1 Atomic Theory

2 Essential Questions What are we made of?How are scientific models developed? Do atoms exist or are they just concepts invented by scientists? What evidence is there in your everyday life for the existence of atoms? How did the understanding of the atom affect historical events? How have historical events affected the model of the atom?

3 Essential Questions What do we think the atom “looks like” now?If the atom is mostly empty space, why doesn’t my butt fall through the chair? How are light and electrons related? How do we “see” where electrons are located in the atom? Why is the location of electrons so important?

4 Historical BackgroundGreek Philosophers Democritus ( BCE) “atomism” Aristotle ( BCE) Earth Air Fire Water No Experiments

5 Historical BackgroundAlchemy (Up to the Middle Ages) Transmutation of other metals into gold Phlogiston Imaginary element Believed to separate from combustible bodies when burned

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7 Historical BackgroundEarly Experimental Chemists Henry Cavendish ( ) (hydrogen – “inflammable air”) Joseph Priestley ( ) (oxygen) Antoine Lavoisier ( ) (oxygen) Karl Wilhelm Scheele ( ) (oxygen) Count Amedeo Avogadro ( ) (gases  mole)

8 Inventions 2. Which of the following were invented a) before 1800?b) between 1800 and 1900? c) after 1900?

9 Before 1800, between 1800 and 1900, or after 1900??Glass Mercury thermometer Barometer Perfume Gunpowder Guns Hot air balloons Telegraph Electric motor Internal combustion engine Car Battery Rechargeable battery Photography X-ray photography Bunsen burner Gas lights Incandescent light bulb Electric lights Glass blowing Sewing machine Cathode ray tube ( TV) Submarine Vacuum technology Asphalt Tin cans Candle Electricity Conduction of electricity Graphite pencil

10 Before 1800 Glass (3000 BC) Perfume (Egypt – BC)Electricity (static electricity 600 BC) Candle (~ 200 BC, China, whale fat) Glassblowing (50 BC) Gunpowder (800s, China) Handguns (1400s) Graphite pencil (1564, England) Muskets (1600s) Barometer (1608) Vacuum technology (1650, Germany) Mercury thermometer (1714) Conduction of electricity (1729, Ben Franklin 1747) Hot air balloons (1783) Gas lights (1792) Battery (1799, Volta)

11 Between 1800 and 1900 Sewing machine (~1800) Tin cans (1810)Asphalt (1824, Paris) Photography (1826) Electric motor (1831) Incandescent light bulb (1835, Scotland) Telegraph (1838) Bunsen burner (1855) Rechargeable battery (1859, Germany) Electric lights (1870s) Cathode ray tube ( TV) (1878) Internal combustion engine (1886; Daimler, Maybach and Benz ) X-ray photography (1895) Car (1897) Submarine (late 1800s, some earlier)

12 Three Laws (late 1700s, ~1800) Law of Conservation of MassLaw of Definite Proportions Law of Multiple Proportions

13 1. Law of Conservation of MassAntoine Lavoisier born 1740 turned to science in his 20’s “father of modern chemistry” invented a balance that read to g

14 Law of Conservation of Mass mass of products = mass of reactantsLavoisier heated tin in air  tin oxide e.g. 2 Sn(s) O2(g)  2 SnO(s) 50.00 g g  g first experimental evidence for law of conservation of mass

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16 2. Law of Definite ProportionsJoseph Proust (1799) Compounds always contain elements in the same proportion by mass, no matter how they are made or where they are found.

17 2. Law of Definite ProportionsSnO will always be 50.00 g Sn x 100% = 88.12% Sn 56.74 g and 6.74 g O x 100% = 11.88% O no matter where SnO is found or how it is made

18 2. Law of Definite ProportionsIn modern, molar terms: 50.00 g Sn x 1 mol Sn = mol Sn g Sn and 6.74 g O x 1 mol O = mol O 16.00 g O

19 3. Law of Multiple ProportionsJohn Dalton (1803) Different compounds made from the same elements: The ratio of mass of an element in the first compound relative to the mass of the same element in a second compound is a fixed whole number.

20 3. Law of Multiple Proportionse.g. SnO2 vs SnO (two oxides of tin) Begin with 50.0 g Sn in both cases: 56.74 g SnO contains 6.74 g O 63.48 g SnO2 contains g O Ratio of the masses of O in these two compounds: 13.48 g O in SnO2 = 2.00 6.74 g O in SnO

21 3. Law of Multiple ProportionsIn modern, molar terms: 6.74 g O x 1 mol O in SnO = mol O 16.00 g O and 13.48 g O x 1 mol O in SnO2 = mol O Mole ratio of O’s: mol in SnO2= 2.000 mol in SnO

22 Last time The 3 laws Law of conservation of mass,Law of definite proportions Law of multiple proportions (Dalton)  Dalton’s theory of the atom

23 Dalton’s Atomic TheoryAll matter is made of extremely small particles called atoms. All atoms of a given element are identical (mass, physical and chemical properties). Atoms of different elements have different masses, and physical and chemical properties. Different atoms combine in simple whole number ratios to form compounds. In a chemical reaction, atoms are combined, separated, or rearranged. Atoms cannot be created, divided into smaller particles, or destroyed.

24 Dalton’s Atomic TheoryAll matter is made of extremely small particles called atoms. All atoms of a given element are identical (mass, physical and chemical properties). Atoms of different elements have different masses, and different physical and chemical properties. Different atoms combine in simple whole number ratios to form compounds. In a chemical reaction, atoms are combined, separated, or rearranged. Atoms cannot be created, divided into smaller particles, or destroyed.

