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1 Chapter 7 Atomic Structure

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2 7.1 Types of EM waves They have different and They have different and n Radio waves, microwaves, infra red, ultraviolet, x-rays and gamma rays are all examples. n Light is only the part our eyes can detect. Gamma Rays Radio waves Frequency decreases Wavelength increases

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3 Parts of a wave Wavelength Frequency = number of cycles in one second Measured in hertz 1 hertz = 1 cycle/second

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4 Frequency = Frequency =

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5 The speed of light c = in a vacuum is x 10 8 m/s c = c = in a vacuum is x 10 8 m/s c = n What is the wavelength of light with a frequency 5.89 x 10 5 Hz? n What is the frequency of blue light with a wavelength of 484 nm? n (YDVD) EMwave

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6 7.2 Nature of Matter n Matter and energy were seen as different from each other in fundamental ways. n Matter was particles. n Energy could come in waves, with any frequency. n Max Planck found that as the cooling of hot objects couldnt be explained by viewing energy as a wave.

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7 Energy is Quantized Planck found E came in chunks with size h Planck found E came in chunks with size h E = n h E = n h n where n is an integer. n and h is Plancks constant n h = x Js These packets of h are called quantum. These packets of h are called quantum.

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8 Einstein is next n Said electromagnetic radiation is quantized in particles called photons. Each photon has energy = h = h c/ Each photon has energy = h = h c/ n Blue color in fireworks has a wavelength of 450 nm. How much energy does a photon of blue light have? A mole of these photons? n Combine this with E = mc 2 n You get the apparent mass of a photon. m = h /( c) m = h /( c)

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9 Which is it? n Is energy a wave like light, or a particle? n Yes n Concept is called the Wave - Particle duality. (Flatman) n What about the other way, is matter a wave? Yes n Photoelectric Effect (YDVD) n (BDVD) Blue & Yellow

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10 Matter as a wave n All matter exhibits both particle and wave properties. De Broglies equation = h/mv De Broglies equation = h/mv n We can calculate the wavelength of an object. n The laser light of a CD is 7.80 x 10 2 m. What is the frequency of this light? n What is the energy of a photon of this light? n What is the apparent mass of a photon of this light? n What is the energy of a mole of these photons?

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11 What is the wavelength? m = h /( c) m = h /( c) n Of an electron with a mass of 9.11 x kg traveling at 1.0 x 10 7 m/s? n Of a softball with a mass of 0.10 kg moving at 35 mi/hr?

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Hydrogen spectrum n Emission spectrum because these are the colors it gives off or emits. n Called a line spectrum. n There are just a few discrete lines showing only certain energies are possible. 410 nm 434 nm 486 nm 656 nm

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13 What this means n Only certain energies are allowed for the hydrogen atom. n Can only give off certain energies. Use E = h = h c/ Use E = h = h c/ n Energy in the in the atom is quantized. n (YDVD) Refraction

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Niels Bohr n Developed the quantum model of the hydrogen atom. n He said the atom was like a solar system. n The electrons were attracted to the nucleus because of opposite charges. n Didnt fall in to the nucleus because it was moving around.

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15 The Bohr Ring Atom n He didnt know why but only certain energies were allowed. n He called these allowed energies energy levels. n Putting Energy into the atom moved the electron away from the nucleus. n From ground state to excited state. n When it returns to ground state it gives off light of a certain energy.

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16 The Bohr Ring Atom n = 3 n = 4 n = 2 n = 1

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17 The Bohr Model n n is the energy level (an integer). n Z is the nuclear charge, which is +1 for hydrogen. n E = x J (Z 2 /n 2 ) n n = 1 is called the ground state. n When the electron moves from one energy level to another. E = E final - E initial E = E final - E initial E = x J Z 2 (1/ n f 2 - 1/ n i 2 ) E = x J Z 2 (1/ n f 2 - 1/ n i 2 )

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18 Examples E = x J Z 2 (1/ n f 2 - 1/ n i 2 ) E = x J Z 2 (1/ n f 2 - 1/ n i 2 ) n Calculate the energy need to move an electron from level n=1 to level n=2. n What is the wavelength of light that must be absorbed by a hydrogen atom in its ground state to reach this excited state. n Calculate the energy released when an electron moves from n= 5 to n=3 in a He +1 ion

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19 The Bohr Model n Doesnt work. n Only works for hydrogen atoms. n Electrons dont move in circles. n The quantization of energy is right, but not because they are circling like planets.

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The Quantum Mechanical Model n A totally new approach. n De Broglie said matter could be like a wave. n De Broglie said they were like standing waves. n The vibrations of a stringed instrument.

