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ATOMIC STRUCTURE AND PERIODICITY CHAPTER 7 WARM-UP QUESTION Be prepared to share out your response to the following questions. What is a photon? What is the source of electromagnetic waves? Is the color spectrum simply a small segment of the electromagnetic spectrum? Defend your answer. ELECTROMAGNETIC RADIATION WARM-UP CONTINUED What is a photon? • A particle of light. • Particle vs Wave Theory • Video 1 • Video 2 What is the source of electromagnetic waves? • Accelerating electric charges Is the color spectrum simply a small segment of the electromagnetic spectrum? Defend your answer. • Yes; the spectrum is also made up of radio waves, IR, UV, Xrays, and gamma rays. CHARACTERISTICS OF WAVES Waves are described according to their Amplitude measures DISPLACEMENT size of the disturbance (from rest to crest) Wavelength distance of a “repeating unit” Also called a cycle Velocity v speed = how fast wave travels FREQUENCY V How often number of wavelengths that pass any point per second measured in wavelengths/second or cycles/second Hertz (Hz) = number of wavelengths in 1 second Frequency is related to velocity: c = v ELECTROMAGNETIC WAVE a transverse wave with an electric component and a magnetic component at right angles to each other How are electromagnetic waves (ex: light) different from mechanical waves (ex: sound and slinky)? micro.magnet.fsu.edu ELECTROMAGNETIC WAVES Electromagnetic waves are special in the fact that they do not need a medium to propagate through. But what is creating the disturbance? What is emitting this energy? © 2003 Mike Maloney 8 ELECTRONS ELECTROMAGNETIC WAVES Electrons in materials are vibrated and emit energy in the form of photons, which propagate across the universe. Photons have no mass, but are pure energy. Electromagnetic Waves are waves that are made up of these “photons”. © 2003 Mike Maloney 9 When these photons come in contact with boundaries, E-M waves interact like other waves would. © 2003 Mike Maloney 10 ELECTROMAGNETIC SPECTRUM SPEED OF E/M WAVES It has been found that the speed of E-M waves and light is --- © 2003 Mike Maloney 11 • 3 x 108 or 300,000,000 m/s • 671,000,000 mph • 186,000 miles per second • We call this value “c” C is constant throughout the universe, as long as light is in a vacuum. When it is in other materials, c can change, but can never be larger than its value in a vacuum. © 2003 Mike Maloney 12 Since “c” is constant, all of E-M waves will have a corresponding frequency to go along with their wavelength. ENERGY IN E-M WAVES Which waves have more energy, Radio waves or gamma waves? The greater the frequency of an E-M wave, the more crests pass a point in a certain amount of time, therefore the more photons pass that point. © 2003 Mike Maloney 13 This means that more energy moves past that point in a certain amount of time or that the wave contains more energy. ELECTROMAGNETIC SPECTRUM “CHECK-UP” True or False… 1. Blue light has a shorter wavelength than red light. 2. X-rays have lower frequencies than radio waves. 3. Microwaves have higher frequencies than gamma rays. 4. Visible radiation composes the major portion of the electromagnetic spectrum. True; False; False; False WAVELENGTH-FREQUENCY RELATIONSHIP EXAMPLE Photosynthesis uses light with a frequency of 4.54x1014s-1. What wavelength does this correspond to? A: 660nm WAVELENGTH-FREQUENCY RELATIONSHIP PRACTICE Calculate the frequency of blue light of wavelength 4.5 x 102nm. Calculate the wavelength of green light of frequency 5.7 x 1014Hz. A:6.7x1014Hz ; 5.3 x 10-7m or 530nm THE NATURE OF MATTER ΔE = hv = hc/λ • ΔE is the change in energy for a system (in Joules per photon) • h is Planck’s constant (6.626 x 10-34J s) • experimentally determined • v is the frequency