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MODERN ATOMIC THEORY Chapter 10 ANCIENT GREEKS’ VIEW OF MATTER About 400 B.C. , Aristotle thought all matter was made of four “elements” : • earth • air • fire • water ANCIENT GREEKS’ VIEW OF MATTER At about the same time another Greek philosopher, Democritus, said that matter was made of tiny, indivisible particles called atoms. Atomos is the Greek word for indivisible. Modern View of the Atom Tiny, dense, positively charged nucleus made up of positive protons and neutral neutrons. Negatively charged electron shells enclose the nucleus and contain negative electrons. Atomic Spectra and Bohr One view of atomic structure in early 20th century was that an electron (e-) traveled about the nucleus in an orbit. + Electron orbit 1. Any orbit should be possible and so is any energy. 2. But a charged particle moving in an electric field should emit energy. End result should be destruction! Electromagnetic Radiation Radiant energy that exhibits wave-like behavior and travels through space at the speed of light in a vacuum. The electromagnetic spectrum. Electromagnetic Radiation wavelength Visible light Amplitude wavelength Ultraviolet radiation Node Waves 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 108 m/s. As the wavelength () decreases, the frequency () increases. Wavelength and frequency can be interconverted. = c/ = frequency (s1, Hz, cyc/s, or waves/s ) = wavelength (m) c = speed of light (m/s) Huygens thought light travels as waves, while Newton believed it travels as particles. Photons Photons -- tiny particle of electromagnetic radiation -- a bundle of light energy. Ground state -- electrons are at their lowest energy state in an atom. Excited state -- electrons have absorbed energy by jumping up to a higher energy state in the atom. Larger energy jumps by electrons produce shorter wavelength (more energetic) light. Line Spectra of Excited Atoms High E Short High Low E Long Low Visible lines in H atom spectrum are called the BALMER series. Line Spectra of Excited Atoms Excited atoms emit light of only certain wavelengths The wavelengths of emitted light depend on the element. Atomic Spectrum of Hydrogen Continuous spectrum: Contains all the wavelengths of light. Bright Line (discrete) spectrum: Contains only some of the wavelengths of light. The diagrams above present evidence for discrete energy levels about a nucleus. Electrons can only be found in certain energy levels with certain energies. Atomic Line Spectra and Niels Bohr Niels Bohr (1885-1962) Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the BRIGHT LINE SPECTRA of excited atoms. Bohr’s Model Bohr’s Model was incorrect. Replaced by QUANTUM or WAVE MECHANICS MODEL. e- can only exist in certain discrete orbitals. e- is restricted to QUANTIZED energy states. e- can not be exactly located--location based upon probability. Quantum or Wave Mechanics L. de Broglie (1892-1987) de Broglie (1924) proposed that all moving objects have wave properties. Quantum or Wave Mechanics E. Schrodinger 1887-1961 Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atoms. Failure of the Bohr Model The Bohr Model of the atom paved the way for the Quantum Mechanical Theory, but current theory is in no way derived from the Bohr Model of the atom. The Bohr Model of the Atom was fundamentally incorrect-atoms do not move in circular orbits about the nucleus. 1s Orbital 2s Orbital p Orbitals A p orbital The three p orbitals lie 90o apart in space 2px Orbital 2py Orbital 2pz Orbital 3px Orbital 3dxy Orbital 3dxz Orbital 3dyz Orbital 3dyz Orbital 2 2 3dx - y Orbital Quantum Numbers (QN) 1. Principal QN (n = 1, 2, 3, . . .) - related to size and energy of the orbital. 2. Angular Momentum QN -- l (s, p, d, & f) relates to shape of the orbital. 3. Magnetic QN -- ml (x, y, or z plane) - relates to orientation of the orbital in space relative to other orbitals. 