2 Electromagnetic Radiation Energy emitted by electrons can be detected at any part of the electromagnetic spectrum
3 Electromagnetic Radiation So, just what is EMR?- an oscillating electric and magnetic fieldwhich travels through spaceOR- a discrete series of “particles” that possess aspecific energy but have no massBOTH!
10 Measuring RadiationAll radiation constantly travels through space at the same velocity (speed) =3.0 x 108 m/s(299,792,458 meters per second)The speed of electromagnetic radiation ---“The speed of light” = c
11 c= lc = 3.0 x108If either the frequency or wavelength is known, the other can be calculated
12 A red light has a wavelength of 728 *10-9 m. What is the speed of the wave in m/s? What is the frequency of the light?
13 A certain blue light has a frequency of 6. 91 x 1014 Hz A certain blue light has a frequency of 6.91 x 1014 Hz. What is the wavelength of the light?
14 Determine the wave length of light with a speed of 50*106 Hz (/sec)
15 If I have a wavelength of 780 *10-9 m and what is the frequency?
16 Microwave ovens often employ radiation with a frequency of 2 Microwave ovens often employ radiation with a frequency of 2.45 x 109 /s. What is the wavelength (in cm) of this radiation?
17 A purple light has a frequency of 7. 42 x 1014 Hz A purple light has a frequency of 7.42 x 1014 Hz. What is its wavelength? What is the energy of one quanta of light
23 Planck’s Quantum Theory In 1900, German Physicist Max Planck proposed: “Radiant energy may only be absorbed or emitted in discrete amounts: quanta.”
24 Electromagnetic spectrum and Energy Planck discovered the energy of a wave or photon of light is constanth = (6.63 x 10-34J/Hz)Planck’s constant (h)
25 Energy and RadiationIf the frequency of a wave is known, then the energy of the wave can be calculated as well in the same mannerE=h OrE= h (c/)
26 Energy and Radiation = 1.035 x 108 Hz E=hv Sample: What is the energy of a wave which has a frequency of x 108 Hz?6.63 x10-34 J1.035 x 108 HzE =6.86 x 10-26J=Hz = x 108 Hz6.63 x10-34 Jh =Hz
27 PhotonsPhotons are a particle of radiation or an individual quantum of lightQuantum is a finite quantity of energy that can be gained or lost by an atom
28 The Nature of Light and Radiation If e- are exposed to energy which matches those levels (a photon), they leap to unstable higher energy levels.As they fall back they emit that energy ( a photon) at a wavelength and frequency which can be detected.
29 Photoelectric EffectAlbert Einstein proposed that light not only behaves as a wave, but as a particle tooAlbert Einstein did not get the Nobel Prize for Relativity. He received it for work that he did between 1905 and 1911 on the Photoelectric Effect.
30 Photoelectric Effect In 1905, Einstein applied this quantum theory to explain the photoelectric effect:
32 Photoelectric Effect-if EMR was absorbed as a wave, then the number of electrons ejected and the energy of the electrons ejected should vary only with the intensity of the light
33 Einstein Photoelectric Effect hvEhv - Ee- = Ionization Energy
34 Photoelectric Effect NUMBER of e-: does vary with EMR intensity ENERGY of e-: vary only with MR frequency AND: no effect if freq is below a threshold value!
35 Photoelectric effectThe reason that certain types of light cause this effect but others do not has to do with threshold energyRemember that electrons must gain energy in certain amounts (E = h)Only certain types of light with “just the right” frequency can cause electrons to become excited
38 Photoelectric Effect View EMR as a collection of particles(called photons), with each photon having the following energy:E = hν
39 Photoelectric Effect Each photon will cause an electron to be ejected IF the energy of the photon is above a minimum (threshold) value. Any energy of the photon above that needed to eject the electron would be transferred to the electron as kinetic energy
40 Photoelectric EffectIncreased EMR intensity translates to an increase in the number of photons (increasing the number of electrons ejected)
41 EinsteinEinstein did not receive his Nobel Prize of the theory of relativity but for his work on the photoelectric effect
43 Hydrogen Line Spectrum Because hydrogen atoms emit only specific frequencies of light indicate that the energy differenced between the atom states are fixed
44 The Bohr AtomThe electron in a hydrogen atom can exist only in discrete orbitsThe orbits are circular paths about the nucleus at varying radii
45 Light leads to Electrons Electrons exist in specific, discrete levels or quanta of energyThe lowest energy of each electron is known as its ground stateIf the right amount (quantum) of energy is applied, then an electron will “leap” to another energy level
46 The Bohr Atom Each orbit corresponds to a particular energy Orbit energies increase with increasing radii
47 The Bohr Atom The lowest energy orbit is called the ground state After absorbing energy, the e- jumps to a higher energy orbit (an excited state)
48 The Bohr AtomWhen the e- drops down to a lower energy orbit, the energy lost can be given off as a quantum of lightThe energy of the photon emitted is equal to the difference in energies of the two orbits involved
49 Bohr’s Model of H Atom Linked an atom’s e- to photon emission Said that e’ can circle the nucleus only is specific paths or orbits.Orbits are separated from the nucleus by large empty spaces (nodes) where the e- cannot exist. (Rungs on a ladder)Remember, E increases as e- are further from the nucleus.
50 Bohr’s Model of H AtomWhile an e- is in orbit, it cannot gain nor lose EIt can, move to a higher E level IF it gains the amount of E = to the difference between the higher E orbit and the lower energy orbit.When falls, it emits a PHOTON = to the difference.Absorption gains E/Emission gives off E
51 Take the good with the bad (Bohr’s Model) Led scientists to believe a similar model could be applied to all atoms.Bad:Yikes!!! Did not explain the spectra of atoms with more than one e’ or chemical behavior of atoms!