25 Dalton’s Model of the Atom

26 Subatomic Particle DiscoveryElectron was the first discovered! J.J. Thomson (1897) Received the Nobel Prize in Physics in 1906 for his work on the electron

27 1. Experiment: Cathode RaysPassed electricity through partially evacuated tube of gas Observed a ray of light passing from one electrode to the other Ray moved a paddle wheel inside the tube

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29 Conclusions Ray must be (-) because it moved toward (+) electrodeRay must be made of particles (moved the paddle wheel) Particles must be in all atoms (same results with different gases)

30 J. J. Thomson’s Model of the Atompositively-charged “dough” embedded with small negatively-charged particles called electrons

31 Nucleons Protons: 1909/1910 Ernest RutherfordReceived the Nobel Prize in Chemistry in 1908 for his work in radioactivity

32 Gold Foil Experiment Alpha particles (He nucleus, + charge) shot through gold foil were deflected in peculiar ways, inconsistent with the plum pudding model of the atom  Rutherford’s idea that there is an area of concentration of positive charge

33 Gold Foil Experiment

34 Gold Foil Experiment Hans Geiger and Marsden performed the experiments; Rutherford interpreted them

35 Comparison between Thomson’s and Rutherford’s models of the atom

36 Warm up – Three Laws and Dalton’s Theory1. Name the three laws and briefly describe what they say. Give examples for each one if you can. 2. What are the five parts of Dalton’s theory? a) Make a table with a column describing his theory, then next to this column, indicate which parts of his theory are still considered to be true, and which ones are no longer considered to be correct. b) Analyze your table – is there a common theme that unites the parts of Dalton’s theory that are no longer considered to be correct? If so, what is the common theme? (describe it)

37 Atomic Theory Video Questions #1 The Earliest Models - Mists of Prehistory1. How do Democritus's theories compare with today’s theories? 2. Who did not believe in atoms? 3. Why did the church oppose atomism? 4. What did alchemists try to do? 5. What did Ben Franklin believe lightning to be? 6. What did Lavoisier show?

38 Atomic Theory Video Questions #1 The Earliest Models - Mists of Prehistory1. How do Democritus's theories compare with today’s theories? pretty well 3/5 2. Who did not believe in atoms? Plato, Aristotle 3. Why did the church oppose atomism? Human spirit claimed to be explained by physical means - atoms 4. What did alchemists try to do? Introduced the idea of observation and experimentation, while trying to change common metals into silver and gold 5. What did Ben Franklin believe lightning to be? A fluid – excess = + charge, deficit = - charge 6. What did Lavoisier show? Mass of reactants = mass of products (Law of Conservation of Mass)

39 Atomic Theory Video Questions #1 Smaller than the SmallestWhat was Dalton trying to figure out when he discovered atomic theory? 2. What did Faraday think of electricity? How did Millikan establish an electron’s charge? 4. What is the raisin bun model?

40 Atomic Theory Video Questions #1 Smaller than the SmallestWhat was Dalton trying to figure out when he discovered atomic theory? Why water absorbs more of one type of gas than another 2. What did Faraday think of electricity? = force of affinity that holds atoms together How did Millikan establish an electron’s charge? Microscopic oil droplets which carry a charge – pass X-rays through plates  free electrons 4. What is the raisin bun model? Positively-charged dough with negative charges sprinkled throughout

41 Atomic Theory Video Questions #1 The Rutherford ModelWhat do electrostatic forces affect? 2. What are the two components of the interaction of charged particles? What does his structure of the atom mimic? 4. What was wrong with Rutherford’s model?

42 Atomic Theory Video Questions #1 The Rutherford ModelWhat do electrostatic forces affect? The path/direction of alpha particles as well as their speed 2. What are the two components of the interaction of charged particles? Perpendicular – change in direction, parallel – change in speed What does his structure of the atom mimic? The solar/planetary system 4. What was wrong with Rutherford’s model? Electrons should lose energy as they accelerate/move around the nucleus, showing a frequency shift but there is no frequency shift

43 Nucleons Neutrons: 1932 James Chadwickreceived the Nobel Prize in Physics in 1935 for the discovery of the neutron (worked with Hans Geiger, then Ernest Rutherford)

44 Nucleons 2. Uncharged particles in the nucleus with mass were pushed out of beryllium when bombarded with alpha particles. These particles accounted for the “missing mass” in the nucleus.

45 Subatomic Particles Particle Location Charge Relative Mass SymbolOutside Nucleus Electron Inside Nucleus Proton Inside Nucleus Neutron

46 Subatomic Particles Particle Location Charge Relative Mass SymbolOutside Nucleus Electron -1 1/1840 e- Inside Nucleus Proton +1 1 p+ Inside Nucleus Neutron 1 no

47 See diagram on board of relative sizes of the parts of the atom record in your notes

48 Atomic Theory Video Questions #2 Introducing the PlayersWhat particle gives the atom its positive charge? 2. What are the three building blocks of atoms? What positively identifies the atom? 4. What defines an isotope? Do isotopes have the same chemical and physical properties? 6. Does the nucleus change in a chemical reaction? 7. Why would the electrons in Rutherford’s model eventually crash into the nucleus?

49 Atomic Theory Video Questions #2 Introducing the PlayersWhat particle gives the nucleus its positive charge? proton 2. What are the three building blocks of atoms? Electrons, protons, neutrons What positively identifies the atom? # protons 4. What defines an isotope? # neutrons Do isotopes have the same chemical and physical properties? Basically yes – only mass differs 6. Does the nucleus change in a chemical reaction? no Why would the electrons in Rutherford’s model eventually crash into the nucleus? An accelerating mass emits EMR; loses energy and crashes into the nucleus

50 Atomic Mass Atomic number = # protons (also = # electrons)Mass number = # protons + # neutrons # neutrons = mass number – atomic number

51 Isotopes 2 Isotope # Protons # Neutrons # Electrons Hydrogen-1 protiumHydrogen-2 deuterium 21H Hydrogen-3 tritium 31H 2

52 Isotopes of Hydrogen and Carbon

53 Isotopes

54 Isotopes WS 1. What do the following symbols represent?a. e- _________________ b. n0 _________________ c. p+ _________________ 2. Which subatomic particles are found in an atom’s nucleus? 3. Which subatomic particle identifies an atom as that of a particular element? 4. Explain why atoms are neutral even though they contain charged particles. 5. What do the numbers, 39, 40, and 41 after the element name potassium refer to? Write the symbolic notation for each of the following isotopes: a. potassium-39 ________________ b. potassium-40 ________________ c. potassium-41 ________________ 7. Write an equation showing the relationship between an atom’s atomic number and its mass number.