of the wave (s-1 or Hz) **Energy can be gained or lost only in integer multiples of hv. (quanta) ENERGY, FREQUENCY, WAVELENGTH EXAMPLE Sodium atoms have a characteristic yellow color when excited in a flame. The color comes from the emission of 589.0nm. • What is the frequency of this radiation? • What is the change in energy associated with this photon? Per mole of photons? ENERGY, FREQUENCY, WAVELENGTH PRACTICE It takes 382 kJ of energy to remove one mole of electrons from gaseous cesium. What is the wavelength associated with this energy? Would we be able to “directly” observe this energy change? Why or why not. THE PHOTOELECTRIC EFFECT Emission of electrons from a metal when light shines on the metal Electromagnetic radiation (light) strikes the surface of the metal ejecting electrons from the metal and causing an electric current, if the frequency was below a certain minimum. Analysis of the kinetic energy and numbers of the emitted electrons led Einstein to suggest that electromagnetic radiation can be viewed as a stream of photons. *Note that the apparent mass of a photon depends on its wavelength. The mass of a photon at rest is thought to be zero, although we never observe it at rest.* BIG IDEAS FROM EINSTEIN AND PLANCK • Energy is quantized. It can occur only in discrete units called quanta. • Electromagnetic radiation, which was previously thought to exhibit only wave properties, seems to show certain characteristics of particulate matter as well. (dual nature of light) WAVE-LIKE BEHAVIOR Diffraction • Light is scattered from a regular array of points or lines. • Constructive interference • In-phase (bright) • Destructive interference • Out-of phase (dim/dark) ATOMIC SPECTRUM OF HYDROGEN Continuous Spectrum • Contains all the wavelengths over which the spectrum is continuous Line Spectrum • Contains certain specific wavelengths that are characteristic of the substance emitting those wavelengths *Hydrogen’s line spectrum shoes that only certain energy transfers are allowed in hydrogen. *Specific energy levels among which the hydrogen electron can shift, thus energy levels are quantized. THE BOHR MODEL 1913 Niels Bohr developed the Quantum Model for the hydrogen atom. • The electron in the hydrogen atom moves around the nucleus only in certain allowed circular orbits. • Hydrogen atom energy levels consistent with the hydrogen emission spectrum. (different wavelength/color associated with the different levels of emission) • Ground state • The lowest possible energy state of an atom or molecule • Excited state • Higher potential energy state than ground state of an atom or molecule Although Bohr’s model fits the energy levels for hydrogen, it is a fundamentally incorrect model for the hydrogen atom. Bohr’s model paved the way for later theories on the quantization of energy in atoms. Electrons do NOT move around the nucleus in circular orbits (planetary model). QUANTUM MECHANICS de Broglie and Schrodinger – wavelike properties of electrons A specific wave function (function of the coordinates x, y, and z of the electron’s position in 3-D space) is often called an orbital. • The wave function corresponding to the lowest energy for the hydrogen atom is called the 1s orbital (no association to the Bohr “orbit”). Nature of an orbital takes into consideration the work of Heisenberg. • Heisenberg uncertainty principle: There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time. 1S ORBITAL The definition most often used by chemists to describe the size of the hydrogen 1s orbital is the radius of the sphere that encloses 90% of the total electron probability. (90% of the time the electron I in this sphere) SUMMARY In the quantum (wave) mechanical model, the electron is viewed as a standing wave. This representation leads to