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Chapter 1

Discovering Chemistry Chapter 1.docx

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Chemistry 1027A/B
Felix Lee

Discovering Chemistry Chapter 1 September 14, 2011 Chapter 1 – Atomic Theory  Atomic Structure o What does an atom consist of?  Inside the nucleus: Protons (positive) and Neutrons (no charge)  Outside the nucleus: Electrons (negative)  Has some amount of energy  We can give it more energy o “excited” state – higher energy level  Eventually it will return to “ground” state o Loses energy it originally gained (emits energy in the form of light)  History o Started by looking at Hydrogen o Wave Nature of Light  Light travels as a wave  Wavelength (lameda) is the distance between two peaks  Frequency, v, is the number of times a peak passes a given point in a given amount of time  Speed of light is given by c: 3.0 x 10 (power 8) m/s o Wave Interference  A) shows constructive interference; two waves “in phase” combine to form a larger wave  B) shows destructive interference; two waves “out of phase” cancel o Particle Nature of Light  Light interacts (may be either absorbed or emitted) as discrete particles with discrete energies  Planck’s constant, h, is the link between wave properties of light and particle properties of light and allows us to calculate the energy of the photon o Hydrogen Emission  Wavelengths of visible light emitted by excited hydrogen are: 410, 434, 486, and 656 nm  See Balmer Equation in book  Light goes from energy level, n, to energy level 2  Energy level 2 is visible light  See Lyman Equation in book  Corresponds with UV Light  Light goes from energy level, n, to energy level 1  Energy level falls further, therefore there is more energy  See Paschen Equation in book  Corresponds with IR light  Light goes from energy level, n, to energy level 3 o Summary of the Bohr Model  Successful at explaining the emission spectrum of hydrogen  Considered the electron as a particle orbiting the nucleus  No explanation for the splitting of emission lines  No explanation of intensities of emission lines  Didn’t work for atoms with more than one electron o Quantum Mechanical Model  Louis de Broglie proposed that electrons (as well as light) could also behave as both waves and particles  Quantum mechanical model takes a wave approach; when bound inside an atom, electrons behave like waves  Bohr’s orbits are explained by standing waves; only certain wave patterns are possible, resulting in only certain frequencies  Principal Quantum Number = # of Nodes + 1  Node: amplitude of the wave is zero  No possibility of finding an electron at a node o What is a Wave Function  A wave function is a variable quantity that mathematically describes the wave characteristics of a particle  By analogy with waves such as those of sound, a wave function, designated by the Greek letter psi may be thought of as an expression for the amplitude of the particle wave (or de Broglie wave), although for such waves amplitude has no physical significance  The intensity of the wave, tells us the probability of finding the particle at that point, and is proportional to psi squared o The Schrodinger Equation  Mathematical representation of de Broglie’s waves  Solution of the Schrodinger Equation gives psi for the wave pattern of the electron  Psi squared gives the probability of finding the electron 95% of the time  This probability is called an orbital (just a probability wave where we find electron most of the time o How is an Orbital Described?  Mathematics defines the wave functions of the orbitals as well as their shapes and sizes  Solution of the Schrodinger Equation gives three numbers which specify the orbital  These numbers (plus one more) describe the orbital of an electron in 3D space; these are called “Quantum Numbers”  No two electrons in an atom can have the same four quantum numbers  If two electrons are in the same orbital at least one number will be different o Quantum Numbers  Principal Quantum Number, n – Energy Level (can be any whole number)  Azimuthal Quantum Number, l – Shape of Orbital (0 up to n-1)  l = 0; s sublevel  l = 1; p sublevel  l = 2; d sublevel  l = 3; f sublevel  Magnetic Quantum Number, ml –  ml = 0, +/- l  if l = 1 (p sublevel); ml = -1, 0, +1  if l = 2 (d sublevel); ml = -2, -1, 0, +1, +2  Spin Quantum Number, ms – the spin of the electron  ms = +1/2 or -1/2  two electrons with opposite spins o Example 1.1.1  Quantum Number Validity
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