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Chapter 4 Spectroscopy & Arrangement of Electrons(e-) Properties of Light • Electromagnetic Radiation- form of energy w/ wavelike properties as it travels through space • Electromagnetic Spectrum- classifies as electromagnetic radiation based on wavelength(λ) and frequency(ν) The Electromagnetic Spectrum Visible Light Violet Blue Green Yellow Orange 400nm 450nm 500nm 550nm 600nm 700nm Gamma Rays X-Rays UV-Rays v Red 1019 Hz 1017 Hz 1015 Hz Infrared Microwave Radio 1014 Hz 109 Hz Long-Wave 108-106 Hz 105 Hz Short Wavelength Long Wavelength High Frequency Low Frequency • Frequency(ν)- # of waves that pass a given point in a specific time (usually 1s) – Measured in waves per second (1/s) or Hz • Wavelength(λ)- the distance between corresponding points on a wave (ex: peak to peak) – Measured in meters (m) Speed of Light(c) • Speed of Light = wavelength x frequency c = λν • Speed of Light is a Constant: c = 3.0 x 108 m/s • Frequency and Wavelength are Inversely Proportional : v Speed of Light Sample Calculations • What is the wavelength of EM Radiation that has a frequency of 1014Hz? Speed of Light Sample Calculations • What is the frequency of microwaves with a wavelength of 0.01m? Dual Wave-Particle Nature of Light • Light has both the properties of waves and of particles – Wave Properties: light can be bent as it passes through objects – Particle Properties: photons have mass and exert force on other objects Photoelectric Effect • When light particles (photons) collide with a metal, the photons knock electrons (e-) loose • These electrons move toward the positive terminal creating an electric current (electricity) Light as Particles • Photons- particles of light carrying a quantum of energy (Max Planck) • A Quantum of Energy- the minimum quantity of energy that can be lost or gained by an atom E = hν E = energy of a photon (J) h = Planck’s constant (6.626 x 10-34Js) ν = frequency (Hz) Energy of a Photon Sample Calculations • What is the energy of violet light with a frequency of 7.50 x 1014 s-1? Energy of a Photon Sample Calculations • What is the frequency of UV light that has an energy of 2.39 x 10-18J? Sample Calculations • What is the energy of EM radiation with a 1.0 x 10-6m wavelength? Emission Spectra • Emission spectra are fingerprints of an atom • Every atom gives off different colors from the visible light spectrum when they release absorbed energy Energy States of Atoms/Electrons • Ground State- lowest energy level of an electron within an atom • Excited State- a higher energy level within an atom that an electron may exist in – Energy must be absorbed for an electron to go from ground to excited state – Energy is given off as visible light when an atom returns to ground state – Every atom gives off a unique spectrum based on the movement of it’s electrons Energy of a Photon • Ephoton = Efinal - Einitial E2 Excited State E2 – E1 = Ephoton = hν E1 Ground State Bohr Model of the Hydrogen Atom • Bohr’s model indicates that as atoms absorb energy their electrons move to higher energy levels • When the absorbed energy is given off as visible light the electrons return to their ground state Bohr Model Drawing Bohr Diagrams • Shows the total number of electrons • The energy levels fill before an e- can be put in the next energy level – 1st shell = 2 e– 2nd shell = 8 e– 3rd shell = 8 e- Practice Drawing Bohr Diagrams • Draw a Bohr Diagram for: – – – – – C H O Ar F Quantum Model of the Atom • The Bohr model was more accurate than previous models but was only completely accurate for Hydrogen, other elements did not behave exactly as Bohr predicted • The Quantum model was later developed based on work of many scientists including Schrodinger, Heisenberg, & Einstein Quantum Model of the Atom Electrons as Waves • Louis deBroglie proved that electrons had wave properties by showing that electrons could produce interference patterns like sound and light waves • Passing electrons through a crystal also caused the stream of electrons to bend like light waves do