flame tests lab Chart: Wavelength (in nanometers) of visible light 10. Take your normal seats. 11. Answer questions at your desks. Turn in lab.

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3 flame tests lab name period lab5. Safety Notice: This lab is exciting, but please be cautious. Wear goggles. Assume all salts are toxic, as are all gases produced. We have all seen the beautiful colors that can form when substances are placed in a flame. Pockets of gas in wood can form green and blue colors when they ignite. What is happening when this occurs? This answer was the key to unlocking the strange behavior of the electron, now known as quantum theory. In this experiment we will observe some of these colors, and will make some initial attempts to explain it. Finally, the color of the emitted light will be used to identify the unknown salts. Procedure. Goggles on please, and go to your stations. The rule all year will be that if the instructor is wearing goggles, you are as well.. Listen to the Bunsen burner lesson 3. Each student should safely light the Bunsen burner and adjust the gas/air mixture. 4. Turn off Bunsen burner 5. Get your set of test tubes and unknowns 6. Dip the paper clip to see what color it turns on its own if you see this color it may be due to the metal in the paper clip. 7. Dip each of solutions using a paper clip and place in flame for less than seconds each. Write down the color of the flame, and estimate the wavelength in nanometers(a color chart will be available). 8. Fill in table during testing. Identify the unknowns based on flame color. 9. Clean up: wet matches in trash. All stations will be inspected, including sinks. Chart: Wavelength (in nanometers) of visible light 0. Take your normal seats.. Answer questions at your desks. Turn in lab. Salts and unknowns: Unknown numbers: Table : Color and wavelength in nanometers of emission spectrum of salts and unknown. unknown # unknown # salt flame color estimated wavelength (nm) Analysis. Each of the known compounds tested contains chlorine, yet each compound produced a flame of a different color. What does this suggest?. We will learn this week that the movement of electrons in atoms produces the colors we observed. What specifically may be going on with the electrons to produce color? (take your best guess). 3

4 name period lab5. spectroscopy lab What is light? How is light created? Is there a relationship between light and matter? This experiment is designed to help you answer these questions. The experiments you will perform are similar to those performed by Niels Bohr and others, and begs the question: what does it all mean? How do the spectral lines relate to the structure of the atom? Procedure: Safety: As before, this lab uses flames and toxic salts. Please wear goggles and work safely.. Put on goggles.. Each group will perform a 5 minute experiment at one of 6 stations, and then proceed to the next. As precisely as you can, draw the component wavelengths observed at each station. Follow the instructions for each station, clean up, and be ready to move to the next station. Station : Sunlight. Each student should point through the spectroscope directly at the sun, and draw the component wavelengths observed. If weather permits, see if the colors are the same when you are not looking through a window. Station : Artificial light Each student should point through the spectroscope directly at the fluorescent lights, and draw the component wavelengths observed: nm Station 3: hydrogen plasma Turn on the spectrum tube and observe the component wavelengths through the spectroscope. You should see individual spectral lines.. Station 4: plasma Turn on the spectrum tube and observe the component wavelengths through the spectroscope. You should see individual spectral lines.. Station 5 plasma Turn on the spectrum tube and observe the component wavelengths through the spectroscope. You should see individual spectral lines nm nm nm nm 4

5 lab5., continued Please help clean up, and answer the following questions at your normal seats:. Describe what you observed at each station:. 5. What were the colors of the individual lines from hydrogen? Which light source provided the simplest spectrum? 6. Now that you have seen a variety of emission spectra, what do you believe causes the lines? Please include details. 3. Which light source provided the most complex or varied spectrum? 4. What were the wavelengths (in nanometers) of the individual lines from hydrogen in nanometers? 7. Describe the color of light at each wavelength: 400 nm 500 nm 600 nm 700 nm 5

6 consider the atomic emission spectra shown below. where do those lines come from? 6

