2. Unit 3: “Atomic Structure” Chemistry: Mr. Blake/Mr. Gower

3. Section 4.1 Defining the Atom The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos”) –He believed that atoms were indivisible and indestructible –His ideas did agree with later scientific theory, but did not explain chemical behavior, and was not based on the scientific method – but just philosophy

4. Dalton’s Atomic Theory (experiment based!) 3)Atoms of different elements combine in simple whole-number ratios to form chemical compounds 4)In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element. 1)All elements are composed of tiny indivisible particles called atoms 2)Atoms of the same element are identical. Atoms of any one element are different from those of any other element. John Dalton (1766 – 1844)

5. Problems with Dalton’s Atomic Theory? 1. matter is composed of indivisible particles Atoms Can Be Divided, but only in a nuclear reaction 2. all atoms of a particular element are identical Does Not Account for Isotopes (atoms of the same element but a different mass due to a different number of neutrons)! Different elements have different atoms. YES! 3. atoms combine in certain whole-number ratios YES! Called the Law of Definite Proportions 4. In a chemical reaction, atoms are merely rearranged to form new compounds; they are not created, destroyed, or changed into atoms of any other elements. Yes, except for nuclear reactions that can change atoms of one element to a different element

6. Sizing up the Atom  Elements are able to be subdivided into smaller and smaller particles – these are the atoms, and they still have properties of that element  If you could line up 100,000,000 copper atoms in a single file, they would be approximately 1 cm long  Despite their small size, individual atoms are observable with instruments such as scanning tunneling (electron) microscopes

7. Section 4.2 Structure of the Nuclear Atom One change to Dalton’s atomic theory is that atoms are divisible into subatomic particles: –Electrons, protons, and neutrons are examples of these fundamental particles –There are many other types of particles, but we will study these three

8. Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron

9. Modern Cathode Ray Tubes  Cathode ray tubes pass electricity through a gas that is contained at a very low pressure. TelevisionComputer Monitor

10. Mass of the Electron 1916 – Robert Millikan determines the mass of the electron: 1/1840 the mass of a hydrogen atom; has one unit of negative charge The oil drop apparatus Mass of the electron is 9.11 x 10 -28 g

11. Conclusions from the Study of the Electron: a)Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. b)Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons c)Electrons have so little mass that atoms must contain other particles that account for most of the mass

12. Conclusions from the Study of the Electron:  Eugen Goldstein in 1886 observed what is now called the “proton” - particles with a positive charge, and a relative mass of 1 (or 1840 times that of an electron)  1932 – James Chadwick confirmed the existence of the “neutron” – a particle with no charge, but a mass nearly equal to a proton

13. Subatomic Particles ParticleChargeMass (g)Location Electron (e - ) 9.11 x 10 -28 Electron cloud Proton (p + ) +1 1.67 x 10 -24 Nucleus Neutron (n o )0 1.67 x 10 -24 Nucleus

14. Nucleus Electron cloud

15. Thomson’s Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model. J. J. Thomson

16. Ernest Rutherford’s Gold Foil Experiment - 1911  Alpha particles are helium nuclei - The alpha particles were fired at a thin sheet of gold foil  Particles that hit on the detecting screen (film) are recorded

17. Rutherford’s problem: In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target? Target #1 Target #2

18. The Answers: Target #1 Target #2

19. Rutherford’s Findings a) The nucleus is small b) The nucleus is dense c) The nucleus is positively charged  Most of the particles passed right through  A few particles were deflected  VERY FEW were greatly deflected “Like howitzer shells bouncing off of tissue paper!” Conclusions:

20. The Rutherford Atomic Model Based on his experimental evidence: –The atom is mostly empty space –All the positive charge, and almost all the mass is concentrated in a small area in the center. He called this a “nucleus” –The nucleus is composed of protons and neutrons (they make the nucleus!) –The electrons distributed around the nucleus, and occupy most of the volume –His model was called a “nuclear model”

21. Atomic Number Atoms are composed of identical protons, neutrons, and electrons –How then are atoms of one element different from another element? Elements are different because they contain different numbers of PROTONS The “atomic number” of an element is the number of protons in the nucleus # protons in an atom = # electrons

22. Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Element# of protonsAtomic # (Z) Carbon66 Phosphorus15 Gold79

23. Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope: Mass # = p + + n 0 Nuclidep+p+ n0n0 e-e- Mass # Oxygen - 10 -3342 - 3115 8 8 18 Arsenic 75 33 75 Phosphorus 15 31 16

