Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Atomic Structure and Periodicity Chapter 7.

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<ul><li> Slide 1 </li> <li> Slide 2 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 1 Atomic Structure and Periodicity Chapter 7 </li> <li> Slide 3 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 2 Overview Introduce Electromagnetic Radiation and The Nature of Matter. Discuss the atomic spectrum of hydrogen and Bohr model. Describe the quantum mechanical model of the atoms and quantum numbers. Use Aufbau principle to determine the electron configuration of elements. Highlight periodic table trends. </li> <li> Slide 4 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 3 Matter and Energy Matter and Energy were two distinct concepts in the 19 th century. Matter was thought to consist of particles, and had mass and position. Energy in the form of light was thought to be wave-like. </li> <li> Slide 5 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 4 Physical Properties of Waves Wavelength ( ) is the distance between identical points on successive waves. Amplitude is the vertical distance from the midline of a wave to the peak or trough. </li> <li> Slide 6 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 5 Properties of Waves Frequency ( ) is the number of waves that pass through a particular point in 1 second (Hz = 1 cycle/s). The speed (v or c) of the wave = x </li> <li> Slide 7 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 6 Electromagnetic Radiation Electromagnetic radiation travels through space at the speed of light in a vacuum. Electromagnetic radiation is the emission and transmission of energy in the form of electromagnetic waves. </li> <li> Slide 8 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 7 Maxwell (1873), proposed that visible light consists of electromagnetic waves. Speed of light (c) in vacuum = 3.00 x 10 8 m/s All electromagnetic radiation x c 7.1 </li> <li> Slide 9 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 8 Electromagnetic Waves Electromagnetic Waves have 3 primary characteristics: 1.Wavelength: distance between two peaks in a wave. 2.Frequency: number of waves per second that pass a given point in space. 3.Speed: speed of light is 2.9979 10 8 m/s. </li> <li> Slide 10 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 9 Wavelength and frequency can be interconverted. = c/ (neu)= frequency (s 1 ) (lamda) = wavelength (m) c = speed of light (m s 1 ) </li> <li> Slide 11 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 10 Electromagnetic Spectrum </li> <li> Slide 12 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 11 Plancks Constant E = change in energy, in J h = Plancks constant, 6.626 10 34 J s = frequency, in s 1 = wavelength, in m n = integer = 1,2,3 Transfer of energy is quantized, and can only occur in discrete units, called quanta. n n </li> <li> Slide 13 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 12 Diffraction X-Ray Diffraction showed also that light has particle properties. </li> <li> Slide 14 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 13 Figure-7.5: (a) Diffraction occurs when electromagnetic radiation is scattered from a regular array of objects, such as the ions in a crystal of sodium chloride. The large spot in the center is from the main incident beam of X rays. (b) Bright spots in the diffraction pattern result from constructive interference of waves. The waves are in phase; that is, their peaks match. (c) Dark areas result from destructive interference of waves. The waves are out of phase; the peaks of one wave coincide with the troughs of another wave. </li> <li> Slide 15 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 14 Energy and Mass Energy has mass E = mc 2 Einsteins Equation E = energy m = mass c = speed of light </li> <li> Slide 16 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 15 Energy and Mass Radiation in itself is quantized </li> <li> Slide 17 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 16 Wavelength and Mass = wavelength, in m h = Plancks constant, 6.626 10 34 J s = kg m 2 s 1 m = mass, in kg v = speed, in ms 1 de Broglies Equation </li> <li> Slide 18 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 17 x = c = c/ = 3.00 x 10 8 m/s / 6.0 x 10 4 Hz = 5.0 x 10 3 m Radio wave A photon has a frequency of 6.0 x 10 4 Hz. Convert this frequency into wavelength (nm). Does this frequency fall in the visible region? = 5.0 x 10 12 nm </li> <li> Slide 19 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 18 Figure 7.6: (a) A continuous spectrum containing all wavelengths of visible light. (b) The hydrogen line spectrum