52 The Bohr Model of the Atom I pictured electrons orbiting the nucleus much like planets orbiting the sun.But I was wrong! They’re more like bees around a hive.WRONG!!!Neils Bohr
53 The Nature of Electrons If light acts as both particles and waves, then so do electrons.If electrons act like waves, then how can we locate them?
54 Atoms are like onions and ogres they have lots of layers plantanswers.tamu.edu
55 OrbitsWe like to draw orbits as circularbut they really aren’t
56 OrbitsTurns out, there's no reason to assume that electron orbits are circular.
57 OrbitsThere are lots of paths an electron can take in order to get around the nucleus.(or through the nucleus)
58 OrbitsIn fact it's very rare for an atom's electron to be in a circular orbit.
59 Electron FactsThe electron moves at different speeds. Fast near the nucleus and slow when it's far from the nucleus.
60 Well we really don’t Because atoms are so small, no one can see them. For this reason, we just can't say exactly where the electron is, as it moves about the nucleus. In those orbit pictures, we know the electron will be somewhere on the white orbit lines.
61 This makes the models involving orbits, whether circles, or ellipses, wrong, because the orbits are pretty specific about the electron's location.
62 So how do we locate electrons in the modern atom? Quantum Numbers
63 Quantum NumbersSpecify the properties of atomic orbitals and the properties of electrons in orbital
64 Quantum NumbersDon’t tell us where the electron is, just where it's most likely to be.The probable location of electrons are described using an address
65 The four quantum numbers their symbols are n, l, m and s.EVERY electron in an atom has a specific,unique set of these four quantum numbers.
66 Quantum Numbers There are 4 parts to each address Principle quantum number (n)Angular quantum number (l)Magnetic quantum number (m)Spin quantum number (s)
67 The Principal Quantum Number The main energy level of an electron
68 Principal Quantum Number (n) Describes the size of the orbital.Since the distance from of an electron from the nucleus is directly proportional to the energy of the electronIt must be a whole number n=1, n=2 …
69 Angular quantum number (l) Azimunthal Describes the shapeThe secondary quantum number divides the shells into smaller groups of orbitals called subshells (sublevels).s p d f g h...
70 Angular (Azimuthal) Momentum Quantum Number LetterShapesSpherepDumbelldCloverleaff“funky”
71 Each suborbital can hold a maximum of 2 electrons per orbital
72 Angular (Azimuthal) Momentum Quantum Number LetterMax number of suborbitalsMax. # of e-s12p36d510f714
73 Quantum NumbersThe more crowded the dots are in a particular region, the better chance you have to finding your electron there.
83 The Magnetic Quantum Number describes the orientation in space of a particular orbital.x, y, zIt is called the magnetic quantum number because the effect of different orientations of orbitals was first observed in the presence of a magnetic field.)
84 Magnetic quantum number (m) The orientation in spaceWhat axis the orbit is located on
85 The Spin Quantum Number allows two electrons of opposite spin (or symmetry) into each orbital.+1/2 or -1/2 or
87 Aufbau (filling up) Principle the number of electrons in an atom is equal to the atomic numbereach added electron will enter the orbitals in the order of increasing energyan orbital can only hold 2 electrons.
88 Hund’s RuleFor orbitals of equal energy, one electron goes into each orbital until all orbitals are half-full.2p42s21s2
89 Aufbau PrincipleLower-energy orbitals of an atom are filled with electrons first.1s 2s 2p 3s 3p 4s 3d 4p …
90 Pauli Exclusion Principle No two electrons in an atom can have an identical set of four quantum numbers. (electrons sharing an orbital have different spins)2s21s2
91 Key TermsGround State Electron Configurations (1s22s2…/Orbital Notations __ __ __ __Aufbau Principle (lowest first/periodic guide)Hund’s Rule: Single e’ before pairing beginsPauli Exclusion (up/down—no 2 e’ same set of quant. #’s)Highest Occupied LevelInner-Shell e’Noble Gas Notation
92 Spin Diagram for Hydrogen Electron configurationfor Hydrogen1s1
93 Spin Diagram for Helium Electron configurationfor Helium1s2
94 Spin Diagram for lithium Electron configurationfor lithium1s22s1
95 Spin Diagram for beryllium Electron configurationfor berylium1s22s2
96 Electron configuration Spin Diagram for boronElectron configurationfor boron1s22s22p1
97 Spin Diagram for carbon Electron configurationfor carbon1s22s22p2
98 Spin Diagram for nitrogen Electron configurationfor nitrogen1s22s22p3
99 Spin Diagram for oxygen Electron configurationfor oxygen1s22s22p4
100 Spin Diagram for fluorine Electron configurationfor fluorine1s22s22p5
101 Electron configuration Spin Diagram for neonElectron configurationfor neon1s22s22p6
102 Standard Notation or Complete electron configuration of Fluorine Number of electronsin the sub level 2,2,51s2 2s2 2p5Main Energy Level Numbers 1, 2, 2Sublevels
103 Shorthand Notation or Noble Gas Configuraiton Use the last noble gas that is located in the periodic table right before the element.Write the symbol of the noble gas in brackets.Write the remaining configuration after the brackets.Ex: Fluorine: [He] 2s2 2p5
107 You broke your big toe! The x ray they take of toe uses waves that have a length 2.19 x 1010m. What is the speed of the wave in m/s? What is the wavelength in nm? What is the frequency of the x ray?