24. Complete Symbols Contain the symbol of the element, the mass number and the atomic number. X Mass number Atomic number Subscript → Superscript →

25. Symbols n Find each of these: a) number of protons b) number of neutrons c) number of electrons d) Atomic number e) Mass Number Br 80 35

26. Symbols n If an element has an atomic number of 34 and a mass number of 78, what is the: a) number of protons b) number of neutrons c) number of electrons d) complete symbol

27. Symbols n If an element has 91 protons and 140 neutrons what is the a) Atomic number b) Mass number c) number of electrons d) complete symbol

28. Symbols n If an element has 78 electrons and 117 neutrons what is the a) Atomic number b) Mass number c) number of protons d) complete symbol

29. Isotopes Dalton was wrong about all elements of the same type being identical Atoms of the same element can have different numbers of neutrons. Thus, different mass numbers. These are called isotopes.

30. Isotopes Frederick Soddy (1877-1956) proposed the idea of isotopes in 1912 Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons. Soddy won the Nobel Prize in Chemistry in 1921 for his work with isotopes and radioactive materials.

31. Naming Isotopes We can also put the mass number after the name of the element: –carbon-12 –carbon-14 –uranium-235

32. Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons. IsotopeProtonsElectronsNeutronsNucleus Hydrogen–1 (protium) 110 Hydrogen-2 (deuterium) 111 Hydrogen-3 (tritium) 112

33. Isotopes Elements occur in nature as mixtures of isotopes. Isotopes are atoms of the same element that differ in the number of neutrons.

34. IONS Ions are atoms or groups of atoms with a positive or negative charge. Taking away an electron from an atom gives a cation with a positive charge Adding an electron to an atom gives an anion with a negative charge. To tell the difference between an atom and an ion, look to see if there is a charge in the superscript! Examples: Na + Ca +2 I - O -2 Na Ca I O

35. A cation forms when an atom loses one or more electrons. An anion forms when an atom gains one or more electrons Mg --> Mg 2+ + 2 e- F + e- --> F -

36. In general metals (Mg) lose electrons ---> cationsmetals (Mg) lose electrons ---> cations nonmetals (F) gain electrons --->nonmetals (F) gain electrons ---> anions

37. Learning Check – Counting State the number of protons, neutrons, and electrons in each of these ions. 39 K + 16 O -241 Ca +2 198 20 #p + ___________________ #n o ___________________ #e - ___________________ 19 20 18 8 8 10 20 21 18

38. One Last Learning Check Write the nuclear symbol form for the following atoms or ions: A. 8 p +, 8 n, 8 e - ___________ B.17p +, 20n, 17e - ___________ C. 47p +, 60 n, 46 e - ___________

39. Charges on Common Ions -2-3 +1 +2 By losing or gaining e-, atom has same number of e-’s as nearest Group 8A atom.

40. Atomic Mass  How heavy is an atom of oxygen?  It depends, because there are different kinds of oxygen atoms.  We are more concerned with the average atomic mass.  This is based on the abundance (percentage) of each variety of that element in nature.  We don’t use grams for this mass because the numbers would be too small.

41. Measuring Atomic Mass Instead of grams, the unit we use is the Atomic Mass Unit (amu) It is defined as one-twelfth the mass of a carbon-12 atom. –Carbon-12 chosen because of its isotope purity. Each isotope has its own atomic mass, thus we determine the average from percent abundance.

42. To calculate the average: Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results. If not told otherwise, the mass of the isotope is expressed in atomic mass units (amu)

43. Atomic Masses IsotopeSymbolComposition of the nucleus % in nature Carbon-12 12 C6 protons 6 neutrons 98.89% Carbon-13 13 C6 protons 7 neutrons 1.11% Carbon-14 14 C6 protons 8 neutrons <0.01% Atomic mass is the average of all the naturally occurring isotopes of that element. Carbon = 12.01

44. - Page 117 Question Solution Answer Knowns and Unknown

45. 3. Calculate the atomic mass of carbon. Isotopes% abundance Atomic mass Carbon-12 98.89% 12.000 amu Carbon-13 1.11% 13.003 amu a. Lithium has two isotopes. If lithium-6 has a mass of 6.015 and 7.42 % occurrence, what is the % abundance and mass of lithium -7? Atomic mass= (%) 1 (mass) 1 +(%) 2 (mass) 2 +….. Atomic mass= (0.9889) 1 (12.000) 1 +(0.0111) 2 (13.003) 2 = 12.0098 amu Periodic Table 6.941= (0.0742)(6.015) +(0.9258)(x) = 7.015 amu