a series of wave functions (orbitals) that describe the possible energies and spatial distributions available to the electron. In agreement with the Heisenberg uncertainty principle, the model cannot specify the detailed electron motions. Instead, the square of the wave function represents the probability distribution of the electron in that orbital. This allows us to picture orbitals in terms of probability distributions, or electron density maps. The size of an orbital is arbitrarily defined as the surface that contains 90% of the total electron probability. The hydrogen atom has many types of orbitals. In the ground state, the single electrons resides in the 1s orbital. The electron can be excited to higher-energy orbitals if energy is put into the atom. QUANTUM MECHANICS HTTP://WWW.META-SYNTHESIS.COM/WEBBOOK/30_TIMELINE/310PX-BOHRATOM-PAR.SVG.PNG Better than any previous model, quantum mechanics does explain how the atom behaves. Quantum mechanics treats electrons not as particles, but more as waves (like light waves) which can gain or lose energy. But they can’t gain or lose just any amount of energy. They gain or lose a “quantum” of energy. A quantum is just an amount of energy that the electron needs to gain (or lose) to move to the next energy level. In this case it is losing the energy and dropping a level. ATOMIC ORBITALS HTTP://MILESMATHIS.COM/BOHR2.JPG Much like the Bohr model, the energy levels in quantum mechanics describe locations where you are likely to find an electron. Remember that orbitals are “geometric shapes” around the nucleus where electrons are found. Quantum mechanics calculates the probabilities where you are “likely” to find electrons. ATOMIC ORBITALS HTTP://COURSES.CHEM.PSU.EDU/CHEM210/QUANTUM/QUANTUM.HTML Of course, you could find an electron anywhere if you looked hard enough. So scientists agreed to limit these calculations to locations where there was at least a 90% chance of finding an electron. Think of orbitals as sort of a "border” for spaces around the nucleus inside which electrons are allowed. No more than 2 electrons can ever be in 1 orbital. The orbital just defines an “area” where you can find an electron. What is the chance of finding an electron in the nucleus? Yes, of course, it’s zero. There are not any electrons in the nucleus. ENERGY LEVELS HTTP://WWW.CHEM4KIDS.COM/FILES/ART/ELEM_PERTABLE2.GIF Quantum mechanics has a principal quantum number. It is represented by a little n. It represents the “energy level” similar to Bohr’s model. Red Orange Yellow Green Blue Indigo Violet n=1 n=2 n=3 n=4 n=5 n=6 n=7 • n=1 describes the first energy level • n=2 describes the second energy level • Etc. Each energy level represents a period or row on the periodic table. It’s amazing how all this stuff just “fits” together. SUB-LEVELS = SPECIFIC ATOMIC ORBITALS Each energy level has 1 or more “sub-levels” which describe the specific “atomic orbitals” for that level. Blue = s block • n = 1 has 1 sub-level (the “s” orbital) • n = 2 has 2 sub-levels (“s” and “p”) • n = 3 has 3 sub-levels (“s”, “p” and “d”) • n = 4 has 4 sub-levels (“s”, “p”, “d” and “f”) There are 4 types of atomic orbitals: • s, p, d and f • Each of these sub-levels represent the blocks on the periodic table. ORBITALS HTTP://MEDIA-2.WEB.BRITANNICA.COM/EB-MEDIA/54/3254-004-AEC1FB42.GIF HTTP://UPLOAD.WIKIMEDIA.ORG/WIKIPEDIA/COMMONS/THUMB/E/E1/D_O RBITALS.SVG/744PX-D_ORBITALS.SVG.PNG s p d In the s block, electrons are going into s orbitals. In the p block, the s orbitals are full. New electrons are going into the p orbitals. In the d block, the s and p orbitals are full. New electrons are going into the d orbitals. What about the f block? QUANTUM NUMBERS Describe the properties of the orbital. Name Symbol Property of the Orbital Related to size and energy of the orbital Range of Values Principal Quantum Number n Integers Angular Momentum Quantum Number l Related to the shape of the orbital Integers from “subshell” n-1 to 0 0 is s; 1 is p; 2 is d; 3 is f; 4 is g; 5 is h Magnetic Quantum Number ml Related to the position of the orbital in space relative to other orbitals 1 to ∞ Integers from -l to 0 to +l DEGENRATE All orbitals having the same value of “n” have the same energy. 