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1Unit 4 –Quantum MechanicsCartoon courtesy of NearingZero.net
2Light and EnergyMuch of what we know about electrons comes from study of lightLight behaves as wavesLight also behaves as particles (photons)
3Wave properties Light waves are part of the electromagnetic spectrum Can be characterized by 5 properties:AmplitudeWavelengthFrequencySpeedEnergy
4Electromagnetic radiation propagates through space as a wave moving at the speed of light. C = speed of light, a constant (3.00 x 108 m/s) = frequency, in units of hertz (hz, /s, s-1) = wavelength, in meters
6Calculating the color of light… Given that a beam of light has frequency of 6.0 x 1014 /swhat is the wavelength of this light?What color is it? –use p.129
7E = h = frequency, in units of hertz (hz, /s, s-1) The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation. This is the energy of a photon of a particular frequency.E = hE = Energy, in units of Joules (kg·m2/s2)h = Planck’s constant (6.63 x J·s) = frequency, in units of hertz (hz, /s, s-1)
8Long Wavelength = Low Frequency Low ENERGY Short Wavelength = Wavelength TableShortWavelength=High FrequencyHigh ENERGY
9Spectroscopic analysis of the visible spectrum… (such as seen from an incandescent light bulb)…produces all of the colors in a “continuous spectrum”
10Spectroscopic analysis of the hydrogen spectrum… (as given off by a hydrogen gas-filled fluorescent light bulb)…produces a “bright line” or “emission spectrum”
11Electron transitions occur when electrons absorb energy in a ‘quantum jump’ to an ‘excited state’.When they ‘fall’ back to their ’‘ground state’they emit photons of light with distinctwavelengths, visible as bands on an ‘emission spectrum’(not all are in the visible range).VISIBLEJUMPS
12The Bohr Model of the Atom I pictured electrons orbiting the nucleus much like planets orbiting the sun.But it turns out they’re more like bees around a hive.WRONG!!!Neils Bohr
13Quantum Mechanical Model of the Atom Mathematical laws were used to identify the regions outside of the nucleus where electrons are most likely to be found.These laws are beyond the scope of this class…But here are two important examples:
14Schrodinger Wave Equation Equation for probability of a single electron being found along a single axis (x-axis)Erwin Schrodinger
15Heisenberg Uncertainty Principle “One cannot simultaneously determine both the position and momentum of an electron.”You can find out where the electron is, but not where it is going.OR…You can find out where the electron is going, but not where it is!WernerHeisenberg
16Electron Energy Level (Shell) Generally symbolized by n, it denotes the average distance of the electron from the nucleus.Number of electrons that can fit in a shell:2n21 holds holds holds 18 etc.
17An orbital is a region within an energy level where there is a probability of finding an electron. This is a probability diagram for the…s orbital… in the first energy level.Orbital shapes are defined as the space thatcontains the electron 90% of the time.
18Quantum Numbers: Electron Addresses! 1st: Energy Level (n = 1, 2, 3, 4 …) Average distance of the electron in the electron cloud from the nucleus.2nd: Sublevel (l = s, p, d, f) Shape of electron cloud.energy levels have n sublevels.3rd: Orbital (ml) Orientation of the cloud in 3-D space.s sublevels have 1 orbital (spherical)p sublevels have 3 orbitals (dumb bell shape)d sublevels have 5 orbitals (varied)f sublevels have 7 orbitals (varied)4th: Spin (ms = +1/2 or -1/2) Direction of electron spin.Each orbital holds 2 electrons, one with each spin.
19Pauli Exclusion Principle: Each electron of an atom has its own unique set of quantum numbers. They may have three that are the same, but never all four.
20Sizes of s orbitals Orbitals of the same shape (s, for instance) grow larger as n increases…Nodes are regions of low probability within anorbital.
21The s orbital the origin of the three axes in space has a spherical shape centered aroundthe origin of the three axes in spaceThere is only one orientation for this shape
22P orbital shape There are three dumbbell-shaped p orbitals in each energy level above n = 1, each assigned toits own axis (x, y and z) in space.
23To remember the shapes, think of “double dumbells” Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3.To remember the shapes, think of “double dumbells”d orbital shapes…and a “dumbellwith a donut”!
24In case you are too curious, here is what the f orbitals look like.
26The Diagonal Rule: Sublevels in order of increasing energy 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f14…The Diagonal Rule: Sublevels in order of increasing energy5g186f147d108p68s24f14 5f143d10 4d10 5d10 6d102p6 3p6 4p6 5p6 6p6 7p61s2 2s2 3s2 4s2 5s2 6s2 7s2Aufbau Principle: Electrons fill the lowest energy position available.
27Energy levels and sublevels on the Periodic Table
28Hund’s Rule: Electrons fill orbitals so that there are a maximum number of orbitals with a single electron.
29Element notation Orbital notation ConfigurationnotationOrbital notationNoble gasLithium1s22s1____ ____ ____ ____ ____1s s p[He]2s1Beryllium1s22s2[He]2s2Boron1s22s22p1[He]2s2p1Carbon1s22s22p2[He]2s2p2Nitrogen1s22s22p31s s p[He]2s2p3Oxygen1s22s22p4[He]2s2p4Fluorine1s22s22p5[He]2s2p5Neon1s22s22p6[He]2s2p6
30Electron configuration of the elements of the first three periods