contains only a few discrete wavelengths. Atomic Spectrum of Hydrogen </li> <li> Slide 20 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 19 Atomic Spectrum of Hydrogen Continuous spectrum: Contains all the wavelengths of light. Line (discrete) spectrum: Contains only some of the wavelengths of light. </li> <li> Slide 21 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 20 1.e - can only have specific (quantized) energy values 2.light is emitted as e - moves from one energy level to a lower energy level Bohrs Model of the Atom (1913) E n = -R H ( ) z2z2 n2n2 n (principal quantum number) = 1,2,3, R H (Rydberg constant) = 2.18 x 10 -18 J 7.3 </li> <li> Slide 22 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 21 E = h </li> <li> Slide 23 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 22 E photon = E = E f - E i E f = -R H ( ) 1 n2n2 f E i = -R H ( ) 1 n2n2 i i f E = R H ( ) 1 n2n2 1 n2n2 n f = 1 n i = 2 n f = 1 n i = 3 n f = 2 n i = 3 </li> <li> Slide 24 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 23 = h/mv = 6.63 x 10 -34 / (2.5 x 10 -3 x 15.6) = 1.7 x 10 -32 m = 1.7 x 10 -23 nm What is the de Broglie wavelength (in nm) associated with a 2.5 g Ping-Pong ball traveling at 15.6 m/s? m in kgh in J su in (m/s) </li> <li> Slide 25 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 24 E photon = 2.18 x 10 -18 J x (1/25 - 1/9) E photon = E = -1.55 x 10 -19 J = 6.63 x 10 -34 (Js) x 3.00 x 10 8 (m/s)/1.55 x 10 -19 J = 1280 nm Calculate the wavelength (in nm) of a photon emitted by a hydrogen atom when its electron drops from the n = 5 state to the n = 3 state. E photon = h x c / = h x c / E photon i f E = R H ( ) 1 n2n2 1 n2n2 E photon = Ignore the (-) sign for and </li> <li> Slide 26 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 25 The Bohr Model Ground State: The lowest possible energy state for an atom (n = 1). Ionization: n f = =&gt; 1/n f 2 = 0 =&gt; E=0 for free electron. Any bound electron has a negative value to this reference state. </li> <li> Slide 27 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 26 E = h = hc P = mv To be well memorized E = mc 2 i f E = R H ( ) 1 n2n2 1 n2n2 E photon = </li> <li> Slide 28 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 27 Schrdinger Wave Equation In 1926 Schrdinger wrote an equation that described both the particle and wave nature of the e - Wave function ( ) describes: 1. energy of e - with a given 2. probability of finding e - in a volume of space Schrdinger's equation can only be solved exactly for the hydrogen atom. Must approximate its solution for multi-electron systems. </li> <li> Slide 29 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 28 Quantum Mechanics Based on the wave properties of the atom = wave function = mathematical operator E = total energy of the atom A specific wave function is often called an orbital. </li> <li> Slide 30 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 29 Schrdinger Wave Equation fn(n, l, m l, m s ) principal quantum number n n = 1, 2, 3, 4, . n=1n=2 n=3 distance of e - from the nucleus </li> <li> Slide 31 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 30 Probability Distribution SQUARE of the wave function: 4 probability of finding an electron at a given position Radial probability distribution is the probability distribution in each spherical shell. </li> <li> Slide 32 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 31 e - density (1s orbital) falls off rapidly as distance from nucleus increases Where 90% of the e - density is found for the 1s orbital </li> <li> Slide 33 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 32 Figure 7.12: (a) Cross section of the hydrogen 1s orbital probability distribution divided into successive thin spherical shells. (b) The radial probability distribution. </li> <li> Slide 34 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 33 Heisenberg Uncertainty Principle x = position mv = momentum h = Plancks constant The more accurately we know a particles position, the less accurately we can know its momentum. </li> <li> Slide 35 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 34 Quantum Numbers (QN) 1.Principal QN (n = 1, 2, 3,...) - related to size and energy of the orbital. 2.Angular Momentum QN (l = 0 to n 1) - relates to shape of the orbital. 3.Magnetic QN (m l = l to l) - relates to orientation of the orbital in space relative to other orbitals. 