2. Atoms: The Building Blocks of Matter Chapter 3

3. The Atom: From Philosophical Idea To Scientific Theory Chapter 3

4. The Atom: From Philosophical Idea To Scientific Theory Remember the idea that matter is made of atoms dates back to 400 BC. Not proven experimentally until 1700’s

5. The Greeks History of the Atom In 400 B.C the Greeks tried to understand matter (chemicals) and broke them down into earth, wind, fire, and air.  

6. Greek Model Greek philosopher Idea of ‘democracy’ Idea of ‘atomos’ Atomos = ‘indivisible’ ‘Atom’ is derived No experiments to support idea Democritus’s model of atom No protons, electrons, or neutrons Solid and INDESTRUCTABLE Democritus

7. Alchemy After that chemistry was ruled by alchemy. They believed that that could take any cheap metals and turn them into gold. Alchemists were almost like magicians. elixirs, physical immortality

8. Alchemy.............. GOLDSILVERCOPPER IRONSAND Alchemical symbols for substances… transmutation: changing one substance into another In ordinary chemistry, we cannot transmute elements. 

9. Contributions of alchemists: Information about elements - the elements mercury, sulfur, and antimony were discovered - properties of some elements Develop lab apparatus / procedures / experimental techniques - alchemists learned how to prepare acids. - developed several alloys - new glassware

10. Timeline 20001000300 AD American Independence (1776) Issac Newton (1642 - 1727) 400 BC Greeks (Democratus ~450 BC) Discontinuous theory of matter ALCHEMY

11. Foundations of Atomic Theory Observations and chemical reactions led to the following scientific laws that describe how compounds are formed.

12. Foundations of Atomic Theory Law of Definite Proportions (Proust) The fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound. Law of Multiple Proportions (Dalton) If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a fixed mass of the first element is always a ratio of small whole numbers. Law of Conservation of Mass (Lavoisier) Mass is neither destroyed nor created during ordinary chemical reactions or physical changes.

13. Conservation of Mass 2 H 2 + O 2 2 H 2 O 4 atoms hydrogen 2 atoms oxygen 4 atoms hydrogen 2 atoms oxygen H H O O O O H H H H H H H2H2 H2H2 O2O2 H 2 O H2OH2O +

14. Legos are Similar to Atoms Legos can be taken apart and built into many different things. H H O O O O H H H H H H H2H2 H2H2 O2O2 H 2 O H2OH2O + Atoms can be rearranged into different substances.

15. 45 g H 2 O? g H 2 O Conservation of Mass Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 204 High voltage Before reaction electrodes glass chamber 5.0 g H 2 80 g O 2 300 g (mass of chamber) + 385 g total H2H2 O2O2 High voltage After reaction 0 g H 2 40 g O 2 300 g (mass of chamber) + 385 g total O2O2 H2OH2O

16. Law of Definite Proportions Joseph Louis Proust (1754 – 1826) Each compound has a specific ratio of elements It is a ratio by mass Water is always 16 grams of oxygen for every 2 grams of hydrogen

17. Law of Definite Proportions 103 g of copper carbonate 53 g of copper 40 g of oxygen10 g of carbon + + Whether synthesized in the laboratory or obtained from various natural sources, copper carbonate always has the same composition. Analysis of this compound led Proust to formulate the law of definite proportions.

18. Law of Multiple Proportions John Dalton (1766 – 1844) If two elements form more than one compound, the ratio of the second element that combines with the first element in each is a simple whole number. H 2 O H 2 O 2 water hydrogen peroxide Ratio of oxygen is 1:2 (an exact ratio)

19. Dalton’s Atomic Theory English chemist in the early 1800’s Dalton stated that elements consisted of tiny particles called atoms He also called the elements pure substances because all atoms of an element were identical and that in particular they had the same mass.

20. Dalton’s Atomic Theory 1.All matter consists of extremely small particles that are indivisible and indestructible called atoms. 2.Atoms of a given element are identical in their physical and chemical properties. Atoms of different elements have different physical and chemical properties. 3.Atoms of different elements combine in simple, whole number ratios to form chemical compounds. 4. Atoms cannot be subdivided, created or destroyed. 5.In chemical reactions, atoms are combined, separated, or arranged. Although some exceptions have been discovered, the theory still stands today, and has been expanded and modified.

21. Daltons’ Models of Atoms Carbon dioxide, CO 2 Water, H 2 O Methane, CH 4

22. Structure of Atoms Scientist began to wonder what an atom was like. Was it solid throughout with no internal structure or was it made up of smaller, subatomic particles? It was not until the late 1800’s that evidence became available that atoms were composed of smaller parts.

