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Chapter 7: Quantum Theory & Atomic Structure

by: Amelia Notetaker

Chapter 7: Quantum Theory & Atomic Structure CH 121

Marketplace > University of Alabama - Huntsville > Chemistry > CH 121 > Chapter 7 Quantum Theory Atomic Structure
Amelia Notetaker
GPA 3.88

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Chapter 7 notes.
GENERAL CHEMISTRY I - 90514 - CH 121 - 02
Pamela D Twigg (P)
Class Notes
Chemistry, twigg, uah, quantum, Theory, atomic, structure, notes
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This 8 page Class Notes was uploaded by Amelia Notetaker on Tuesday December 1, 2015. The Class Notes belongs to CH 121 at University of Alabama - Huntsville taught by Pamela D Twigg (P) in Fall 2015. Since its upload, it has received 20 views. For similar materials see GENERAL CHEMISTRY I - 90514 - CH 121 - 02 in Chemistry at University of Alabama - Huntsville.


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Date Created: 12/01/15
Chapter 7 Lecture Notes  Chemical bonding o Electron orbitals  Electron magnetic radiation o Most subatomic particles behave as particles and obey physics of waves  EM radiation curves o Wavelength (lambda) - distance from peak to peak o Frequency (nu) - number of cycles per unit time o Amplitude- height from starting point to wave peak  EM radiation o Consists of an oscillating electric and magnetic field o Waves have a frequency o Use nu for frequency = lambda x nu = c  Wavelength (frequency): constant, speed of light  C= velocity of light: 3.00 x 10^8 m/s o Long wavelength --> smaller frequency o Short wavelength --> higher frequency  EM spectrum o Long wavelength --> low frequency, low energy o Short wavelength --> high frequency, high energy o Red light lambda = 700 nm --> calculate frequency  Nano = 10^-9 m  Lambda x nu = c  700nm 10^-9 m 1 nm  = 7.00 x 10^-7 m  Nu = c/lambda  3.00x10^8 m/s 7.00 x 10^- 7 m  = 4.29 x 10^4 sec^-1  Quantization of energy o Wave model of light can explain some behavior, but not all (early 1900's) o Heated objects emit light (blackbody radiation) o Emission of electrons from metal surfaces on which light shines (photoelectric effect) o Emission of light from electronically excited gas atoms (emission spectra) o Max Planck solved by quantization - theory that energy is absorbed or released by atoms in discrete "chunks" o Only certain wavelength allowed o An object can gain or lose energy in quanta o Energy of red is proportional to frequency  E = h x v  H= 6.626 x 10^-34 J(s) o Matter is allowed to emit and absorb energy in only whole number multiples of hu (3 hu = 3 quanta) o Light with large lambda (small nu) has small E o Light with short lambda (large nu) has high E  Photoelectric effect o Demonstrates the particle nature of light o There is a minimum frequency of light until electrical current occurs  Amount of current was proportional with intensity of light shown on it o Classical theory said that E if ejected e- should increase with rise is light intensity, it was not observed o No e- observed until light of a minimum E was used o Light consists of particles called photons of discrete energy o Calculate E of 100 mol of photons of red light lambda = 700 nm nu = 4.29 x 10^14 sec^-1  Lambda = (6.626 x 10^-34 J(s)) (4.29 x 10^14 s^-1)  = 2.85 x 10^-19 J per photon  2.85 x 10^- 6.022 x 10^23 19 J photons/mol  = 171.6 kJ/mol  Enough energy to break bonds  Excited gases and atomic structure o Electricity off --> colorless o Electricity on --> excited electrons = color  Atomic emission spectra and Niels Bohr o Bohr's greatest contribution was building a simple model of the atom  Based off of an understanding the sharp emission spectra  Line emission spectra and excited atoms o Excited atoms emit light of only certain wavelengths o Wavelength of emitted light depend on element  Visible series in H atom spectrum are called Balmer series  Rydberg equation o Balmer and rydberg worked on a math equation to calculated the wavelengths of all spectrial lines of H+  1/lambda = (R_h)((1/n_1)^2 - (1/n_2)^2)  R_h is constant = 1.096676 x 10^7 m^-1  N_1 and N_2 are positive integers, n_2 > n_1  Atomic spectra and Bohr o Early 20th century was that e- traveled about the nucleus in orbit  Any orbit should be possible and so is any energy  But changed particles moving in electric field should emit energy and orbit will eventually decay  End result is destruction o Bohr: classical view is wrong  Quantum or wave mechanics o E- can only exist in certain discrete locations (stationary states) o E- restricted to quantized energy states  Fixed amount of E  E_n = -R_h (h x c/n^2) = (-2.18 x 10^-18 J) (1/n^2) o E of state = -c/n^2  N = quantum number o