4.Electron Spin QN (m s = + 1 / 2, 1 / 2 ) - relates to the spin states of the electrons. </li> <li> Slide 36 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 35 = fn(n, l, m l, m s ) angular momentum quantum number l for a given value of n, l = 0, 1, 2, 3, n-1 n = 1, l = 0 n = 2, l = 0 or 1 n = 3, l = 0, 1, or 2 Shape of the volume of space that the e - occupies l = 0 s orbital l = 1 p orbital l = 2 d orbital l = 3 f orbital Schrdinger Wave Equation </li> <li> Slide 37 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 36 m l -lto +l </li> <li> Slide 38 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 37 Each orbital can take a maximum of two electrons and a minimum of Zero electrons. Zero electrons does not mean that the orbital does not exist. </li> <li> Slide 39 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 38 Degenerate </li> <li> Slide 40 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 39 Pauli Exclusion Principle In a given atom, no two electrons can have the same set of four quantum numbers (n, l, m l, m s ). Therefore, an orbital can hold only two electrons, and they must have opposite spins. </li> <li> Slide 41 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 40 Figure 7.13: Two representations of the hydrogen 1s, 2s, and 3s orbitals. </li> <li> Slide 42 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 41 Figure 7.14: Representation of the 2p orbitals. (a) The electron probability distributed for a 2p orbital. (b) The boundary surface representations of all three 2p orbitals. </li> <li> Slide 43 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 42 m l = -1 m l = 0m l = 1 2p Degenerate Orbitals </li> <li> Slide 44 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 43 Figure 7.16: Representation of the 3d orbitals. </li> <li> Slide 45 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 44 m l = -2m l = -1m l = 0m l = 1m l = 2 3d </li> <li> Slide 46 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 45 How many 2p orbitals are there in an atom? 2p n=2 l = 1 If l = 1, then m l = -1, 0, or +1 3 orbitals How many electrons can be placed in the 3d subshell? 3d n=3 l = 2 If l = 2, then m l = -2, -1, 0, +1, or +2 5 orbitals which can hold a total of 10 e - </li> <li> Slide 47 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 46 Figure 7.20: A comparison of the radial probability distributions of the 2s and 2p orbitals. Probability to be in the nucleus Zero probability to be in the nucleus P orbital is more diffuse </li> <li> Slide 48 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 47 Figure 7.21: (a) The radial probability distribution for an electron in a 3s orbital. (b) The radial probability distribution for the 3s, 3p, and 3d orbitals. </li> <li> Slide 49 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 48 Energy of orbitals in a single electron atom Energy only depends on principal quantum number n E n = -R H ( ) 1 n2n2 n=1 n=2 n=3 </li> <li> Slide 50 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 49 Energy of orbitals in a multi-electron atom Energy depends on n and l n=1 l = 0 n=2 l = 0 n=2 l = 1 n=3 l = 0 n=3 l = 1 n=3 l = 2 </li> <li> Slide 51 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 50 Fill up electrons in lowest energy orbitals (Aufbau principle) H 1 electron H 1s 1 He 2 electrons He 1s 2 Li 3 electrons Li 1s 2 2s 1 Be 4 electrons Be 1s 2 2s 2 B 5 electrons B 1s 2 2s 2 2p 1 C 6 electrons ?? </li> <li> Slide 52 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 51 Figure 7.25: The electron configurations in the type of orbital occupied last for the first 18 elements. </li> <li> Slide 53 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 52 C 6 electrons The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hunds rule). C 1s 2 2s 2 2p 2 N 7 electrons N 1s 2 2s 2 2p 3 O 8 electrons O 1s 2 2s 2 2p 4 F 9 electrons F 1s 2 2s 2 2p 5 Ne 10 electrons Ne 1s 2 2s 2 2p 6 </li> <li> Slide 54 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 53 Order of orbitals (filling) in multi-electron atom 1s &lt; 2s &lt; 2p &lt; 3s &lt; 3p &lt; 4s &lt; 3d &lt; 4p &lt; 5s &lt; 4d &lt; 5p &lt; 6s </li> <li> Slide 55 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 54 Figure 7.26: Electron configurations for potassium through krypton. </li> <li> Slide 56 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 55 Figure 7.27: The orbitals being filled for elements in various parts of the periodic table. </li> <li> Slide 57 </li> <li> Copyright2000 by Houghton Mifflin Company. All rights reserved. 56 What is the electron configuration of Mg? Mg 12 electrons 1s &lt; 2s &lt; 2p &lt; 3s &lt; 3p &lt; 4s 1s 2 2s 2 2p 6 3s 2 2 + 2 + 6 + 2 = 12 electrons 7.7 Abbreviated as [Ne]3s 2 [Ne] 1s 2 2s 2 2p 6 What are the possible quantum numbers for the last (outermost) electron in Cl? Cl 17 electrons1s &lt; 2s &lt; 2p &lt; 3s &lt; 3p &lt; 4s...</li></ul>

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