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Week 4 Notes

by: Alexa Johnson

Week 4 Notes CHEM 111

Alexa Johnson
GPA 4.0
General Chemistry I (GT-SC2)
Claire M Filloux

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General Chemistry I (GT-SC2)
Claire M Filloux
Class Notes
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This 3 page Class Notes was uploaded by Alexa Johnson on Sunday September 20, 2015. The Class Notes belongs to CHEM 111 at Colorado State University taught by Claire M Filloux in Fall 2015. Since its upload, it has received 12 views.


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Date Created: 09/20/15
CHEM 111 General Chemistry Week 4 Notes gt 3134 The Nature of Lioht and Atomic Spectra gt 31 Waves of Light In order to understand the stability of atoms structure we must rst study how atoms interact with a form of energy Eectromagnetic radiation or radiant energy Any form of radiant energy in the electromagnetic spectrum Visibe light colors are the most common form of electromagnetic spectrum we know but there are more with which we are likely familiar Radio waves infrared xrays etc Eectromagnetic spectrum a continuous range of radiant energy that includes gamma rays Xrays ultraviolet radiation visible light infrared radiation and radio waves in that order from left to right Al of these forms of radiation are considered electromagnetic because of James Clerk Maxwell39s theory that electromagnetic radiation move through space as waves with of an electric eld and a magnetic eld Waves of radiation have a characteristic wavelength 2 which is the distance from one wave crest to the next Each also has a characteristic frequency which is the number of crests that pass a stationary point in space per second in a unit often called hertz Hz which is equal to 1 wave per second Wavelengths and frequency are inversely related one goes up other goes down and vice versa The relationship for radiant energy traveling through a vacuum where it is called the speed of light is lvc l wavelength vfrequency cspeed of light2998 x 10quot8 ms The reciprocal equation for this relationship is vlc gt 32 Atomic Spectra Observations about interactions of electromagnetic radiation started over 200 years ago Wiliam Wollaston studied light through glass prisms and found that they were not completely continuous Later Joseph von Fraunhofer mapped the wavelengths of over 500 lines that we call Fraunhofer lines or dark lines in the otherwise continuous solar spectrum Using an instrument called a spectroscope it was discovered that light was given off or emitted by elements Bunsen and Kirchhoff discovered that the lines in atomic emission spectra of certain elements exactly matched the wavelengths of some of Fraunhofer lines of sunlight This showed that at very high temperatures the atoms of elements emit a characteristic spectrum On the other hand the atoms of elements in the gaseous state absorb electromagnetic radiation when illuminated by an external source of radiation which is called atomic absorption spectra gt 33 Particles of LightQuantum Theory In 1900 Max Plank proposed that objects emit electromagnetic radiation only in integral multiples of an elementary unit or quantum of energy de ned by the equann Ehv vfrequency hPlanck39s constant 6626 x 10quot34 5 Thus created the equation of quantum theory Ehcv For example if you were to walk up a set of stairs you cannot stop at a height between each You must stand on the steps On the steps height is quantized meaning that the discrete changes in height model Planck s hypothesis that energy is released down the stairs or absorbed up the stairs in packets of energy caed quanta Quantized restricts values to having wholenumber multiples of a speci c base value Today we call the tiny packets of energy photons a quantum of electromagnetic radiation energy The photoelectric effect is the release of electrons from a material as a result of electromagnetic radiation striking it gt 34 The Hydrogen Spectrum and Bohr Model In 1888 Johannes Rydberg created an equation for predicting wavelengths of hydrogen s spectral lines 10 RH1n1quot21n2quot2 where n1 and n2 are positive integers and RH is the constant In 1913 Neils Bohr proposed a model that would explain 1why hydrogen atoms lose and gain discrete quanta 2why their electrons do not spiral into their nuclei In his model the electrons revolve around the nucleus in a series of orbits representing an energy level n1 is the lowest energy and closest to nuclei higher energy the farther away According to his model orbits have larger values of n and higher energies based on this equation E 2178 X 10quot18J 1nquot2 The higher the n integer the less negative the value becomes and the closer it gets to in nity At in nity energy approaches 0 Another equation of Bohr39s accurately explains the observations of previous others and how energy levels change E 2178 x 10quot18J 1n nalquot2 1ninitialquot2 When the electron of a hydrogen atom is the lowest energy level it is said to be in ground state but if the energy level of a hydrogen atom is above that n1 it is said to be in an excited state The movement of the electron between two energy levels is called electron transition gt 37 The Sizes and Shapes of Atomic Orbitals Atomic orbitals are three dimensional shapes that are graphical representations There are three main kinds of orbitals s p and d S is the a singular spherical representation of the electrons near the nucleus Shells with ngt or 2 have a subshell containing three P orbitals with two lobes oriented around the nuclei Shells with principal quantum numbers of 3 or more have ve D orbitals The shapes and purpose for the orbitals will be discussed further later on


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