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!Hcomb of benzoic acid (C6H5CO2H) is -3227 kJ/mol. When a 2.442-g sample of benzoic acid is burned in a calorimeter, all of the resulting heat is transfered to 2.5 kg of water (c=4.184 J/g.K). What is the temperature change? Calculate !Hrxn for 2 NOCl N2 + O2 + Cl2 given the following set of reactions N2 + O2 2NO + Cl2 2NO 2NOCl !H = 180.6 kJ !H = -77.2 kJ Combustion of octane C8H18 produces carbon dioxide and water. C8H18 + O2 CO2 + H2O (a) Balance this reaction and calculate !Hocomb for 1 mol of octane, assuming water forms as a gas; !Hof = -208.45 kJ/mol for octane !Hof = -393.5 kJ/mol for carbon dioxide !Hof = -241.8 kJ/mol for H2O (g) (b) How many liters of CO2 (at 25 oC and 1 atm, R = 0.0821 atm.L/mol.K) are formed when 1 kg of octane is burned? How much work is done by the expanding CO2 as 1 kg of octane is burned (again, at 25 oC and 1 atm). (Hint, 1 J = 9.87.10-3 atm.L). What is !E for the reaction? (Hint, the definition of H = E + PV, assume PV is done by carbon dioxide only (ignore water and octane)) Atomic Structure: Quantum Theory nuclear model from Ratherford’s experiments Problem with Classic Structure accelerating charged particle loses energy The Nature of Light ! frequency (") ! amplitude #= c " Electromagnetic Spectrum Diffraction and Interference diffraction interference here, light behaves as a wave Blackbody Radiation However: 0K 1000 K 1500 K Max Planck, 1900 E = nh" n is a positive whole number, quantum number h is a constant " is the frequency of light emitted Photoelectric Effect • threshold frequency • no time lag Einstein, ~1905 Ephoton = h" > 2000 K Photoelectric Effect 7.96. Electric power is typically stated in units of watts (1W = 1J/s). About 95% of the power output of an incandescent bulb is converted to heat and 5% to light. If 10% of that light shines on your chemistry text, how many photons per second shine on the book from a 75-W bulb? (assume a wavelenght of 550 nm) The Nature of Light • electromagnetic radiation travels in waves • at the speed of light (in vacuum, c = "#) • intensity is either amplitude or number of photons/second • energy of light is quantized, E = h" • the energy of atom is also quantized, E = nh" h = 6.626$10-34 J•s Atomic Spectra H He Na sunlight Ba K Atomic Spectra (Neon) Atomic Spectra (H) Rydberg equation 1 =R ! 1 22 _ 1 n2 The Bohr Model of the Hydrogen Atom Niels Bohr proposed these postulates: 1. The H atom has fixed energy levels: stationary states 2. The atom does not radiate energy while in one of those states 3. Change to another state: !E = !Ephoton = h" The Bohr Model of the Hydrogen Atom Niels Bohr proposed these postulates: E = h! ground state: the lowest energy The Bohr Model of the Hydrogen Atom E = h! the charge of the nucleus -18 E = -2.18x10 J Z2 n2 energy levels for the H atom De Broglie: electrons are waves matter is wavelike: #= h mv electrons within atoms behave like waves The Quantum-Mechanical Model Schrödinger equation: d2% d2% d2% 8&2me [E - V(x,y,z)]%(x,y,z)=0 + + + dx2 dy2 dz2 h2 or H% = E% %2 is the probability that the electron is in a certain region of space The Quantum-Mechanical Model Schrodinger equation: H% = E% solutions (energy states) are associated with a given %, called an atomic orbital Hydrogen Atom “point” probability radial probability Hydrogen Atom 32% 74% 99% Atomic Orbitals different energy states are associated with atomic orbitals atomic orbitals are characterized by 3 quantum numbers n (1,2,3, etc.) is the size (level) principal quantum number l (fromangular 0 to n-1) is the shape (sublevel) momentum quantum number ml (from -l to +l) is the orientation magnetic quantum numbber Atomic Orbitals Sample Problem 7.5 Specify possible orbitals (sets of n, l, m) for n=3. Atomic Orbitals n (level) 1,2,3, etc l (sublevel, shape) s (l=0), p (l=1), d (l=3), f(l=4), etc S orbital - m equals 0 Atomic Orbitals (s) 1 s orbital 3 s orbital Atomic Orbitals: s energy increases 1s n=1 l=0 2s n=2 l=0 3s n=3 l=0 4s n=4 l=0 S orbitals means l=0 (darker color indicates highest radial probability %2) Atomic Orbitals: p energy increases 1p n=1 l=1 2p n=2 l=1 3p n=3 l=1 P orbitals means l=1 2p Atomic Orbitals: p 2px 2py n=2 l=1 m = -1 n=2 l=1 m=0 2pz n=2 l=1 m = +1 Atomic Orbitals: d d orbitals: start with n = 3; l = 2 Energy Energy Levels: Hydrogen 3s 3p 2s 2p 1s 3d Energy Energy Levels: Multielectron Systems 3s 3p 2s 2p 3d 1s Practice! Problems 7.49. How many orbitals in an atom can have each of the following designations: (a) 1s (b) 4d (c) 3p (d) 2f (e) n = 3 Do 7.50 Practice! Problems 7.73. Carbon-carbon bonds form the backbone of nearly every organic and biological molecule. The average bond energy of the C-C bond is 347 kJ/mol. Calculate the frequency and wavelength of the least energetic photon that can break this bond. In what region of electromagnetic spectrum is this radiation? Practice! Problems *7.76. A ground state H atom absorbs a photon of wavelength 94.91 nm, and its electron attains a higher energy level. The atom then emits two photons: one of wavelength 1281 nm to reach an intermediate level, and a second to return to the ground state. (a) What level did the electron reach? (b) What intermediate level did the electron reach? (c) What was the wavelength of the second photon emitted? Practice! Problems *7.76. A ground state H atom absorbs a photon of wavelength 94.91 nm, and its electron attains a higher energy level. The atom then emits two photons: one of wavelength 1281 nm to reach an intermediate level, and a second to return to the ground state. (a) What level did the electron reach? (b) What intermediate level did the electron reach? (c) What was the wavelength of the second photon emitted? Bohr and Schrödinger: !E = -2.18x10-18J 1 n 2 final 1 n2initial