3s; 3p; 3d Energy is required to transfer an electron to a higher-energy orbital (excited state). **In polyelectronic atoms we find that the s, p, and d have different levels of potential energy. THE 4TH QUANTUM NUMBER The electron spin quantum number. Electron spin • Two spin states + ½ and – ½ • Produce two oppositely directed magnetic moments Pauli Exclusion Principle • In a given atom no two electrons can have the same set of four quantum numbers (n, l, ml, ms) • Thus, an orbital can hold only TWO electrons, and they must have opposite spins. PRACTICE WITH QUANTUM NUMBERS Which of the following quantum numbers are allowed? For each that is incorrect state why. Principal, Angular Momentum, Magnetic Quantum Numbers (n, l, ml) a. 1, 0, 1 b. 2, 2, 1 c. 5, 3, 2 d. 6, -2, 2 e. 6, 2, -2 QUANTUM NUMBERS AND LEVELS OF ORBITALS Table 7.2 on page 294 in text Energy Level Sublevels Total Orbitals Total Electrons Total Electrons per Level n=1 s 1 (1s orbital) 2 2 n=2 s p 1 (2s orbital) 3 (2p orbitals) 2 6 8 n=3 s p d 1 (3s orbital) 3 (3p orbitals) 5 (3d orbitals) 2 6 10 18 Complete the chart in your notes as we discuss this. first has an s orbital. n = The 4 s level (n=1) 1 (4s orbital) 2 It has only 32 1. There pare no other energy level. 3 (4p orbitals orbitals) in the first 6 d 5 (4d orbitals) 10 We call this orbital the 1s orbital. f 7 (4f orbitals) 14 WHERE ARE THESE ORBITALS? HTTP://WWW.BIOSULF.ORG/1/IMAGES/PERIODICTABLE.PNG 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 5d 6p 7s 6d 7p 4f 5f ELECTRON CONFIGURATIONS What do I mean by “electron configuration?” The electron configuration is the specific way in which the atomic orbitals are filled. Think of it as being similar to your address. The electron configuration tells me where all the electrons “live.” RULES FOR ELECTON CONFIGURATIONS HTTPS://TEACH.LANECC.EDU/GAUDIAS/SCHEME.GIF In order to write an electron configuration, we need to know the RULES. 3 rules govern electron configurations. • Aufbau Principle • Pauli Exclusion Principle • Hund’s Rule Using the orbital filling diagram at the right will help you figure out HOW to write them • Start with the 1s orbital. Fill each orbital completely and then go to the next one, until all of the elements have been acounted for. Each line represents an orbital. 1 (s), 3 (p), 5 (d), 7 (f) High Energy FILL LOWER ENERGY ORBITALS FIRST HTTP://WWW.META-SYNTHESIS.COM/WEBBOOK/34_QN/QN3.JPG The Aufbau Principle states that electrons enter the lowest energy orbitals first. The lower the principal quantum number (n) the lower the energy. Within an energy level, s orbitals are the lowest energy, followed by p, d and then f. F orbitals are the highest energy for that level. Low Energy NO MORE THAN 2 ELECTRONS IN ANY ORBITAL…EVER. HTTP://WWW.FNAL.GOV/PUB/INQUIRING/TIMELINE/IMAGES/PAULI.JPG Wolfgang Pauli, yet another German Nobel Prize winner The next rule is the Pauli Exclusion Principal. The Pauli Exclusion Principle states that an atomic orbital may have up to 2 electrons and then it is full. The spins have to be paired. We usually represent this with an up arrow and a down arrow. Since there is only 1 s orbital per energy level, only 2 electrons fill that orbital. Quantum numbers describe an electrons position, and no 2 electrons can have the exact same quantum numbers. Because of that, electrons