Ch 6 Atomic Spectra. Masterson & Hurley

Transcription

1 Ch 6 Atomic Spectra Masterson & Hurley

2 1 Joule = 1 kg m 2 s 2

3 Ch 6.1 Light, Photon Energies, & Atomic Spectra

4 What scientists know about light, scientists are able to explain the structure of the atom. Light is radiant energy and is characterized by two variables: 1. Wavelength ( ) distance between two peaks, measured in meters (m) 2. Frequency ( ) number of waves (cycles) per second, measure in a hertz (Hz), which is 1/sec.

5 1 second V 1 = 6 cycles/second = 6 Hz V 2 = 16 cycles/second = 16 Hz

6 1 second V 1 = 6 cycles/second = 6 Hz Wavelength and frequency are related to one another by: = c Speed, where (c) = speed of light = 3.00 x 10 8 m/s Note how wavelength and frequency are inversely related, the bigger the wavelength the smaller the frequency and viceversa. ( = c/ ) V 2 = 16 cycles/second = 16 Hz

7 Red fireworks are seen by the emission of light with wavelengths around 650 nm when strontium salts are burned. What is the frequency? Rearrange: = c

8 = c V = 3.00 x 10 8 m/s 650 nm Notice, the speed of light is in meters and the wavelength in in nm, you need to convert nm to m 650 nm x 1 m = 6.50x10-7 m 1 x 10 9 nm V = 3.00 x 10 8 m/s 6.50 x 10-7 m v = 4.62 x Hz

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10 What is radiant energy? It consists of an electrical and magnetic component. Thus it is called electromagnetic radiation.

11 Different wavelengths of energy affect matter differently

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14 The ionization of rubidium is 403 kj/mol. Do X- rays with a wavelength of 85 nm have sufficient energy to ionize rubidium?

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16 1410 kj/mol > 403 kj/mol. YES, X-rays with a wavelength of 85 nm have sufficient energy to ionize rubidium?

17 A continuous spectrum contains all wavelengths of visible light--note your defraction grating, look through at the projector as a light source. The defraction grating acts a prism.

18 A line spectrum or emission spectrum contains only a few discrete wavelengths. Scientist theorize that the electron is the portion of the atom that absorbs and releases energy emitting electromagnetic energy.

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20 We will use an instrument known as a spectroscope to view the continuous and line spectrums. It is similar to the defraction grating that you are using, but it has a number scale on it, so that we are able to view the wavelength emitted by the atom (element). Each element emits it s own unique line spectrum and people are able to identify what element is absorbing and giving off energy. Note the next slide

21 Line Spectrum of Selected Elements

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24 Notice how each element emits it s own unique line spectrum when an element is absorbing and giving off energy. What is the theory behind this phenomenon? Note the next slide

25 6.2 The Hydrogen Atom

26 German Physicist Max Planck ( ) studied the radiation given off by heated, solid bodies. He postulated that energy can be gained or lost only in whole-number multiples of a quantity. He called these small packets of energy a quantum. If we throw a waded piece of paper across the room, we see the energy of the paper moving through the air as continuous. But, the energy of matter is not continuous, but is transferred in whole quanta--or in a step-like transfer of energy. We are not able to see it with our eye, because the quantum is so tiny.

27 Plank calculated a constant known as Plank s Constant, h = Planck s constant = 6.626x10-34 J s The energy gained or lost only in whole-number multiples of a quantity is equal to (h ). These small packets of energy is called a quantum. The energy of matter is not continuous, but is transferred in whole quanta. Albert Einstein ( ) suggested that radiation can be viewed as a stream of particles called photons--they act as particles and waves. The change in energy for a system: E = nh h = Planck s constant = 6.626x10-34 J s n = is an integer (1, 2, 3, ) = frequency of radiation absorbed or emitted E photon = nh = nhc

28 Ok, this is the idea An electron behaves like light (electromagnetic radiation) An element not absorbing energy has it s electrons in the ground state. When an element absorbs energy, the electron absorbs a set amount of energy (a quanta). Now the electron is in the excited state. When it gets there, then it falls back down to it s original state, emitting a unique line spectrum And it does this over and over again, until there is no more energy applied to the element.

29 When atoms receive energy from a source they become excited, (a source can be electricity or heat). When the electron falls back down to the ground state it releases energy.

30 The excited state of an atom is the atom absorbing energy--the highest energy state. The ground state is the lowest possible energy state of an atom.

31 When an atom absorbs energy it goes from a higher to a lower energy state. When it falls back down to its lower energy state (the ground state) it can release energy in the form of visible light. The more energy it absorbs the greater the frequency of the visible light seen.

