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General Chemistry 1314 Chapter 7

by: Morgan Walker

General Chemistry 1314 Chapter 7 Chem 1314

Marketplace > Oklahoma State University > Chemistry > Chem 1314 > General Chemistry 1314 Chapter 7
Morgan Walker
OK State
GPA 3.2

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These notes cover Quantum Mechanics from Chapter 7.
General Chemistry
Dr. Jimmie Weaver
Class Notes
25 ?




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This 7 page Class Notes was uploaded by Morgan Walker on Saturday March 26, 2016. The Class Notes belongs to Chem 1314 at Oklahoma State University taught by Dr. Jimmie Weaver in Winter 2016. Since its upload, it has received 20 views. For similar materials see General Chemistry in Chemistry at Oklahoma State University.


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Date Created: 03/26/16
Chapter 7 Quantum and Mechanical Model of the Atom Quantum mechanics forms the foundation of chemistry. It explains the periodic table, behavior of elements and the behavior of chemical bonds. Behavior of the very small  Electrons, their behavior determines much of an atoms behavior  Directly observing electrons changes their behavior o Shining light on them changes their behavior. Quantum Mechanical Model Theory  Quantum-mechanical model- explains the manner in which electrons exist and behave in atoms  it helps us understand and predict the properties of atoms Nature of Light  Light- form of electromagnetic radiation o Perpendicular oscillating waves, one for electric and one for magnetic  Electric field- area where an electrically charged particles feels force  Magnetic field- area where a magnetic particle feels force  Electromagnetic waves move at the same speed o 3 X 10 m/s (speed of light) Waves  Amplitude- wave height o Distance from node to crest or from node to trough o Measure of intensity, Larger the amplitude the brighter the light and vice versa  Wavelength (λ)- distance covered by the wave o From crest to crest or trough to trough  Frequency (v)- number of waves that pass a point in a given period of time o Number of waves = number of cycles o Measured in Hertz (Hz) is defined as one cycle o Frequency is directly proportional to wavelength  When given wavelength one can find frequency or vice versa, speed of light (c) is constant o v= c/ λ Color  determined by wavelength (red-long violet-short)  white light is all the colors of visible light o it is a spectrum (visible light spectrum) o Red, Orange, Yellow, Green, Blue, Indigo, Violet  Objects absorb some of the wavelengths and reflect others o Red apple absorbs all wavelengths BUT RED  Different wavelength, different color  Different amplitude, different brightness o High amplitude= bright , low amplitude= dim Electromagnetic Spectrum  All wavelengths of light  Shorter wavelength, higher frequency, more energy  Radio wave- low frequency, long wavelength, low energy  Gamma wave- high frequency, short wavelength, high energy Interference  Interaction between waves  Constructive- when two waves of equal amplitude and are in phase interact and form a bigger wave  Destructive- when two waves are out of phase interact and cancel each other out Diffraction  When a wave comes across an obstacle that is comparable to its size in wavelength it bends (diffracts) around it Photo Electron Effect  If you shine a light on a metal it will emit an electrode  Classic wave theory  transfer of energy from the light to the metal, which results in a release of an electron o If light is brighter it will emit more electrodes  Threshold frequency- a minimum frequency is needed before electrons would be emitted regardless of the intensity  Albert Einstein proposed that light energy must come in packets (also called photons or quantums)  The amount of energy depends on frequency according to this equation o E=hv -34  h is planks constant h= 6.626x10 J * s Binding Effect  1 photon at the threshold frequency gives just enough energy for it to escape the atom o Excess energy becomes kinetic energy of the electron Spectra  When atoms or molecules absorb energy the energy is often released as light  Emission Spectrum- when a ray of light is passed through a prism it emits a pattern of particular wavelengths of light o A pinkish purple light is shown through a prism and the resulting rays are red, blue and violet. These colors show up on a photographic film. Ryndberg Analysis  Analyzed the Hydrogen spectrum and found that it could be described with an equation that involves an inverse square of intervals o 1/λ = R( 1/m – 1/n ) Nuclear Model  Nucleus- the atoms dense center o Makes up essentially the entire mass of the atom o Positively charged o Positive balances out the negative o Electron cloud around the nucleus  Electrons are charged particles  Moving particles give off energy, therefore electrons should give off energy Bohr Model  Energy of the atom was quantized o Amount of energy in the atom is related to the electrons position in the atom o Quantized- an atom can only have specific amounts of