7 observation: light passing through slits can create multiple lines.: this is superposition introduction to the electron what do you know about light? how can this be? light travels in waves. s = w f Speed of light = wavelength x frequency 3 x 0 8 m/s per second = s - = Hz Short wavelength High frequency dangerous what is the wavelength of violet light in nanometers; f = 7.3 x 0 4 s -? s = wf s w = f = 8 3 x 0 m/s x 0 s - = 4.5 x 0-7 m = 45 nm Long wavelength Low frequency safe What is the frequency of green light, which has a wavelength of 4.90 x 0-7 m? 6. x 0 4 s - 7

8 the electromagnetic spectrum so much you don t see safe often dangerous constan t 656 nm 8

9 where are the electrons?: the story of bohr s epiphany Bohr hydrogen emission spectrum w nm = 656, 486, 434, 40 what number is next?? Balmer elements emit unique sharp lines w nm = inner Rydberg constant Balmer formula outer Try it for 3 outer w = 656 nm inner Rydberg 9

10 bohr sees the connection between light and the electron w = n electron emission creates light 0

11 434 nm (5-) -e absorption: no light -e -e hydrogen emission: it all fits 3- emission: 656 nm +P 486 (4-)nm -e 3 -e nm 400 nm 700 nm

12 energy of hydrogen photon emission Planck found the energy of a photon is proportional to it s frequency: where his Planck s constant frequency is related to wavelength where s is the speed of light E = hf s = wf h = 6.66 x 0-34 J s s = 3 x 0 8 m/s so hydrogen photon energy can be re-expressed by wavelength and since the wavelength of hydrogen electron emission is known: the energy of hydrogen photon emission can be calculated directly: E = hs/w w nm = E hydrogen =.8 x 0 8 joules.0097 ( inner outer ) inner outer

13 shell atomic orbital theory (to argon) # electrons total atomic orbitals: subshells of paired electrons paired electrons = orbital electron configuration level s 8 s and p 3 s s orbital s orbital p x orbital p y orbital p zorbital 3s orbital orbital s s p 6 3s # e s igor mikhailovskij bohr shells nucleus after: schrodinger, mikhailovskij, others 3 and p 3p x orbital 3p y orbital 3p z orbital 3p 6 orbitals 3

14 metal nonmetal aufbau order is a powerful tool s s p 6 3s 3p 6 4s 3d 0 4p 6 5s 4d 0 5p 6 6s 4f 4 5d 0 6p 6 7s 5f 4 6d 0 s s 3s 4s 5s 6s 7s valence electron + Alkali metals Group H hydrogen.0 lithium 6.94 Na sodium.99 + Alkaline earth metals Group (H is a nonmetal) 3 Li 4 Be beryllium potassium K Ca calcium Rb rubidium Cs cesium Fr francium 3.0 valence electrons , Mg Al Si P S Cl Ar magnesium 3s Transition metals: valence electrons 3p 3p 6 aluminum silicon phosphorus sulfur chlorine argon 4.3 Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9 Group 0 Group Group Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr scandium titanium vanadium chromium manganese iron cobalt nickel copper zinc gallium germanium arsenic selenium bromine 4s 3d 3d 0 4p 6 4p krypton Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I 54 Xe strontium yttrium zirconium niobium 4d molybdenum technetium ruthenium rhodium palladium silver cadmium indium tin Antimony) tellurium iodine xenon p barium Ra radium 6.0 5d 6d 7 Lu Lutetium Lr lawrencium 6. 7 Hf hafnium Rf rutherfordium Ta tantalum Db dubnium W tungsten Sg seaborgium Re rhenium Bh bohrium Os osmium Hs hassium Ir iridium Mt Meitnerium (68) - halogens helium 4.00 Group 3 Group 4 Group 5 Group 6 Group B C N O F Ne p boron carbon nitrogen s p 6 oxygen fluorine neon Pt Au Hg Tl Pb Bi Po At Rn platinum gold 5d 0 mercury thallium lead bismuth polonium astatine p 6p 6 radon (0) (0) (0) 0 Ds Rg Darmstadtium roentgenium (8) (7) Uub Ununbium (85) 7p 3 Uut ununtrium (84) Valence electrons: 8 4 Uuq ununquadium (89) 5 Uup ununpentium (88) 6 Uuh ununhexium (89) 7 Uus ununseptium (95) Noble gases Group 8 5s 4d 0 5p 6 56 Ba 6s 7s 6d He 8 Uuo ununoctium (93) s 4f 5f La Ce lanthanum cerium Ac actinium Th thorium Pr praseodymium Pa protactinium Nd neodymium U uranium Pm promethium Sm samarium Eu europium Np Pu Am neptunium plutonium americium Gd gadolinium Cm curium (47) 65 Tb terbium Bk berkelium (49) 66 Dy dysprosium Cf californium (5) 67 Ho Holmium Es einsteinium (54) 68 Er erbium Fm fermium Tm thulium Md mendelevium (56) 70 Yb ytterbium No nobelium (54) Atomic number 4f 4 to 7 to 03 Sc scandium Symbol: Solid Liquid Gas Manmade name Atomic mass 5f 4 metal nonmetal metalloid conclusion: the periodic table is based on this pattern. 4