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33 Here is an electron moving from a ground state to an excited state, releasing energy a firework!

34 The numbers 1, 2, 3, etc. are integers known as principal quantum numbers The bigger the number, the further away you are from the nucleus and thus an increase in energy.

35 Let s absorb a low energy wavelenght, red at 657 nm Notice the e- is going from: n = 3 --> n = 2 The electron is going from an excited state to a ground state--it has already absorbed energy and is now releasing energy, emitting light in the visible region of 657 nm, our eye sees red

36 Now if it were going from: n = 2 --> n = 3 Then the electron is going from a ground state to an excited state--it is absorbed energy and is NOT emitting light.

37 The Quantum Model for Hydrogen was developed by Niels Bohr ( ). He proposed that the hydrogen s electron moves around the nucleus in certain allowed circular orbits. He calculated the hydrogen atom energy levels atch?v=qi50gbuj48s

38 E = E inner - E outer E = -R H 1-1 n 2 inner - n 2 outer n = an integer (the larger the value the larger the orbit radius) R H = Rhydberg constant = x J

39 Strategy and example What is the energy of a hydrogen electron when n = 3 --> n = 2? Let s see The e- is going from higher quantum energy to lower. The e- is going from an excited to a ground state. The e- is thus emitting energy (not absorbing) Ok, now let s apply Bohr s theory to calculate the energy emitted and if it is within the visible spectrum

40 E = -R H 1-1 n 2 inner - n 2 outer E = (-2.18 x J) 1-1 n 2 inner - n 2 outer E 3 = (-2.18 x J) = x J E 2 = (-2.18 x J) 1 = x J 2 2 Inner orbit - Outer orbit Inner orbit = 2 and outer orbit = 3 (-5.45 x J) - (-2.42 x J) = x J

41 = hc E = (6.63 x10-34 J/s) (3.00x10 8 m/s) = 3.03 x J (get rid of (-) it just means energy released) = 6.56 x10-7 m Wavelength = 656 nm

42 When Bohr s model was applied to other atoms it did not work. Therefore, this applies only to the hydrogen atom. It was later concluded that orbits are neither circular or elliptical, as we shall see later. But his model paved the way for later theories. But this helps show that electrons behave as light particles

43 But can particulate matter exhibit wave properties? Louis de Broglie ( ) verified this relationship using electrons: = h mv m = mass v = velocity

44 Particles have wavelength, exhibited by diffraction patterns. When electromagnetic radiation, such as X rays are directed onto a crystal areas on a photographic plate produces: bright spots: constructive interference waves in phase dark areas: destructive interference waves out of phase

45 In 1927 a beam of electrons were directed at a crystal and a pattern was seen similar to the diffraction of X rays

46 De Broglie helped scientists understand why our eye cannot see the jumps in a quantized amount of energy. This has to do with the mass of the electron Remember Louis de Broglie ( ) verified this relationship using electrons, a relationship between the wave-like nature (wavelength) of the electron and the mass of the electron. = h mv m = mass v = velocity

47 Compare the wavelengths Wavelength of an electron m = 9.11x10-31 kg speed = 1.0x 10 7 m/s = h mv an electron to a ball Wavelength of a ball m = 0.10 kg speed= 35 m/s = h mv e = 6.626x10-34 kg m 2 /s (9.11x10-31 kg)(1.0x 10 7 m/s) b = 6.626x10-34 kg m 2 /s (0.10 kg)(35 m/s) e = 7.27 x m small particles have very long wavelengths!!! b = 1.9 x m Large particles have very short wavelengths!!!

48 To Summarize All matter exhibits both particulate and wave properties. Large pieces of matter, such as baseballs, predominately exhibit particulate properties. Very small bits of matter, such as photons, predominately exhibit wave properties and some particulate properties. Intermediate masses, such as electrons, clearly show particulate and wave properties.

49 Lab The Hydrogen Atom

50 For Procedure A: Emission Spectrum of H: You should see 4 colors: Note, pg 133 of Report Sheet Data Table: Write these down NOW!!!! Color Assignment Violet n = 6 --> n =2 Blue n = 5 --> n = 2 Green n = 4 --> n =2 Red n = 3 --> n = 2

51 For your lab Part A. Emission Spectrum of H Accepted Wavelength (pg 134 of data)

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53 That was a lot again Now you need to complete the Pre-Lab: Summarize the procedure. Answer the pre-lab questions. Attach the data sheets to your lab notebook.

