1. UEQ Where does the concept and understanding of the atom originate?

2. LEQ Who are the main players in the development of the concept of the atom?

3. Lecture Notes Chapter 4-The Structure of the Atom Early Ideas of the Atom: Democritus: atomos: small cells Aristotle: Rejected the concept of the atom. All things came from earth, wind and fire. John Dalton: English teacher.

4. Dalton’s Atomic Theory: 1. All matter is made-up of tiny indivisible particles called atoms. 2. The atom is the smallest form of matter. 3. Atoms are represented by symbols called elements. 4. Atoms are the foundation of the universe. 5. All atoms of the same elements will be the same. 6. All atoms of different elements will be different. 7. Atoms of different element can combine in definite proportions to form more compounds

5. Question? Dalton was right about most of his theory. What part of Dalton’s Theory is not correct with today’s knowledge? Answer: Atoms are divisible into small subatomic particles.

6. Sir William Crookes Developed the use of the cathode ray tube in the study of the atom. a. cathode ray tube: vacuum tube with a (+) end called the anode and a (-) end called the cathode. b. Crookes study show movement from the cathode (-) to the anode (+). From -  +

7. JJ Thomson Used the cathode ray tube to show the charge to mass ratio of the stream of light in the tube. a. Compare the charge to mass ratio to other known values b. Concluded that the charged particles are less than the mass of the lightest atom known. c. Identified the first subatomic particle called the electron. d. Concluded the atom formed a ‘plum pudding’ model.

8. Cathode Ray Tube

9. Robert Millikan Use Thomson’s charge to mass ratio to define the mass of an electron. Mass of electron = 1/1840 the mass of hydrogen or one electron = 9.1 x 10 -28 g or 9.1 x 10 -31 Kg

10. Milliken Oil Drop Experiment

11. Ernest Rutherford Gold Foil Experiment. a. The atom has a dense center called the nucleus. b. The nucleus has a positive charge. c. The positive charge is due to the protons. d. The atom is mostly empty space. (diagram: Link)

12. Rutherford Gold Fold Experiment

13. James Chadwick Discovered the neutron a. Chadwick was a student of Rutherford. b. Discovered that the nucleus is heavier than could be account for with just protons. c. The extra mass called the neutron and had NO charge. No charge = neutral = neutrons

14. LEQ What is the structure of the atom?

15. Conclusion: Structure of the Atom Within the nucleus: *Protons: (+) charge; half the mass of the atom; number is fixed, unchanging (Rutherford) *Neutrons: NO charge; the other half of the mass of the atom (Chadwick) Outside the nucleus: * Electrons: (-) charge; in the electron cloud about the nucleus; NO mass (JJ Thomson)

16. Conclusion: The Atom 1. Every atom of the same atom has the same number of protons. THE NUMBER OF PROTONS NEVER CHANGES. 2. The number of neutrons can change from atom to atom of the same element. This is called an isotope, (same atom but different number of neutrons). Note: Cesium-131, Cesium-134 3. The number of protons NEVER changes for the same atom. Found in the nucleus.

17. 4. Elements are identified by the number of protons. 5. In a neutral atom (net charge = 0), the # of protons = # of electrons 6. Ion form when # of electrons changes (gain electron(s) yields (-) ion ; loss electron(s) yields (+) ion 6. All the mass of the atom is in the nucleus. 7. The # of neutrons + the # of protons = the mass number. 8. The # of protons = the atomic number 9. Periodic Table (Robert Mosley)

18. Reading the Periodic Table 50 Sn Tin 118.69

19. ElementAtomic NumberProtonsElectron (neutral atom) Fe 16 19

20. ElementsAtomic Number Atomic Mass Mass # Number of Neutrons Cl Mg Zn La

21. ElementsAtomic Number Atomic Mass Mass Number ProtonsNeutrons Br 55 26.9815

22. Questions ElementSymbolAtomic #Atomic Mass #e- for neutral atom Oxygen 20 Br Magnesium 83 Iodine 11 1.

23. Each atom is a neutral atom ElementSymbolAtomic #Mass Number#p + #n o #e - 92 Si 17 Tin 16 55 Al Note, the #e- will indicate for a neutral atom.

24. Indicate the magnitude of the ion ElementSymbolAtomic #Mass Number Atomic Mass#p + #n o #e - Ion? 9 10 Zinc 65 65.39 38 36 Argon 39.948 K 39 18 Manganese 54.938 15 18 3. Note, the number of #e- may indicate a neutral atom or an ion. Indicate the magnitude of the ion for those elements indicated.

25. All atoms are indicated as neutral Atomic #Atomic Mass#p+#no#e-Element 37 26 79 36 80 5 21 All the atoms listed here are neutral.

26. LEQ What is the significant role played by each part of the atom?

27. Ions Formed when there is a different number of electrons to protons (# of protons for an atom NEVER changes Charged form of an atom brought about by the gain of electrons (-) or loss of electrons (+) Ions are formed: 1. response to another atom 2. metals form (+) ions 3. non-metal form (-) ions 4. unlike charges attract; like charges repel.