energy  Electrons travel in orbits that are a fixed distance from the nucleus o Stationary states o Energy of electron is proportional to the distance the orbital is from the nucleus  Electrons emit radiation when they jump from a high energy orbit to low energy o Emitted radiation (photon of light) o Distance between the orbits determines the energy of the photon produced Wave behavior of Electron  Proposed particles can have wavelike characters  Wavelength of a particle was inversely proportional to its momentum Electron Diffraction  Electrons behave like waves since they show a diffraction property instead of going in a straight path to a specific spot, like particles would do Complimentary Properties of Electrons  Wave nature- interference pattern o Ripples in a lake  Particle nature- position, which slit it is going through o Ray of sun shining through a window  Electron has both natures, they are inversely related Uncertainty Principle  uncertainty in position and speed of a particle are inversely proportional to its mass  the more we know about one the less we know about another o the more we know about velocity the less we know about mass  Heisenberg’s Principal- Δx * mΔv ≥ h/4π Determinacy and indeterminacy  Particles move in a path determined by the particles velocity, position and forces acting on it  We do not know the position and velocity of electrons therefore we cannot predict its path Electron Energy  Electrons energy and position are complementary  Many properties of the atom come from the electrons energy Schrödinger’s Equation  Allows us to calculate the probability of finding an electron with a particular amount of energy at a particular location in an atom  Psi (ѱ)- wave function, it describes the location of the electron o Ѱ represents an orbital, a position probability map of an electron  Calculations show the size, shape and orientation in space of an orbital are determined by three integers in the wave function o n- the principal quantum number, distance from nucleus o l-the angular quantum number o m-lmagnetic quantum number Principal Quantum number (n)  characterizes the energy of an electron in a particular orbital  n can be any integer ≥1  larger the n greater the energy  larger the n larger the orbital Angular Quantum Number  determines the shape of the orbital  l can be integer values from 0 to (n-1)  l=0 s orbital- circular shape  l=1 p orbital- 2 balloons tied at the ends  l=2 d orbital- four balloons tied at the ends  l=3 f orbital- eight balloons tied at the ends Magnetic Quantum number  an integer that specifies the orientation in space of an orbital  - lto + l o (-3,-2,-1,0,1,2,3)  each set of n, l and mldescribes one orbital  orbitals with the same n value are in the same principal energy level  orbitals with the same value of mnlre said to be in the same sublevel  the number of sublevels within a level= n  the number of orbitals = 2l Quantum Mechanical Explanation of Atomic Spectra  each wavelength in the spectrum of an atom corresponds to an electron transition between orbitals  when excited electron goes from a low energy level (orbital) to high  when it relaxes, electrons transitions from high energy level (orbital) to a lower one o a photon of light is released and its energy equals the energy difference between the orbitals Electron Transition  to go from low to high it must gain the correct amount of energy to transfer  in high state they are unstable and will drop down Predicting the Hydrogen Spectrum  wavelengths of the emission spectrum can be predicted by calculate the difference in energy between any 2 states  electron in an energy state n there are n-1 energy states it can transition to  both Bohr and the quantum mechanical models can predict very accurately  energy in the photon releases is equal to the difference in energy of the orbital the electron is jumping to o subtract energy of initial state from energy of final state o ΔE electronEfinalΔE initial ΔE photon - Δelectron 2  Ѱ is the probability of finding an electron o Represents total probability of finding and electron particle at a particular point in space o S orbital its max is the nucleus  Node= 0 probability o The neutral point in a wave o v1.0/section_10/d7d6077cda061c03591398212effebdf.jpg Radial distance function  total probability of finding an electron within a thin spherical shell at a distance r from the nucleus  basically what is the chance of finding an electron at a certain radius from the nucleus Shapes l=0 s orbital- circular shape  each principal level has 1 s orbital  nodes equals n-1 o n=3 nodes= 2 l=1 p orbital- 2 balloons tied at the ends  principal energy state above n=1 has 3 p orbitals  m l -1,0,1 l=2 d orbital- four balloons tied at the ends  Each energy state above n=2 has 5 d orbitals  m l -2,-1,0,1,2 l=3 f orbital- eight balloons tied at the ends  principal energy state above n=3 has 7 f orbitals  m l -3,-2,-1,0,1,2,3 Phase of Orbital  wave function can be positive or negative o positive is in phase o negative is out of phase


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