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(Electron Configurations) Scientists began to understand how electrons acted by observing the way that light interacts with matter. In your notebook, answer the question: How would you describe what light is to someone else? Electromagnetic Radiation-form of energy that exhibits wave-like behavior as it travels through space. Electromagnetic Spectrum-ordered arrangement by wavelength or frequency for all forms of electromagnetic radiation. Wavelength-lambda (λ) The distance between corresponding points on adjacent waves. Units: m, nm, cm, or Å Frequency-nu (ν) The number of waves passing a given point in a definite amount of time. Units: hertz (Hz) or cycles/s = 1/sec = s-1 When an electric field changes, so does the magnetic field. The changing magnetic field causes the electric field to change. When one field vibrates—so does the other. RESULT-An electromagnetic wave. EM waves do not require a medium to propagate through Today you will need: Notes Calculator Periodic Table Agenda: Review Properties of Waves and EM Spectrum Flame Test Demo Frequency and Wavelength Calculations Goal: Know the relationship between λ and v and how to calculate them! Wavelength-lambda (λ) The distance between corresponding points on adjacent waves. Units: m, nm, cm, or Å Frequency-nu (ν) The number of waves passing a given point in a definite amount of time. Units: hertz (Hz) or cycles/s = 1/s = s-1 Electromagnetic Radiation-form of energy that exhibits wave-like behavior as it travels through space. Electromagnetic Spectrum-ordered arrangement by wavelength or frequency for all forms of electromagnetic radiation. c = λ∙ν λ = wavelength (m) ν = frequency (Hz) c = speed of light= 3.0 x 108 m/s (constant) λ and ν are _______________ related. Calculate the frequency for violet light. Truck-mounted helium-neon laser produces red light whose wavelength (λ ) is 633 nanometers. Determine the frequency (v). *Remember that c=3.0x108m/s. *Use the formula c= λ . v c =3.0x108 m/s c= λ . v v=c / λ λ = 633nm= 6.33x10-7m v = 3.0x108 m/s = 0.47x 1015s-1 = 4.7x1014 s-1 6.33x10-7m Frequency = 4.7x1014 Hz (cycles per second) EX: Find the frequency of a photon with a wavelength of 434 nm. GIVEN: WORK: =c =? = 434 nm = 4.34 10-7 m = 3.00 108 m/s -7 m 8 4.34 10 c = 3.00 10 m/s = 6.91 1014 Hz Calculate the frequency for the yelloworange light of sodium. Calculate the frequency for violet light. E = h∙ν E = energy (joule) h = Planck’s constant = 6.63 x 10-34 J∙s ν = frequency (Hz) E and ν are ______________ related. Calculate the energy for the yellow-orange light for sodium. Calculate the energy for the violet light. Welcome! Please turn into the box: -Chp. 4 Outline -EM Spectrum Practice (if you didn’t Fri) -Energy Calculations (if you didn’t Fri) Chp. 4 Reading Quiz Pass Back Chp. 3 Quiz Wavelength, Frequency, Energy Problem Evidence that light is a wave! If Time: Evidence that light is a particle! Pick your favorite (or any, really) radio station. Knowing that radio station frequencies are measured in Mega Herz and that 1 MHz = 106Hz, calculate: A)The wavelength (in m) of your radio wave B) The Energy associated with your station Discuss: What do you know/notice about the relationship between wavelength and energy? Today! Understand that there is experimental evidence that supports the idea that light is both a wave AND a particle. Constructive Interference: Before During Destructive Interference: Before During 2 problems that could not be explained if light only acted as a wave. 1.) Emission of Light by Hot bodies: Characteristic color given off as bodies are heated: red yellow white If light were a wave, energy would be given off continually in the infrared (IR) region of the spectrum. 2.) Absorption of Light by Matter = Photoelectric Effect Light can only cause electrons to be ejected from a metallic surface if that light is at least a minimum threshold frequency . The intensity is not important. If light were only a wave intensity would be the determining factor, not the frequency! Please turn your photoelectric effect assignment into the box. Tomorrow is the Flame Test Lab. We will meet in Rm. 126 To participate, you MUST be wearing proper lab attire. When an object loses energy, it doesn’t happen continuously but in small packages called “quanta”. “Quantum”-a definite amount of energy either lost or gained by an atom. “Photon”-a quantum of light or a particle of radiation. Excited State: Higher energy state than the atom normally exists in. Ground State: Lowest energy state “happy state” Line Spectrum: Discrete wavelengths of light emitted. 2 Types: 1.) Emission Spectrum: All wavelengths of light emitted by an atom. 2.) Absorption Spectrum: All wavelengths of light that are not absorbed by an atom. This is a continuous spectrum with wavelengths removed that are absorbed by the atom. These are shown as black lines for absorbed light. Continuous Spectrum: All wavelengths of a region of the spectrum are represented (i.e. visible light) Hydrogen’s spectrum can be explained with the wave-particle theory of light. Niel’s Bohr (1913) 1.) The electron travels in orbits (energy levels) around the nucleus. 