At_Struct_1 - 1 CHEMICAL CHEMICAL BONDING BONDING Chemical...

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Unformatted text preview: 1 CHEMICAL CHEMICAL BONDING BONDING Chemical Bonding 2 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? Cocaine 4 ATOMIC STRUCTURE ELECTROMAGNETIC RADIATION Periodic Table & Chemistry Li, 3eLi+ Na, 11eNa+ Mg, 12eMg2+ Al, 13eAl3+ C, 6eCH4 Si, 14eSi4+, SiH4 5 6 Electromagnetic Radiation Electromagnetic Radiation • Most subatomic particles behave as PARTICLES and obey the physics of waves. Page 1 3 7 Electromagnetic Radiation Electromagnetic Radiation wavelength 8 Electromagnetic Radiation Electromagnetic Radiation Visible light Electromagnetic Radiation Electromagnetic Radiation 9 Long wavelength --> small frequency Short wavelength --> high frequency • Waves have a frequency ν • Use the Greek letter “ nu”, , for frequency, ”, and units are “cycles per sec” Amplitude λν=c wavelength • All radiation: • All where c = velocity of light = 3.00 x 10 8 m/sec • Long wavelength --> small frequency • Short wavelength --> high frequency Node Ultaviolet radiation 10 Electromagnetic Radiation Electromagnetic Radiation Red light has λ = 700 nm. Calculate the Red frequency. Quantization of Energy Max Planck (1858-1947) Solved the “ultraviolet catastrophe” 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 = 4.29 x 10 14 sec increasing frequency increasing wavelength 11 12 Quantization of Energy Quantization of Energy Energy of radiation is proportional to frequency E = h•ν E = h•ν -1 An object can gain or lose energy by absorbing or emitting radiant energy in QUANTA. Page 2 h = Planck’s constant = 6.6262 x 10 -34 J•s 13 Quantization of Energy Quantization of Energy 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. Light Light with a short λ ((large ν)) has a large E. Light with a short λ large ν has a large E. Light Energy of Radiation Energy of Radiation Energy of 1.00 mol of photons of red light. Photoelectric Effect Photoelectric Effect 14 PROBLEM: Calculate the energy of 1.00 mol PROBLEM: Calculate the energy of 1.00 mol PROBLEM: mol off photons of red light. o photons of red light. λ = 700. nm λ = 700. nm ν = 4.29 x 1014 ssec-1 ν = 4.29 x 1014 ec-1 17 . Atomic Line Emission Spectra and Niels Bohr E = h•ν 10-34 1014 15 Understand experimental observations if light consists of particles called PHOTONS of particles PHOTONS of discrete energy. 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 minimum E is used. • Number of e - ejected depends on light intensity. 16 Photoelectric Effect Photoelectric Effect 18 Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the sec-1) = (6.63 x 10 J•s)(4.29 x 10 sec ) = 2.85 x 10-19 J per photon E per mol = (2.85 x 10-19 J/ph)(6.02 x 1023 ph/mol) = 171.6 kJ/mol This is in the range of energies that can break bonds. Niels Bohr (1885-1962) Page 3 SHARP LINE EMISSION SPECTRA of excited of atoms. 19 Line Emission Spectra of Excited Atoms 20 Line Emission Spectra of Excited Atoms • Excited atoms emit light of only certain wavelengths • The wavelengths of emitted light depend on the element. High E Short λ Short High ν High 21 The Electric Pickle • Excited atoms can emit light. • Here the solution in a pickle is excited electrically. The Na+ ions in the pickle juice give off light characteristic of that element. Low E Long λ Long Low ν Low Visible lines in H atom spectrum are called the BALMER series. 22 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. + 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! 23 24 Atomic Spectra and Bohr Atomic Spectra Bohr Atomic Spectra and Bohr Atomic Spectra Bohr Bohr said classical view is wrong. Need a new theory — now called QUANTUM or WAVE MECHANICS. or WAVE e- can only exist in certain discrete orbits — called stationary states. stationary e- is restricted to QUANTIZED energy QUANTIZED energy states. Energy of quantized state = - C/n2 Energy of state = - C/n2 where n = quantum no. = 1, 2, 3, 4, .... Page 4 • Only orbits where n = integral no. are permitted. • Radius of allowed orbitals = n2 • (0.0529 nm) • But note — same eqns. come But same from modern wave mechanics approach. • Results can be used to explain atomic spectra. If e-’s are in quantized energy states, then ΔE of states can have only certain values. This explain sharp line spectra. E = -C (1/2 2) E = -C (1/1 2) E = -C (1/22) E = -C (1/12) n=2 n=1 Calculate ΔE for e- “falling” from high energy level (n = 2) to low energy level (n = 1). n=2 ΔE = Efinal - Einitial = -C[(1/1 2) - (1/2)2] ΔE = -(3/4)C n=1 Note that the process is EXOTHERMIC Note EXOTHERMIC 26 Atomic Atomic Spectra Spectra and Bohr and Bohr ENERGY Atomic Atomic Spectra Spectra and Bohr and Bohr Atomic Spectra and Bohr Atomic Spectra Bohr ENERGY 25 E = -C (1/22) E = -C (1/12) n=2 n=1 ΔE = -(3/4)C C has been found from experiment (and is now called R, the Rydberg constant) called Rydberg R (= C) = 1312 kJ/mol or 3.29 x 10 15 cycles/sec so, E of emitted light = (3/4)R = 2.47 x 10 15 sec-1 and λ = c/n = 121.6 nm and 121.6 This is exactly in agreement with experiment! 28 Atomic Line Spectra and Atomic Line Spectra and Niels Bohr Niels Bohr Niels Bohr (1885-1962) 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 So, QUANTUM WAVE MECHANICS Page 5 27 ...
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This note was uploaded on 10/19/2011 for the course CHM 2210 taught by Professor Romanmanetsch during the Summer '08 term at University of South Florida - Tampa.

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