Chapters 4-5: The Atom

1. Chapters 4-5: The Atom Chemistry

2. What is this a painting of? Henri Matisse, THE BEES, 1948

3. Seurat, 1884

4. Cans Seurat, 200760x92"Depicts 106,000 aluminum cans, the number used in the US every thirty seconds. Pointillism

5. Ferris Bueller’s Day Off

6. The Atom • Indivisible particle of matter • The smallest part of an element sample • Atoms combine to form molecules & compounds Al Foil ?

7. Democritus of Abdera, 100 Greek Drachma (1967) This note is now obsolete. Theories Involving Matter & Atoms • Democritus (400 BC) • “atomos” = indivisible • atoms are the smallest particles possible • there are different kinds of atoms • a very general theory • not supported by evidence

8. Aristotle (384-322 BC) • Matter is Continuous • “Hyle”: a continuous substance making up all matter • A philosophical theory

9. Antoine Lavoisier(1743-1794) • Law of Conservation of Mass: • Matter cannot be created nor destroyed in a reaction • Mass Before = Mass After • Determined formula for water (H2O) • Established the Metric System • French Revolution: Killed in Guillotine

10. Joseph Proust(1754-1826) Law of Definite Composition • Elements combine in definite ratios to form compounds Ex: Table salt always contained 1.5 times as much chlorine as sodium Water always contained eight times as much oxygen as hydrogen

11. John Dalton (1766-1844) • British • Quaker • School Teacher • Recorded daily weather and barometric pressure • Color blind (called “Daltonism”) • Child Prodigy

12. Dalton used the work of Lavoisier and Proust to form… Dalton’s Atomic Theory (1807) 1) All matter is composed of atoms 2) Atoms of an element are alike 3) Each atom has its own weight 4) Atoms unite in definite ratios to form compounds and…

13. Dalton used the work of Lavosier and Proust to form… i.e. NO 1:1 NO2 1:2 NO3 1:3 Law of Multiple Proportions Atoms may combine in more than one ratio when forming compounds Different compounds with entirely different properties

14. Which Law is Illustrated? Law of Definite Composition ____ a) Water is 89% oxygen by mass ____ b) 3.10 g Mg combine with 2.04 g of oxygen to form ___ g of magnesium oxide ____ c) H2O (water) & H2O2 (hydrogen peroxide) Law of Conservation of Mass 5.14 Law of Multiple Proportions

15. Example from HW

16. Parts of the atom subatomic particles Electron (e-) • negative charge (-1) • discovered by JJ Thomson (1897)—Nobel Prize for Physics in 1906 • smallest unit of electrical charge • mass = 1/1837 of a proton J.J. Thomson (1856-1940)

17. Ernest Rutherford 100 New Zealand Dollars Proton (p+) • positive charge (+1) • discovered by Ernest Rutherford (J.J. Thomson’s student at Cavendish Labs) • Mass = 1 amu (atomic mass unit) • Atomic Number = Z = p+

18. James Chadwick (1891-1974) Neutron (no) • discovered in 1932 by Chadwick (a student of Rutherford) • Awarded Nobel Prize for Physics in 1935 • no charge –neutral • Mass of neutron  mass of proton

19. Nucleons • Particles in the nucleus: protons & neutrons • Atomic Mass = #protons + #neutrons • Mass # = Atomic Weight rounded to the nearest whole number

20. 24 12 12 12 24 12 12 10 24 12 12 13 Example

21. 9 9 10 9 18.998 19 Try This for Fluorine (F) Z = __________ #p+ = _________ #no = _________ #e- = __________ Atomic Weight = ________ Mass # = ________

22. Isotopes Atoms of the same element having different masses same number of ________ different number of ________ protons neutons

23. 40 Atomic Mass (#p+ + #no) 22 2+ Ne Ca 10 20 Atomic Number (Z) (#p+) charge Not always the same as the P. T. !!! Symbolizing

24. 6 6 6 6 6 8 6 6 12 14 Try This: C C

25. Which is the most common isotope for Neon? Ne-22 or Ne-21 or Ne-20 How do you know? Average Atomic Weight for Ne is 20.2 g/mol

26. Why are atomic masses not whole numbers? Atomic masses are averages of the isotopes present

27. Think About It! • An atoms weight comes from the protons and neutrons. The electrons are so tiny they don’t matter for weight! • Like a sumo wrestler weighing in with ten lady bugs on his shoulder. (the electrons are like the lady bugs, they don’t change his mass very much).