2.) The orbits closest to the nucleus are lowest in energy, those further out are higher in energy. 3.) When energy is absorbed by the atom, the electron moves into a higher energy orbit. This energy is released when the electron falls back to a lower energy orbit. A photon of light is emitted. 1. 2. 3. Bohr models are used to predict reactivity in elements. Reactivity refers to how likely an element is to form a compound with another element. When looking at Bohr models, we look at its valence electrons (the electrons on the last energy level) to determine reactivity. 1. 2. 3. 4. 5. Draw the nucleus. Write the number of neutrons and the number of protons in the nucleus. Draw the first energy level. Draw the electrons in the energy levels according to the rules (on the next slide). Make sure you draw the electrons in pairs. Keep track of how many electrons are put in each level and the number of electrons left to use. Each electron shell can hold a certain number of electrons Electron shells are filled from the inside out Noble Gases have full outer electron shells All other elements have partially filled outer electron shells Electron Shell 1 Number of Electrons 2 2 8 3 8 4 18 5 18 6 32 7 32 In order to draw Bohr models of these elements, you must first determine the number of protons, neutrons, and electrons. Once you have found this information, follow the directions to draw your model. 6 C Carbon 12.011 6 6 6 Protons: _____ Neutrons: _____ Electrons: ______ 2 How many energy shells will this have? ____ 4 How many valence (outer) electrons does this element have? ____ Bohr Model: 16 S Sulfur 32.066 16 16 16 Protons: _____ Neutrons: _____ Electrons: ______ 3 How many energy shells will this have? ____ 6 How many valence (outer) electrons does this element have? ____ Bohr Model: 3 Li Lithium 6.941 3 4 3 Protons: _____ Neutrons: _____ Electrons: ______ 2 How many energy shells will this have? ____ 1 How many valence (outer) electrons does this element have? ____ Bohr Model: 10 Ne Neon 20.180 10 10 10 Protons: _____ Neutrons: _____ Electrons: ______ 2 How many energy shells will this have? ____ 8 How many valence (outer) electrons does this element have? ____ Bohr Model: 15 P Phosphorus 30.974 15 Protons: _____ 16 15 Neutrons: _____ Electrons: ______ 3 How many energy shells will this have? ____ 5 How many valence (outer) electrons does this element have? ____ Draw the Bohr Model: 11 Na Sodium 22.990 11 Protons: _____ 12 11 Neutrons: _____ Electrons: ______ 3 How many energy shells will this have? ____ 1 How many valence (outer) electrons does this element have? ____ Draw the Bohr Model: Lyman Series-electrons falling to the 1st orbit, these are highest energy, _____ region. Balmer Series- electrons falling to the 2nd orbit, intermediate energy, _______ region. Paschen Series-electrons falling to the 3rd orbit, smallest energy, ______ region. En = (-RH) 1/n2 En = energy of an electron in an allowed orbit (n=1, n=2, n=3, etc.) n = principal quantum number (1-7) RH = Rydberg constant (2.18 x 10-18 J) When an electron jumps between energy levels: ΔE =Ef – Ei By substitution: ΔE = hν = RH(1/ni2 - 1/nf2) When nf > ni then ΔE = (+) When nf < ni then ΔE = (-) DeBroglie (1924)-Wave properties of the electron was observed from the diffraction pattern created by a stream of electrons. Schrodinger (1926)-Developed an equation that correctly accounts for the wave property of the electron and all spectra of atoms. (very complex) DO NOT WRITE THIS IN YOUR NOTES! Rather than orbits we refer to orbitals. These are 3-dimensional regions of space where there is a high probability of locating the electron. Heisenberg Uncertainty Principle-it is not possible to know the exact location and momentum (speed) of an electron at the same time. Quantum Numbers-4 numbers that are used to identify the highest probability location for the electron. 1.) Principal Quantum Number (n) States the main energy level of the electron and also identifies the number of sublevels that are possible. n=1, n=2, n=3, etc. to n=7 2.) Azimuthal Quantum Number (l) Values from 0 to n-1 Identifies the shape of the orbital l=0 l=1 l=2 l=3 s p d f sphere dumbbell 1 orbital 3 orbitals 4-4 leaf clovers & 1-dumbbell w/doughnut5 orbitals very complex 7 orbitals 3.) Magnetic Quantum Number (ml) Values from –l l States the orientation in space (x, y, z) ml = 0 ml = -1, 0, +1 ml = -2,-1,0,+1,+2 ml = -3,-2,-1,0,+1+2,+3 s p d f only 1 orientation 3 orientations 5 orientations 7 orientations 4.) Spin Quantum Number (ms) Values of +1/2 to -1/2 States the spin of the electron. Each orbital can hold at most 2 electrons with opposite spin.