55 Isotopes WS 1. What do the following symbols represent?a. e- ___electron_____ b. n0 ___neutron_______ c. p+ ___proton_______ 2. Which subatomic particles are found in an atom’s nucleus? protons, neutrons 3. Which subatomic particle identifies an atom as that of a particular element? protons 4. Explain why atoms are neutral even though they contain charged particles. # protons = # electrons 5. What do the numbers, 39, 40, and 41 after the element name potassium refer to? mass number Write the symbolic notation for each of the following isotopes: a. potassium-39 _____3919K______ b. potassium-40 _____4019K______ c. potassium-41 _____4119K______ 7. Write an equation showing the relationship between an atom’s atomic number and its mass number. Mass # = atomic # + # of neutrons

56 Warm Up Compare the atomic models of Dalton, Thomson, and Rutherford:For each model, 1. draw a diagram, 2. use words to describe the model, 3. use words to describe the experiment that provided the evidence to change the preceding model into the next one, 4. explicitly state the connection that prompted Thomson and Rutherford made to propose the changes from the previous model to their new model

57 Last Time Mass number Atomic number Number of protonsNumber of neutrons Number of electrons WHOLE NUMBERS

58 Average Atomic Mass 1 atomic mass unit (amu)=1/12 mass of a C-12 atomaverage Atomic Mass = weighted average (by abundance) shown on Periodic Table

59 Average Atomic Mass Practice MarblesMass = 1.59 g, 1.51 g, 1.76 g Average Mass = (1.59 g g g) 3 = 1.62 g

60 Average Atomic Mass Practice MarblesThe same calculation can be done as 4.86 g x … = 1.62 g.

61 Average Atomic Mass Practice Marbles79 are 1.59 g, 10 are 1.51 g, and 11 are 1.76 g 79 x 1.59 g = g 10 x 1.51 g = g 11 x 1.76 g = g 160.5 g divided by 100 = 1.60 g

62 Average Atomic Mass Practice Marblesor, using relative abundances: 0.79 x 1.59 g =+1.26 g 0.10 x 1.51 g = g 0.11 x 1.76 g = g 1.60 g

63 Average Atomic Mass Practice AtomsWhat is the average atomic mass of antimony? The isotopes of antimony and their percent abundances are Sb-121 ( amu, 57.21%) and Sb-123 ( amu, 42.79%) Use your periodic table to check your answer.

64 Average Atomic Mass Practice AtomsWhat is the average atomic mass of vanadium? The isotopes of vanadium and their percent abundances are V-50 (49.95 amu, 0.250%) and V-51 (50.94 amu, %). Use your periodic table to check your answer.

65 Warmup (from Atomic Theory WS #1)2. Give the number of protons, electrons, and neutrons in each of the following atoms. a Ag b Ca c Na 3. Name each isotope, and write it in symbolic notation. a. Atomic number 26; mass number 56 b. Atomic number 29; mass number 64 c. Atomic number 17; mass number 37

66 Warmup KEY (from Atomic Theory WS #1)2. Give the number of protons, electrons, and neutrons in each of the following atoms. a Ag p+, 47 e-, 61 n0 b Ca p+, 20 e-, 20 n0 c Na p+, 11 e-, 12 n0 3. Name each isotope, and write it in symbolic notation. a. Atomic number 26; mass number 56 iron-56, 5626Fe b. Atomic number 29; mass number 64 copper-64, 6429Cu c. Atomic number 17; mass number 37 chlorine-37, 3717Cl

67 Warmup (from Atomic Theory WS #1)How many protons, electrons and neutrons are in each of the following isotopes? a. Uranium-235 b. Hydrogen-3 c. Silicon-29 Additional problem: How many protons, electrons and neutrons are in the following ion isotopes? a As3- b W6+ 14. An element has three naturally occurring isotopes Isotope 1 has a mass of amu, 90.48% abundance Isotope 2 has a mass of amu, 0.27% abundance Isotope 3 has a mass of amu, 9.25% abundance a. Calculate the (average) atomic mass of the element. b. Identify the element, using the periodic table.

68 Warmup (from Atomic Theory WS #1)How many protons, electrons and neutrons are in each of the following isotopes? a. Uranium p+, 92 e-, 143 n0 b. Hydrogen p+, 1 e-, 2 n0 c. Silicon p+, 14 e-, 15 n0 Additional problem: How many protons, electrons and neutrons are in the following ion isotopes? a As p+, 36 e-, 42 n0 b W p+, 68 e-, 106 n0 14. An element has three naturally occurring isotopes Isotope 1 has a mass of amu, 90.48% abundance Isotope 2 has a mass of amu, 0.27% abundance Isotope 3 has a mass of amu, 9.25% abundance a. Calculate the (average) atomic mass of the element amu b. Identify the element, using the periodic table. neon

69 Atomic #, Mass # and Average Atomic Mass# p+ #p+ + #no Weighted average of all isotopes whole # Decimal, limited by sfs Found on PT NOT on PT

70 Radioactivity Nuclear Reaction  change in the identity of the elements Radioactivity = radiation emitted by atoms with an unstable n0:p+ ratio

71 Radioactivity Smaller Elements (atomic # < 20) Larger ElementsStable ratio = 1 n0:1 p+ i.e. Mass # = 2 x atomic # Larger Elements - Stable ratio = 1.5 n0:1p+ - All elements with atomic # > 83 are radioactive

72 Radioactivity Unstable nuclei  emit radiation and change their identities This is called radioactive decay

73 Historical figures Wilhelm Roentgen ( ) - discovered X-rays –1895 X-ray of his wife’s hand Nobel Prize in Physics, 1901

