Understanding Periodicity & Atomic Structure: From Mendeleev to Quantum Mechanics, Study notes of Chemistry

Download Understanding Periodicity & Atomic Structure: From Mendeleev to Quantum Mechanics and more Study notes Chemistry in PDF only on Docsity! Chapter 5 Periodicity and Atomic Structure McMurray and Fay M. Todd Tippetts, Ph.D. Periodic Table • Periodic repeated properties based on observation • Now scientific basis is understood Predictions for Unknown Elements TABLE 5.1 A Comparison of Predicted and Observed Properties for Gallium (eka-Aluminum) and Germanium (eka-Silicon) Mendeleev’s Observed Element Property Prediction Property Atomic mass 68 amu 69.72 amu , Density 5.9 gicm® 5.91 g/cm? Gallium Melting point Low 29.8 °C (eka-Aluminum) 9 point . Formula of oxide X,0, Ga,0, Formula of chloride XCl, GaCl, Atomic mass 72 amu 72.61 amu . Density 5.5 gicm? 5.35 g/cm? Germanium Color Dark gra’ Light gra (cka-Silicon) gray gnt gray Formula of oxide xO, GeO, Formula of chloride XCl, GeCl, Table 5-1 Chemistry, 5/e © 2008 Pearson Prentice Hall, Inc. Explaining Periodicity • Mendeleev’s work bases purely on observation • Basis not explained until 1920’s • Explanation came from interaction of matter with electromagnetic radiation, light. • Electromagnetic energy (“light”) is characterized by wavelength, frequency, and amplitude. Wavelength () and Frequency () • Wavelength ()– the distance light travels to complete one cycle • Frequency () – the number of wave cycles in one second • units are cycles per second (cps), Hertz (Hz) • Hz = s-1 Notice: As wavelength decreases the frequency increases 400 nm 800 nm Violet light Infrared radiation (v = 7.50 X 1014 5~1) (v = 3.75 X 10571) Oe What we perceive as different kinds of electromagnetic energy are waves with different wavelengths and frequencies. Figure 5-4 Chemistry, 5/e © 2008 Pearson Prentice Hall, Inc. Wavelength and Electromagnetic Spectrum The familiar visible region accounts for only a small portion near the middle of the spectrum. Ore a rl 4 it Wavelength a) Atom Virus Bacteria Dust Pinhead Fingernails Humans in meters 10—12 10-10 10-8 10-6 10-4 107? 1 | | l | l | | | I | | Gamma rays X rays Ultraviolet Infrared Microwaves Radio waves I T T T T T T I T T T T T 1020 1018 1016 10'4 1012 10'° 10° Frequency (v) in hertz Visible 780 nm 380nm 500 nm 600nm 700 nm 7.8 X1077m 3.8 X10-7m Waves in the X-ray region have a length that is approximately the same as the diameter of an atom (107 '° m). Figure 5-3 Chemistry, 5/e © 2008 Pearson Prentice Hall, Inc. • Substances vary in how they absorb light • Absorbance spectra for chlorophyll shows how violet and red light are absorbed for photosynthesis, but green light is reflected from the leaves. This is the green color of the light that we perceive. Line Spectrum of Hydrogen • Discrete or Line Spectra from excited atoms have only a few energies represent. Other energies are missing Mathematical Prediction of Line Spectra Johannes Rydberg later modified the equation to fit every line in the entire spectrum of hydrogen.  1 = R n2 1 m2 1 - Johann Balmer in 1885 discovered a mathematical relationship for the four visible lines in the atomic line spectra for hydrogen.  1 = R n2 1 22 1 - R (Rydberg Constant) = 1.097 x 10-2 nm-1 Calculate the wavelength for the Balmer series when n=2 and m=4  1875.010097.1 4 1 2 11 12 22        nmXR  nm 2.486 10057.2 1 3  X   1 = R n2 1 m2 1 - R = 1.097 x 10-2 nm-1 Light has Properties of a Particle • Einstein explained the photoelectric effect used light as particles “photons” of light, each having a discrete energy • One particle could dislodge just one electron, no matter how energetic the particle was • The only thing that gave more electrons would be more particles 21 Energy And Light Waves • The energy of a wave is proportional to the wave frequency, E=hν • h= Planck’s constant, 6.626 × 10-34 J•s/photon • Calculate the energy of a photon of uv light (320 nm)and a photon of red light (700 nm)? Why is ultraviolet light more damaging? J/photon6.2X10 ) sec 1038.9 s/photon)(J6.626X10( sec 1038.9 10320 sec/1000.3 19- 14 34 320 14 9 8      X hE X mX mXc    • Red light =700 nm UV lightRed light J/photonX10 48.2 ) sec 1028.4 s/photon)(J6.626X10( sec 1028.4 mX10 700 m/secX10 3.00 19- 14 34 700 14 9 8      X hE Xc    E=6.20 X 10-19 