Chemistry 111 Week 5 Atomic Structure
Chemistry 111 Week 5 Atomic Structure CHEM 111 - 02
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This 5 page Class Notes was uploaded by Makayla Richardson on Sunday September 25, 2016. The Class Notes belongs to CHEM 111 - 02 at New Mexico State University taught by Dr. Antonio Lara in Fall 2016. Since its upload, it has received 20 views. For similar materials see General Chemistry I in Chemistry and Biochemistry at New Mexico State University.
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Date Created: 09/25/16
Chapter 3 Sections 14: The Atomic Structure 1) Wavesof Light Electromagnetic Radiation (Radiant Energy): Forms of radiant energy in the electromagnetic spectrum. Electromagnetic Spectrum: Continuous range of radiant energy. Includes gamma rays, X rays, ultraviolet radiation, visible light, infrared radiation, and radio waves. Gamma Rays: The highest energy with the shortest wavelength. Extremely dangerous to the human skin, is very radioactive. XRays: Also high in energy. The photons are absorbed by calcium, which is why they are used in hospitals. Ultraviolet Light: Types A,B, and C. Very energetic photons, found in the 2 atmosphere. O slows down ultraviolet photons. Visible Light: The light and colors that you can see with the naked eye. Infrared Radiation: vibrating molecules, moving electron clouds, less than red light. Microwaves: All depends on the size and electrons of the electron in how they move. Microwaves affect H O m2lecules. Radio Waves: low frequency with flipping molecules. Violet light has the highest energy and shortest wavelength. Red light is the lowest frequency and longest wavelength. James Clerk Maxwell developed the theory that electromagnetic radiation moves as waves of oscillating electric and magnetic fields. Electric waves move up and down. Magnetic waves move side to side. Wavelength ( λ ): The distance from crest to crest of wave. Frequency ( ν ): number of crests of a wave that passes in a specific time per second. per second: s 1 1 Hertz (Hz): SI unit of frequency with units of reciprocal seconds. 1 Hz= 1s = 1 cycle cycle per second (cps) The shorter the wavelength, the higher the frequency. Speed of light (c): how fast photons travel through a transparent medium. c λν=c∨ν= λ 8 value of c=2.998x10 m/s A) Calculating Frequency from Wavelength What is the frequency of yelloworange light ( λ=589nm ) produced by sodium vapor street lights. (c=2.998x10 m/s) I) Convert nanome9rs to meters 1 nm=10 m is the conversion factor Frequencies of visible light are about 10 Hz or 10 s4 1 c Substitute into ν= λ II) Substitute and Solve 8 2.998x10 m/s =5.09x10 s14 −1 10 m 589nm x 1nm Cross out all of the meters and nanometers Divide Round up 14 1 Add the hertz per second 10 s *Remember wavelength and frequency are reciprocal; so one increases and one decreases. 2) Atomic Spectra Fraunhofer Lines: dark lines in an otherwise continuous solar spectrum. Atomic Emissions Spectra: patterns of bright lines produced when an atom vaporizes when excited. Atomic Absorption Spectra: patterns of dark lines produced after radiation passes through atoms. 3) Particles of Light Quantum Theory Kirchkoff studied the continuous emission spectra with heated metals. Blackbody Radiations: sources of radiant energy that absorb all light that strikes them. A) Photons of Energy Max Planck, threw out the laws of classical physics and discovered that elements emitted electromagnetic radiation based on the quantum. Quantum: Smallest quantity from a form of energy Defined by: E=h ν ν = frequency of radiation = the Planck Constant: 6.626x10 J*s h E= hc ν The module has become known as the Quantum Theory: the idea that energy is absorbed and emitted in discrete quantities. Quantized: values restricted to whole numbers of a certain value. Ex: By taking an elevator you can really only stop on the specific floors, not between the floors. But if you take the stairs you can stop between the floors When you walk up the stairs energy is being absorbed. When you walk down the stairs energy is released. Photons: a quantum of electromagnetic radiation. They are the building blocks of electromagnetic radiation. I) Calculating Energy of a Photon What is the energy of a photon of a red light that has a wavelength of the 656 electromagnetic spectrum? a) Calculate energy of a photon starting with wavelength: E=hv h=6.626x10 J*s c=2.998x10 m/s b) Convert nanometers to meters so that distance cancels out −34 2.998x10 m 8 (6.626x10 J∗s)( s ) c) Ε= hc = =3.03x10 −19J ν 10 m 656nmx 1nm B) The Photoelectric Effect Won Albert Einstein a Noble Prize. Photoelectric Effect: when electrons are released from an element as a result of electromagnetic radiation strikes it. Threshold Frequency (v )o minimum frequency of light required to produce the photoelectric effect. If the radiation of frequencies is less than the threshold frequency of that element, then it will not emit any photoelectrons, regardless of the radiation intensity. If the radiation produces at least a few photoelectrons then it was equal to or greater than the threshold frequency. Work Function ( ϕ¿ : minimum quantity of energy needed to emit photoelectrons from an element. ϕ=hv o Is related to the strength of the attraction of the element nuclei and the element electrons. You can also calculate the work function of a target metal by: ɸ=hv−¿ KE electrons *Kinetic Energy Electrons I) Using the Work Function 19 The work function of mercury is 7.22x10 J. What is the minimum frequency required to get mercury to release photoelectrons? Can visible light produce the photoelectric effect in mercury? a. Rearrange the work function and solve for (v o. ov = hɸ b. Plug in the work function of mercury and Planck’s Constant. ϕ 7.22x10 −19J 15 −1 V o = −34 =1.09x10 s h 6.626x10 J∗s Cross out joules Divide numerator by denominator Take the exponents, and subtract. 19 34= 15 c. The frequencies of visible light are between 10 and 10 Hz. 4) The Hydrogen Spectrum and the Bohr Model Johann Balmer studied the emission spectrum of hydrogen atoms. He created an empirical equation. To express wavelength in nanometers and units most involved in visible light; use the equation: 2 364.54m = 2 2 m −n mis greater than 2. Values of 3,4,5, and 6 wavelengths of the four brightest hydrogen lines. n is equal to 2. Try Balmer’s equation and plug in m=3. Your answer should be 656.21 nm. Another version of finding hydrogen emission wavelengths is the Rydberg empirical equation: 1 1 1 (H 2 − 2) λ n1❑ n2❑ n1 nd n 2are positive integers, 2 is greater than1n RH s the Rydberg Constant. This depends on the uints expressed for wavelength. It could be: 1.097x10 m 1 5 1 1.097x10 cm 1.097x10 nm 1 Generally the nanometer conversion is used to determine the wavelength. Now try out the Rydberg Emerpical Equation and let n =7 a2d n =2. 1our answer should be λ=397.0nm . A) The Bohr Model Niels Bohr made huge advancements in the way we find emission wavelengths. The Bohr Model looks like: 1 1 ΔΕ=−2.178x10 −18J( − ) n final2 ninitial2 Δ this symbol is called delta and means change. n final where the electron ended up after absorbing energy. n initialwhere the electron started. *If the initial value is larger than the final value then it means the electron loses energy. *If the initial value is smaller than the final value it means that the electron gains energy. * The edge of the atom is marked at 0, anything closer to the nucleus is a negative number, where anything outside of the nucleus is a positive number. Ground State: the most stable, and lowest energy state of an electron particle. Excited State: any energy state above ground state. Electron Transition: movement of electrons between energy levels. Ionization: completely getting rid of the electron from the atom. n= ∞
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