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by: Lyna Nguyen

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CHEM 101 Chapter 6 Chem 101

Marketplace > Texas A&M University > Chemistry > Chem 101 > CHEM 101 Chapter 6
Lyna Nguyen
Texas A&M

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textbook + lecture notes
COURSE
General Chemistry 1
PROF.
Dr. Daniel Collins
TYPE
Class Notes
PAGES
8
WORDS
CONCEPTS
Chemistry, Chem, CHEM 101, tamu
KARMA
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This 8 page Class Notes was uploaded by Lyna Nguyen on Friday January 29, 2016. The Class Notes belongs to Chem 101 at Texas A&M University taught by Dr. Daniel Collins in Fall 2015. Since its upload, it has received 19 views. For similar materials see General Chemistry 1 in Chemistry at Texas A&M University.

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Date Created: 01/29/16
10/20/15­10/22/15 Chemistry 101 Chapter 6  Electromagnetic Radiation o James Clerk Maxwell (1831-1879)  Proposed that visible light consists of electromagnetic waves  Electromagnetic radiation emission and transmission of energy in the form of electromagnetic waves  Developed a mathematical theory to describe light and other forms of radiation in terms of oscillating or wave-like, electric an magnetic field o Electromagnetic Radiation: visible light, microwaves, television, and radio signals, x-rays, and other forms of radiation  Can be categorized by its wavelength and frequency  Wavelength (λ): distance between a given point on a wave and the corresponding point in the next cycle of the wave o Between crests or troughs  Amplitude: the vertical distance from the midline of a wave to the peak or trough  Frequency (v): number of waves that pass a given point in some unit of time o Unit: hertz (1/s)=1 cycle/s  Wave length and frequency are related by c= λ*v  For all electromagnetic radiation  Short wavelength, high frequency  Long wavelength, low frequency  Red light: higher wavelength, lower frequency  Blue light: lower wavelength, higher frequency  Red to blue: more harmful, increase in wavelength 1 10/20/15­10/22/15 o Speed of visible light and all other forms of electromagnetic radiation in a vacuum is a constant 8  C=3 x 10 m/s  Consists of oscillating electric and magnetic disturbances  Electrons can interact with fields and interactions allow scientists to probe matter at the atomic and molecular level  X-rays < UV wavelength < visible light < infrared  Wavelength increase, frequency decrease  Frequency increase, energy increase  Quantization Planck, Einstein, Energy, and Photons o Planck’s Equation  Max Planck  Theory: the electromagnetic radiation emitted originated in vibrating atoms (called oscillators) in heated objects o Each oscillator had a fundamental frequency (v) of oscillation and oscillators could only oscillate at this frequency or whole number multiples (nv)  Radiant energy emitted by an object at a certain temp depends on its wavelength  Energy (light) is emitted or absorbed in discrete units (quantum)  E=nhv  Quantization: only certain energies are allowed  H: Planck’s constant=6.63 x 10 -3Js  If oscillator changes from higher to lower, energy is emitted as electromagnetic radiation  ΔE = ΔE higherΔElower= Δnhv  Einstein and the Photoelectric Effect (1905) o Albert Einstein: incorporated Planck’s ideas into the photoelectric effect o Photoelectric effect: electrons are ejected when light strike the surface of a metal, but only if the frequency of the light is high enough  If light with a lower frequency is used, no electrons are ejected regardless of intensity  If frequency is at or above a minimum, critical frequency, increasing the light intensity causes more electrons to be ejected o Photons: packets of energy; massless  Energy of each photon is proportional to the frequency the radiation as defined by Planck’s Equation  “Particle” of light  Distinct packaging of energy  KE e : mass increase, velocity decrease  Work function: depends on how strongly electrons are held in the metal o Given and metal specific o hv=KE+W 2 10/20/15­10/22/15 o Photons striking atoms on a metal surface will cause electrons to be ejected only if photons have high enough energy o Electromagnetic radiation: wave-particle duality  Wave nature and particle nature  Atomic Line Spectra and Particles o If a high