must have opposite spins from each other in order to “share” the same orbital. HUND’S RULE HTTP://INTRO.CHEM.OKSTATE.EDU/AP/2004NORMAN/CHAPTER7/LEC111000.HTML Don’t pair up the 2p electrons until all 3 orbitals are half full. Hunds Rule states that when you get to degenerate orbitals, you fill them all half way first, and then you start pairing up the electrons. What are degenerate orbitals? Degenerate means they have the same energy. So, the 3 p orbitals on each level are degenerate, because they all have the same energy. Similarly, the d and f orbitals are degenerate too. Paramagnetic unpaired electrons 2p Diamagnetic all electrons paired 2p APPLICATION NOW that we know the rules, we can try to write some electron configurations. Remember to use your orbital filling guide/PERIODIC TABLE to determine WHICH orbital comes next. Lets write some electron configurations for the first few elements, and let’s start with hydrogen. H; Li; B; N; F; Na; K; Fe ELECTRON CONFIGURATIONS Element Configuration Element Configuration H Z=1 1s1 He Z=2 1s2 Li Z=3 1s22s1 Be Z=4 1s22s2 B Z=5 1s22s22p1 C Z=6 1s22s22p2 N Z=7 1s22s22p3 O Z=8 1s22s22p4 F 1s22s22p5 Ne Z=10 1s22s22p6 (2p is now full) Na Z=11 1s22s22p63s1 Cl Z=17 1s22s22p63s23p5 K Z=19 1s22s22p63s23p64s1 Sc Z=21 1s22s22p63s23p64s23d1 Fe Z=26 1s22s22p63s23p64s23d6 Br Z=35 1s22s22p63s23p64s23d104p5 Z=9 Note that all the numbers in the electron configuration add up to the atomic number for that element. Ex: for Ne (Z=10), 2+2+6 = 10 CONCEPTUAL CHECK One last thing. Look at the previous slide and look at just hydrogen, lithium, sodium and potassium. Notice their electron configurations. Do you see any similarities? Since H and Li and Na and K are all in Group 1A, they all have a similar ending. (s1) ELECTRON CONFIGURATIONS Element Configuration H Z=1 1s1 Li Z=3 1s22s1 Na Z=11 1s22s22p63s1 K Z=19 1s22s22p63s23p64s1 This similar configuration causes them to behave the same chemically. It’s for that reason they are in the same family or group on the periodic table. Each group will have the same ending configuration, in this case something that ends in s1. NOBLE GAS NOTATION… “SHORT CUT” Be Al Br Mo Ag ORBITAL NOTATION Be Al N Br Mo Ag ION ELECTRON CONFIGURATION Be2+ Al3+ Br- Ag1+ PERIODIC TRENDS Atomic Radii Ionic Radii Electronegativity Ionization Energy Electron Affinity ATOMIC RADIUS 2r Definition: Half of the distance between nuclei in covalently bonded diatomic molecule. Radius decreases across a period • Increased effective nuclear charge due to decreased shielding Radius increases down a group • Each row on the periodic table adds a “shell” or energy level to the atom IONIZATION ENERGY Definition: the energy required to remove an electron from an atom **Increases for successive electrons taken from the same atom. (First Ionization Energy is lower than the second…and that lower than the third…etc) Tends to increase across a period (general pattern-valence e-) • Electrons in the same quantum level do not shield as effectively as electrons in inner levels. • Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove Tends to decrease down a group • Outer electrons are farther from the nucleus and easier to remove IONIZATION ENERGY AND ORBITAL NOTATION Phosphorus vs. Sulfur Potassium 1st Ionization vs. 2nd ELECTRON AFFINITY Definition: the energy change associated with the addition of an electron Affinity tends to increase across a period Affinity tends to increase up in a group • Electrons farther from the nucleus experience less nuclear attraction • Some irregularities due to repulsive forces in the relatively small “p” orbitals ELECTRONEGATIVITY Definition: A measure of the ability of an atom in a chemical compound to attract electrons from another atom. Electronegativity tends to increase across