charge of +2 (2 protons) • 2 neutrons 42He • beta () rays/particles • NO mass • Charge of -1 0-1e

45. C. Balancing Nuclear Equations • Two rules: 1. the sum of the MASS NUMBERS (superscripts) on both sides of the equation must be EQUAL 2. the sum of the ATOMIC NUMBERS (subscripts) on both sides of the equation must be EQUAL • General equation: alpha emission ____ → 4He + ____ beta emission ____ → 0e + ____ 2 -1

46. 4He 2 0 e -1 C. Balancing Nuclear Equations • examples: Thorium-223 (Th) is an alpha emitter. Write an equation for this nuclear reaction and identify the product. Iodine-139 (I) is a beta emitter. Write an equation for this nuclear reaction and identify the product. 219  + 223Th Ra 90 88 139  + 139I Xe 53 54

47. Natural Decay Video

48. Nuclear half-life • when a nucleus gives off radiation, it is said to decay • half-life, t1/2: the time it takes for one half of any sample of radioactive material to decay • in theory, it would take an infinite amount of time for all of a radioactive sample to disappear • in reality, most of the radioactivity disappears after 5 half-lives (only 3% of the original remains)

49. Nuclear half-life • INDEPENDENT of temperature and pressure (and all other conditions) • we do not know of any way to speed up or slow down radioactive decay! In assessing long-range health effects of atomic bomb damage or nuclear power plant accidents (Three Mile Island, PA in 1979 or Chernobyl, Soviet Union in 1986), LONG nuclear half-lives are more important • In medical imaging or therapy, SHORT-lived isotopes are more useful because they disappear faster

50. Nuclear half-life • some half lives are short, some are long! Hydrogen-3 3H 12.26 years Beta Carbon-14 14C 5730 years Beta Phosphorus-28 28P 0.28 second Positrons Phosphorus-32 32P 14.3 days Beta Potassium-40 40K 1.28 x 109 years Beta + gamma Scandium-42 42Sc 0.68 second Positrons Polonium-210 210Po 138 days Alpha Radon-205 205Rn 2.8 minutes Alpha Radon-222 222Rn 3.8 days Alpha Uranium-238 238U 4 x 109 years Alpha 1 6 15 15 19 21 84 86 86 92