voltage is applied to atoms of an element in the gas phase at low pressure, the atoms energy absorb energy and are said to be excited  Can emit light o Line emission spectrum: spectrum obtained by passing a beam of light from excited neon or hydrogen atoms through a prism  Only a few colors are seen  Every element has a unique emission spectrum  Characteristic lines can be used in chemical analysis, both to identify and how much is present  More lines, more transition  Johann Balmer (1825-1898) and Johannes Rydberg (1854-1919) o Mathematical approach o Balmer Equation: used to calculate the wavelength of the red, green and blue lines in the visible emission spectrum of hydrogen  E n-R (n/n )2 -18  Rydberg constant=2.18x10 J o 4 visible lines in the spectrum of hydrogen atoms are known as Balmer series o e can only have specific (quantized) energy levels - o Light is emitted as e waves moves from one energy level to a lower one o Larger n, farther away from nucleus  Bohr Model of the Hydrogen Atom o Early 20 century, Niels Bohr proposed a model for the electronic structure of atoms o Planetary structure for the hydrogen atom where electrons moved in a circular orbit around the nucleus  Had to contradict laws of classical physics o Postulated that there are certain orbits corresponding to particular energy levels  As long as an electron is in one of these energy levels, the system is stable o Introduced quantization into description of electronic structure o Principle quantum number: positive, limitless integer  Defines the energies of the allowed orbits o Energy of an electron in orbit has a negative value because has lower energy than when it is free  0 energy when n=infinite o Ground: an atom with its electrons in the lowest possible levels o States for the H atom with higher energies (n>1) are excited states o Because energy is dependent on 1/n , energy levels are progressively closer as n increases 3 10/20/15­10/22/15 o An electron in the n=1 orbit is closest to the nucleus and has the lowest (most negative) energy  As n increases, distance of an electron from the nucleus increases, and energy becomes higher  Bohr Theory and the Spectra of Excited Atoms o Theory describes electrons as having only specific orbits and energies. o If an electron moves, energy must be absorbed or evolved o To move from n=1 to an excited state, energy must be transferred  Only exact amount will cause transition  Consequence of quantization o Energy is provided from an electric discharge or by heating o Electrons can go back to their original state by releasing energy  Energy of Emission Line 2 o E =fR (1hn ) o Lyman series: ultraviolet region of emission lines  Results from electrons going from n>1 to n=1 o Balmer series: visible light  Results from electrons going from n>2 to n=2  Particle-wave duality: prelude to quantum mechanics o Louis Victor de Brogie (1924) proposed a free electron of mass m moving with a velocity of v should have an associated wavelength  λ=h/mv or 2*pi*r=n*λ  V=velocity of e -  M=mass of e -  M*V=momentum  More velocity, less mass=more like a wave o Linked particle properties of the electron (m & v) with wave property (wavelength) o CJ Davisson and LH Gremer: diffraction (property of waves) was observed when a beam of electrons was directed at a thing sheet of metal foil o Only possible to view wavelike properties of particles with extremely small mass and extremely high velocity  Species with massive mass need extremely large velocity to make it look like a wave  The Modern View of Electronic structure: wave or quantum mechanics o Erwin Schrodinger: comprehensive theory of the behavior of electrons  Quantum mechanics or wave mechanics: model for electrons o Standing wave: only certain matter waves are possible for an electron in an atom  Ex: string plucked o Schrodinger Wave function (ψ) ONLY FOR HYDROGEN  Equation that described both the particle and wave nature of the e -  Describes: - o Energy of e with a given - o Probability of finding e in a volume of space 4 10/20/15­10/22/15  When solved for energy  Only certain wave functions are found to be acceptable and each is associated with an allowed energy value o Energy of an electron is quantized  In 3-D depend on n, l and m (quantum numbers) l o Only certain combinations are possible  Physical significance (Max Born’s interpretation)  