a period • As the radius decreases, electrons get closer to the bonding atom’s nucleus Electronegativity tends to increase up a group or remain the same. • As radius increases, electrons are farther from the bonding atom’s nucleus IONIC RADII Cations • Positively charged ions formed when an atom loses one or more electrons • SMALLER that the corresponding atom Anions • Negatively charged ions formed when an atom gains one or more electrons • LARGER than the corresponding atom ANALYZING PERIODIC TABLE TRENDS WITH PEERS Explain why atomic radius increases going down a group and right to left in a period. Rank the following from largest to smallest. Explain why electronegativity, ionization energy, and electron affinity increase from left to right and from bottom to top. Rank the following from largest to smallest. SPECTROSCOPY Photoelectron Spectroscopy (PES) • Provides data for ionization energy trends and applications Mass Spectrometry • Provides atomic/molar mass data as it ionizes SPECTROSCOPY GROUPS Group 1: Group 2: Group 3: Group 4: Group 5: Group 6: PLEASE BRING YOUR GREEN LAB MANUAL TOMORROW!! PHOTOELECTRON SPECTROSCOPY (PES) Ephoton = hv Atom Monochromatic Beam of X-Rays IEelectron = Ephoton - KE KE = mv2 2 e- PHOTOELECTRON SPECTRUM Relative Intensity = 2 Each peak is relative to the others. This indicates the relative number of electrons. If the peak is twice as big, there are twice as many electrons. Relative Intensity = 1 20 MJ/mol 10 MJ/mol 0 MJ/mol PHOTOELECTRON SPECTRUM Valence 19.3 MJ/mol 1.36 MJ/mol RI = 2 RI = 2 Gap is due to increased energy of the orbital (decreased amount of energy to remove the electron) 20 MJ/mol 10 MJ/mol 0.80 MJ/mol RI = 1 0 MJ/mol PHOTOELECTRON SPECTRUM Boron (Z=5) 19.3 MJ/mol RI = 2 Analysis: 1) Valence has 2 values: 2) RI is 2 to 1 in valence: 3) Closest core has RI 2 not 6: 4) s2s2p1 must be 1s22s22p1 RI = 2 0.80 MJ/mol RI = 1 2s2 1s2 20 MJ/mol 1.36 MJ/mol 10 MJ/mol Inner orbitals require the most energy 2p1 0 MJ/mol Valence orbitals require the least energy PHOTOELECTRON SPECTRUM Depending on the size of the table, 1s may be intentionally cut out of view because it’s too far away and makes the graph too long Remember IE is about REMOVING electrons, which means they are removed from the OUTSIDE to the INSIDE, and NOT in reverse order of energy! For example, 4s is removed BEFORE 3d. PHOTOELECTRON SPECTRUM 2p6 3.67 MJ/mol 1s2 2s2 104 MJ/mol Sodium (Z=11) 3s1 6.84 MJ/mol 0.50 MJ/mol {} 8 MJ/mol 4 MJ/mol 0 MJ/mol Online PES Resources http://www.chem.arizona.edu/chemt/Flash/photoelectron. html https://www.youtube.com/watch?v=NRIqXeY1R_I https://www.youtube.com/watch?v=vANbxozsRSA From the AP Sample Questions… From the AP Sample Questions… Which peaks in the photoelectron spectrum are representative of the binding energy of p orbital electrons? a. C only c. C and E b. D only d. B, C and D Mass Spectrometry • Mass spectrometry gives the mass to charge ratio • Like PES, the relative size of the peaks indicates the relative number of particles • Separates isotopes according to mass • Used to find relative abundance and atomic/molar mass of unknown samples Mass Spectrometry From the AP Sample Questions… The elements I and Te have similar average atomic masses. A sample that was believed to be a mixture of I and Te was run through a mass spectrometer, resulting in the data above. All of the following statements are true. Which one would be the best basis for concluding that the sample was pure Te? From the AP Sample Questions… a. Te forms ions with a -2 charge, whereas I forms ions with a -1 charge. b. Te is more abundant that I in the universe. c. I consists of only one naturally occurring isotope with 74 neutrons, whereas Te has more than one isotope. d. I has a higher first ionization energy