J/photonE=2.84 X 10-19 J/photon De Broglie Wave Equation • 1924 proposes that matter has wavelike properties • For heavy objects, the wavelength is so short that wave-like properties are not observable   m h  Calculate the de Broglie wavelength for tennis ball (m=6.0 X 10-2 kg) and electron (m=9.1 X 10-31 kg) each traveling at 62 m/s. m10 X 8.1 m/s) 62(kg)10 X 0.6( Js10 X 6.63 34- 2- 34   tb   m h  nm 12,000m10 X 2.1 m/s) 62(kg)10 X 1.9( Js10 X 6.63 5- 31- 34    e 1 atom diameter = 10-10 m Infrared light Electron Microscope • Light microscopes cannot visualize anything smaller than shortest visible wave (400 nm) • Electrons traveling at high speed have wavelength around 0.001 nm, so they can be used to visualize much smaller objects First electron microscope, Berlin, 1933 Pollen granules Heisenberg Uncertainty Principle • We do not know the detailed pathway of an electron. • 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. Δx = uncertainty in a particle’s position Δ(m) = uncertainty in a particle’s momentum h = Planck’s constant  m 4      h x • Calculate the uncertainly in position of an electron in a hydrogen atom • mass = 9.11X10-31 kg • velocity = 2.2X106 m/s ± 0.2 X106 m/s  m 4      h x )(4  m h x   pm 300103 )102.0)(1011.9()1416.3(4 10626.6 10 631 34 2       mXx XkgX X x s m s mkg Wave Functions and Quantum Numbers • A wave function is characterized by three parameters called quantum numbers: n, l, ml. • These three quantum numbers describe an orbital, a three dimensional region of space where we will most likely find the electron Probability of finding electron in a region of space (2) Wave equation Wave function or orbital () solve Wave Functions and Quantum Numbers Magnetic Quantum Number (ml ) • Defines the spatial orientation of the orbital • There are 2l + 1 values of ml and they can have any integral value from -l to +l • If l = 0 then ml = 0 • If l = 1 then ml = -1, 0, or 1 • If l = 2 then ml = -2, -1, 0, 1, or 2 • etc. TABLE 5.2 Allowed Combinations of Quantum Numbers n, /, and m, for the First Four Shells Orbital Number of Number of nil my Notation Orbitals in Subshell Orbitals in Shell 1 °0 0 1s 1 1 2 0 0 2s 1 4 1 -1, 0, +1 2p 3 0 0 3s 1 3 1 -1, 0, +1 3p 3 9 2 —2, -1, 0, +1, +2 3d 5 0 0 4s 1 4 7 -1, 0, +1 4p 3 i. 2 —2,-1, 0, +1, +2 4d 5 3 -3 -2,-1, 0, +1, +2, +3 4f 7 Table 5-2 Chemistry, 5/e © 2008 Pearson Prentice Hall, Inc. Orbital Energies for Multi-electron Atom Within each energy level or shell there are an increasing number of subshells, each having different energies Each subshell can be described by a principal and angular momentum quantum number n=1, l =0 Shapes of the Orbitals • s orbitals • Occur on all energy levels • Always spherical in shape • Only single orientation (ml=0) allowed • 2s has two regions of high probability • 3s has three regions of high probability Nodes • Inherent property of waves • Region where amplitude of wave = 0 • Positive and negative values for wave shown by blue/red colors in orbital figures Shapes of the Orbitals * p orbitals Each p orbital has two lobes of high electron probability separated by a nodal plane passing through the nucleus. Nodal Nodal \ 7 7 < Nodal plane = | ae « The different colors of the lobes represent different algebraic signs, analogous to the different phases of a wave. Figure 5-12 Chemistry, 5/e © 2008 Pearson Prentice Hall, Inc. Quantum Mechanics and Line Spectra • Electrons only have certain orbitals available to them • Can be excited and jump to a higher energy orbital, but not stop in between orbitals • Energy emitted as electrons fall back to ground state – lower energy orbital • Further they fall, the more energy emitted Quantum Mechanics and Line Spectra  1 = R n2 1 m2 1 - n: shell the transition is from (outer-shell) m: shell the transition is to (inner-shell) More complex for multi- electron atoms since all sublevels with an energy level no longer have identical energies • Calculate the energy (kJ/mol) emitted when an electron emitted from n=2 falls to n=1 in a hydrogen atom.   nm 121 1003.875.010097.1 2 1 1 11 1312 22             nmXnmX R     kJ/mol 989 J/mol10 X 9.89 mol 10 X 022.6 m 10 X 121 10 X 3.00 s)J(6.626X10 5 1-23 9 s m8 34-                   AN c hE  50 • Shielding: occurs when core electrons block the valence electrons from experiencing