Value of wave function at a given point (x, y, z) is the amplitude of the electron matter wave o Has magnitude and a sign (+/-)  Square of wave function is related to the probability of finding an electron in a tiny region of space o Probability density 2  Wave function * volume o Heisenberg’s Uncertainty Principle: If we choose to know the energy of an electron in an atom with only a small uncertainty, the we must accept a correspondingly large uncertainty in its position  Wemer Heisenberg  Another definition: there is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time  Δx*Δ(mv)>= h/4*pi  x: uncertainty o Quantum Numbers and Orbitals  Quantum numbers is a function of ψ (n, l, m, m ) l s  Atomic Orbital: described by wave function  N: the principle quantum number (n=1,2,3,4,5,6,7) (period number)  Primary factor in determining the energy of an orbital o Most important  Defines the size of the orbitals o Greater the n, greater the size  2 or more electrons may have the same n value o Said to in electron shells  Distance of e from the nucleus o Easier to remove the larger the n number o More accessible to steal or share  L: the orbital angular momentum number (0,1,2,3…n-1)  Subshells: orbital groups of a given electron shell  Defines the characteristic shape and type of the column of - space that the e occupies  L=0, s orbital (1A, 2A)  L=1, p orbital (3A-7A)  L=2, d orbital (trench of doom)  L=3, f orbital (Actinide, Lanthanide) 5 10/20/15­10/22/15  M:lthe magnetic quantum number (0, +/- 1, +/-2, +/-3, +/-l)  Related to the orientation of space of the orbitals within a subshell o +/- L  Number of values of m=2l+1 o Specifies number of orbitals in the subshell  Shells and Subshells  N= number of subshells  2l+1=number of orbitals in a subshell  n =number of orbitals in a shell  First Electron shell, n=1  L=0  M=l  Only one subshell exists and only one orbital, 1s  Second Electron shell, n=2  N=2  L= 0 or 1 o 2 subshells in the second shell  2s and 2p  m=l1,0,1 o 3 2p orbitals exist  Third Electron shell, n=3  L=0,1,2  N=3 o 3s and 3p and 3d  m=l2, -1,0,1,2 o 5d orbitals in d subshell  Fourth Electron shell, n=4  N=4  L=0,1,2,3 o 4s, 4p, 4d, and 4f  7 orbitals exist because 7 m vllues  The Shapes of Atomic Orbitals o S orbital  Holds only 2 electrons  Found near the nucleus  Denser around the nucleus: electron cloud  Surface densi2y plot or radial2distribution plot  4*pi*r *wave function v. r 6 10/20/15­10/22/15  For s orbital, equation=0  Spherical in shape  Boundary surface: easier to draw pictures  Where 90% of the e density is found  All s orbitals are spherical in shape  Increase in size as n increases  1s is more compact than 2s which is more than 3s  Important features:  There is not an impenetrable surface which electrons are contained  Probability of finding the electron is not the same throughout the volume enclosed by the surface  “Electron cloud” and “electron distribution” imply that electron is a particle o Basic premise in quantum mechanics is that the electron is treated as a wave, not a particle o P orbital  Only holds 6 electrons  Shape like a weight lifter’s “dumbbell”  Has a nodal surface and spread out density  A surface which the probability of finding the electron is zero  Result of wave function  3 p orbitals in a subshell  Drawn along x, y, z axis o D orbitals  Holds 10 electrons 7 10/20/15­10/22/15  Value of l is equal to number of nodal surfaces slicing through the nucleus  5 d orbitals have 2 nodal surfaces; 4 regions  Look like 4 leaf clovers o F orbitals  Only holds 14 electrons  3 nodal surfaces cause the electron density to lie in up to 8 regions of space o Sum of all orbitals is a sphere  One More Electron Property: Electron Spin o Otto Stern and Waltner Gerlach  Probed the magnetic behavior of atoms by passing a beam of silver atoms in the gas phase through a magnetic field, split electrons half attracted, half repelled  Best interpreted by imagining the electron has a spin and behaves as a tiny magnet that can be attracted or repelled by another magnet  Electron spin is quantized o Electron spin quantum number (m ) s  One orientation is associated with a value of m of s and the other with m os -1/2  Pretty sure of location  Can not have another electron have the same 4 quantum numbers  Must be flipped 8

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