28. A practical use for Isotopes and Mass Differences • To build an atom bomb scientists needed pure U-235, a tiny fraction of natural uranium samples. • Separated using mass differences.

29. Isotope ReviewHeavy Water • Isotopes of hydrogen: • Hydrogen • Deuterium • Tritium Deuterium Water = D2O and a molecular weight of 20 g/mol

30. Remember John Dalton’s Atom • Dalton thought of atoms as indivisible spheres • His proof for this was the Law of Definite Composition (water is H2O not H1.5O) • He of course was wrong… • Atoms can be split into parts • nop+ and e-

31. Models of the Atom—Thomson Model electrons J.J. Thomson sphere of positive charge • Discovered the existence of the electron (1897) • Proved Dalton’s idea that the atom was indivisible was wrong • Thought electrons were imbedded in the atom • Called the Plum Pudding Model

32. Shortcoming of Thomson’s model: • No Nucleus

33. Rutherford Model • Gold Foil Experiment (1911) • The atom has a dense nucleus • The atom is made of mostly empty space click picture for video turn down volume

34. Major Shortcomings of Rutherford’s Model: • Electrons cannot orbit the nucleus; they would lose energy and the atom would collapse

35. Bohr Model • e-s move at different levels around the nucleus n = 4 n = 3 n = 2 n = 1 Niels Bohr 1885-1962 Click picture for video (after Lab 16 demo)

36. Continuous Spectrum ROY G BIV Bright-Line Spectrum Bohr Model of the Atom • Energy (heat or electricity) causes e-s to move to higher level/excited states • When e-s move back, energy is released in packets called QUANTA

37. Shortcomings of Bohr • Only worked for the Hydrogen atom 500 Danish Kroner

38. Quantum Theory—Modern Model of the Atom • Cannot be visualized • A mathematical description of atom • Energy is transferred in units (quanta) • Electron position described by probability and quanta (packets) of energy

39. n = 4 n = 3 n = 2 n = 1 Energy Levels of the Atom Levels • Principal Quantum Number (n) • Also called: shell, quantum, energy level • Horizontal Rows on the diagram on the back of your PT n = 1 is ground state (lowest energy level) n > 1 is an excited state

40. n = 2 n = 3 n = 4 n = 1

41. Maximum #e-s per energy level = 2n2 2(1)2 = 2 2(2)2 = 8 2(3)2 = 18 2(4)2 = 32

42. n = 2 n = 3 n = 4 n = 1 2 32 18 8

43. Sublevels (l) • Shape of the probability region of the e- • Each energy level has n sublevels • Circles on the diagram on the back of your PT n = 1 1 sublevel s n = 2 2 sublevels s, p n = 3 3 sublevels s, p, d n = 4 4sublevels s, p, d, f

44. n = 2 n = 3 n = 4 n = 1 2 32 18 8 s s, p s, p, d s, p, d, f

45. s – p –

46. # orbitals in that energy level = n2 Orbitals (m) The s, p, d, f circles on the diagram on the back of your PT • Orientation of the sublevel • Maximum of 2 e-s per orbital • s sublevels contain 1 orientation • p sublevels contain 3 orientations • d sublevels contain 5 orientations • f sublevels contain 7 orientations s sublevel p sublevel f sublevel d sublevel

47. Spin • 2 e-s in the same orbital have opposite spin (opposite North and South Poles) • Discovered in 1928 • The lines you draw in the circles on the diagram

48. Electron Configurations 3 Simple Rules • Electrons fill lowest energy level first • Maximum number of 2 e- per orbital • Orbitals half-fill before they fill (Each orbital gets one before an orbital can get seconds) NOT

49. 6p 5d 4f 6s 5p 4d 5s 4p 3d 4s 3p 3s 2p 2s 1s Si = 14 e- 1s2 2s2 2p6 3s2 3p2

50. 6p 5d 4f 6s 5p 4d 5s 4p 3d 4s 3p 3s 2p 2s 1s Br = 35 e- 1s2 4p5 2s2 2p6 3s2 3p6 4s2 3d10