Development of Modern Atomic Theory

1. Lesson 5 Atomic Theory Anything in black letters = write it in your notes (‘knowts’)

2. Section 1 – Defining the Atom Section 2 – The Spectrum of Light Section 3 – Development of Modern Atomic Theory Section 4 – The Nucleus Section 5 – Electrons in Atoms

3. 5.3 – Development of Modern Atomic Theory Wilhelm Roentgen Discovered x-rays in 1895.

4. Crooke’s Tube Roentgen noticed a glow on a nearby fluorescent screen when the Crooke’s tube was on. Even with the tube covered in black cardboard, the glow was still present. No known ray of light could do this – these unknown rays were called x-rays; x for unknown.

5. First x-ray taken When roentgen placed his hand between the crookes’ tube and the fluorescent screen, he saw the bones of his own hands.

6. Henri Becquerel Discovered radioactivity in 1896. Becquerel was studying phosphorescence in uranium containing compounds. He noticed that when these compounds were wrapped in black paper and left in the dark they still marked photographic plates. This indicated that the uranium compound was giving off energy at all times.

7. 3 Types of Radioactivity Alpha () – helium nucleus; heavy and slow Beta () – electron; fast and light Gamma () – high energy light wave

8. J. J. Thomson Discovered the electron in 1897 Thomson performed experiments that involved passing electric current through gases at low pressure. The result was a glowing beam, or cathode ray, that traveled from the cathode to the anode.

9. Thomson found that a cathode ray is deflected by electrically charged metal plates. Thompson knew that opposite charges attract and like charges repel, so he hypothesized that a cathode ray is a stream of tiny negatively charged particles moving at high speed; now called electrons. To test his hypothesis, Thompson set up an experiment to measure the ratio of an electron’s charge to its mass. Also, the charge-to-mass ratio of electrons did not depend on the kind of gas in the cathode-ray tube or the type of metal used for the electrodes.

10. A cathode ray can also be deflected by a magnet.

11. The discovery of the electron led to the “Plum Pudding” model of the atom. Dalton Democritus Thomson’s Plum Pudding Model

12. In 1911, Ernest Rutherford and others performed the Gold Foil Experiment to test the plum pudding model Ernest Rutherford

13. The Gold Foil Experiment

14. The results… It was expected that alpha particles would pass through the plum pudding model of the gold atom undisturbed. Expected It was observed that a small portion of the alpha particles were deflected, indicating a small, concentrated positive charge (the nucleus!) Actual

15. “It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you. On consideration, I realized that this scattering backward must be the result of a single collision, and when I made calculations I saw that it was impossible to get anything of that order of magnitude unless you took a system in which the greater part of the mass of the atom was concentrated in a minute nucleus. It was then that I had the idea of an atom with a minute massive center, carrying a charge.” — Ernest Rutherford

16. Nucleus – tiny positively charge core of an atom Rutherford’s Nuclear Model of the Atom Is this the current model of the atom? NO… • If an atom were the size of a football stadium, the nucleus would be about the size of a marble

17. Proton – positively charged particle in the nucleus of all atoms. Neutron – particle with no charge in the nucleus of all atoms except 1H

18. Lesson 5 Quick Quiz 1. Identify the idea or discovery these people are credited with. Democritus, Dalton, Roentgen, Becquerel, Thomson, Rutherford 2. Place the above ideas or discoveries in chronological order. 3. Describe the setup of the gold foil experiment. Also describe the expected and actual results of this experiment. What conclusion was made about the atom from the gold-foil experiment? 4. Name the three subatomic particles. Also give their symbol and charge. What particles are present in the nucleus of the atom? 5. What is an atom?

19. 5.4 – The Nucleus Atomic Number (Z) - The number of protons in an atom; identifies the element. Atoms have no net charge: # p+ = # e-

20. Mass Number (A) - The number of protons (p+) plus the number of neutrons (n0) in an atom. The mass number is NOT the atomic mass. 1 1 1 ??? 8 8 8 ??? 20 20 20 ??? The number of n0 depends on the mass number of the isotope

21. Isotopes - Atoms of the same element that have a different number of neutrons.

22. Chemical Symbols for Isotopes A is the superscript Z is the subscript

23. a. Be b. Ne c. Na 9 4 20 10 23 11 Determining the Composition of an Atom How many protons, electrons, and neutrons are in each atom? p+ e- n0

24. Naturally Occurring Isotopes of Neon Percent Abundance in Nature 90.48% 0.27% 9.25%

25. The masses of atoms are rarely expressed in grams. The C-12 isotope has been given a mass of exactly 12 atomic mass units (amu) The masses of all other elements are based on the mass of the C-12 isotope.