15 electrons aufbau order is infinite s s p 6 3 d 0 4 f 5 g p 4 d 5 f 6 g s 4 p 5 d 6 f s 5 p 6 d 7 f s 6 p 7 d 6 s 7 p 8 d s 8 p 6 8 s s 9 p 0 s 6h 7g8 8f 4 9d 0 0p 6 s s s p 6 3s 3p 6 4s 3d 0 4p 6 5s 4d 0 5p 6 6s 4f 4 5d 0 6p 6 7s 5f 4 6d 0 5

16 electron configuration with orbital notation tells us where the electrons are in an atom in great detail 3Li: s s try it for carbon: 6C: orbital notation electron configuration Pauli Principle:: please give the electron configuration with orbital notation for sulfur Electrons pair up into atomic orbitals with opposite spins 6S: no! yes! X s s p s s p s s p 6 3s 3p 4 hund s rule: electrons spread out within orbital groups 6

17 heisenberg s uncertainty proposal heisenberg: einstein: electrons are part matter, part light lets say where they probably are God does not play dice with the cosmos Heisenberg is wrong 7

18 aufbau order; shorthand notation [He] [Ne] [Ar] [Kr] [Xe] [Rn] shorthand notation s s p 6 3s 3p 6 4s 3d 0 4p 6 5s 4d 0 5p 6 6s 4f 4 5d 0 6p 6 7s 5f 4 6d 0 Mg: s s p 6 3s [Ne]3s Sc: s s p 6 3s 3p 6 4s 3d [Ar]4s 3d 8

19 principles and rules of electron configuration principle or rule bad good heisenberg (e-position uncertain) aufbau (build up) hund s rule (spread out) pauli (opp. spins) s p s s s s p s s p s s 9

