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# chem1030 chapter 6 notes: electronic structure of atoms

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Department
Chemistry
Course Code
CHEM 1030
Professor
Marla Spergel

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Chapter 6: Electronic Structure of Atoms  Waves o The distance between corresponding points on adjacent waves is the wavelength (). o The number of waves passing a given point per unit of time is the frequency (). o For waves traveling at the same velocity, the longer the wavelength, the smaller the frequency.  Electromagnetic Radiation o All electromagnetic radiation travels at the same velocity: the speed of light (c), 8  3.00  10 m/s. o Therefore, c =   The Nature of Energy o Max Planck explained it by assuming that energy comes in packets called quanta. o Einstein used this assumption to explain the photoelectric effect. o He concluded that energy is proportional to frequency:  E = h  where h is Planck’s constant, 6.626  104J-s. o Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light:  c =   E = h o For atoms and molecules one does not observe a continuous spectrum, as one gets from a white light source. Only a line spectrum of discrete wavelengths is observed. o Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:  Electrons in an atom can only occupy certain orbits (corresponding to certain energies).  Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom.  Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by E = h  The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation:   where R is the Rydberg constant, 2.18  1018J, and n and n are the H i f initial and final energy levels of the electron.  The Wave Nature of Matter o Louis de Broglie posited that if light can have material properties, matter should exhibit wave properties. o He demonstrated that the relationship between mass and wavelength was   The Uncertainty Principle o Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known:  o In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself!  Quantum Mechanics o Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated; known as quantum mechanics. o The wave equation is designated with a lower case Greek psi (). 2 o The square of the wave equation,  , gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time.  Quantum Numbers o Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies.  Each orbital describes a spatial distribution of electron density.  An orbital is described by a set of three quantum numbers.  Principal Quantum Number (n) o The principal quantum number, n, describes the energy level on which the orbital resides. o The values of n are integers ≥ 1. 
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