26. Atomic Mass – Weighted average of all the naturally occurring isotopes of the element. 12.000 amu 98.93 % 13.003 amu 1.07 % 14.003 amu 0.0000000001 % Atomic Mass of Carbon = 12.011 amu

27. 12.000 amu 98.93 % 13.003 amu 1.07 % 14.003 amu 0.0000000001 % Atomic Mass of Carbon = 12.011 amu No atom of carbon actually weighs 12.011 amu. But a typical carbon atom averages 12.011 amu. Atomic masses are weighted averages.

28. There are 2 stable isotopes of silver Silver-107; 106.905097 amu; 51.84% Silver-109; 108.904752 amu; 48.16% Calculate the atomic weight of silver. Atomic Weight of Silver = 107.868 amu

29. Weighted Averages are NOT just for Atomic Masses Your Chemistry grade is calculated as a weighted average. 50% (0.50) Quizzes/Tests, 20% (0.20) Homework, 15% (0.15) Lab Reports, 15% (0.15) Quarter Project, Calculate the un-weighted average grade Calculate the weighted average grade 4/10 = 40% 50/50 = 100% 70/75 = 93.3% 45/50 = 90% 169/185 = 91.35% 67.50%

30. Mass Spectrometer – separates isotopes by mass differences.

31. Mass Spectrum for Cadmium

32. Mass Spectrum for Zinc; Atomic Weight = 65.395

33. Despite differences in the number of neutrons, isotopes of an element are chemically similar. Neutrons do not determine chemical reactivity; the electrons do.

34. Example… Calculate The average mass of zinc The weighted average mass of zinc

35. Terms to Know for Section 3 & 4 Atom Atomic Number (Z) Mass Number (A) Proton, Neutron, Electron Nucleus Isotope Atomic Mass (a.m.u.) Weighted Average Natural Abundance

36. Things to know for Section 3 & 4 Roentgen, Becquerel, Thomson, Rutherford 3 types of radiation (, , ) Explain Gold Foil Experiment (expected and actual results, conclusion) How to determine the number of p+, n0, e- How to write isotope symbols (Sodium-22 or 22Na) How to calculate weighted averages Basically know how to do the Practice Quiz!

37. 5.5 – Electrons in Atoms Rutherford’s Planetary Model of the Atom Electrons moving around a tiny nucleus

38. Problems with Rutherford’s Model • Did not explain the chemical properties of the elements. • Did not explain atomic spectra (…later…) • e- would spiral into the nucleus, but they don’t

39. Low energy ( = 700 nm) High energy ( = 380 nm) Frequency  (s-1) 3 x 106 3 x 1012 3 x 1022 10-8 10-14 102 Review….. The Electromagnetic Spectrum

40. White light produces a rainbow of colors. Screen Light bulb Slit Prism A prism separates light into the colors it contains.

41. Light from a helium lamp produces discrete lines. Screen Slit Prism Helium lamp

42. The lines that result are unique for each element and are called its atomic emission spectrum.

43. For years, nobody had an explanation for atomic spectra…

44. The Bohr Model ~1913 Electrons are found only in specific locations (or energy levels) around the nucleus. To move from one energy level to another, an e- must gain or lose a quantum of energy. Niels Bohr

45. The energy levels in atoms are unequally spaced, like the rungs in this unusual ladder. The higher energy levels are closer together.

46. Bohr’s Model explained the emission spectra When an atom absorbs energy an electron jumps to a higher energy level (excited state). The electron returns to the lower energy level, emitting a photon with a definite energy. The photon’s energy shows up as a line in the emission spectrum.

48. Electron cloud The Quantum Mechanical Model The modern description of e- in atoms. Similar to Bohr Model except the exact location of an electron is impossible to know Electrons are likely to be found in electron ‘clouds’ around the nucleus We will use the Bohr model in this class…

49. Atomic orbital – Most probable place for e- to be. Orbitals can hold 2 e- maximum. The orbitals are named s, p, d & f

50. S (1 type) p (3 types) d (5 types) There are 7 types of f orbitals Don’t worry about these shapes…