20 metal nonmetal it s all about the valence electrons the big idea: atoms want full outer shells. (almost always 8 electrons) neon: stable sodium: unstable: will lose lose s s 3s 4s 5s 6s 7s valence electron + Alkali metals Group H hydrogen.0 (H is a nonmetal) 3 Li lithium 6.94 Na sodium.99 9 K potassium Rb rubidium Cs cesium Fr francium 3.0 lose valence electrons + Alkaline earth metals 4 Be beryllium 9.0 Mg magnesium Ca calcium Sr strontium 56 Group 87.6 Ba barium Ra radium 6.0 3d 4d 5d 6d Sc scandium Y yttrium Lu Lutetium Lr lawrencium 6. Ti titanium Zr zirconium 9. 7 Hf hafnium Rf rutherfordium 6. 3 V 4 vanadium Nb niobium Ta tantalum Db dubnium 6. (outer shell) Transition metals: valence electrons 4 Cr chromium Mo molybdenum W tungsten Sg seaborgium 63. lose. complex: unfilled inner shells Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9 Group 0 Group Group 5 Mn manganese Tc technetium Re rhenium Bh bohrium Fe iron Ru ruthenium Os osmium Hs hassium Co cobalt Rh rhodium Ir iridium Mt Meitnerium (68) 8 Ni nickel Pd palladium Pt platinum Ds Darmstadtium (8) 9 Cu copper Ag silver Au gold Rg roentgenium (7) 30 Zn zinc Cd cadmium Hg mercury Uub Ununbium (85) p 3p 4p 5p 6p 7p Al aluminum Ga gallium In indium Tl thallium or share lose 4 or gain 4 lose Uut ununtrium (84) Noble Valence electrons: 8 gases Group , - halogens Group 3 Group 4 Group 5 Group 6 Group B C N O F boron carbon nitrogen oxygen fluorine.0 4 Si silicon Ge germanium Sn tin Pb lead Uuq ununquadium (89) gain 3 gain P phosphorus As arsenic Sb Antimony) Bi bismuth Uup ununpentium (88) S sulfur Se selenium Te tellurium Po polonium (0) 6 Uuh ununhexium stable gain (89) Cl chlorine Br bromine I iodine At astatine (0) 7 Uus ununseptium (95) 0 He helium Ne neon Ar argon Kr krypton Xe xenon Rn radon (0) 8 Uuo ununoctium (93) 4f 5f 57 La lanthanum Ac actinium Ce cerium Th thorium Pr praseodymium Pa protactinium Nd neodymium U uranium Pm promethium Np neptunium Sm samarium Pu plutonium Eu europium Am americium Gd gadolinium Cm curium (47) 65 Tb terbium Bk berkelium (49) 66 Dy dysprosium Cf californium (5) 67 Ho Holmium Es einsteinium (54) 68 Er erbium Fm fermium Tm thulium Md mendelevium (56) 70 Yb ytterbium No nobelium (54) to 7 to 03 Atomic number Sc scandium metal metalloid Symbol: Solid Liquid Gas Manmade name Atomic mass nonmetal chlorine unstable: will gain one or share 0

21 electron dot structures: a quick look at valence electrons Ne Li Be X Be no always spread out valence electrons try H,O,N,C H O N C valence electrons are the key to understanding: Chemical reactivity fortunately, it is nicely categorized in our next topic: The Periodic Table

22 The electron: fact sheet exists is located: Outside the nucleus In shells In subshells (s p d f) In orbitals With opposite spins is an elementary particle. doesn t have much mass (0-8 g; 836x lighter than a proton)

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24 name period wavelength worksheet: introduction to waves ws5. introduction to waves If you look down from Diamondhead in Hawaii, you will see waves rolling in at a steady rate. Some days they are nicely spread apart, meaning they have a long wavelength. Other days they come in more frequently; this is more dangerous for the surfers. The surfers prefer the long wavelength days. They know that as the wavelengths get shorter, their frequency gets higher, and there is more energy- more danger to the high frequency waves. (They are also paying close attention to the height, or amplitude of the wave, but that is another story. This is summarized in the diagram: waves o light are similar to waves in the ocean Light travels in the same way. It travels at a steady rate: about 300,000,000 meters per second, or 3 x 0 8 m/s. And as the wavelength decreases, the frequency must increase if their velocities don t thance. wavelength formula s = wf s= speed of light = 3 x 0 8 m/s w = wavelength in meters (m) f = frequency in waves per second (Hz, or s - ) wavelength chart Our eyes are really important to us, but they are kind of lame when you consider the tiny portion of light from the electromagnetic spectrum that they can detect: We can use the wavelength formula and the chart on the previous page to understand things like radio stations, visible light, and sunburns (due to ultraviolet light). Our ultimate goal is to make the connection between light and the electron. 4