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21 Whats possible? n You can only have a standing wave if you have complete waves. n There are only certain allowed waves. n In the atom there are certain allowed waves called electrons. n 1925 Erwin Schroedinger described the wave function of the electron. n A lot of math, but what is important are the solutions.

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22 Schroedingers Equation The wave function ( is a (x, y, z) function. The wave function ( is a (x, y, z) function. n Solutions to the equation are called orbitals. n These are not Bohr orbits. n Each solution is tied to a certain energy. n These are the energy levels.

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23 The Heisenberg Uncertainty Principle. n There is a fundamental limitation on how precisely we can know both the position and momentum of a particle at a given time. n MATHEMATICALLY: x · (mv) > h /4 x · (mv) > h /4 x is the uncertainty in the position. x is the uncertainty in the position. (mv) is the uncertainty in the momentum. (mv) is the uncertainty in the momentum. the minimum uncertainty is h /4 the minimum uncertainty is h /4

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24 What does the wave Function mean? n Nothing. It is not possible to visually map it. n The square of the function is the probability of finding an electron near a particular spot. n Best way to visualize it is by mapping the places where the electron is likely to be found.

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25 Defining the size n The nodal surface. n The size that encloses 90% to the total electron probability. n NOT at a certain distance, but a most likely distance. n For the first solution it is a a sphere. This is of course and s orbital.

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Quantum Numbers n There are many solutions to Schroedingers equation n Each solution can be described with quantum numbers that describe some aspect of the solution. n Principal quantum number (n) - size and energy of of an orbital. n Has integer values > 0

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27 Quantum numbers n Angular momentum quantum number l. n shape of the orbital. n integer values from 0 to n - 1 n l = 0 is called s n l = 1 is called p n l =2 is called d n l =3 is called f n l =4 is called g

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28 Quantum numbers n Magnetic quantum number (m l ) n integer values between - l and + l n tells direction in each shape. n Electron spin quantum number (m s ) n Can have 2 values. n either +1/2 or -1/2 n Q: What are the possible quantum numbers for the last electron of a sulfur atom? Iron?

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29 Example: Describe the electrons defined by the following quantum numbers: n l m s electron p electron d electron not allowed (l must be < n) not allowed (ml must be between -l and l)

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30 1. Wha t are the shapes of s, p, and d orbitals respectively? 2. How many 1s orbitals are there in an atom? 4p orbitals? 4d orbitals? 3. What is the maximum number of orbitals with: n = 4 l = 1 n = 2 l = 2 n = 3 l = 2 n = 5 l = 1 ml = Which orbitals cannot exist? 2p 3p 4d 3f 6s 2d 5. Write a set of quantum numbers for a 4f orbital.

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31 1. Wha t are the shapes of s, p, and d orbitals respectively? s= spherical p = dumbbell d = cloverleaf 2. How many 1s orbitals are there in an atom? 4p orbitals? 4d orbitals? 1s: 1 4p: 3 4d: 5 3. What is the maximum number of orbitals with: n = 4 l = 1 3 (the 4p orbitals) n = 2 l = 2 none (l must be < n) n = 3 l = 2 5 (the 3d orbitals) n = 5 l = 1 ml = -1 1 (3 q.n. define a unique orbital) 4. Which orbitals cannot exist? 2p 3p 4d 3f 6s 2d 3f and 2d 5. Write a set of quantum numbers for a 4f orbital. n = 4 l = 3 ml = 3, 2, 1, 0, -1, -2, -3

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32 S orbitals (BDVD) orbitals

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33 P orbitals

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34 D orbitals

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35 F orbitals

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36 Pauli Exclusion Principle n In a given atom, no two electrons can have the same set of four quantum numbers ( n, l, m l, m s ). n Therefore, an orbital can hold only two electrons, and they must have opposite spins.

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Polyelectronic Atoms n More than one electron. n Three energy contributions. n The kinetic energy of moving electrons. n The potential energy of the attraction between the nucleus and the electrons. n The potential energy from repulsion of electrons.

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38 Polyelectronic atoms n Electron shielding occurs. This affects the properties of these atoms. n Electrons are attracted to the nucleus. n Electrons are repulsed by other electrons. n Electrons would be bound more tightly if other electrons were not present.

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The Periodic Table n Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870s). n Didnt know much about atom. n Put in columns by similar properties. n Predicted properties of missing elements.

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Aufbau Principle n Aufbau is German for building up. n As the protons are added one by one, the electrons fill up hydrogen-like orbitals. n Fill up in order of energy levels.