4. Electron Spin QN -- ms (+1/2, 1/2) - relates to the spin states of the electrons-- clockwise or counterclockwise. Electron Arrangement Level Sublevel # Orbitals # electrons 1-7 s 1 2 2-7 p 3 6 3-7 d 5 10 4-7 f 7 14 Energy Levels and Orbitals • n = the number of the energy level. • n2 = the number of orbitals in an energy level. • 2n2 = the number of electrons in an energy level. Pauli Exclusion Principle In a given atom, no two electrons can have the same set of four quantum numbers (n, l, ml, ms). Therefore, an orbital can hold only two electrons, and they must have opposite spins. Aufbau Principle As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen-like orbitals. Electron Filling Order -Aufbau Hund’s Rule The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals. Orbitals halffill before they completely fill. Writing Atomic Electron Configurations Two ways of writing configs. One is called the electron configuration notation. Electron configuration notation for H, atomic number = 1 1 1s value of n Electron-dot symbol is H. no. of electrons value of l Writing Atomic Electron Configurations Two ways of writing configs. Other is called the orbital box notation. ORBITAL BOX NOTATION for He, atomic number = 2 Arrows 2 depict electron spin 1s 1s Quantum numbers are an energy address instead of a positional address. Electron-dot symbol is He: Lithium Group 1A Atomic number = 3 1s22s1 ---> 3 total electrons 3p 3s Li. 2p 2s 1s Beryllium Group 2A Atomic number = 4 3p 3s 1s22s2 ---> 4 total electrons 2p 2s 1s Be: Boron Group 3A Atomic number = 5 3p 1s2 2s2 2p1 ---> 3s 2p 2s 1s 5 total electrons . B: Carbon Group 4A Atomic number = 6 1s2 2s2 2p2 ---> 6 total electrons 3p 3s 2p 2s 1s . .C : Here we see for the first time HUND’S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible. Nitrogen Group 5A Atomic number = 7 1s2 2s2 2p3 ---> 3p 3s 2p 2s 1s . 7 total electrons .N : . Oxygen Group 6A Atomic number = 8 3p 1s2 2s2 2p4 ---> 3s 2p 2s 1s .. 8 total electrons .O : . Fluorine Group 7A Atomic number = 9 3p 1s2 2s2 2p5 ---> 3s 2p 2s 1s 9 total electrons .. :F: . Neon Group 8A Atomic number = 10 1s2 2s2 2p6 ---> 10 total electrons 3p 3s 2p 2s 1s .. Note that we have reached the end of the 2nd period, and the 2nd shell is full! : Ne : .. Electron Dot Filling Order 63 4 7 X 58 2 1 Sodium Group 1A Atomic number = 11 1s2 2s2 2p6 3s1 or “neon core” + 3s1 Na. [Ne] 3s1 (uses rare gas notation) Note that we have begun a new period. Aluminum Group 3A Atomic number = 13 1s2 2s2 2p6 3s2 3p1 3p 3s [Ne] 3s2 3p1 . Al : 2p 2s 1s Phosphorus Group 5A Atomic number = 15 1s2 2s2 2p6 3s2 3p3 [Ne] 3s2 3p3 3p . .P : 3s 2p 2s . 1s Calcium Group 2A Atomic number = 20 6 4s2 1s2 2s2 2p6 3s2 3p Ca : [Ar] 4s2 Valence Electrons The electrons in the outermost principle quantum level of an atom. Inner electrons are called core electrons. Relationship of Electron Configuration and Region of the Periodic Table Green = s block Yellow = p block Lt. Blue = d block Med. Blue = f block Broad Periodic Table Classifications Representative Elements (main group): filling s and p orbitals (Na, Al, Ne, O) Transition Elements: filling d orbitals (Fe, Co, Ni) Lanthanide and Actinide Series (inner transition elements): filling 4f and 5f orbitals (Eu, Am, Es) Transition Metals All transition elements have the configuration (n-1) d. D orbitals are always one behind their period. Chromium Iron Copper Transition Element Configurations 3d orbitals used for Sc - Zn Lanthanides and Actinides Rare earth elements always have the configuration (n-2) f. F orbitals are always two behind their period. Cerium [Xe] 6s2 5d1 4f1 Uranium [Rn] 7s2 6d1 5f3 Lanthanide Element Configurations 4f orbitals used for Ce - Lu and 5f for Th - Lr Properties of Metals • malleable • ductile • good conductors of heat & electricity • tend to lose electrons--oxidation • left of zigzag line on periodic table • most active metal in lower left corner (Fr) Properties of Nonmetals • not malleable or ductile • brittle • nonconductors of heat & electricity • tend to gain electrons -- reduction • right of zigzag line on periodic table • most active nonmetal in upper right corner (F) Properties of Metalloids • properties intermediate between metals and nonmetals • found bordering zigzag line on periodic table • B, Si, Ge, As, Sb, & Te ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY Atomic Size SIZE Size goes UP on going down a group. Because electrons are added further from the nucleus, there is less attraction. Size goes DOWN on going across a period, the addition of protons pulls electrons tighter. Atomic Radii Trends in Atomic Size Radius (pm) 250 K 1st transition series 3rd period 200 Na 2nd period Li 150 Kr 100 Ar Ne 50 He 0 0 5 10 15 20 25 Atomic Number 30 35 40 Sizes of Transition Elements 3d subshell is inside the 4s subshell. 4s electrons feel a more or less constant Z*. Sizes stay about the same and chemistries are similar! Ion Sizes F,64 pm 9e and 9p F- , 136 pm 10 e and 9 p Forming an anion. ANIONS are LARGER than the atoms from which they come. The electron/proton attraction has gone DOWN and so size INCREASES. Ion Sizes + Li,152 pm 3e and 3p Li + , 60 pm 2e and 3 p Forming a cation. . CATIONS are SMALLER than the atoms from which they come. The electron/proton attraction has gone UP and so size DECREASES. Trends in Ion Sizes Ionization Energy IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ ---> Mg+ (g) + eMg+ (g) + 1451 kJ ---> Mg2+ (g) + e- Trends in Ionization Energy 1st Ionization energy (kJ/mol) 2500 He Ne 2000 Ar 1500 Kr 1000 500 0 1 3 H Li 5 7 9 11 Na 13 15 17 19 K 21 23 25 27 29 31 Atomic Number 33 35 Trends in Ionization Energy IE increases across a period because Z* increases. Metals lose electrons more easily than nonmetals. Metals are good reducing agents. Nonmetals lose electrons with difficulty. Trends in Ionization Energy IE decreases down a group Because size increases. Reducing ability generally increases down the periodic table. Electron Affinity A few elements GAIN electrons to form anions. Electron affinity is the energy involved when an anion loses an electron. A-(g) ---> A(g) + e- E.A. = DE Trends in Electron Affinity Affinity for electron increases across a period (EA becomes more positive). Affinity decreases down a group (EA becomes less positive). Atom EA F +328 kJ Cl +349 kJ Br +325 kJ I +295 kJ Trends in Electron Affinity F Cl Br 35 0 30 0 S Si 20 0 Se 15 0 S4 10 0 Ge P S3 Period 50 S2 S1 0 K 1 2 3 4 5 Group 6 7 Ele ctron a ffinity (kJ/mol) C H 25 0 O electronegativity, ionization energy, ionic radii, electron affinity atomic radii ionization energy, electron affinity, & electronegativity Noble gases 02_29 Alkaline 1 earth metals Halogens 1A 1 Alkali metals H 8A 2 13 14 15 16 17 2A 3A 4A 5A 6A 7A 2 He 3 4 5 6 7 8 9 10 Li Be B C N O F Ne 13 14 15 16 17 18 Al Si P S Cl Ar 11 12 Na Mg 4 3 5 6 9 8 7 Transition metals 10 11 12 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 Cs Ba La* Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 104 105 106 107 108 109 110 111 87 88 Fr Ra 89 Ac† Unq Unp Unh Uns Uno Une Uun Uuu *Lanthanides † Actinides ionic & atomic radii 18 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 90 91 92 93 94 95 96 97 98 99 100 101 102 103 Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Increasing Periodic Trends Similarity of Elements Elements are grouped together in vertical columns (Groups) that have similar properties. Alkali Metals -- Li, Na, K, Rb, & Cs Halogens -- F2, Cl2, Br2, & I2 Noble Gases -- He, Ne, Ar, Kr, Xe, & Rn Noble gases 02_29 Alkaline 1 earth metals Halogens 1A 1 Alkali metals H 18 8A 2 13 14 15 16 17 2A 3A 4A 5A 6A 7A 2 He 3 4 5 6 7 8 9 10 Li Be B C N O F Ne 11 12 13 14 15 16 17 18 Na Mg Al Si P S Cl Ar 3 4 5 6 7 8 9 Transition metals 10 11 12 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 Cs Ba La* Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 104 105 106 107 108 109 110 111 87 88 Fr Ra 89 Ac† Unq Unp Unh Uns Uno Une Uun Uuu *Lanthanides † Actinides 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 90 91 92 93 94 95 96 97 98 99 100 101 102 103 Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Periodic Table of the Elements