46. The Periodic Table: A Preview  A “periodic table” is an arrangement of elements in which the elements are separated into groups based on a set of repeating properties  The periodic table allows you to easily compare the properties of one element to another

47. The Periodic Table: A Preview  Each horizontal row (there are 7 of them) is called a period  Each vertical column is called a group, or family  Elements in a group have similar chemical and physical properties  Identified with a number and either an “A” or “B”

48. Section 6.1 Organizing the Elements A few elements, such as gold and copper, have been known for thousands of years - since ancient times Yet, only about 13 had been identified by the year 1700. As more were discovered, chemists realized they needed a way to organize the elements.

49. Section 6.1 Organizing the Elements Chemists used the properties of elements to sort them into groups. In 1829 J. W. Dobereiner arranged elements into triads – groups of three elements with similar properties –One element in each triad had properties intermediate of the other two elements

50. Mendeleev’s Periodic Table By the mid-1800s, about 70 elements were known to exist Dmitri Mendeleev – a Russian chemist and teacher Arranged elements in order of increasing atomic mass Thus, the first “Periodic Table”

51. Mendeleev He left blanks for yet undiscovered elements –When they were discovered, he had made good predictions But, there were problems: –Such as Co and Ni; Ar and K; Te and I

52. A Better Arrangement In 1913, Henry Moseley – British physicist, arranged elements according to increasing atomic number The arrangement used today The symbol, atomic number & mass are basic items included-textbook page 162 and 163

54. The Periodic Law When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties. Horizontal rows = periods –There are 7 periods Vertical column = group (or family) –Similar physical & chemical prop. –Identified by number & letter (IA, IIA)

55. Areas of the Periodic Table Three classes of elements are: 1) metals, 2) nonmetals, and 3) metalloids 1) Metals: electrical conductors, have luster, ductile, malleable 2) Nonmetals: generally brittle and non- lustrous, poor conductors of heat and electricity

56. Areas of the Periodic Table Some nonmetals are gases (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br) Notice the heavy, stair-step line? 3) Metalloids: border the line-2 sides –Properties are intermediate between metals and nonmetals

57. Squares in the Periodic Table The periodic table displays the symbols and names of the elements, along with information about the structure of their atoms: Atomic number and atomic mass Black symbol = solid; red = gas; blue = liquid (from the Periodic Table on our classroom wall)

59. Groups of Elements - Family Names Group IA (1) – alkali metals –Forms a “base” (or alkali) when reacting with water (not just dissolved!) Group 2A (2)– alkaline earth metals –Also form bases with water; do not dissolve well, hence “earth metals” Group 7A (17) – halogens –Means “salt-forming” Group 8A (18) – noble gases –Nonreactive because of their electron configuration

60. ELEMENTS THAT EXIST AS DIATOMIC MOLECULES Remember: HOFBrINCl These elements only exist as PAIRS. Note that when they combine to make compounds, they are no longer elements so they are no longer in pairs!

61. Nuclear Chemistry Chemistry – Unit 4 Chapter 25

62. Mass Defect Difference between the mass of an atom and the mass of its individual particles. 4.00260 amu4.03298 amu

63. Nuclear Binding Energy Energy released when a nucleus is formed from nucleons. High binding energy = stable nucleus. E = mc 2 E:energy (J) m:mass defect (kg) c:speed of light (3.00×10 8 m/s)

64. Nuclear Binding Energy Unstable nuclides are radioactive and undergo radioactive decay.

65. Types of Radiation Alpha particle (  ) –helium nucleus paper 2+ Beta particle (  -) –electron 1- lead Positron (  +) –positron 1+ Gamma (  ) –high-energy photon 0 concrete

66. Nuclear Decay Alpha Emission parent nuclide daughter nuclide alpha particle Numbers must balance!!

67. Nuclear Decay Beta Emission electron Positron Emission positron

68. Nuclear Decay Electron Capture electron Gamma Emission –Usually follows other types of decay. Transmutation –One element becomes another.

69. IQ# 1 1.Balance the following equations:

70. Nuclear Decay Why nuclides decay… –need stable ratio of neutrons to protons

71. Belt of Stability and Radioactive Decay

72. Half-life Half-life (t ½ ) –Time required for half the atoms of a radioactive nuclide to decay. –Shorter half-life = less stable.

