1. Atoms: The Building Blocks of Matter Chapter 3

2. Development of Atomic Theory

3. In the Beginning… Democritus—matter is composed of small, indivisible particles Atoms – “indivisible” Greeks—four “elements” make up all substances: earth, water, fire, and air

4. Slow Progress No experimental evidence for atoms—even though Democritus was right, everyone believed Aristotle. Next 2 000 years--alchemy Discoveries: elements and preparation of mineral acids

5. “Modern” Chemistry 17 th century—chemists worked with metals and used elements & compounds for medicine. Robert Boyle—first quantitative experiments—pressure/volume relationships in gases Evidence led to support for atomic theory of matter

6. New Ideas Boyle—An element is any substance that cannot be reduced to a simpler substance. Greek idea of elements died. Joseph Priestly—discovered oxygen, “dephlogisticated air”— led to study of combustion

7. Fundamental Laws

8. Lavoisier Law of Conservation of Mass—mass is neither created nor destroyed in a chemical reaction

9. Proust Law of Definite Proportions—a given compound always contains the same proportion of elements by mass H 2 0 NaCl NaOH AlCl 3

10. Example: H 2 O is always 2 parts hydrogen and 16 parts oxygen by mass. Example: CO 2 is always 12 parts carbon and 32 parts oxygen by mass.

11. Dalton Law of Multiple Proportions—When two elements form a series of compounds, the ratios of the masses of he second element that combines with 1 g of the first element can always be reduced to small whole numbers.

12. Law of Multiple Proportions (Cont’d) Examples: CO CO 2 Always a 1:2 ratio in the amount of oxygen used to make CO & CO 2

13. Dalton’s Atomic Theory 1. Each element is made up of tiny particles called atoms. 2. Atoms of a given element are identical in size, mass,and other properties. Atoms of different elements are different. 3. Atoms cannot be subdivided, created, or destroyed.

14. Dalton’s Atomic Theory 4. Atoms of different elements combine in simple, whole- number ratios to form compounds. 5. Chemical reactions involve the reorganization of atoms, but the atoms themselves are not changed.

15. Dalton’s Atomic Theory By relating atoms to mass (something that could be measured) Dalton developed Democritus’ idea into a theory. Notice, not all of Dalton’s points are correct!

16. Modern Definition of an Atom An atom is the smallest particle of an element that retains the chemical properties of that element.

17. Structure of the Atom

18. Discovery of Electrons J.J. Thomson—experiments with cathode ray tubes (CRT’s)

19. Cathode Ray Experiment Applying voltage causes a glow to travel from the negative to positive end glass tube in which a partial vacuum exists. What can we learn from a ray??

20. Cathode Ray Observations and Conclusions Ray is deflected by a magnetic field. Has electromagnetic properties Ray travels away from negative and toward positive probably is negatively charged

21. Deflection of cathode rays by an applied electric field.

22. Cathode Ray Observations and Conclusions Any metal will produce a ray. All atoms must contain these particles. Atoms are electrically neutral. Some positive particle must be present to balance the negative charge.

23. Charge vs. Mass Thomson was able to calculate the charge to mass ratio 1.7 x 10 11 Coulombs / kg VERY LARGE charge for a very small mass

24. New Model Plum pudding—atoms are spherical masses of positive charge with electrons scattered throughout Since electrons have a small mass, the mass must come from something else.

25. Charge of an Electron Robert Millikan’s oil drop experiment Oil drops can be suspended in an electric field by adjusting voltage of charged plates

26. Millikan’s Experiment

27. Millikan’s Conclusion The charge produced on the oil drop was always a whole number multiple of an electron’s charge.

28. Gold Foil Experiment Question: How will positive alpha particles behave when passing through gold foil? Background: “Plum Pudding” model of the atom Hypothesis: All particles should crash through the foil—slight deflection.

29. Expected Results

30. Experiment:  -particle bombardment of metal foil.

31. Actual Results

32. Gold Foil Experiment Result: Most particles went through, but some were deflected and some were reflected. Conclusion: “Plum Pudding” model cannot be correct. Atoms must have a dense, positively charged nucleus.