25 name period wavelength worksheet: problems Much of what we know about the electron comes from wondering about light. Be sure you have read the previous page before answering the questions below. You will need to refer to the wavelength chart occasionally. Also, be sure you have a scientific calculator out and ready to use. Finally, refer to the solved example below to assist you. And don t forget to show your work as was done in the example. Example. What is the frequency of green light, which has a wavelength of 4.90 x 0-7 m? 8 s 3 x 0 m/s 4 - solution: s = wf; f = = = 6. x 0 s please show your work and circle your answers -7 w 4.90 x 0 m. An X-ray has a wavelength of.5 x 0-0 m. What is its frequency? ws5. (continued). The speed of light is always meters per second, and since the speed of light equals wavelength x frequency, we can rearrange this formula: wavelength = and frequency =. 5. Cable television operates at a wavelength of about 300 nanometers. What is the frequency of that wave, and what region of the electromagnetic spectrum is it in? Is it dangerous? (Any wave more frequent than visible light is considered dangerous). 3. What is the speed and wavelength of an electromagnetic wave that has a frequency of 7.80 x 0 6 Hz? 4. A popular radio station broadcasts with a frequency of 94.7 megahertz (MHz). What is the wavelength of the broadcast? ( MHz = x 0 6 Hz =,000,000 Hz) 6. Which is more dangerous, a radio wave or ultraviolet light? 7. The moon is 34,000 miles from earth. Light travels at 3 x 0 8 meters per second, and there are.609 kilometers in a mile. If a flashlight is shined at the moon, how long does it take for the flashlight s light to reach the moon s surface? The smallest particle of light is the photon. Max Planck discovered that the energy of light can be calculated, where it is simply equal to a constant number multiplied by the frequency of the light: Use the light energy formula to answer the next two questions light energy formula e = hf e is the energy of the light in joules h = Planck s Constant = 6.66 x 0-34 joules. seconds f = the frequency of light in Hz (which is /seconds) 8. What is the energy of a photon of green light? (See example at the top of the page for the frequency of green light) 9. What is the energy of a photon of light with a wavelength of.00 meters? 0. Since s =wf, and e = hf, can we calculate energy using wavelength, by combining the two formulas? Please show the combined formula, solving for energy. (Hint: note that f appears in both formulas). 5

26 Name Period the bohr model of the atom worksheet 5. Prior to the work of Niels Bohr, it was known that electrons existed outside of the nucleus, but beyond that very little was known. It was known, however, that elements, when heated, and ionized to form plasmas, gave off light of various measurable colors. Strontium emitted a beautiful magenta color, sodium produced white light, and is still used or streetlights. It was also known that if you placed a prism in front of theses elemental light forms, they divided into individual sharp wavelengths. The bigger elements gave off in some cases hundreds of lines of light, but hydrogen only gave off four visible lines. Hydrogen gas emits sharp lines of light what does it mean about the atom? To use this formula, one needs to use a whole number for n, and it will reveal the wavelength of visible light in nanometers that is in the hydrogen emission spectrum. Eventually more spectrum lines were discovered outside of the visible spectrum. The formula needed only to be generalized to account or this new data: w nm the Balmer formula general form = inner outer This more general form of the Balmer formula predicts all wavelengths of the hydrogen emission spectrum, and represents the first working model of the atom. Johann Balmer was able to see a pattern or these numbers; the balmer formula visible form w nm = n The Balmer formula exactly predicts the visible hydrogen emission spectrum wavelengths, and includes whole numbers what s up with that? Originally inner and outer were just variables, but it was Bohr who, in a moment of epiphany, realized that these were shells, ( st shell, nd shell, etc.)which he visualized as rings like a target, around the nucleus where the electron existed. As the electrons lost energy they moved from outer to inner shells and emitted packets of light of exactly the wavelengths that are observed in the hydrogen emission spectrum. In other words, his simple target model of the hydrogen atom was in good agreement with the hydrogen emission spectrum. nucleus st shell nd shell 3 rd shell Bohr s conception of what the hydrogen atom might look lie. What was the observation that Bohr based his research on? After reading the above, answer the questions below.. Solve the visible formula of the Balmer formula shown above for n = 4. What does that answer have to do with the electron? 3. The heart of Bohr s discovery was that he was able to come up with real meaning to this formula. Draw a hydrogen atom with several energy levels ( shells ) around it and show electronic emission from the fourth shell to the second shell. Indicate the movement of the electron with an arrow. 4. Draw diagrams indicating atomic emission and absorbance. Label each diagram. 6