74 Historical figures Henri Becquerel ( ) - discovered radioactivity in U Nobel Prize in Physics, 1903, shared with the Curies, for his discovery of spontaneous radioactivity

75 Historical figures Ernest Rutherford – identified different types of radiation, and explored their properties (beg. 1898) Nobel Prize in Chemistry in 1908

76 Historical figures Pierre ( ) and Marie Curie ( ) - discovered radium and polonium – 1898; first used the term “radioactivity” Nobel Prize in Physics 1903, Pierre and Marie, with Henri Becquerel Nobel Prize in Chemistry 1911 (Marie only) for discoveries of radium and polonium

77 Types of Radiation Alpha radiation Alpha particles = 2 p+ + 2 n0(most common in elements with atomic # > 83,  increase the number of neutrons) Alpha particles = 2 p+ + 2 n0 (He nucleus, 42He, a) with 2+ charge e.g Ra  Rn He (+ energy)

78 Types of Radiation Beta radiation(most common in elements with high n0:p+ ratio  decrease the number of neutrons) Beta particles = 1 e- (0-1b) with 1- charge Neutron  proton + beta particle 10n  11p b e.g. 146C  147N b (+ energy)

79 Types of Radiation Note:The sum of the mass #s and atomic #s on both sides of the equation are the same

80 Types of Radiation 3. Gamma radiationGamma rays = high-energy radiation with no mass and no charge (00g) usually accompany alpha and beta radiation e.g U  Th He g

81 Types of Radiation Nuclei with lower neutron:proton ratios than optimal: Positron Emission (most common in lighter elements with low n0:p+ ratio)  more neutrons by converting a proton into a neutron Positron = particle with same mass as an e-, but opposite charge Proton  neutron + positron 11p  10n b e.g. 116C  115B b

82 Types of Radiation Nuclei with lower neutron:proton ratios than optimal: Electron Capture (most common in elements with a high n0:p+ ratio)  more neutrons by pulling in an e- which combines with a proton to form a neutron Proton + electron  neutron 11p e  10n e.g. 0-1e Rb  8136Kr + X-ray photon

83 Radioactive Particles WSPositron same mass as e-’s 0+1b / Electron capture electrons 0-1e / (Added to the reactants side)

84 Radioactive Particles WSWhich radioactive emission has the greatest mass? Least mass? 2. Why do you think gamma rays are drawn as wavy lines? 3. Which charged plate are the alpha particles attracted to? Explain. 4. Which charged plate are the beta particles attracted to? Why do the beta particles have a greater curvature than the alpha particles? 5. Explain why the gamma rays do not bend toward one of the electrically charged plates.

85 Radioactive Particles WSWhich radioactive emission has the greatest mass? Least mass? alpha – greatest; gamma – no mass Why do you think gamma rays are drawn as wavy lines? Gamma rays have no mass and are EMR, which is often drawn as wavy lines. Which charged plate are the alpha particles attracted to? Explain. To the negatively-charged plate, as the alpha particles are positively-charged. Which charged plate are the beta particles attracted to? Why do the beta particles have a greater curvature than the alpha particles? to the positively-charged plate, as the beta particles are negatively-charged. They have a smaller mass, so are more greatly influenced by the electric field. Explain why the gamma rays do not bend toward one of the electrically charged plates. Gamma rays have no charge, therefore, they are not attracted to either plate.

86 Nuclear Fission Nuclear fission = the splitting of a nucleus into smaller, more stable fragments, accompanied by a large release of energy e.g. Uranium-235: 23592U + 10n  23692U  9236Kr Ba n + energy (unstable) The new neutrons (10n)  fission of more U-235 (= chain reaction, a self-sustaining process)

87 Nuclear Fission Chain reaction requires a critical mass (= minimum amount of starting material to maintain a chain reaction) supercritical mass may  violent nuclear explosion results in radioactive waste Practical examples = nuclear power plant, atomic bomb

88 Nuclear Fusion Nuclear Fusion = the process of binding smaller atomic nuclei into a single larger and more stable nucleus, requiring a huge amount of energy to initiate, followed by a large release of energy

89 1. Creation of Natural ElementsElements are created by nuclear reactions a. Hydrogen, other light elements - from the Big Bang

90 Creation of Natural Elementsb. Elements #2-92 (except Fr, Pr, Te, At) Nuclear fusion occurs in stars (naturally) Occurs in hydrogen bomb (artificially) > 2 x 107oC The sun converts 3 x 1014 g of H into He every second. 4 11H  42He + energy Mass is not conserved. Mass is converted into energy via E = mc2

91 Creation of Natural ElementsOther fusion reactions occur in the sun: 42He He  84Be + g (gamma ray) 42He Be  126C + g

92 2. Synthetic Elements a. Nuclear bulletsi. Bombard nuclei of elements with small particles such as p+, n , 42He (a particles) & e- (0-1b particles) ii. Elements #

93 2. Synthetic Elements iii. 1919 first experiment:147N He  178O + 11H (Rutherford)

94 2. Synthetic Elements b. Crashing nucleii. Accelerators hurl nuclei into each other at very high speeds. e.g. 126C Cm  No n carbon curium nobelium neutron ii. Elements beyond #100 iii. These elements are very unstable: e.g. Element 109 existed for only 3.4 x 10-3 sec (3 atoms)

95 2. Synthetic Elements Superheavy elements (“transuranium” elements)Stability of nucleus of atom depends on filling "shells" within nucleus with alternating p+ and n. The more filled shells, the more stable it would be. e.g. Element 114 24494Pu Ca  Fl n (1999, Russia)

96 Nuclear equations Complete the following equations :21483Bi  42He + _____ 23993Np  23994Pu + ______

97 Production of Transuranium Elements

98 Production of Transuranium Elements WS1. Does the diagram illustrate a natural transmutation reaction or an induced transmutation reaction? 2. What is the name and nuclear symbol of the isotope produced in the reaction? 3. Write a nuclear equation to show how dubnium-263, lawrencium-262, and seaborgium-266 can be produced from a nuclear reaction of neon-22 and americium-244. 2210Ne Am  Db n 2210Ne Am  5. Each of the radioisotopes in the table decays within 20 seconds to 10 hours. Write a nuclear equation for each decay. 26695Am  42He + 263105Db  0-1b + 262103Lr + 0-1e-  266106Sg  42He + 6. Which, if any, of the four isotopes listed in the table would you expect to find at Earth’s surface? Why?

