1. Atomic Theory & Atomic Structure

2. Tro, Chapter 4 & 9 Sections 4.1 – 4.4, 4.8, 4.9; 9.2 – 9.9

3. Document BIG IDEAS about: Atomic structure –Electrons (mass, size, position) –Protons and neutrons (mass, position) –Isotopes Changes in (MODERN) thought –Dalton –Thomson, Rutherford and Bohr Quantum theory (CONTEMPORARY)

4. Early Atomic Theories

5. Democritis (400 BCE) First to propose idea of atom Atom = “ a ” + “ tomos ” = cannot be cut Based solely on logic; not supported by experiments

6. Alchemy (12-1500 CE) Modern word ‘ chemistry ’ came from Arabic ‘ alkimiya ’ recognized importance of experimentation Responsible for developing lab equipment & procedures still used today NOTE: Alchemy is a field, NOT a person…

7. Galileo (~1600 CE) Birth of modern science - combining logic, experimenting, publishing results

8. Lavosier & Priestly (1700 ’ s) Quantitative analysis of chemicals Law of Conservation of Mass: Matter can neither be created nor destroyed

9. Proust (1700 ’ s) Developed Law of Definite Proportions Law of Definite Proportions: Different samples of the same compound always contain its constituent elements in the same proportions by mass

10. Law of Definite Proportions Copper carbonate always contains –5.3 parts copper –4 parts oxygen –1 part carbon by mass

11. Dalton (1800 ’ s) School teacher that proposed the first modern-day idea of atoms Law of Multiple Proportions: If 2 elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in small whole # ratios

12. Law of Multiple Proportions

13. Dalton ’ s Atomic Theory - 1808 All matter is composed of atoms which cannot be subdivided Atoms of same element are identical (size, mass, reactivity) Atoms combine to form compounds in simple, whole # ratios Chemical reactions involve the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction

14. Modern Atomic Theories

15. General Principle #1 Electric Charges Objects with an equal amount of positive and negative charge are said to be electrically neutral + – positive negative

16. General Principle #2 Forces between Charges Objects with like charge repel Objects with opposite charge attract + + + + – + – +

17. Forces between Charges Electrostatic force becomes greater with more charge Electrostatic force becomes smaller the greater the distance between the charges

18. Thomson ’ s Atomic Model (1904)

19. Cathode Ray Experiments Any metal worked for anode Negative electric field repelled beam Object placed in path of glow blocked beam

20. J.J. Thomson ’ s Contribution Discovered the electron (1897) Plum Pudding model Determined the charge-to-mass ratio of an electron using data from cathode ray tube experiments

21. Evidence & Conclusions cathode rays consisted of subatomic particles from atoms of anode cathode rays are negatively charged  must also be positive charge Millikan (oil drop experiment, 1909) calculated electron ’ s mass to be 9.11 x 10 -31 kg

22. Modern View of Atomic Structure ParticleSymbol Relative Charge Mass (kg) protonp+p+ +1 1.6726 x 10 -27 neutronn0n0 0 1.67510 x 10 -27 electrone-e- 9.1096 x 10 -31 + 0 nucleons

23. Modern View of Atomic Structure Particle Relative Charge Mass (kg) Relative mass (amu) p+p+ +1 1.6726 x 10 -27 ~1 n0n0 0 1.67510 x 10 -27 ~1 e-e- 9.1096 x 10 -31 ~0 + 0

24. Rutherford ’ s Problems How is nucleus held together? Why don ’ t electrons collapse into nucleus? H atom has 1 proton & He atom has 2 protons,  mass ratio should be 2:1; instead the ratio is 4:1 …there must be another particle

25. The Gold Foil Experiment: Hypothesis The α -particles will pass straight through the atoms What is an (  ) alpha particle? It is a positively charged Helium nucleus

26. Rutherford’s Gold Foil Experiment

27. The Gold Foil Experiment: Outcome

28. What ’ s happening?

29. The Gold Foil Experiment: Conclusions Atoms : must be mostly space must have a very small, dense area of + charge Protons have same charge as e -, but almost 2000x more mass!