Bohr showed that energy possessed by a single e- in one nth orbit or energy level of the H atom was given o R_h = ryberg's constant, h = planck's constant, c = speed of light o Only orbit where n = integral A are permitted o Radius allowed orbitals is proportional to n^2 o If e-'s are in quantized energy states, the delta E of states can have only certain values: explain sharpness of the line spectra o Calculate delta E = E final = E initial = -C [(1/n_f^2) - (1/n_i^2)]  N=2 and to n=1  Delta E = -C [(1/1 ^2) - (1/2 ^2)]  =-3/4 C  EXOTHERMIC PROCESS  Delta E= -3/4 C C= 2.18 x 10^-18 J --> E of emitted light  -3/4(2.18 x 10^-18 J)  -1.64 x 10^-18 J  E = h x nu --> 1.64 x 10^-18 J = (6.626 x 10^-34 J(s)) nu  Nu = 2.47 x 10^15 s^-1  Lambda = c/nu --> 1.12 x 10^-7 m x c = 121 nm o Bohr's theory was accomplished  Noble prize 1922  Problems with theory  Only successful for H  Introduced quantum idea artifically  Quantum/wave mechanics o L.De Broglie (1924) proposed that all moving objects have wave properties o For light: E = mc^2, E= h(nu) = hc/lambda  Therefore, mc^2 = hc/lambda --> mc = h/lambda o More generally for particles with mass (m) moving at velocity (v)  Mass x velocity = h/lambda  Or lambda = h/mv o Wave particle duality: an e- has properties of both o E Schroedinger applied idea of e- behavior as a wave to the problem of e- in atoms --> develops wave equation o Solution gives in expression called wave function (Y) o Each describes an allowed energy state of e-  Quantization into naturally  Uncertainty principle o W Heisenberg: problem with defining nature of e- in atoms o Can't simultaneously define the position and momentum of e- o Define e-'s energy exactly but accept the limitation  Wave Function (Y) o Each Y corresponds to orbital --> region of space within an e- is forced o Doesn't describe exact location --> proportional  Quantum numbers and atomic orbitals o An allowed electron energy state (orbital) is defined by 3 quantum numbers o The principal quantum number (n) is a positive integer  The value of n indicates the relative size of the orbital and therefore its relative distance from the nucleus o The angular momentum quantum number (l) is an integer from 0 to (n- 1)  The value of l indicates the shape of the orbital o The magnetic quantum number (m_1) is an integer with values from -l to +l  The value of m_1, indicates the spatial orientation of the orbital  Allowed values of quantum numbers o N = the number of subshells in a shell ( l=0 to n-1)  N=1 l=0  N=2 l= 0,1  N-3 l=0,1,2 o 2l+1 = the number of orbitals in a given subshell  = the number of values of m_1 (m_1 = -l,..0,…+l)  L=0, m_1=0  L=1, m_1= -1,0,+1  L=2, m_1= -2,-1,0,+1,+2 o N^2 = the total number of orbitals in a shell (all orbitals in all subshells) o Name, symbol (property) Allowed values Principal, n (size, energy) Positive integer (1,2,3…) Angular momentum, l 0 to n-1 (shape) Magnetic, m_1 -l…0…+l (orientation)  Subshells = different atomic orbital types o Orbital names come from descriptions of line spectra: s= "sharp" p= "principal", d= "diffuse", f= "fundamental"  Naming subshells (sublevels) o L = 0 is an s subshell o L = 1 is a p subshell o L = 2 is a d subshell o L = 3 is a f subshell o L = 4 is a g subshell  Subshells = different atomic orbital types o Sublevel "name" = n value + letter designation  3s or 2p  Shells and subshells o When n= 1, then l o M_l has a single value --> 1 orbital  This subshell is labeled 1s o Every shell has 1 orbital labeled s, and it is a spherical in shape  S orbitals o All s orbitals are spherical in shape  P orbitals o When n=2 then l=0 and 1 o Therefore, in n=2 shell there are 2 types pf orbitals - 2 subshells o For l=0, m_l=0  This is a subshell o For l= 1, m_1 = -1, 0, +1  This is a p subshell with 3 orbitals o When l=1, there is a single planar node through the nucleus o The three orbitals lie 90 degrees apart in space  D orbitals o When n=3, what are the values of l?  L = 0, 1, 2  So there are 3 subshells in the shell o For l=0, m_l=0  S subshell with single orbital o For l=1, m_l=-1, 0, +1  P subshell with 3 orbitals o For l=2, m_l= -1, -1, 0 ,1, 2  D subshell with 5 orbitals o S orbitals have mo planar node (l=0) and so are spherical o P orbitals have l=1, and have 1 planar node  Dumbbell shaped o This mean d orbitals (with l =2) have 2 planar nodes  F orbitals o When n=4, l=0, 1, 2, 3 so there are 4 subshells in the shell o For l=0, m_l=0  S subshell with single orbital o L=1, m_l= -1, 0, 1  P subshell with 3 orbitals o L=2, m_l= -2, -1, 0, 1, 2  D subshell with 5 orbitals o L=3 m_l= -3, -2, -1, 0, 1, 2, 3  F subshell with 7 orbitals o The F_xyz orbital, one of seven f orbitals


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