99 Production of Transuranium Elements WS1. Does the diagram illustrate a natural transmutation reaction or an induced transmutation reaction? Induced transmutation 2. What is the name and nuclear symbol of the isotope produced in the reaction? Dubnium-266; Db 3. Write a nuclear equation to show how dubnium-263, lawrencium-262, and seaborgium-266 can be produced from a nuclear reaction of neon-22 and americium-244. 2210Ne Am  Db n 2210Ne Am  Lr + 42He 2210Ne Am  Sg + 0-1b

100 Production of Transuranium Elements WS5. Each of the radioisotopes in the table decays within 20 seconds to 10 hours. Write a nuclear equation for each decay. 24495Am  42He Np 263105Db  0-1b Sg 262103Lr + 0-1e-  No 266106Sg  42He Rf 6. Which, if any, of the four isotopes listed in the table would you expect to find at Earth’s surface? Why? None – they all have very short half-lives.

101 Nuclear equations Complete the following equations :21483Bi  42He + _21081Tl_ 23993Np  23994Pu + __0-1b____

102 Nuclear equations Write a balanced nuclear equation for the alpha decay of americium-241. Write a balanced nuclear equation for the beta decay of bromine-84.

103 Nuclear equations Write a balanced nuclear equation for the alpha decay of americium-241. 24195Am  42He Np Write a balanced nuclear equation for the beta decay of bromine-84. 8435Br  0-1b Kr

104 Nuclear equations Complete the following equations: 21483Bi  42He +23993Np  23994Pu + 24195 Am  42He + 8435Br  _______ b

105 Nuclear equations (KEY)Complete the following equations: 21483Bi  42He Tl 23993Np  23994Pu + 0-1b 24195 Am  42He Np 8435Br  8436Kr b

106 Warmup – Nuclear Equations, including Nuclear Fusion94Be H  42He + _______ U He  2 10n + _______ 156C n  _______ Cs  ________ b

107 Warmup – Nuclear Equations94Be H  42He Li U He  2 10n Pu 156C n  166C Cs  Ba b

108 Next Steps - Properties of ElectronsWave nature of light – EMR (James Maxwell, 1864) Particle nature of light – quantum (Max Planck, late 1800s) Emission of light and other EMR from heated elements  emission spectra

109 Electromagnetic RadiationEMR = energy that exhibits wave-like behavior as it travels through space James Maxwell (1864) – unified electric and magnetic forces into electromagnetic force

110 Electromagnetic RadiationUnified the electric and magnetic forces  electromagnetic force (emf) James Maxwell (1864) – unified electric and magnetic forces into electromagnetic force

111 Electromagnetic RadiationSpeed of EMR always the same c = 3.00 x 108 m/s Examples include: microwaves, TV, Radio, X-rays ←Memorize

112 Electromagnetic RadiationWavelength = λ (lambda) usually in nm Frequency = n (nu) or f Waves per second = Hz (Hertz) = cycles/s or s-1 Speed = c, measured in m/s c = λ n Note the inverse relationship between λ and n

113 Electromagnetic Spectrum

114 Demo EMR speed is the same, while frequency and wavelength change – red vs. blue light

115 EMR Spectrum

116 EMR Spectrum What kinds of waves have the longest wavelength? What kinds of waves have the shortest wavelength? Which waves have the lowest frequency? Which has a higher frequency: microwaves or X rays? Which waves can be seen by the eye? Sequence the different segments of the visible spectrum in order from shortest wavelength to longest wavelength. Sequence the following types of waves from lowest frequency to highest frequency: ultraviolet rays, infrared rays, gamma rays, radio waves, and green light. Compare the wavelengths and frequencies of each kind of wave. What is the relationship between frequency and wavelength? 8. What is the wavelength of a radio station emitting its signal at 95.5 MHz? Estimate your answer to the nearest power of ten.

117 EMR Spectrum What kinds of waves have the longest wavelength? Radio waves 2. What kinds of waves have the shortest wavelength? Gamma rays 3. Which waves have the lowest frequency? Radio waves 4. Which has a higher frequency: microwaves or X rays? X-rays 5. Which waves can be seen by the eye? Visible portion of the spectrum

118 EMR Spectrum 6. Sequence the different segments of the visible spectrum in order from shortest wavelength to longest wavelength. Violet, Indigo, Blue, Green, Yellow, Orange, Red 7. Sequence the following types of waves from lowest frequency to highest frequency: radio waves, infrared waves, green light, ultraviolet waves, gamma rays 8. Compare the wavelengths and frequencies of each kind of wave. What is the relationship between frequency and wavelength? Inversely proportional 9. What is the wavelength of a radio station emitting its signal at 95.5 MHz? Estimate your answer to the nearest power of ten. About 3 m, or 3 x 100 m

119 EMR Practice Problems from Textbook (pp. 121, 124)c = lx n Microwaves are used to transmit information. What is the wavelength of a microwave having a frequency of 3.44 x 109 Hz? (8.72 x 10-2 m)

120 EMR Practice Problems c = lx nWhat is the frequency of green light which has a wavelength of 4.90 x 10-7 m? (6.12 x 1014 s-1) An X-ray has a wavelength of 1.14 x m. What is its frequency? (2.63 x 1018 s-1) What is the speed of an electromagnetic wave that has a frequency of 7.8 x 106 Hz? (3.00 x 108 m/s) 4. A popular radio station broadcasts with a frequency of 94.7 MHz. What is the wavelength of the broadcast? (1 MHz = 106 Hz) (3.17 m)