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56 Scientists have taken the first 3-D picture of a molecule

57 6.2 The Quantum Mechanical Model of the Atom

58 Werner Heisenberg ( ), Louis de Broglie, ( ), and Erwin Schrodinger ( ) developed wave mechanics or quantum mechanics. Schrodinger and de Broglie Worked out a model to describe the behavior of an electron, that it behaves as a standing wave. A specific wave function is often called an orbital. The wave function for the lowest energy for the hydrogen atom is the 1s orbital. Remember, the orbital is not circular, so how then is the electron moving?

59 This movie will describe the quantum mechanical model of the atom

60 Heisenberg Orbitals are areas of probability for locating electrons, Heisenberg Uncertainty Principle: The more accurately we know the position of any particle, the less accurately we can know its momentum, and vice-versa. Probability distribution is used to indicate the probability value near a given point in space by using color intensity, (electron density map)

61 So in summary (so far): Electrons are found in orbitals Orbitals are areas of the probability of finding an electron Heisenberg Uncertainty Principle: If we know it s exact position, then we will not know it s exact momentum and vice-versa

62 6.4 Atomic Orbitals; Shapes & Sizes

63 s Orbitals spherical shape = s = sharp Labeling system arose from early spectral line studies n = 2 or greater s orbitals become larger as the value of n increases Nodes are areas of zero probability

64 p Orbitals peanut shaped (having two lobes) = p = principle Are labeled n px, n py, n pz according to the axis at which the lobe lies n = 2 or greater

65 d Orbitals daisy shaped (two p orbitals) = d = diffuse Note the axis at which each lobe lies n = 3 or greater

66 d Orbitals Watch out when filling electrons

67 f Orbitals fancy shaped (highly complex shape) = fundamental n = 4 greater Not involved in bonding in most compounds

68 There is a g orbital, but I could not find any pictures showing a g orbital

69 Orbital Energies All orbitals with the same value of n have the same energy degenerate orbitals Hydrogen only Three p orbitals in the same energy level are degenerate or equal in energy The lowest energy state is called ground state Excited state when the atom absorbs energy and electrons move to higher energy orbitals

70 6.5 Electron Configurations In Atoms 6.6 Orbital Diagrams

71 Before we get into the nitty-gritty of wave mechanics or quantum mechanics let s just figure out how to write an electron configuration, which is the distribution of electrons among the orbitals of an atom. Below are examples of electron configurations

72 First you need to understand that electrons arrange themselves in orbitals and these orbitals take-on various shapes: s-shaped, remember spherical shaped p-shaped, remember peanut shaped d-shaped, remember daisy shaped f-shaped, remember fancy shaped

73 You need to remember that the s-orbital only has one orientation in space, since it is circular. Therefore, it will have one subshell Electrons like to hang-out in groups of two, and each subshell can hold a maximum of two electrons. Note the s-configurations shown below

74 p-orbitals have three orientations in space, x, y, and z. Therefore, it will have three subshells. p-orbitals can hold a maximum of 6 electrons. Note the p-configurations shown below

75 The d-orbital has five orientations in space. Therefore, it will have five subshells d-orbitals can hold a maximum of ten electrons. Note the d- configurations shown below

76 The d-orbital has five orientations in space. Therefore, it will have five subshells d-orbitals can hold a maximum of ten electrons.

77 To help you remember this, you get to use various cheat sheets to help you complete electron configurations. The first is your periodic table. The other will help you add the electrons in the proper order.

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79 Let s view an animation on filling orbitals er/electronconfnew.html

80 Let s do a couple electron configurations together. Complete the e-configuration for magnesium Fill at lowest energy level, fill sublevel one at a time and then go back and pair-up, with opposite spin.

81 Complete the e-configuration for P and F

82 7.8 Electron Spin and the Pauli Principle

83 Pauli Exclusion Principle An orbital can hold only two electrons, and they must have opposite spins

84 Spin Quantum Number (m s ): covered in section /2 and 1/2 the electron has two possible spins

85 7.11 The Aufbau Principle and the Periodic Table

86 Aufbau Principle Aufbau is German for building up. As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogenlike orbitals. Add e- to the lowest energy orbital (1s)

87 Hund s Rule One e- in each separate orbital with parallel spins before pairing e-.

88 Shorthand Notation of e- Configurations

89 6.3 Quantum Numbers

90 Quantum numbers describe various properties of the orbital, there are four quantum numbers: 1. Principle Quantum Number (n): 2. Angular Momentum (l): 3. Magnetic Quantum Number (m l ): 4. Spin Quantum Number (m s ): covered in section 7.8