23. The Structure of the Atom Chapter 3

24. Structure of Atoms The cathode ray tube led to the discovery of electrons, small, negatively charged particles at the turn of the century. J. J. Thomson - English physicist Made a piece of equipment called a cathode ray tube. It is a vacuum tube - all the air has been pumped out.

25. A Cathode Ray Tube Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 58

26. Thomson’s Experiment + - vacuum tube metal disks voltage source

27. Thomson’s Experiment + - voltage source OFF ON Passing an electric current makes a beam appear to move from the negative to the positive end

28. Thomson’s Experiment + - voltage source OFF ON

29. Thomson’s Experiment + - voltage source OFF ON + - By adding a magnetic field… he found that the moving pieces were negative.

30. J.J. Thomson He proved that ALL atoms of any element must contain these negative particles. He knew that atoms did not have a net negative charge and so there must be balancing the negative charge. J.J. Thomson

31. Plum Pudding Model In 1910 proposed the Plum Pudding model Negative electrons were embedded into a positively charged spherical cloud. Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 56 Spherical cloud of Positive charge Electrons

32. Television Picture Tube Fluorescent screen Shadow mask Glass window Blue beam Green beam Red beam Electron gun Electron beam Deflecting electromagnets

33. Ernest Rutherford (1871-1937) Learned physics in J.J. Thomson’ lab. Noticed that ‘alpha’ particles were sometime deflected by something in the air. Gold-foil experiment

34. Rutherford ‘Scattering’ In 1909 Rutherford undertook a series of experiments He fired  (alpha) particles at a very thin sample of gold foil According to the Thomson model the  particles would only be slightly deflected Rutherford discovered that they were deflected through large angles and could even be reflected straight back to the source particle source Lead collimator Gold foil  

35. Rutherford’s Apparatus beam of alpha particles radioactive substance gold foil circular ZnS - coated fluorescent screen Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 120

36. What 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 California WEB

37. What he expected… California WEB

38. What he got…

39. Density and the Atom Since most of the particles went through, the atom was mostly empty. Because the alpha rays were deflected so much, the positive pieces it was striking were heavy. Small volume and big mass = big density This small dense positive area is the nucleus California WEB

40. Rutherford’s Experiment Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 56

41. The Rutherford Atom Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 57 n +

42. Size of an atom Atoms are incredibly tiny. Measured in picometers (10 -12 meters) Hydrogen atom, 32 pm radius Nucleus tiny compared to atom Radius of the nucleus near 10 -15 m. Density near 10 14 g/cm If the atom was the size of a stadium, the nucleus would be the size of a marble. California WEB

43. Atoms All atoms have similar structure protons and neutrons cluster together to form a nucleus, or central core electrons orbit the space surrounding the nucleus then things change depending on the element

44. Counting Atoms Chapter 3

45. Atoms Each element has a characteristic number of protons Si- silicon 14 protons H- hydrogen 1 ptoton Ag- silver 47 protons ↓ Atomic number

46. Atomic Particles Protons are constant for an element, but electrons and neutrons can vary When electrons vary, the charge of the atom changes When neutrons vary, you have a different isotope of the atom Isotope- one of 2 or more atoms having the same number of protons but different numbers of neutrons

47. Isotopes Hydrogen 3 isotopes protium deuterium tritium 1 proton 1 proton 1 proton 0 neutrons 1 neutron 2 neutrons isotopes have very similar chemical properties Deuterium or hydrogen-2

48. Isotopes Contain the symbol of the element, the mass number and the atomic number X Mass number Atomic number # protons + # neutrons mass number

49. Isotopes Atomic Number = number of protons # of protons determines kind of atom/element Atomic Number = number of electrons in a neutral atom Mass Number = the number of protons + neutrons California WEB

50. Mass Number mass # = protons + neutrons always a whole number NOT on the Periodic Table! + + + + + + Nucleus Electrons Nucleus Neutron Proton Carbon-12 Neutrons 6 Protons 6 Electrons6

51. Isotopes Find the number of protons number of neutrons number of electrons Atomic number Mass number F 19 9 = 9 = 10 = 9 = 19 +

52. Isotopes Find the –number of protons –number of neutrons –number of electrons –Atomic number –Mass number Br 80 35 = 35 = 45 = 35 = 80

53. Isotopes Find the number of protons number of neutrons number of electrons Atomic number Mass number Na 23 11 Sodium atom

54. Isotopes Find the number of protons number of neutrons number of electrons Atomic number Mass number Na 23 11 1+ Sodium ion = 11 = 12 = 10 = 11 = 23

