At_Struct_1 - 1 CHEMICAL CHEMICAL BONDING BONDING Cocaine...

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Unformatted text preview: 1 CHEMICAL CHEMICAL BONDING BONDING Cocaine Chemical Bonding How is a molecule or polyatomic ion held together? Why are atoms distributed at strange angles? Why are molecules not flat? Can we predict the structure? How is structure related to chemical and physical properties? How is all this connected with the periodic table? 2 Periodic Table & Chemistry Li, 3eLi, Li+ Mg, 12eMg2+ Al, 13eAl3+ C, 6eCH4 3 Na, 11eNa, Na+ Si, 14eSi, Si4+, SiH4 4 ATOMIC STRUCTURE ELECTROMAGNETIC RADIATION 5 6 Electromagnetic Radiation Electromagnetic Radiation • Most subatomic particles behave as PARTICLES and obey the physics of waves. Page 1 7 8 9 Electromagnetic Radiation Electromagnetic Radiation wavelength Electromagnetic Radiation Electromagnetic Radiation • Waves have a frequency • Use the Greek letter “nu”, , for frequency, “ nu”, and units are “cycles per sec” • All radiation: • where c = velocity of light = 3.00 x 10 8 m/sec m/sec • Long wavelength --> small frequency • Short wavelength --> high frequency Electromagnetic Radiation Electromagnetic Radiation Long wavelength --> small frequency Short wavelength --> high frequency Visible light ν Amplitude λν=c wavelength Node Ultaviolet radiation increasing frequency increasing wavelength 10 Electromagnetic Radiation Electromagnetic Radiation Red light has λ = 700 nm. Calculate the 700 nm. frequency. Quantization of Energy Max Planck (1858-1947) Solved the “ultraviolet catastrophe” 11 12 Quantization of Energy Quantization of Energy Energy of radiation is proportional to frequency 1 x 10 -9 m 700 nm • = 7.00 x 10-7 m 1 nm Freq = 3.00 x 10 8 m/s 7.00 x 10 -7 m E = h•ν E = h•ν h = Planck’s constant = 6.6262 x 10 -34 J•s J•s An object can gain or lose energy by absorbing or emitting radiant energy in QUANTA. = 4.29 x 10 14 sec -1 Page 2 13 14 15 Quantization of Energy Quantization of Energy Photoelectric Effect Photoelectric Effect A. Einstein (1879-1955) • Experiment demonstrates the particle nature of light. (Figure 7.6) • Classical theory said that E of ejected electron should increase with increase in light intensity—not observed! • No e- observed until light of a certain observed minimum E is used. • Number of e - ejected depends on light ejected intensity. Photoelectric Effect Photoelectric Effect Understand experimental observations if light consists of particles called PHOTONS of PHOTONS of discrete energy. PROBLEM: Calculate the energy of 1.00 mol PROBLEM: Calculate the energy of 1.00 mol of photons of red light. of photons of red light. λ = 700. nm λ = 700. nm 700. ν = 4.29 x 1014 sec-1 ν = 4.29 x 1014 sec -1 4.29 sec E = h•ν E = h•ν Light with large λ (small ν)) has a small E. Light with large λ ((small ν has a small E. small Light with a short λ (large ν)) has a large E. Light with a short λ ((large ν has a large E. large 16 17 Energy of Radiation Energy of Radiation Energy of 1.00 mol of photons of red light. . Atomic Line Emission Spectra and Niels Bohr 18 E = h•ν = (6.63 x 10 J•s)(4.29 x 10 sec ) (6.63 J•s)(4.29 = 2.85 x 10-19 J per photon per E per mol = (2.85 x 10-19 J/ph)(6.02 x 1023 ph/mol) J/ph)(6.02 ph/mol) = 171.6 kJ/mol kJ/mol This is in the range of energies that can break bonds. 10-34 1014 sec-1) Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the SHARP LINE EMISSION SPECTRA of excited of Niels Bohr (1885-1962) atoms. Page 3 19 20 21 Line Emission Spectra of Excited Atoms • Excited atoms emit light of only certain wavelengths • The wavelengths of emitted light depend on the element. Line Emission Spectra of Excited Atoms High E Short λ High ν Low E Long λ Low ν The Electric Pickle • Excited atoms can emit light. • Here the solution in a pickle is excited electrically. The Na+ ions in the pickle ions juice give off light characteristic of that element. Visible lines in H atom spectrum are called the BALMER series. 22 23 24 Atomic Spectra and Bohr Atomic Spectra Bohr One view of atomic structure in early 20th century was that an electron (e-) traveled about the nucleus in an orbit. Atomic Spectra and Bohr Atomic Spectra Bohr Bohr said classical view is wrong. Need a new theory — now called QUANTUM or WAVE MECHANICS. MECHANICS. e- can only exist in certain discrete orbits — called stationary states. stationary e- is restricted to QUANTIZED energy states. Atomic Spectra and Bohr Atomic Spectra Bohr Energy of quantized state = - C/n2 • Only orbits where n = integral no. are permitted. • Radius of allowed orbitals = n2 • (0.0529 nm) (0.0529 nm) • But note — same eqns. come eqns. from modern wave mechanics approach. • Results can be used to explain atomic spectra. + 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! Energy of state = - C/n2 where n = quantum no. = 1, 2, 3, 4, .... Page 4 25 If e-’s are in quantized energy states, then ∆E of states can have only certain values. This explain sharp line spectra. E = -C (1/2 2) n=2 ENERGY ENERGY Atomic Spectra and Bohr Atomic Spectra Bohr Atomic Atomic Spectra Spectra and Bohr and Bohr E = -C (1/22) n=2 26 E = -C (1/12) n=1 Atomic Atomic Spectra Spectra and Bohr and Bohr E = -C (1/22) n=2 27 E = -C (1/12) n=1 Calculate ∆E for e- “falling” from high energy level (n = 2) to low energy level (n = 1). ∆E = Efinal - Einitial = -C[(1/1 2) - (1/2)2] -C[(1/1 E = -C (1/1 2) n=1 ∆E = -(3/4)C Note that the process is EXOTHERMIC EXOTHERMIC ∆E = -(3/4)C C has been found from experiment (and is now called R, the Rydberg constant) R (= C) = 1312 kJ/mol or 3.29 x 10 15 cycles/sec kJ/mol cycles/sec so, E of emitted light = (3/4)R = 2.47 x 10 15 sec-1 sec and λ = c/n = 121.6 nm c/n 121.6 This is exactly in agreement with experiment! 28 Atomic Line Spectra and Atomic Line Spectra and Niels Bohr Niels Bohr Bohr’s theory was a great accomplishment. Rec’d Nobel Prize, 1922 Problems with theory — • theory only successful for H. • introduced quantum idea artificially. • So, we go on to QUANTUM or WAVE MECHANICS Niels Bohr (1885-1962) Page 5 ...
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