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Transcript
Radiant Energy • The light form that travels through space • what makes you feel hot in bright sunlight • has properties of both waves and particles • A form of electromagnetic radiation Electromagnetic Wave Electromagnetic Wave • Electric and magnetic fields oscillating at right angles to each other and to the direction of the motion of the wave Electric field Magnetic field Properties of a Wave • Amplitude - height of a wave from the origin to the crest (peak); determines brightness or intensity of light amplitude Direction of wave travel Properties of a Wave • Wavelength - distance between two successive crests or troughs; distance the wave travels as it completes one full cycle of up & down movement wavelength 1 Properties of a Wave • Visible light wavelengths range from 400-700 nm Properties of a Wave • Frequency - tells how fast a wave is oscillating up and down (cycles per second, 1/s, s-1, Hz) • Speed of light is always 3.00 x 108 m/s • THE SHORTER THE WAVELENGTH THE GREATER THE FREQUENCY!!!!!!!!!!!!!!!! Lambda λ = wavelength Nu ν = frequency C = speed of light Electromagnetic Spectrum Electromagnetic Spectrum Electromagnetic Spectrum Quantum Theory • Visible spectrum - continuous; each color fades into the next. Each color has its own wavelength and frequency. A very small part of the spectrum but the only part we can see! • Infrared - radiant heat (ex: heat lamp) • Microwaves - transfer energy to the moisture in food • 1900 - Max Planck predicted how the spectrum of radiation emitted from an object changes with temperature • There is a restriction on the amounts of energy an object emits or absorbs and that each piece of energy is a quantum 2 Quantum Theory • E = amount of energy E = hν • h = 6.6262 x 10-34 J·s (Planck’s constant) KNOW!! • ν = frequency • Can use to determine the temperatures of planets, stars, etc. Photoelectric Effect • 1905 - Einstein - used Planck’s constant to explain the photoelectric effect • Electrons are ejected from the surface of a metal when light shines on the metal Photoelectric Effect • When a photon strikes the metal, the energy is transferred to the metal’s electrons • The greater the frequency, the more energy • If the photon has enough energy, the electron receives the energy and is ejected from the metal’s surface Quantum Energy or • Quantized energy is like stair steps • We don’t feel this because the “steps” are so small!! (equal to h) • A small change on our scale = HUGE change on atomic scale! Photoelectric Effect • Einstein proposed that light consists of photons • Photon - a quanta of energy that behaves like a tiny particle • Each photon carries an amount of energy given by E=hν • Photons can transfer energy to a substance when they strike the surface Photoelectric Effect • Photons behave like particles traveling at the speed of light, with frequencies AND wavelengths! • X-rays = photons with high frequency, high energy; can cause damage to cells! **Thus, light has a dual nature - it is like a particle AND a wave (waveparticle duality)** 3 Line Spectrum Line Spectrum • Spectrum with only certain colors or wavelengths • All elements emit light when heated or when electricity is passed through them when they are in their gas state • When elements absorb energy, they re-emit that energy as light with only certain wavelengths • Gives each element a unique atomic emission spectrum (“fingerprint”) Line Spectrum Line Spectrum • 1911 - Niels Bohr said that to get spectral lines, energy of electrons must be quantized • Electrons are only allowed certain orbits corresponding to different amounts of energy Quantum electrons! • Bohr gave each energy level (orbit) a quantum number • Lowest energy level = ground state, closest to nucleus, n = 1 • Electron plus right amount of energy = EXCITED STATE! • Excited states have quantum numbers n=2, n=3, n=4…..larger orbits are farther from the nucleus Quantum electrons! • When energy (radiation) is absorbed - electron jumps from ground state to excited state! • Radiation is emitted when electron falls back down, resulting in visible emissions of light! 4 More Quantum Things • Light travels through space like a wave, but behaves like a stream of particles when it hits matter • 1924 - Louis de Broglie said wavelike behavior of particles = matter waves; predicted that all moving objects have a wavelike behavior! More Quantum Things • 1927 - Heisenberg said that the position and momentum of a moving object cannot be simultaneously measured • You can’t pin down an electron, but you have a good idea of where it will probably be! A New Approach to the Atom Quantum-Mechanical Model • Known: • Explains the properties by treating electrons as a wave with quantized energy • Electron density density of an electron cloud; high probability = high density, low probability = low density – Electrons can only have certain amounts of energy because energy is quantized (only certain orbits correspond to certain amounts of energy) – Electrons have wave-like behavior – The exact position and momentum of an electron is unknown Atomic Orbitals Atomic Orbitals • Atomic orbital- a region around the nucleus of an atom where an electron with a given energy is likely to be found • Each has a specific shape, size, and energy. The amount of energy in an electron determines the orbital in which it is located • S-orbital: spherical • P-orbital: dumbbell • D- and f- orbitals: complex! • Go here to play with orbitals! http://www.d.umn.edu/~pkiprof/ChemW ebV2/AOs/index.html 5 Principal Energy Levels • Main energy levels given the principal quantum numbers n=1, n=2, etc • Energy increases as n increases • Each energy level has one or more sublevels (s, p, d, f) • Number before sublevel = quantum number (ex: 4s) • Superscript = number of electrons in sublevel (ex: 4s2) Sublevels • • • • S sublevel = 1 s-orbital P sublevel = 3 p-orbitals D sublevel = 5 d-orbitals F sublevel = 7 f-orbitals • Pauli Exclusion Principle each orbital of an atom can hold only 2 electrons with opposite spins, thus determining how many electrons are allowed in each principle quantum energy level and sublevel • Paired = 2 electrons, unpaired = 1 electron • So principal energy levels are made out of sublevels, which are made out of orbitals • The higher the number and the orbital, the more energy it has • The 2p sublevel has 3 “p-orbitals” of equal energy (2px, 2py, 2pz) Electron Spin Increasing energy! • EACH ORBITAL HOLDS 2 ELECTRONS! Electron Spin Principal Energy Levels • Electrons spin! (but not really!) • Clockwise (north pole points up) • Counterclockwise (north pole points down) • Creates a magnetic field; a pair with opposite spins cancel each other; a pair with parallel spins = magnetic effect! Electron Configuration • The distribution of electrons among orbitals describing where electrons are found and what energies they have • Determined by distributing among energy levels, sublevels, and orbitals 6 Electron Configuration • Aufbau principle - electrons are added one at a time to the LOWEST ENERGY orbitals until all electrons of the atom have been accounted for • Aufbau - german for “construction”, theory came from Bohr • Exceptions - Cr, Cu, inner transition elements Electron Configuration • Hund’s Rule - electrons occupy equal energy orbitals so that a maximum number of UNpaired electrons results Orbitals fill up in this order! Electron Configuration • Orbital diagrams - combination of all information about electron placement! • Know how to write these! Which is the correct orbital diagram for oxygen? Must follow Aufbau, Pauli exclusion, and Hund’s rule! Electron Configuration • Compact form - does not show orbitals and electron spins • Principal quantum energy number, sublevels, electrons (superscripts) ex: oxygen : 1s22s22p4 • Sum of superscripts = atomic number of element! Summary • Energy levels - n=1, n=2, n=3, etc • Each energy level has SUBLEVELS (s, p, d, f) • S sublevel - 1 orbital - 2 electrons • P sublevel - 3 orbitals - 6 electrons • D sublevel - 5 orbitals - 10 electrons • F sublevel - 7 orbitals - 14 electrons 7