Journey through the Atom: From Ancient Greece to Modern Science

1. Chapter 4 Atoms and their structure

2. History of the atom • Not the history of atom, but the idea of the atom. • Original idea Ancient Greece (400 B.C.) • Democritus and Leucippus- Greek philosophers.

3. History of Atom • Smallest possible piece? • Atomos - not to be cut • Looked at beach • Made of sand • Cut sand - smaller sand

4. Another Greek • Aristotle - Famous philosopher • All substances are made of 4 elements • Fire - Hot • Air - light • Earth - cool, heavy • Water - wet • Blend these in different proportions to get all substances

5. Who Was Right? • Did not experiment. • Greeks settled disagreements by argument. • Aristotle was a better debater - He won. • His ideas carried through middle ages. • Alchemists tried to change lead to gold.

6. Who’s Next? • Late 1700’s - John Dalton- England. • Teacher- summarized results of his experiments and those of others. • Elements substances that can’t be broken down • In Dalton’s Atomic Theory • Combined idea of elements with that of atoms.

7. Dalton’s Atomic Theory • All matter is made of tiny indivisible particles called atoms. • Atoms of the same element are identical, those of different atoms are different. • Atoms of different elements combine in whole number ratios to form compounds. • Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed.

8. Parts of Atoms • J. J. Thomson - English physicist. 1897 • Made a piece of equipment called a cathode ray tube. • It is a vacuum tube - all the air has been pumped out. • A limited amount of other gases are put in

9. Voltage source Thomson’s Experiment - + Metal Disks

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

11. Voltage source Thomson’s Experiment + - • By adding an electric field

12. Voltage source Thomson’s Experiment + - • By adding an electric field he found that the moving pieces were negative • By adding an electric field

13. Thomson’s Experiment • Used many different metals and gases • Beam was always the same • By the amount it bent he could find the ratio of charge to mass • Was the same with every material • Same type of piece in every kind of atom

14. Thomsom’s Model • Found the electron. • Couldn’t find positive (for a while). • Said the atom was like plum pudding. • A bunch of positive stuff, with the electrons able to be removed.

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

16. Rutherford’s experiment • When the alpha particles hit a florescent screen, it glows. • Here’s what it looked like (pg 72)

17. Flourescent Screen Lead block Uranium Gold Foil

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

19. What he expected

20. Because

21. Because, he thought the mass was evenly distributed in the atom

22. Because, he thought the mass was evenly distributed in the atom

23. What he got

24. + How he explained it • Atom is mostly empty. • Small dense, positive piece at center. • Alpha particles are deflected by it if they get close enough.

25. +

26. Modern View • The atom is mostly empty space. • Two regions. • Nucleus- protons and neutrons. • Electron cloud- region where you might find an electron.

27. Density and the Atom • Since most of the particles went through, it was mostly empty. • Because the pieces turned so much, the positive pieces were heavy. • Small volume, big mass, big density. • This small dense positive area is the nucleus.

28. Other pieces • Proton - positively charged pieces 1840 times heavier than the electron. • Neutron - no charge but the same mass as a proton. • Where are the pieces?

29. Subatomic particles Actual mass (g) Relative mass Name Symbol Charge Electron e- -1 1/1840 9.11 x 10-28 Proton p+ +1 1 1.67 x 10-24 Neutron n0 0 1 1.67 x 10-24

30. Structure of the Atom • There are two regions. • The nucleus. • With protons and neutrons. • Positive charge. • Almost all the mass. • Electron cloud- most of the volume of an atom. • The region where the electron can be found.

31. Size of an atom • Atoms are small. • Measured in picometers, 10-12 meters. • Hydrogen atom, 32 pm radius. • Nucleus tiny compared to atom. • IF the atom was the size of a stadium, the nucleus would be the size of a marble. • Radius of the nucleus is near 10-15m. • Density near 1014 g/cm3.