121 Quantum Max Planck ( )– Nobel Prize in Physics, 1918, for his discovery of energy quanta Revolutionary concept in physics

122 The Idea of the Quantum Quantum = the smallest discrete amount of energy that can exist independently, esp. as EMR 1 quantum = 1 photon E = hn, where h = a constant The amount of energy in EMR increases with increasing frequency

123 Demo Photons – glow in the dark

124 EMR Practice Problems E = hnPlanck’s constant = h = x J·s Example: Tiny water drops in the air disperse the white light of the sun into a rainbow. What is the energy of a photon from the violet portion of the rainbow if it has a frequency of x 1014 s-1? (4.79 x J)

125 EMR Practice Problems E = hnPlanck’s constant = h = x J·s 5. What is the energy of each of the following types of EMR? a x 1020 s-1 (4.19 x J) b x 1013 Hz (6.29 x J) c x 1016 s-1 (6.96 x J) 6. Name the types of radiation in each part of #5.

126 EMR Practice Problems 1. What is the frequency of EMR with a wavelength of 235 pm? What type of EMR is this? 2. What is the frequency of EMR with a wavelength of cm? What type of EMR is this? 3. What is the wavelength of EMR with a frequency of 8,512 Hz? What type of EMR is this? 4. What is the wavelength of EMR with a frequency of 625 x 1017 Hz? What type of EMR is this? 5. If the speed of light is 3.00 x 108 m/s, calculate the wavelength of the electromagnetic radiation whose frequency is x 1012 Hz. 6. Determine the frequency of light with a wavelength of x 10-7 cm. 7. For the following sources, calculate the missing member of the wavelength/frequency pair. a) FM radio waves with a frequency of 94.7 Hz. b) A laser with a wavelength of 1064 nm. c) An X-ray source, emitting X-rays with a wavelength of pm. 8. How long would it take a radiowave with a frequency of 7.25 x 105 Hz to travel from Mars to Earth if this distance between the two planets is approximately 8.00 x 107 km?

127 EMR Practice Problems 1. What is the frequency of EMR with a wavelength of 235 pm? What type of EMR is this? (1.28 x 1018 s-1; X-rays) What is the frequency of EMR with a wavelength of cm? What type of EMR is this? (4.89 x 1010 s-1; microwaves) What is the wavelength of EMR with a frequency of 8,512 Hz? What type of EMR is this? (3.52 x 104 m; radio waves) What is the wavelength of EMR with a frequency of 625 x 1017 Hz? What type of EMR is this? (4.80 x m; X-rays, gamma rays)

128 EMR Practice Problems radiation whose frequency is 7.500 x 1012 Hz.5. If the speed of light is 3.00 x 108 m/s, calculate the wavelength of the electromagnetic radiation whose frequency is x 1012 Hz. (4.00 x 10-5 m) Determine the frequency of light with a wavelength of x 10-7 cm. (7.05 x 1016 s-1) For the following sources, calculate the missing member of the wavelength/frequency pair. a) FM radio waves with a frequency of 94.7 Hz. (3.17 x 106 m) b) A laser with a wavelength of 1064 nm. (2.82 x 1014 s-1) c) An X-ray source, emitting X-rays with a wavelength of pm. (1.71 x 1018 s-1) How long would it take a radiowave with a frequency of 7.25 x 105 Hz to travel from Mars to Earth if this distance between the two planets is approximately 8.00 x 107 km? (Note: v is not required for the calculation.) (2.67 x 102s)

129 Warmup - EMR Which color has the shorter wavelength (l) – blue or red?2. Which color has the higher frequency (n) – blue or red? 3. What is the wavelength of light with a frequency of 4.90 x 1016 s-1)? (A: 6.12 x 10-9 m) 4. What is the frequency of EMR with a wavelength of 5.26 mm? What type of EMR is it? (A: x 1010 s-1) 5. Determine the energy, in joules, of a photon whose frequency is 3.55 x 1017 Hz. (h = x J s) (A: 2.35 x J)

130 Warmup - Honors 5. When sodium is heated, a yellow spectral line whose energy is 3.37 x J per each photon is produced. a. What is the frequency of this light? (A = 5.09 x 1014 s-1) b. What is its wavelength? (A = 5.89 x 10-7 m)

131 Planetary Model – Neils Bohr( ) (Danish physicist) Studied with Thomson and Rutherford Refined Rutherford’s model in 1913 Received the Nobel Prize in 1922 for his work on the structure of atoms

132 Planetary Model – Neils Bohr( ) (Danish physicist) Incorporated Planck’s idea of quanta of energy Provided an explanation for the spectral lines of hydrogen

133 Bohr Model Electrons… nucleus, “orbits”.1. are arranged in circular paths around nucleus, “orbits”. have fixed energy levels to prevent them from falling into nucleus. Electrons closest to the nucleus have lowest Etotal = KE + PE (most stable).

134 Bohr’s Model + As e- goes further E3 - E1 + 1 2 3 41 = lowest energy 4 = highest energy E3 - E1 = a quantum of energy in the form of EMR + As e- goes further away from the nucleus, it increases in potential energy As e- goes further

135 Bohr Model Electrons… must gain or lose energy to change energy levels. EMR is emitted from the atom when electrons fall down to a lower energy level. 4. in different energy levels are not the same distance apart. A “quantum” = amount of energy needed to make the leap between energy level.