27 Name: Period: WS5.3 electron configuration with orbital notation worksheet Directions: Using your periodic tables, Draw the electron configurations with orbital notation for each of the following atoms.. Scandium: 6S: Example: Here is the electron configuration of Sulfur with orbital notation. s s p 6 3s 3p 4 orbital notation electron configuration. Gallium: 3. Californium 8. Write the electron configuration using shorthand notation of the following elements: a. sodium b. An oxygen anion, O - 9. Two substances that have the same number of electrons are isoelectronic. For example, both the fluorine anion F - and neon have ten electrons, they are isoelectronic. a. The bromine anion is isoelectronic b. Argon is isoelectronic with which monocation? with what uncharged element? 7

28 Name Period WS 5.4 In this unit we have seen how the electrons are organized around the nucleus. It is a very detailed view of the electrons location, and various rules to help keep it all straight have been devised, and are shown below. In each problem below, the electron configuration and/or the orbital notation is incorrect. Fix it, and explain what law or principle (not Principal!) was violated. incorrect s hydrogen electron configuration NOT! worksheet example fixed Law Violated: Aufbau Principle s hydrogen Principles and rules of electron configuration Principle or rule Heisenberg (e-position uncertain) Aufbau (build up) Hund s Rule (spread out) Pauli (opp. spins) Unit 5 electrons Dr. B. s ChemAdventure Bad s p Good s s s s p s s p s s use these rules and principles to answer each question find what is wrong with the electron configuration and/or the orbital notation for each example. 7Cl Circle the incorrect region and describe the law of principle violated. s p 6 3s s 3p 5. 40Zr Circle the incorrect region and describe the law of principle violated. s s p 6 3s 3p 6 4s 3d 0 4p 6 5s 4d 8

29 3. 8 Oxygen electron configuration NOT! worksheet (page ) Circle the incorrect region and describe the law of principle violated. WS 5.4 (continued) s s p Seaborgium s s p 6 3s 3p 6 4s 3d 0 4p 6 5s 4d0 5p 6 6s 4f 4 5d0 6p 6 7s 5f 4 6d 4 Circle the incorrect region and describe the law of principle violated. 9

30 Name: Period: WS 5.5 electron configuration and orbital notation self test Chemical behavior is determined by electron position. It s a simple statement, but it says a lot. Another way of saying it is Chemistry is all about where the electrons are. That s why we ve been spending the last week focusing on electrons. However, somehow it always seems to bog down in some weird world of s s p 6, and the Pauli Principle, and we forget our goal: if we know where the electrons are we know how the substance will behave. Why Neon is stable, and sodium is very unstable, and in fact why all the elements and the substances they form behave the way they do. Let s pick an element. We know that an uncharged oxygen atom contains protons. And since it is not charged, it contains electrons. We know that of the electrons occupy the first shell, and the other six are in the second shell. We know that the first shell consists of a s orbital that holds electrons, and so we say that the electron configuration of that first shell is s. For the second shell we have six electrons, and we have learned that the first two will occupy a orbital, and the next four go into p orbitals. Thus the electron configuration of oxygen is. We can go into more detail, and show the exact orbitals that the electrons are in, which even show the relative directions the electrons are spinning in. An atomic orbital is simply a of electrons, and the Pauli Principle tells us that electrons prefer to pair up with spins. The first shell of oxygen contains one orbital, which we draw with a box like this:, showing that the electrons are paired up with opposite spins. The second shell begins with one more orbital for the two electrons of the s subshell, for a total of four electrons so far. We have more electrons in oxygen, and they will occupy the three p orbitals. We remember to apply rule and spread these electrons out as far as possible in those three boxes. Thus we can draw the electron configuration of oxygen with its orbital notation right above it: Note that this tells us that oxygen has four electrons in its outer (second) shell, and the two of them are unpaired.we also know from HONC that oxygen likes to form two bonds a coincidence?? Let s work out the electron configuration of nitrogen and see if we get three unpaired electrons: Nitrogen has electrons, so the electron configuration with orbital notation is (be sure to spread out your p electrons): 30