33. Nuclear Structure Nucleus—protons (+) and neutrons (0) Atomic number (Z)--# of protons (element) Atomic mass (A)—protons + neutrons (different isotopes)

34. How can the positive protons stay together in the tiny nucleus? Nuclear forces—strong forces that only act when particles are VERY close together Much stronger than electrostatic forces that would cause repulsion

35. Size Atomic radii: 40-270 picometers (billionths of a meter) Nuclear radii: 0.001 picometer Nucleus is very small compared to the whole atom, but most of the mass is there—VERY dense

36. Counting Atoms

37. Atomic Number The number of protons in the nucleus of an atom Sometimes written as Z Makes the element what it is Equal to number of electrons if atom is neutral.

38. Isotopes Atoms with the same number of protons but different numbers of neutrons and thus different masses Examples: protium, deuterium, tritium (isotopes of hydrogen with masses of 1, 2, and 3)

39. Mass Number Total number of protons & neutrons in the nucleus of an atom # Neutrons = mass # - atomic #

40. Writing Isotopes Hyphen notation—symbol followed by mass number H-1, H-2, H-3, C-14, U-235 Nuclear symbol—symbol with atomic number as subscript and mass number as superscript H H H C U 1 11 2 1 3 6 12235 92

41. In Your Scientist’s Notebook: 1. How many protons and neutrons are in Pb-208? 2. How many protons and neutrons are in I? 3. Write a nuclear symbol for Ca-42. 132 53

42. Relative Atomic Mass Because measuring atomic mass in grams would be cumbersome, relative mass is measured in atomic mass units (amu) 1 amu = 1/12 mass of carbon-12 Mass of proton & neutron is very close to 1 amu

43. Examples Find atomic masses for copper, chlorine, and helium. Find formula masses for sulfur dioxide, magnesium chloride, and ammonium sulfate

44. In Your Scientist’s Notebook: 4. What is the atomic mass of: Sodium (Na) Iron (Fe) Nitrogen (N) 5. What is the formula mass of: Sodium chloride—NaCl Sulfuric acid—H 2 SO 4 Aluminum hydroxide Al(OH) 3

45. Average Atomic Mass Weighted average of naturally occurring isotopes To calculate:  (relative abundance) x (mass)

46. Examples: Calculate the relative atomic mass of copper if 69.17% of the isotopes are Cu-63 (62.930 amu) and 30.83% are Cu-65 (64.928 amu).

47. 69.17% Cu-63 (62.930 amu) & 30.83% Cu-65 (64.928 amu)

48. Examples: Calculate the average atomic mass of carbon if 98.90% is C- 12 (12.000 amu) and the remainder is C-14 (14.003 amu).

49. 98.90% C-12 (12.000 amu) & C-14 (14.003 amu)

50. In your Scientist’s Notebook: 6. Calculate the average atomic mass of oxygen if 99.76% is O-16 (15.995 amu), 0.04% is O-17 (16.999 amu), and 0.20% is O-18 (17.999 amu)

51. Counting Demo

52. The Mole SI unit for amount of a substance The amount of a substance that contains the same number of particles as exactly 12 g of C-12 6.022 x 10 23 of anything

53. Avogadro’s Number 6.022 x 10 23

54. Molar Mass Related to atomic mass because 1 mole of carbon weighs 12 g. 12 amu per carbon atom : 12 grams per mole of carbon 1 amu : 1 gram per mole

55. Molar Mass One mole of any element has a mass in grams equal to its atomic mass. The mass in grams of one mole of any compound is the sum of the atomic masses of its elements.

56. Conversions Use factor label (dimensional analysis) to convert between grams, moles & number of particles

57. Examples: How many grams of carbon are in 3.2 moles?

58. Examples: How many atoms of carbon are in 3.2 moles?

59. Examples: How many moles are there in 60.0 grams of Ca?

60. Examples: How many grams of water are in 1 mole of H 2 O?

61. In Your Scientist’s Notebook: 7. How many moles of Al in 13.50 g? 8. How many atoms of Xe in.0087 mol? 9. How many atoms of Sn in 43.23 g?