CHEM 101 Chapter 6
CHEM 101 Chapter 6 Chem 101
Popular in General Chemistry 1
Popular in Chemistry
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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