91 Principle Quantum Number (n): 1. Integral values: 1, 2, 3 2. Indicates probable distance from the nucleus: 1. Higher numbers = greater distance (larger orbital) 2. Greater distance = less tightly bound = higher energy

92 Angular Momentum (l): 1. Integers from 0 to (n-1) 2. Is related to the shape of the atomic orbital Value of l Letter used s p d f g Example, for principal quantum level n = 5, determine the number of allowed subshells (different values of l ), and give the designation of each. If n = 5, then n 1 = 5-1 = 4 values run from 0 to 4 l = 0, l = 1, l = 2, l = 3, l = 4 5s 5p 5d 5f 5g

93 Magnetic Quantum Number (m l ): 1. Integers from l to + l 2. Related to position of orbital in space relative to other orbitals Value of l Letter Used s p d f g # of Sublevels Total Number of Electrons In Sublevel

94 Magnetic Quantum Number (m l ):

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97 Spin Quantum Number (m s ): covered in section /2 and 1/2 the electron has two possible spins

98 Magnetic Properties: Paired and Unpaired Electrons

99 When substances are placed in a magnetic field two possible types of behavior occur: 1. Diamagnetism: If all the electrons in the substance are paired, the substance is weakly repelled by a magnetic field 2. Paramagnetism: If the substance has at least one unpaired electron, it is attracted by the magnetic field. The paramagnetism of unpaired electrons have a much stronger effect than the weak diamagnetism Each electron in an atom acts like a tiny magnet. A pair of electrons with opposite spins cancel each other.

100 The next slide will exhibit liquid oxygen s ability to display paramagnetism

101 6.8 Periodic Trends In The Properties of Atoms

102 Internal Atomic Energies 1. Kinetic energy of moving electrons 2. Potential energy of attraction between nucleus and electrons--nuclear electron attraction 3. Potential energy of repulsion between electrons--electron-electron repulsion

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105 The movie will show how nuclear-electron attraction effects the atom

106 This movie will review if there is a great nuclearelectron attraction then this causes a high electron affinity (a BIG negative number)

107 The Electron Correlation Problem 1. Electron pathways are not known, so electron repulsive forces cannot be calculated exactly 2. We approximate the average repulsions of all other electrons

108 Screening or Shielding (same thing) On past AP Exams Consider the outermost e- and the forces it feels. It is attracted to the nucleus, however it also feels the repulsions caused by the other 10 e-. The net effect is that the e- is not bound to the nucleus if the other e- were not present. Thus the e- is shielded from the nuclear charge by the repulsion of the other e-.

109 Variations in Energy Within the Same Quantum Level Atoms other than hydrogen have variations in energy for orbitals having the same principal quantum number Electrons fill orbitals of the same n value in preferential order E ns < E np < E nd < E nf e- density profiles show that s e- penetrate to the nucleus more than other orbital types Closer to nucleus = lower energy

110 Ionic Radius

111 Ionization Energy

112 How and why do atoms lose electrons? Atoms have a tendency to lose electrons due to ionization energy Metals tend to be more stable when they lose electrons--to obtain an octet

113 The ionization energy is the energy required to remove an electron from an atom or ion Note your table of elements. What is the trend for ionization energy?

114 Ionization Energy What is the charge of Al? Why the large jump in energy needed to remove the 4th electron?

115 How Do Ions Form?

116 Metals and Nonmetal Charges Metals tend to lose electrons (low ionization energy) Easier to remove electrons Larger radii & thus greater electron-electron repulsion Nonmetals tend to gain electrons (high ionization energy) More difficult to remove electrons Smaller radii & thus greater nuclear-electron attraction

117 Ions that lose electrons are called cations (positively (+) charged ions)

118 Ions that gain electrons are called anions (negatively charged ion)

119 Electron Affinity

120 The electron affinity is the energy required to add an electron to an atom

121 What is the periodic trend for electron affinity?

122 Metals and Nonmetal Charges Metals tend to lose electrons (low electron affinity--low affinity for electrons) B/C of larger radii and increased shielding Nonmetals tend to gain electrons (high electron affinity--high affinity for electrons) B/C of smaller radii and increased nuclear-electron attraction

123 Let s Review Metals tend to lose e- Low ionization energy --the less energy is needed to remove electrons Low electron affinity electron affinity because they do not want to acquire electrons Nonmetals tend to gain e- High ionization energy --the more energy needed to remove electrons High electron affinity because they acquire electrons

124 Chapter 7 Summary