Interference Patterns Heisenburg Uncertainty Principle • This principle states that it is impossible to determine simultaneously both the position and velocity of an electron • This theory led to the concept of the electron cloud Quantum Theory • This theory describes mathematically the wave properties of electrons based on probability – Electrons are not in set energy levels – Electrons are in 3D orbits around the nucleus called orbitals Orbitals • Orbitals are 3D regions around the nucleus of an atom that indicate the probable location of an electron Quantum #’s • Quantum #’s are used to specify the properties and location of electrons in orbitals around the nucleus • There are 4 quantum #’s, each is represented by a letter : n, l, m, & s Principle Quantum # (n) • The principle quantum number indicates the main energy level of an electron and it’s distance from the nucleus • n=1 : e- close to nucleus • n=7 : e- further from nucleus * As the n value increases so does the energy of the e-, and the distance from the nucleus Angular Momentum Quantum #(l) • The angular momentum quantum # indicates the shape of an orbital (s, p, d, f) • Each l value has a corresponding shape l = 0 : s-shape, sphere orbital around nucleus - an s-orbital can hold up to 2 electrons l = 1 : p-shape, 2 lobes on either side of nucleus - an atom can have 3 p-orbitals, one in each plane (x, y, z) - p-orbitals can hold up to 6 electrons Angular Momentum Quantum #(l) • l = 2 : d-shape, 4 lobes “clover” around nucleus – An atom can have up to 5 d-orbitals – d-orbitals can hold up to 10 electrons Angular Momentum Quantum #(l) • l = 3 : f-shape – f-orbitals can hold up to 14 electrons Magnetic Quantum # (m) • The magnetic quantum # indicates the orientation of an orbital around the nucleus – Ex. Would indicate if your p orbital is on the x, y, or z plane Spin Quantum # (s) • The spin quantum number indicates the direction that an electron is spinning, either clockwise or counterclockwise • s = +1/2 : clockwise spin • s = -1/2 : counterclockwise spin * 2 electrons in the same orbital must have opposite spins Quantum #’s n – principle quantum # distance from nucleus / main energy level l – angular momentum quantum # shape of orbital m – magnetic quantum # orientation of orbitals around nucleus s – spin quantum # direction of e- spin around nucleus Pauli Exclusion Principle • The Pauli Exclusion principle states that no 2 electrons in the same atom can have the same combination of 4 quantum #’s • This means that no 2 electrons could be in the same place at the same time Orbital Notation • Orbital notation uses boxes and arrows to indicate electrons in orbital by energy level • Aufbau Principle- an electron will occupy the lowest possible energy level that can hold it • Hund’s Rule- orbitals of equal energy will each receive one electron before they receive a second electron e Configurations 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p Order of Atomic Sublevels • Orbital Notation: • P: • F: * Each p-orbital gets 1 ebefore it gets a 2nd e- Examples • N: • Si : • Fe : • Mg : Electron Configuration Notation • 1s22s22p63s23p64s23d104p65s24d105p66s2… Principle Quantum # or Energy Level Angular Momentum Quantum # or Orbital Shape • Ex: Carbon = 6 electrons • C = 1s22s22p2 • Ex: Sodium = 11 electrons • Na = 1s22s22p63s1 # of Electrons in Orbital Examples • N: • Si : • Fe : • Mg : Noble Gas Notation (Shortcut) • To eliminate repetitive electron configuration for elements with large #’s of electrons the symbol of a Nobel Gas can be substituted for a portion of the electron configuration • Ex: K = 1s22s22p63s23p64s1 [Ar]4s1 • Ex: Zn = 1s22s22p63s23p64s23d10 [Ar]4s23d10 Examples • N: • Si : • Fe : • Mg : END OF CHAPTER 4 NOTES !!! Flame Test Lab Compound NaCl Na(NO3) Sr(NO3)2 Ca(NO3)2 Color Description Unknown A B C Ba(NO3)2 K(NO3) D Cu(NO3)2 E Cu(SO4) F Li(NO3) Color Compound Description Calculations Compound Flame Color λ (nm) ν (Hz) Sr(NO3)2 CuSO4 NaCl General Equations: Speed of Light: c c = 3.0 x 108 m/s =λν Energy: E =hν h = 6.626 x 10-34 Js E (J)