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Chapter 4 Electron Configurations

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Early thoughts Much understanding of electron behavior comes from studies of how light interacts with matter. Early belief was that light traveled as a wave, but some experiments showed behavior to be as a stream of tiny fast moving particles

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Waves Today scientists recognize light has properties of waves and particles Waves: light is electromagnetic radiation and travels in electromagnetic waves.

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4 Characteristics of a wave: 1) amplitude - height of the wave. For light it is the brightness 2) Wavelength ( )– distance from crest to crest. For light – defines the type of light Visible light range – 400- 750nm

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Properties continued 3) Frequency ( )– measures how fast the wave oscillates up and down. It is measured in number per second. Hertz = 1 cycle per second FM radio = 93.1 MHz or 93.1 x 10 6 cycles per second Visible light = 4 x 10 14 cycles per second to 7 x 10 14 cycles per second 4) speed – 3.00 x 10 8 m/s

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Short wavelength, high frequency Long wavelength, low frequency Visible Spectrum ROY G BIV Longer wavelength shorter wavelength

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Electromagnetic spectrum (meters) 10 -11 gamma 10 -9 x-rays 10 -8 UV 10 -7 visible light 10 -6 infrared 10 -2 microwave 1 TV

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Wavelength and frequency Wavelength and frequency are inversely related!! = c/ Where is the wavelength, c is the speed of light and is the frequency Speed of light = Constant = 3.00 x 10 8 m/sec

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Example Example: An infrared light has a wavelength of 2.42 x 10 -6 m. Calculate the frequency of this light. = c/ = 3.0 x 10 8 m/sec = 2.42 x 10 -6 m = 1.2 x 10 14 waves/sec

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Wavelength and frequency ****Remember and are inverse. Therefore short wavelength = high frequency!!

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Quantum Theory Needed to explain why certain elements when heated give off a characteristic light (certain color) 1900- Max Planck – idea of quantum

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Quantum Theory - The amount of energy (electromagnetic radiation) an object absorbs/emits occurs only in fixed amounts called quanta (quantum) - Quanta – finite amount of energy that can be gained or lost by an atom.

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1905 Einstein’s theory Einstein explained photoelectric effect by proposing that light consists of quanta of energy called PHOTONS Photon = discrete bit of energy Consider light traveling as photons

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Energy equation Amount of energy of a photon described as E = h Where E = energy = frequency h = Planck’s constant = 6.6262 x 10 -34 J s Joule = SI unit for energy

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Photoelectric effect – electrons are ejected from the surface of a metal when light shines on the metal. The frequency determines the amount of energy. The higher the frequency, the more energy per photon.

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Photoelectric effect Ex. X-rays have a high frequency; therefore can damage organisms while radio waves have a low frequency.

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Dual nature of radiant energy Photons act both like particles and waves. Line Spectra: A spectrum that contains only certain colors, or wavelengths

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Question: How are electrons arranged in atoms? Note: All elements emit light when they are vaporized in an intense flame or when electricity is passed through their gaseous state.

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How are electrons arranged in atoms? Explanation: Bohr atom: 1911 postulated that atoms have energy levels in which the electrons orbit energy levels and/or orbits are labeled by a quantum number, n. lowest energy level = ground state n=1

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How are electrons arranged in atoms? when an electron absorbs energy, it jumps to a higher level (known as the excited state) n = 2,3 or 4 Bohr model of an atom Hydrogen Red- e falls from 3 to 2 Blue e falls from 4 to 2 Violet e falls from 6 to 2

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1924 – Louis DeBroglie – If waves of light can act as a particle, then particles of matter should act like a wave. Found to be true.

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DeBroglie Matter waves = wavelike behavior of particles. From equation – wave nature is inversely related to mass therefore we don’t notice wave nature of large objects. However, electrons have a small mass so they have a larger wave characteristic

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Schroedinger’s wave equation predicted probability of finding an electron in the electron cloud around nucleus. Gave us four quantum numbers to describe the position.

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Heisenberg’s Uncertainty Principle The position and momentum of a moving object cannot simultaneously be measured and known exactly. Cannot know where it is and where its going at the same time.

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Quantum mechanical model of an atom – Treats the electrons as a wave that has quantized its energy Describes the probability that electrons will be found in certain locations around the nucleus.

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Orbitals Orbitals – atomic orbital is a region around the nucleus of an atom where an electron with a given energy is likely to be found (high probability) Have characteristic shapes, sizes and energies. Four different kinds of orbitals s, p, d, f.

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S and p orbitals

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Orbitals and energy Principal Energy Level = quantum number n. Principal Quantum number Value of n = 1,2,3,4,5,6,7 Tells you the distance from the nucleus

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sublevels Each level is divided into sublevels # of sublevels = value of n N=1 1 sublevel n=2 2 sublevels s,p,d or f shows the shape

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Orbital shape Each sublevel has a certain number of orbitals which are directed three dimensionally S = 1 sphere P = 3 figure 8 (along the x,y or z axis) D = 5 figure 4-27 p. 145 F = 7

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Each electron in an orbital will have a spin – 2 options clockwise, vs. counter clockwise. Electron configuration Pauli Exclusion Principle – each orbital in an atom can hold at most 2 electrons and their electrons must have opposite spin.

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Aufbau Principle- electrons are added one at a time to the lowest energy orbitals available Hund’s Rule – electrons occupy equal energy orbitals so that the maximum number or unpaired electrons result. Occupy singly before pairing

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Diagonal Rule: for order of sublevels: must remember 1s, 2s, 2p, 3s, 3p, 4s Exceptions to Aufbau Principle Cr Z=24 Cu Z = 29

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Energy levels Number of electrons per energy level 1st = 2 2nd = 8 3rd = 18 4th = 32

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Number of electrons per orbital 2 Difference between paired and unpaired electrons Paired = 2 electrons in the same orbital Unpaired = 1 electron in the orbital

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