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41 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f He with 2 electrons

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42 Details n Valence electrons - the electrons in the outermost energy levels (not d). n Core electrons - the inner electrons. n Hunds Rule - The lowest energy configuration for an atom is the one have the maximum number of of unpaired electrons in the orbital. n C 1s 2 2s 2 2p 2

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43 Details n Elements in the same column have the same electron configuration. n Put in columns because of similar properties. n Similar properties because of electron configuration. n Noble gases have filled energy levels. n Transition metals are filling the d orbitals

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44 Exceptions n Ti = [Ar] 4s 2 3d 2 n V = [Ar] 4s 2 3d 3 n Cr = [Ar] 4s 1 3d 5 n Mn = [Ar] 4s 2 3d 5 n Cu = [Ar] 4s 1 3d 10 n Half filled orbitals. n Scientists arent sure of why it happens

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Periodic Trends in Properties n Ionization Energy - The energy necessary to remove an electron from a gaseous atom. n The ionization energy for a 1s electron from sodium is 1.39 x 10 5 kJ/mol. n The ionization energy for a 3s electron from sodium is 4.95 x 10 2 kJ/mol. n Demonstrates shielding.

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46 Shielding n Electrons on the higher energy levels tend to be farther out. n Have to look through the other electrons to see the nucleus. n They are less effected by the nucleus.

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47 How orbitals differ n The more positive the nucleus, the smaller the orbital. n A sodium 1s orbital is the same shape as a hydrogen 1s orbital, but it is smaller because the electron is more strongly attracted to the nucleus. n The helium 1s is smaller as well. n This provides for better shielding.

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48 Periodic Trends n Highest energy electron removed first. n First ionization energy (I 1 ) is that required to remove the first electron. n Second ionization energy (I 2 ) - the second electron

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49 Trends in ionization energy n for Mg I 1 = 735 kJ/moleI 1 = 735 kJ/mole I 2 = 1445 kJ/moleI 2 = 1445 kJ/mole I 3 = 7730 kJ/moleI 3 = 7730 kJ/mole n The effective nuclear charge increases as you remove electrons. n It takes much more energy to remove a core electron than a valence electron because there is less shielding.

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50 Explain this trend n For Al I 1 = 580 kJ/moleI 1 = 580 kJ/mole I 2 = 1815 kJ/moleI 2 = 1815 kJ/mole I 3 = 2740 kJ/moleI 3 = 2740 kJ/mole I 4 = 11,600 kJ/moleI 4 = 11,600 kJ/mole

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51 Across a Period n Generally from left to right, I 1 increases because there is a greater nuclear charge with the same shielding. n As you go down a group I 1 decreases because electrons are farther away.

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52 It is not that simple n Z eff changes as you go across a period, so will I 1 n Half filled and filled orbitals are harder to remove electrons from. n Heres what it looks like.

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53 First Ionization energy Atomic number

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54 First Ionization energy Atomic number

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55 First Ionization energy Atomic number

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56 The information it hides n Know the special groups n It is the number and type of valence electrons that determine an atoms chemistry. n You can get the electron configuration from it. n Metals lose electrons have the lowest IE n Non metals- gain electrons most negative electron affinities.

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57 Electron Affinity & Atomic Radii n Electron affinity - The energy change associated with the addition of an electron to a gaseous atom. X(g) + e - X - (g) X(g) + e - X - (g) n Affinity tends to increase across a period and decrease as you go down a group. n Atomic Radii: n Decrease going from left to right across a period and increase going down a group.

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The Alkali Metals n Doesnt include hydrogen - it behaves as a non-metal. n Decrease in IE. n Increase in radius. n Decrease in density. n Decrease in melting point. n Behave as reducing agents.

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59 Reducing ability n Lower IE = better reducing agents. n Cs>Rb>K>Na>Li n Works for solids, but not in aqueous solutions. n In solution Li>K>Na n Why? n Its the water - there is an energy change associated with dissolving.

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60 Hydration Energy n It is exothermic. n for Li kJ/mol n for Na kJ/mol n for K kJ/mol n Li is so big because of it has a high charge density, a lot of charge on a small atom. n Li loses its electron more easily because of this in aqueous solutions.

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61 The reaction with water n Na and K react explosively with water. n Li doesnt. Even though the reaction of Li has a more negative H than that of Na and K. Even though the reaction of Li has a more negative H than that of Na and K. n Na and K melt. H does not tell you speed of reaction. H does not tell you speed of reaction. n More in Chapter 12.

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