136 Bohr Model Bohr’s model did not explain the line spectra of atoms with >1 electron.

137 Definitions related to SpectraSpectrum = whole range of related qualities [Latin: appearance, from specere – to view] Electromagnetic spectrum = all EMR arranged according to l

138 Definitions related to SpectraEmission = any radiation of energy by EM waves [Latin: emitto – to send out, to utter, to hurl, to set free] Emission spectrum = the spectrum into which light or other EMR from any source can be separated Continuous spectrum = a spectrum whose source emits light of every l in a continuous band Bright-line spectrum = pattern of bright lines on a dark background. Source = glowing gas that radiates in special l’s characteristic of the chemical composition of the gas

139 Continuous white light spectrum

140 Line-Emission Spectrumexcited state ENERGY IN PHOTON OUT ground state

141 Comparison of Spectra

142

143 Comparison of continuous, line and absorption spectra

144 Infrared Paschen SeriesThe Hydrogen Spectrum E2 Each photon emitted has a characteristic λ which contributes a line to the spectrum E1 Visible Balmer Series UV Lyman Series Infrared Paschen Series

145 The Hydrogen Spectrum

146 Atomic Theory Video Questions #3 The Bohr AtomWhat does a change in color indicate? 2. What did Max Planck propose about how matter emits energy? 3. Which has more energy: an electron near the nucleus or an electron far from the nucleus? Why? 4. What element did Bohr use in his experiments? 5. How much energy must be carried by a free electron in order to knock another electron out of orbit? What happens if an electron of more than 13.6 eV hits another electron? 7. What did Bohr propose happens when an electron goes down an energy level?

147 Atomic Theory Video Questions #3 The Bohr Atom KEYWhat does a change in color indicate? Change in frequency 2. What did Max Planck propose about how matter emits energy? In packets called quanta 3. Which has more energy: an electron near the nucleus or an electron far from the nucleus? Why? Far – more potential energy 4. What element did Bohr use in his experiments? hydrogen 5. How much energy must be carried by a free electron in order to knock another electron out of orbit? > 10.2 eV What happens if an electron of more than 13.6 eV hits another electron? Liberates it from the nucleus  ionized What did Bohr propose happens when an electron goes down an energy level? Emits photon equivalent to amount of energy

148 Atomic Theory Video Questions #3 SpectraWhat is the orbit of an electron closest to the nucleus called? 2. What region of the spectrum of excited hydrogen gas did a. Balmer predict? ________________ b. Paschen predict? ________________ c. Lyman predict? ________________ What was wrong with Balmer’s formula? 4. What made Bohr’s mathematical model so special? 5. Does the electron emit radiation when it is bumped up an energy level or when it falls back down? 6. What level (n) does the electron fall to in order to produce: a. the Balmer series? _______________ b. the Paschen series? _______________ c. the Lyman series? _______________

149 Atomic Theory Video Questions #3 Spectra KEYWhat is the orbit of an electron closest to the nucleus called? ground state 2. What region of the spectrum of excited hydrogen gas did a. Balmer predict? visible b. Paschen predict? infrared c. Lyman predict? UV What was wrong with Balmer’s formula? Nobody knew why it worked 4. What made Bohr’s mathematical model so special? Based on a possible structure of the atom 5. Does the electron emit radiation when it is bumped up an energy level or when it falls back down? when it falls 6. What level (n) does the electron fall to in order to produce: a. the Balmer series? 2 b. the Paschen series? 3 c. the Lyman series? 1

150 Modern (Quantum) TheoryWave nature of the electron Louis de Broglie received the Nobel Prize for Physics in 1929 for his discovery of the wave nature of electrons

151 Modern (Quantum) TheoryQuantum Mechanics Erwin Schrödinger received the Nobel Prize for Physics in 1933 for his work in atomic theory Wave equation: electrons as waves (1926) Foundation of the quantum theory of the atom

152 Modern (Quantum) TheoryOrbital

153

154 Modern (Quantum) TheoryQuantum Mechanics Werner Heisenberg, received Nobel Prize for Physics in 1932 for quantum mechanics Heisenberg’s Uncertainty Principle We cannot simultaneously measure an electron’s position and its velocity

155

156 Orbitals and Quantum #’sOrbital in the shape of an “electron cloud” contains all (90%) locations of an electron Each e- in an atom has its own set of four quantum #’s — n, l, m, and s

157 Orbitals and Quantum #’sPrincipal energy levels = Bohr’s orbits Total # of e- in one principal energy level = 2n2 Sublevels (l), magnetic position (m), and spin (s)—additions to classify e- energies

158 Atomic Orbitals: predict 90% probability of location of electrons (electron cloud)Each orbital can contain a maximum of 2 electrons, spinning in opposite directions. f d p s

159 Orbitals

160 Orbitals WS What is the shape of an s orbital?2. What is the relationship between the size of an s orbital and the principal energy level in which it is found? What is the shape of a p orbital? How many p orbitals are there in a sublevel? How many electrons can each orbital hold? Look at the diagrams of the p orbitals. What do x, y and z refer to? How many d orbitals are there in a given sublevel? How many total electrons can the d orbitals in a sublevel hold? Which d orbitals have the same shape? What point in each diagram represents an atom’s nucleus? 9. How likely is it that an electron occupying a p or a d orbital would be found very near an atom’s nucleus? What part of the diagram supports your conclusion?

161 Orbitals WS What is the shape of an s orbital? spherical2. What is the relationship between the size of an s orbital and the principal energy level in which it is found? Size increases with increasing principal energy level What is the shape of a p orbital? Dumbbell How many p orbitals are there in a sublevel? 3 How many electrons can each orbital hold? 2 Look at the diagrams of the p orbitals. What do x, y and z refer to? 3 perpendicular axes How many d orbitals are there in a given sublevel? 5 How many total electrons can the d orbitals in a sublevel hold? 10 Which d orbitals have the same shape? 4 out of the 5 What point in each diagram represents an atom’s nucleus? The origin – where the x, y and z axes intersect 9. How likely is it that an electron occupying a p or a d orbital would be found very near an atom’s nucleus? Very unlikely What part of the diagram supports your conclusion? The shapes of the orbitals come to a point at the intersection of the three axes, making the possibility of an electron being found there very unlikely.

