Electrons in AtomsChapter 5 Chemistry 2
The Development of Atomic Models 5.1 • Rutherford: model of electrons moving like planets around the Sun • Did not explain: colors of flames, glowing objects at high temp = chemical properties
The Bohr Model 5.1 • Niels Bohr (1885 – 1962) – Danish Physicist • Rutherford’s student • Electrons move in orbits around nucleus • Has energy levels • To move from one level to another, must gain/lose energy • The higher the energy level, the farther from the nucleus (normally) • Quantum Energy – amount of energy required to move electrons to different energy levels • Levels are not evenly spaced • Higher energy levels = closer together = less energy to change energy levels
The Quantum Mechanical Model 5.1 • Erwin Schrodinger (1887 – 1961) – Austrian Physicist • Mathematical solution for electron movement • Determines allowed energies can have and likely locations • Probability – how likely electron is in a particular location • Compare to plane propeller • More dense area in cloud = high probability • Less “ “ “ “ = low “
Atomic Orbitals 5.1 • Region of space which there is high probability of finding electron • Labeled by principal quantum numbers (n) – denoted by numbers • For each level = could be several orbitals with different shapes and energy levels • Each sublevel = different shape • Denoted by letters • Sublevels • s = spherical • p = dumbbell • 3 types with different orientations in space • d = 4 of the 5 are clover leaf shape • f = 7 types that are complicated
Electron Configuration 5.2 • The way electrons are arranged in various rbitals around the nuclei • 3 rules • Aufau Principle – e- occupy lowest energy first • Pauli Exclusion – atomic orbital may have at most 2 e- • To occupy same orbital must have opposite spin – clockwise or counterclockwise • Hund’s Rule – electrons occupy orbital of the same energy in a way that the e- spin in the same directions large as possible
Lets look at Hydrogen H ______ or 1s1Let's take a look at helium. (Z=2) • He _____ or 1s2Now let's take a look at Li (Z = 3) • Li ______ ______ or 1s2 2s1Now take a look at boron (Z = 5) 1s 2s 2p B O O OOO or 1s2 2s2 2p1One final example before you get turned loose to do some yourself: Ne (Z = 8) 1s 2s 2p Ne O O OOO or 1s2 2s2 2p6Both the orbital diagram and shorthand have uses. The orbital diagram is best for showing how bonding takes place. While the shorthand method is used on the periodic table to show the electronic configuration of the outermost orbitals.One last thing. The orbitals would seem to be filled in the following order:1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s etc. • BUT THIS IS NOT SO.The electron order of filling is in fact: 1s 2s 2p 3s 3p 4s 3d 4p 5a 4d 5p and so on.
Exceptional Electron Configurations 5.2 • Can get correct configurations to Vanadium (atomic # 23) w/ Aufbau principle • Exceptions due to e- - e- interaction w/ similar energy orbitals • Don’t worry about exceptions now, just know they exist • Refer to handout on exceptions!!!
Light 5.3 • Isaac Newton (1642 – 1727) • Tried to explain behavior of light by particle movement
Light 3.5 • Wavelength =Greekletterlambda (λ) • Frequency = v - # of wave cycles that pass a point in a specific time • Usually cycles per second • Measured in Hertz (Hz) • c = λv • c = constant = speed of light • λ & v are inversely proportional (1 increases, the other decreases) • Electromagnetic Radiation – radio waves, microwaves, infrared waves, visible light, uv waves, xrays, and gamma rays • All waves travel at 2.998 x 108m/s in vacuum • Sunlight = various wavelengths • In prism, sunlight separates into different frequencies to form a spectrum • Rainbow – droplet of water acts a prism
Atomic Spectra 5.3 • pass electric current thru neon tube (absorbe energy) energize electrons electrons move to higher energy levels lose energy move back to lower energy level emit light • No elements have the same emission spectrum • Like people’s fingerprints
Exploration of Atomic Spectra • Ground State – lowest possible energy of electron • Electronic transition • The quantum of energy in the form of light is emitted when electron drops back t o ground state after excited state • Quantum of energy related to frequency • E = hv • h = 6.626 x 10-34 J.s • Hydrogen Spectrum • Lyman series – transition to the n = 1 level – Ultraviolet • Balmer series – transition to the n = 2 level – Visible light • Paschen series – transition to the n= 3 level = Infrared
Notice – spectral lines become closer together with increase value of n b/c energy levels are closer
Quantum Mechanics 5.3 • Einstein explain that light could be described as quanta of energy called photons • Louis de Brogilie asked: Given that light behaves as waves and particles, can particles of matter behave as waves? • Proposal = yes! • Accepted b/c of scientific experiment from Davisson and Germer • Bombardment of metals with beams of e- • e- bounced back as waves • Theory: all moving particles have wavelike behavior • Mass must be small enough to be big enough to observe • Use beams of e- in e- microscope to magnify objects
Quantum Mechanics 5.3 • Classical mechanics – older theory • Describes motion of bodies much larger than atoms • Quantum Mechanics – newer theory • Describes motions of subatomic particles and atoms • Heisenberg uncertainty principle – impossible to know exactly both the velocity and the position of a particle at the same time • Matters for small particles, not big
Technology & SocietyLasers at Work • Laser – produce intense beam of light where waves that have crests that coincide • Used to read CD-ROMS, DVDs, scanning bar codes, cut metal, surgery • Surgeon can reshape the cornea by removing some tissue so patient doesn’t need glasses • Flashlight vs. Laser beam • Flashlight light travels in all directions • Laser light travels parallel to one another = beam http://science.howstuffworks.com/laser.htm 8 minutes