than Te does. Based on the mass spectrum of atom Y, which of the following statements is false? a. peak A and peak D come from atoms that have the same number of electrons b. there are seven isotopes of atom Y c. peak C comes from the most abundant isotope of atom Y d. peak D comes from an atom with 4 more protons than the atom that gave peak B PREPARATION OF A STANDARD SOLUTION A chemist whishes to prepare 1.00L of a 0.200 M sodium hydroxide solution. Describe the steps, with calculations, necessary to complete this task starting with solid sodium hydroxide and distilled water. DILUTION OF SOLUTIONSM1V1 = M2V2 (a) A Measuring Pipet (b) A Volumetric (transfer) Pipet You’ve been asked to prepare 150 ml of a 0.035M solution of sodium hydroxide from the 0.200M stock sodium hydroxide solution prepared earlier. Detail the steps necessary to complete this task. BEER- LAMBERT LAW Relates the amount of light being absorbed to the concentration of the substance absorbing the light A=abc A = measured absorbance a = molar absorptivity constant (a characteristic of the substance being monitored). b = path length through which the light must pass. c = Molar concentration of the absorbing substance. BEER’S LAW SAMPLE PROBLEMS 1. A solution with a concentration of 0.14M is measured to have an absorbance of 0.43. Another solution of the same chemical is measured under the same conditions and has an absorbance of 0.37. What is its concentration? 2. The following data were obtained for 1.00 cm samples of a particular chemical. What is the concentration of a 1.00 cm sample that has an absorbance of 0.60? Conc. Abs. (M) 3. The absorptivity of a particular chemical is 1.5/M·cm. What is the concentration of a solution made from this chemical if a 2.0 cm sample has an absorbance of 1.20? 0.50 0.69 0.40 0.55 0.30 0.41 0.20 0.27 BEER’S LAW SAMPLE PROBLEMS 4. Using the data from the graphing example in question #2, what are the concentrations of solutions with absorbances of 0.20, 0.33, and 0.47? 5. A solution is prepared to be 0.200M. A sample of this solution 1.00 cm thick has an absorbance of 0.125 measured at 470nm and an absorbance of 0.070 measured at 550nm. Calculate the concentrations of the following solutions: Sample Absorbance Wavelength Path length 1 0.055 470nm 1.00cm 2 0.155 470nm 1.00cm 3 0.120 550nm 1.00cm 4 0.048 550nm 5.00cm GATORADE LAB/INVESTIGATION What is the relationship between the concentration of a solution and the amount of transmitted light through the solution? Video Spectrophotometer See Pre-lab and Report Sheet for support. **Dilution calculations MUST be done prior to data collection!!** (Due TOMORROW!) –complete first 3 columns in data table prior to entering lab. LAB PARTNERS AND DILUTION ASSIGNMENTS Dilution Assignments (mL of stock/ mL of water) Group 1: • 10mL/0mL; 6mL/4mL • 9mL/1mL; 3mL/7mL Group 2: • 5mL/5mL; 0mL/10mL • 8mL/2mL; 1mL/9mL Group 3: • 2mL/8mL; 4mL/6mL, • 0mL/10mL, 7mL/3mL Group 4: • 10mL/0mL; 6mL/4mL • 9mL/1mL; 3mL/7mL Group 5: • 5mL/5mL; 0mL/10mL • 8mL/2mL; 1mL/9mL Group 6: • 2mL/8mL; 4mL/6mL, • 0mL/10mL, 7mL/3mL Will share class data for plots. UNIT 3 EXAM ANALYSIS 20 points total 10 points Multiple Choice (MC) 10 points Free Response Questions (FRQ) -Analyze areas of strength and those in need of assistance. -Utilize Learning Objectives and Topic Sheets (tan colored) -Reflect on *labs *practice packets *readings/notes UNIT 3 EXAM ANALYSIS 30 minutes individual review. 10 minutes peer review….for half credit (based on reflection)! Learning is a PROCESS! For credit you must correct your answer and provide a brief reflection of why your answer choice was incorrect and how you corrected it. (3 to 4 sentences per question) Question Number (MC or FRQ) Incorrect Answer Corrected Answer REFLECTION of LEARNING