32. Counting the Pieces • Atomic Number = number of protons • # of protons determines kind of atom. • the same as the number of electrons in the neutral atom. • Mass Number = the number of protons + neutrons. • All the things with mass. • NOT on the periodic table

33. Symbols • Contain the symbol of the element, the mass number and the atomic number. Mass number X Atomic number

34. Symbols • Find the • number of protons • number of neutrons • number of electrons • Atomic number • Mass Number • Name 24 Na 11

35. Symbols • Find the • number of protons • number of neutrons • number of electrons • Atomic number • Mass Number • Name 80 Br 35

36. Symbols • if an element has an atomic number of 34 and a mass number of 78 what is the • number of protons • number of neutrons • number of electrons • Complete symbol • Name

37. Symbols • if an element has 91 protons and 140 neutrons what is the • Atomic number • Mass number • number of electrons • Complete symbol • Name

38. Symbols • if an element has 78 electrons and 117 neutrons what is the • Atomic number • Mass number • number of protons • Complete symbol • Name

39. Unit 3 notes Ions – charged atoms Most atoms in their natural state are not stable. In order to understand stability, we have to look at the last energy level the electrons fill. This shell is called the valence shell. If that shell is full the atom is happy, if not, then the atom will go out and react to make itself full. If the atom needs to steal electrons to become stable it will, if it needs to give up electrons to become stable it will, and some have the ability to do both. Rule is this, the atom will do whatever is easiest.

40. Unit 3 notes Ex. Flourine :2 in the first level 7 in the second. That second level can hold 8, so it needs one more, so it will go out and steal one from something to make itself happy. Now the PNE will change from 9, 10, 9 to 9, 10, 10, now the protons and electrons are not equal to each other and so the overall charge is not zero but now -1. So the flouride ion has an F-1 overall charge.

41. Unit 3 notes Where does it get this extra electron from? If you look at Lithium, the PNE for lithium is 3, 4, 3 2 electrons in the first level and 1 in the second. If lithium wants to fill its second level it would need to steal 7 electrons, very difficult. But if it could dump it to flourine, then the first level (which is now the only level) would be full and everything could be stable. Lithium ion would have a PNE of 3, 4, 2 and because it gave up an negative particle would have a charge of Li +1 These go hand in hand, you have to have one thing giving it up if you want to have something take it in. It is the backbone for every chemical reaction.

42. Unit 3 notes If the atom gives up an electron we say that it has been oxidized (overall charge increases) If the atoms takes in an electron we say that it has been reduced (overall charge decreases) These come from Benjamin Franklin’s names of oxidation and reduction during a chemical reaction. Cation- + ions Anion - - ions Something is unique about C though, it can do either gain 4 or lose 4, so it can have a charge of + or – 4 depending on whatever it needs to do.

43. Isotopes • Dalton was wrong. • Atoms of the same element can have different numbers of neutrons. • different mass numbers. • called isotopes.

44. Naming Isotopes • Put the mass number after the name of the element. • carbon- 12 • carbon -14 • uranium-235

45. Relative Abundance The atomic mass given to you is the overall average of all of the isotopes found. It is based on the relative abundance of each of the isotopes found in nature. For example, Cu has two isotopes: Cu-63 and Cu-65 Because the average is closer to 63, then Cu-63 is more abundant in nature.

46. Relative Abundance To calculate the average of all the isotopes, use the atomic mass and the decimal form of the percent abundance to find the mass contributed by each isotope and then add them together.

47. Relative Abundance What??????

48. Relative Abundance Take the mass of one isotope and multiply it by its percentage that exists in nature (just use the decimal form) Ex. You have element X-10 that is 19.91% abundant in nature and element X-11 that is 80.80% abundant in nature. 10 (.1991) = 1.991 11 (.8080) = 8.888 Now add them up to get the average: 1.991 + 8.888 = 10.879 amu

49. Relative Abundance Now you try it! Calculate the average atomic mass of strontium. Here are the relative abundances: Sr-84 .960% Sr-86 9.86% Sr-87 7.10% Sr-88 82.08%

50. Relative Abundance Bromine has two isotopes Br- 79 and Br- 81. If the average mass of bromine is 79.9, calculate the relative abundances?