the full attraction of the nucleus • Effective nuclear charge (Zeff): the amount of positive charge “felt” by outer electrons in atoms other than hydrogen • Zeff=Zactual-shielding electrons • Shielding electrons are all those on lower energy levels • Zeff is lower than the atomic number because of shielding Shielding And Effective Nuclear Charge Shielding and Orbitals • Inner core electrons are very effective at shielding outer electrons from positive charges in nucleus • Effectiveness of shielding depends on the orbital the outer electron occupies Electron Configurations • Electron Configuration: A description of which orbitals are occupied by electrons. • Degenerate Orbitals: Orbitals that have the same energy level. For example, the three p orbitals in a given subshell. • Ground-State Electron Configuration: The lowest-energy configuration. • Aufbau Principle (“building up”): A guide for determining the filling order of orbitals. Electron Configuation and Orbital Diagram N: H: He: Li: Orbital-Filling Diagram 1s 1s 2s1s 1s 2p2s Electron Configuration 1s2 2s2 2p3 1s2 2s1 1s2 1s1 Write Electron Configuration and Orbital diagrams for O, F and Ne • Remember the relative energies of different sublevels • Note: 4s is lower energy than 3d, therefore fills first • Pattern is repeated for 5s, 6s etc. Some Anomalous Electron Configurations [Ar] 4s1 3d 5Cr: Cu: [Ar] 4s1 3d 10 Actual Configuration Expected Configuration [Ar] 4s2 3d 4 [Ar] 4s2 3d 9 Electron Configurations and Periodic Table • Atoms within same main group column have similar chemical properties • They also have the same number of valence electrons • Na [Ne] 3s1 • K [Ar] 4s1 • Rb [Kr] 5s1 TABLE 5.3 Valence-Shell Electron Configurations of Main-Group Elements Valence-Shell Electron Group Configuration 1A ns! (1 total) 2A ns? (2 total) 3A ns?np' (3 total) 4A ns?np? (4 total) 5A ns?np> (5 total) 6A ns?np* (6 total) 7A ns?np° (7 total) 8A ns? np® (8 total) Table 5-3 Chemistry, 5/e © 2008 Pearson Prentice Hall, Inc. Use the periodic chart to predict the abbreviated electron configuration for: • Ruthenium (Ru) [Kr] 5s2 4d6 • Strontium (Sr) [Kr] 5s2 • Antimony (Sb) [Kr] 5s2 4d10 5p3 • Polonium (Po) [Xe] 6s2 4f14 5d10 6p4 Electron Configurations 22 23 24 25 26 27 28 29 30 Ti v Cr Mn Fe Co Ni cu in asad? | astad? | astad® | 4s’3d° | as7ad® | 4s"3d7 | 45730 | 4s'3d'° | 4s73d'? 40 4 42 43 44 45 46 47 43 Zr Nb | Mo Tc Ru Rh Pd Ag Cd stad? | Sslad* | Sstad? | SstaP | Ss'ad’ | Sslad* ad? | sstad'® | sstag!? 55 56 57 72 73 74 75 76 77 78 Cs Ba la Hf Ta Ww Re Os Ir Pt bs! 6s? 6575! |6s7a9*5a? |6574f'5d° | 6574 1'*5df* | 6s?4F'*5d°| 657A s'*5cl | 6s74t"*5a" |6s'at 45a los'aF 45a" 87 88 89 104 105 106 107 108 109 WW Fr Ra Ac Rf Db Sg Bh Hs Mt Rg 7s*5#6a" 735 eu" Figure 5-17 Chemistry, 5/e © 2008 Pearson Prentice Hall, Inc. 58 59 60 61 62 63 64 65 66 67 68 69 70 71 Ce Pr Nd | Pm | Sm | Eu | Gd Tb Dy | Ho Er | Tm | Yb lu ostafisd'| 6staP | bs?af | ostaf | ostae | Gstar? |esarsa'| 6s"? | esta? | ostart! | stag? | estat? | ostas'* |os?aritsa! 90 91 92 93 94 95 96 7 98 99 100 101 102 103 Th Pa uU Np Pu Am | Cm Bk Cf Es Fm | Md No ir T#6d* | 75°5f6d' | 7s75P6d' | 75*5f'6d" | 7s'5° | 755 |7s'sfed'| 75751 | 7s*5#'° | 7s?5f"' | 7s'5#? | 7575f° | 75*5#'* |7s75f'*od" Atomic radius (pm) Electron Configurations and Periodic Properties: Atomic Radii Maxima occur for atoms of group 1A elements (Li, Na, K, Rb, Cs, Fr). 300 . ow |S Fr column radius 250 200 150 100 F 50 Minima occur for atoms of group > . 7A elements (F, Cl, Br, I). row rad lus 0 0 10 20 30 40 50 60 70 80 90 100 Atomic number Z., and Atomic Radius 200 180 = a °o Atomic radius (pm) NB ° o 100 Atomic As Zogf increases, the valence-shell radius —_ electrons are attracted more strongly to the nucleus... \ ...and the atomic radius therefore decreases. Zeff 80 10 Figure 5-20 Chemistry, 5/e 11 12 13 14 15 16 17 18 Na Mg Al Si P s cl Ar Atomic number © 2008 Pearson Prentice Halll, Inc. Zeff Which should be the larger atom? Why? Na Cl Which should be the larger atom? Why? Na Cs • Choose the larger atom in each of the following pairs: a) Mg or Ba b) Ta or Pt c) Si or Sn d) Ge or Se Spectra From Common Sources of Visible Light Relative Energy 500 Wavelength (Nanometers) Figure 3 • Output of several commercial CFL • Varying mixture of phosphor will vary composition of light