31 Howtoaceitunit5 how to ace the electrons test In this unit our goal was to determine where the electrons are in atoms. To find out, we performed two experiments that revealed the sharp lines that ionized elements produced. We then analyzed this data from a historical perspective, beginning with the work of Niels Bohr. For this we needed to review the properties of light, including frequency, wavelength, energy, and, common types. This involved the use of the speed of light equation (s = wf), the energy of an electron (e = hf), and an understanding of the electromagnetic spectrum. We then showed how the key mathematical solutions of Balmer and Rydberg allowed Bohr to put it all together to postulate energy levels, where atomic emission explains light, and produces the spectral lines observed for all elements. This was followed by a detailed look at the electron around the nucleus. We found that not only do electrons reside in shells, there are also subshells and orbitals within each shell. We observed how they spread out within an orbital group (Hund s Rule), and even how they spin when near each other (the Pauli Principle). We learned the configurations of electrons for all elements follow the Aufbau Order, and we write it all down by electron configuration, and orbital notation. This can rapidly tell us how many electrons are in each shell and subshell, the spin of each electron, and the number of unpaired electrons. There may be limits to what can be known about the electron. This is the controversial Heisenberg Uncertainty Principle, which states that it is impossible to measure the position and velocity of an electron simultaneously, due to the extreme sensitivity of the electron, and perhaps due to the fact that an electron is so small it is partly matter and partly energy. Finally, we showed how valence is easy to determine using the periodic table, and that valence may be drawn using electron dot formulas, also known as Lewis Dot Formulas. We emphasized that nearly all chemical behavior is simply a result of those outer shell electron.. Draw the symbols for Democritus, Aristotle, Ghazali, Lavoisier, Dalton, Thomson, Rutherford, and Bohr. What is the significance of each symbol? Try to assign one or two key words for each symbol. 3. What are the dangerous wavelengths of light? 4. How does light relate to electrons? 5. What is wavelength? Units? 6. What is frequency? Provide 3 equivalent units. 7. Rearrange the speed of light equation to show what frequency is equal to. 8. The electromagnetic spectrum: what is it? The range of frequencies emitted by electrons when they move around 9. Frequency: how does it relate to energy and safety? 0. Wavelength- how does it relate to frequency?. Energy: which rays have the highest energy?. Safety: why are radio waves generally considered safe? 3. Where is the visible region of the electromagnetic spectrum? 4. Convert 45 nanometers to meters (0 9 nm = m) 5. Use s = wf to find the frequency of 45 nm light. 6. (Level one only) The Balmer formula. Find it in your notes: 7. Significance: 8. Solve for the n= 3 to n = transition: 9. Atomic Emission Spectra: How did we observe it? 0. Emission vs. absorbance- what is the difference?. The Bohr model of the atom- draw a model.5 What is the difference between electron configuration, and orbital notation? During this study we found that the periodic table is well designed to show the number of valence electrons for any element. In our next unit we will apply this to our understanding of the periodic table. To dominate this test, review all of the material in his packet: The lessons, the labs, and the worksheets. Here are some questions you should answer to help prepare for the test: 3

32 . Know the electron names up to zirconium. For example, manganese has the symbol. Here are some others: Mg B Be Na K Y Sr Cr Howtoaceitunit5 3. Be able to provide the electron configurations for all elements. Do iodine using noble gas notation. 4. L only: Be able to provide the orbital notation for all elements. Do silicon. Include the number of valence electrons and the number of unpaired electrons. 5. The Heisenberg Uncertainty Principle. State what it is and why briefly. 6. L only: Sub-shells: s, p, d, and f how many electrons for each? How many orbitals for each? 7. L only: Aufbau principle. Give an example where it is broken, and fix it. 8. L only: Pauli exclusion principle. Break it and fix it. 9. L only: Hund s Rule. Break it and fix it. 30. Lewis Dot Structures, also known as electron dot structures. Draw oxygen, for example 3. Valence Electrons. Do each column in the periodic table. 3. Why is it important to use scientific references, rather than websites, when writing a scientific paper? 33. Where are the electrons in an atom? Be prepared to answer this question in detail, with evidence to support your answer. To help, arrange the following from a elementary to a deep level of understanding: a. the atom exist, b. electrons exist in shells, c. electrons exist, electrons exist in subshells, d. electrons spin in opposite directions in orbitals, e. electrons exist in orbitals, f. electrons exist outside the nucleus. The best answers would include evidence to support each claim. 3