73. Half-life m f :final mass m i :initial mass n:# of half-lives

74. Half-life Fluorine-21 has a half-life of 5.0 seconds. If you start with 25 g of fluorine-21, how many grams would remain after 60.0 s? GIVEN: t ½ = 5.0 s m i = 25 g m f = ? total time = 60.0 s n = 60.0s ÷ 5.0s =12 WORK : m f = m i (½) n m f = (25 g)(0.5) 12 m f = 0.0061 g

75. Example: How much of a 500. g sample of Uranium-235 would be left after five half-lives? Example: A 16.00 mg sample of Radon-222 decays to 0.250 mg after 24 hours. Determine the half-life. 16 → 8 → 4 → 2 → 1 → 0.5 → 0.250 = 6 half lives (n = # of half-lives) M i = 500 g n = 5 M f = ? m f = m i (½) n m f = (500 g)(0.5) 5 m f = 15.6 g

76. Example: The half-life of molybdenum-99 is 67 hours. How much of a 1.000 mg sample is left after 335 hours? M i = 1.000 mg Half-life = 67 h Rxn time = 335 h M f = ? n = 335 / 67 = 5 m f = m i (½) n m f = (1.000 mg)(0.5) 5 m f = 0.03125 mg

77. Learning Check! The half life of I-123 is 13 hr. How much of a 64 mg sample of I-123 is left after 39 hours? M i = 64 mg n = 3 M f = ? m f = m i (½) n m f = (64 mg)(0.5) 3 m f = 8.0 g

78. Half Life and Radioactivity Lab Work in groups of 2 at your table. Each cup has 1 penny in it which will be shaken and then GENTLY emptied on the table. For the first trial, shake the penny out 100 times on the table. Record the number of times that it came up heads. For the next trail, you will shake out the penny the number of times that it landed on heads in the last round. The same procedure will follow until no more of the pennies have landed on “heads” (tails = decayed). Record all data in the lab book following the example on page 809. Answer questions 1-4 and be sure to follow the graphing rules (R74 and in the “Math Review” handout from the beginning of the year).

79. Graphing the Results Important !! Graph directly on lab book Title every graph and label each axis Graph at least 2/3 page Use a ruler Circle all data points Use a best-fit line (no “connect the dots”!) 5) Find the average half-life (in # of trials) of your sample by interpolating your curve at exactly 50, 25, and 12.5 flips)

81. F ission splitting a nucleus into two or more smaller nuclei 1 g of 235 U = 3 tons of coal

82. F ission chain reaction - self-propagating reaction critical mass - mass required to sustain a chain reaction

83. Fusion combining of two nuclei to form one nucleus of larger mass thermonuclear reaction – requires temp of 40,000,000 K to sustain 1 g of fusion fuel = 20 tons of coal occurs naturally in stars

84. Fission vs. Fusion 235 U is limited danger of meltdown toxic waste thermal pollution fuel is abundant no danger of meltdown no toxic waste not yet sustainable FISSIONFISSION FUSIONFUSION

85. Nuclear Power Fission Reactors Cooling Tower

86. Nuclear Power Fission Reactors

87. Nuclear Power Fusion Reactors (not yet sustainable)

88. Nuclear Power Fusion Reactors (not yet sustainable) Tokamak Fusion Test Reactor Princeton University National Spherical Torus Experiment

89. Synthetic Elements Transuranium Elements –elements with atomic #s above 92 –synthetically produced in nuclear reactors and accelerators –most decay very rapidly

90. Radioactive Dating half-life measurements of radioactive elements are used to determine the age of an object decay rate indicates amount of radioactive material EX: 14 C - up to 40,000 years 238 U and 40 K - over 300,000 years

91. Nuclear Medicine Radioisotope Tracers –absorbed by specific organs and used to diagnose diseases Radiation Treatment –larger doses are used to kill cancerous cells in targeted organs –internal or external radiation source Radiation treatment using  -rays from cobalt-60.

92. Nuclear Weapons Atomic Bomb –chemical explosion is used to form a critical mass of 235 U or 239 Pu –fission develops into an uncontrolled chain reaction Hydrogen Bomb –chemical explosion  fission  fusion –fusion increases the fission rate –more powerful than the atomic bomb

93. Others Food Irradiation –  radiation is used to kill bacteria Radioactive Tracers –explore chemical pathways –trace water flow –study plant growth, photosynthesis Consumer Products –ionizing smoke detectors - 241 Am