55. Isotopes If an element has an atomic number of 23 and a mass number of 51 what is the –number of protons –number of neutrons –number of electrons –Complete symbol V 51 23 = 23 = 28 = 21 +2

56. Isotopes If an element has 60 protons and 144 neutrons what is the –Atomic number –Mass number –number of electrons –Complete symbol Nd 204 60 = 60 = 204 = 60

57. 56 Can atoms be counted or measured? atomic mass- the mass of an atom in atomic mass units atomic mass unit- 1/12 of the mass of the carbon-12 isotope Atoms too small to measure in grams. Created new unit. Carbon-12 atom was assigned a value of 12 atomic mass units (amu).

58. 57 22.990 is the average atomic mass of all the isotopes a weighted average

59. Isotopes The percent natural abundances for mercury isotopes are: Hg-196 0.146% Hg-196 0.146% Hg-198 10.02% Hg-198 10.02% Hg-199 16.84% Hg-199 16.84% Hg-200 23.13% Hg-200 23.13% Hg-201 13.22% Hg-201 13.22% Hg-202 29.80% Hg-202 29.80% Hg-204 6.85% Hg-204 6.85% (0.00146)(196) + (0.1002)(198) + (0.1684)(199) + (0.2313)(200) + (0.1322)(201) + (0.2980)(202) + (0.0685)(204) = x 0.28616 + 19.8396 + 33.5116 + 46.2600 + 26.5722 + 60.1960 + 13.974 = x x = 200.63956 amu Hg 200.59 16

60. 59 Example: Si 92.21% of Si atoms in nature have mass of 27.98 amu 4.70% of Si atoms in nature have mass of 28.98 amu 3.09% of Si atoms in nature have mass of 29.97 amu What is the average of these masses, taking into consideration how often they are found in nature (abundance)? (92.21)(27.98) + (4.70)(28.98) + (3.09)(29.97) 100 = 28.09 amu

61. 60 Example: Cl 2 isotopes occur naturally, chlorine-35, which has a natural abundance of 75% and a mass of 34.969 amu and chlorine-37 which occurs only 25% of the time and has a mass of 36.966 amu. What is the average atomic mass of chlorine? chlorine-35 75% 34.969 amu chlorine-37 25% 36.966 amu (75)(34.969) + (25)(36.966) = 35.453 amu 100

62. 61 Mole discussing tiny particles and tiny amounts cannot create and work with in lab chemists derived a new unit as a bridge between the microscopic and macroscopic worlds fundamental SI unit used to measure the amount of a substance

63. 62 collection of 6.022137 X 10 23 particles or usually 6.022 X 10 23 (Avogadro’s number) 1 mole of any substance contains 6.022 X 10 23 particles 1 mole of oxygen contains 6.022 X 10 23 atoms 1 mole of water contains 6.022 X 10 23 molecules Just like a dozen is always 12 pieces!!!

64. 63 molar mass- the mass in grams of 1 mole of a given substance molar mass = average atomic mass in grams C → 1 mol = 6.022 X 10 23 C atoms = 12.011 g of C Fe → 1 mol = 6.022 X 10 23 Fe atoms = 55.85 g of Fe Mo → 1 mol = 6.022 X 10 23 Mo atoms = 95.94 g of Mo So…

65. 64 1 mol = 6.022 X 10 23 atoms for any substance can be written as conversion factor 1 mol or 6.022 X 10 23 6.022 X 10 23 1 mol

66. 65 How many atoms in 2.5 mols of Si? 2.5 mol Si 1 mol = 6.022 X 1023 atoms ? = # of atoms 2.5 mol Si 6.022 X 10 23 atoms 1 mol = 1.5 X 10 24 atoms Si

67. 66 How many mols is 9.7 X 10 24 atoms of Cu? 9.7 X 10 24 atoms 1 mol = 6.022 X 10 23 atoms 9.7 X 1024 atoms 1 mol 6.022 X 10 23 atoms = 16 mol Cu

68. 67 Determine the mass in grams of 3.5 mols of Cu. 3.5 mol Cu molar mass of Cu = 63.55g, which means 1 mol = 63.55g of Cu mass of Cu = ? 3.5 mol Cu 63.55 g Cu = 222 g Cu 1 mol

69. 68 What is the mass of a single atom of Si? molar mass of Si = 28.09g 1 mol = 6.022 X 10 23 atoms mass of 1 atom = ? 28.09g Si 1 mol 1 mol 6.022 X 10 23 atoms = 4.665 X 10 -23 g/atom