30. The Neutron Discovered by James Chadwick in 1932. Neutron is electrically neutral & has slightly greater mass than a proton Mystery solved.

31. Atomic theory timeline

32. Updating Dalton ’ s Atomic Theory 3 major differences between modern atomic theory & Dalton ’ s atomic theory: Atoms are NOT indivisible – they are made up of protons, neutrons, and electrons Atoms of the same element are NOT exactly alike – they can have different masses (isotopes) Atoms CAN be changed from one element to another, but not by chemical reactions (nuclear reactions)

33. Atomic Structure & Isotopes

34. Atomic Mass Unit (amu) defined as a more convenient unit for reporting mass of small numbers of atoms 12 C is used as the reference 1 amu is defined as exactly 1/12 of a 12 C atom

35. Getting Information from the Periodic Table 6 C 12.0111 Atomic # = # p + in nucleus Elemental symbol Atomic mass (more on this later)

36. Isotopic Notation Atomic number (Z) = # of p + in the nucleus Mass number (A) = sum of # p + & n 0 in nucleus For a neutral atom, # e - = # p +

37. H 1111 He 4242 C 12 6 O 16 8 Zn 63 30 Mass number (A) Examples Atomic number (Z)

38. Isotopes All atoms in an element have the same atomic number However, 2 atoms of the same element can have different mass numbers – called isotopes Isotopes have: –Same # of p + –Different # of n o

39. Some Common Isotopes HHHHHH 1111 2121 3131 CCCCCC 12 6 13 6 14 6 UUUU 235 92 238 92

40. Relative Abundance

41. Mass Spectrometry Technique used to determine atomic mass e-e- Atom bombarded by stream of high energy electrons e-e- e - collides with atom, “bounces” off, but transfers some energy to it e-e- + Atom dissipates excess energy by expelling an electron

42. Mass Spectrometry, cont. Ions are accelerated through a magnetic field Amount of deflection depends on the ion ’ s mass Highest mass deflected least Lowest mass deflected most N S + + + + + + + +

43. Mass Spectrometry, cont. Mass (amu) Sample mass spec for chlorine Relative abundance of each isotope can be determined from relative peak heights 35 37

44. Relative Abundance & Atomic Mass Relative isotopic abundance is then used to calculate atomic mass Atomic mass is the weighted average of the mixture of isotopes

45. Example average atomic mass = (atomic mass 35 Cl)(fraction 35 Cl) + (atomic mass 37 Cl)(fraction 37 Cl) = (34.968 amu)(0.7577) + (36.965 amu)(0.2423) = 35.45 amu Calculate the atomic mass of Cl given the relative abundances of its isotopes: 35Cl – 75.77% 37Cl – 24.23%

46. Practice Copper, a metal known since ancient times, is used in electrical cables & pennies, among other things. The atomic masses of its 2 stable isotopes, 63 Cu (69.09%) and 65 Cu (30.91%) are 62.93 amu and 64.9278 amu, respectively. Calculate the average atomic mass of copper – the relative abundances of each ion is given in parentheses. Answer: 63.54 amu

47. The Bohr Model

48. Electromagnetic Spectrum

49. Light c =  c = speed of light (3.0 x 10 8 m) = wavelength  = frequency

50. Frequency vs. Wavelength

51. Light Energy  as frequency  Energy  as wavelength  Light behaves like a particle (photon) as well as a wave c = 

52. Emission Spectrums When electricity is run through a sample of hydrogen gas, hydrogen atoms gain energy H atoms loose that energy by emitting photons Resulting spectrum is discontinuous continuous discontinuous

53. What ’ s happening?

54. Bohr Model Electrons move in circular orbits around the nucleus Only certain energy levels are “ permitted ” (this explains the discrete lines for the emission spectrum of hydrogen)

55. Schroedinger/Heisenburg Experiments used mathematics (probability) to predict behavior of electrons –Schroedinger equation approximated the probability of finding a single electron for H within a region close to the nucleus –Heisenburg [Uncertainty Principle] reinforces the idea that we just don’t know!

56. Math in Context: Blackbody Experiments