162 Origin of Orbital Namess – sharp (or use sphere) p – principal (peanut) d – diffuse (daffodil or daisy) f – fundamental (funky) Names come from the spectrum analysis, e.g. the hyperfine splitting of the d-line of the sodium spectrum

163 shape of probability cloudPrincipal Energy Level (n) (or “Shell”) distance from nucleus Sublevel (l) (or “Subshell”) shape of probability cloud 3-D (m) (position in space) x, y, z Spin (s) + ½ - ½ Total # e- 2n2 1 s only 1 orientation 1 e ½ 1 e ½ 2(1)2 = 2 2 s, p 3 orientations for p: px, py, pz 2(2)2 = 8 3 s, p, d etc. 2(3)2 = 18 4 s, p, d, f 2(4)2 = 32 5 6 s, p, d, (f) 7 s: holds 2 e- x orbital = 2 e- total p: holds 2 e- x orbitals = 6 e- total d: holds 2 e- x orbitals = 10 e- total f: holds 2 e- x orbitals = 14 e- total

164 n = # of sublevels per leveln2 = # of orbitals per level Sublevel sets: 1 s, 3 p, 5 d, 7 f

165 1. The p orbitals px py pz

166 2. The d orbitals

167 Electron spin An orbital can hold 2 electrons that spin in opposite directions.

168 Three Rules For Filling OrbitalsAufbau Principle Fill in order of increasing energy levels Hund’s Rule Fill all orbitals at same energy level with at least 1 e-, before adding the second e- Pauli Exclusion Principle only 2 e- per orbital, of opposite spin

169 Electron Configuration

170 Electron Configuration WSWhat does each small box in the diagram represent? How many electrons can each orbital hold? How many electrons can the d sublevel hold? Which is associated with more energy: a 2s or a 2p orbital? Which is associated with more energy: a 2s or a 3s orbital? According to the Aufbau Principle, which orbital should fill first, a 4s or a 3d orbital? Which orbital has the least amount of energy? What is the likelihood that an atom contains a 1s orbital? 9. Sequence the following orbitals in the order that they should fill up according to the Aufbau Principle: 4d, 4p, 4f, 5s, 6s, 3d, 4s.

171 Electron Configuration WS KEYWhat does each small box in the diagram represent? An orbital How many electrons can each orbital hold? 2 How many electrons can the d sublevel hold? 10 Which is associated with more energy: a 2s or a 2p orbital? 2p Which is associated with more energy: a 2s or a 3s orbital? 3s According to the Aufbau Principle, which orbital should fill first, a 4s or a 3d orbital? 4s Which orbital has the least amount of energy? 1s What is the likelihood that an atom contains a 1s orbital? 100% 9. Sequence the following orbitals in the order that they should fill up according to the Aufbau Principle: 4d, 4p, 4f, 52, 62, 3d, 4s: 4s, 3d, 4p, 5s, 4d, 6s, 4f

172 Electron ConfigurationOrder of Filling—Aufbau Diagram s s p s p d s p d f Oxygen (8 e-) Orbital Diagram Electron Configuration 1s2 2s p4

173 Practice with orbital diagrams and electron configurations:Draw your own Aufbau diagram, then write orbital diagrams and electron configurations for oxygen calcium gallium

174 Practice with complete and condensed electron configurations:Write electron configurations (only) for: a) helium b) carbon a) neon b) aluminum a) argon b) bromine a) krypton b) palladium a) xenon b) lead a) radon b) uranium c) Write all b)’s in noble gas (condensed) configuration d) Show how to jump into an Aufbau diagram e) s,p,d,f blocks in the periodic table

175 Atomic Theory Video Questions # 4 Electron ArrangementWhat 2 parts of Bohr’s model did Schrödinger and Heisenberg keep? How did they change Bohr’s model? What is the probability distribution of the first energy level called? 4. What is the second type of probability distribution shaped like? What is it called? 5. How many orientations do s orbitals have? p orbitals? d orbitals? How many electrons can be in any one orbital? 7. Which orbitals are filled with electrons first? What is the second “rule” about filling orbitals? Which electron orbitals have the most bearing on the chemical properties of a particular atom? 10. What practical application do the periodic table groupings provide to working chemists?

176 Atomic Theory Video Questions # 4 Electron Arrangement KEYWhat 2 parts of Bohr’s model did Schrödinger and Heisenberg keep? Positively-charged nucleus, energy levels for electrons How did they change Bohr’s model? No definite orbits, but rather probability electron clouds called orbitals What is the probability distribution of the first energy level called? s 4. What is the second type of probability distribution shaped like? Dumb-bells What is it called? p 5. How many orientations do s orbitals have? p orbitals? d orbitals? 5 How many electrons can be in any one orbital? 2, of opposite spin 7. Which orbitals are filled with electrons first? Lowest energy What is the second “rule” about filling orbitals? 1 e- per orbital of same energy, then add second e- Which electron orbitals have the most bearing on the chemical properties of a particular atom? Last ones to be filled What practical application do the periodic table groupings provide to working chemists? Helps predict chemical reactions

177 Atomic Theory Scientists

178 Atomic Theory and the Periodic Tables p 1 2 3 4 5 6 7 d (n-1) f (n-2) 6 7 © 1998 by Harcourt Brace & Company

179 E. Periodic Patterns Period # A/B Group # Column within sublevel blockenergy level (subtract for d & f) A/B Group # total # of valence e- Column within sublevel block # of e- in sublevel

180 1s1 Periodic Patterns 1st column of s-block 1st Period s-blockExample - Hydrogen 1s1 1st column of s-block 1st Period s-block

181 Periodic Patterns p s d (n-1) f (n-2) Shorthand ConfigurationCore e-: Go up one row and over to the Noble Gas. Valence e-: On the next row, fill in the # of e- in each sublevel. s d (n-1) f (n-2) p

182 Periodic Patterns Example - Germanium [Ar] 4s2 3d10 4p2

183 Stability 1+ 2+ 3+ NA 3- 2- 1- Ion FormationAtoms gain or lose electrons to become more stable. Isoelectronic with the Noble Gases. 1+ 2+ 3+ NA 3- 2- 1-