28. Oxidation Numbers Ions will allow you to assign oxidation numbers to each element. **Number of protons (+) NEVER changes **Number of electrons (-) changes. No mass **Difference in protons to electrons = oxidation numbers Mg: 12 p+ and 12 e- = ______ oxidation number Mg: 12 p+ and 10 e- = ______ oxidation number

29. Element Atomic Number Number of Protons Number of electrons Oxidation Number O 2- 2018 F1- 562+

30. Isotopes Same atoms with different number of neutron in the nucleus: **Mass of one neutron = mass of one proton **Accounts for different masses for the same atom. **amu: atomic mass units, is the average of the mass in one mole. **Accounts for some radioactivity

31. Germanium-70Germanium-73Germanium-76 Protons Neutrons Electrons

32. Calculating the Average Atomic Mass NOTE: 1. The relative abundance is the decimal form of the % abundance. % AbundanceRelative Abundance 56%  0.56 7%  0.07

33. Relative Abundance How much is present based on the amount of the sample given. -It is the decimal form of the percentage -Used to find the amu.

34. Calculating the Average Atomic Mass 2. Mass of Isotope times the relative abundance will equal the mass contributed by that isotope to the total mass 3. The sum of each contributed mass will equal the amu.

35. Calculate the Average Atomic Mass IsotopePercent AbundanceMass Cr-504.35 %49.946 Cr-5283.79 %51.941 Cr-539.5 %52.941 Cr-542.36 %53.939

36. Sampler In a given sample of an element, 81.3% has a mass of 32.45 g and 18.7% has a mass of 30.39 g. What is the amu?

37. Sampler 1. Copper has two isotopes. Cu-63 has an amu of 62.930 occurring 69.17% and Cu-65 has an amu of 64.928 occurring 30.83% of the time is a sample. Find the atomic mass.

38. 2. Find the element by calculating the amu. IsotopeAbundance (%)Atomic Mass X-7021.2369.924 X-7227.6671.922 X-737.7372.923 X-7435.9473.921 X-767.4475.921

39. 3. Chlorine has two natural occurring isotopes. Cl- 35 with an atomic mass of 34.969 occurring 75.77% of the time, and Cl-37 that occurs 24.23% of the time. What is the atomic mass of Cl-37?

40. Radioactivity The spontaneous emission of radioactive (high energy) particles. Radioactivity is caused by a proton to neutron ratio. The greater the difference between the p + and n o, the more unstable. An unstable atom will emit energy until a more stable form is reached. This is called radioactive decay.

41. Types of Radiation, page 122 Alpha Decay Alpha radiation: (α), contains two protons and two neutrons giving a 2+ charge and a change in the mass of 4. Note that the mass and the atomic number changes. For Rn-222: 222 Rn 86  218 Po 84 + 4 He 2 226 Ra 88  210 Po 84 

42. Beta Decay Beta radiation: (β), contains a single electron and carries a 1 - charge. Note the mass does not change and the number of protons increase to balance the 1- charge. For Bi- 241 241 Bi 83  241 Po 84 + 0 e 1- 228 Ra 88  14 C 6 

43. Gamma Decay Gamma radiation: (γ), Very high energy particle with no mass and no charge. Accompanies a alpha or a beta decay. For U-238 238 U 92  234 Th 90 + 4 He 2 + 0 γ 0 244 Pu 94  200 U 92 + 209 Bi 83  209 Po 84 +

44. Radioactive Questions 1.Why are some atoms radioactive while some are not? 2.Write the symbols used to denote alpha, beta, and gamma radiation. Write the mass and charge of each. 3.What is the primary factor that determines if a nucleus is stable or unstable? 4.Boron-10 emits alpha particles and Cesum-137 emits beta particles. Write a balance equation.

45. Half Life Half Life is the time intervals it takes for ½ the amount of material to undergo radioactive decay. 100 % sample  50%  25%  12.5%  6.25% ……… Two ways to find Half Life Intervals (n): 1. n = t / t 1/2 t = time of the decay t 1/2 = half life 2. Divide the original amount by 2 until the amount decay a and the end of time ‘t’ is reached.

46. Half Life To calculate the amount of material decayed: A = A o / 2 n where A = amount that remains at the end of time ‘t’. A o = original or initial amount of radioactive material n = number of half life intervals

47. Calculating Half Life 1.How much of a 1 kg sample of K-39 remains after 151 years of decay? The half life for K-39 is 30.2 years.t = t 1/2 = A = A o = 2.Given 64 g of an unknown sample undergoes radioactive decay for 12.5 hours to yield an amount of 2 g. What is the half life of this sample? t = t 1/2 = A = A o =

48. Calculating Half Life 3.Cobalt-60 has a half life of 5.26 years. How much time is needed for a 16 g sample of Co-60 to decay to 1 g? 4. How much of Nb-259 will remain after 290 min if you begin with 125 g. The half life of Nb-259 is 58 minutes.