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# Chem 103 electrons and light CHEM 103 - 002

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This 9 page Class Notes was uploaded by Karlee Nelsen on Friday November 6, 2015. The Class Notes belongs to CHEM 103 - 002 at University of Wisconsin - Eau Claire taught by Sanchita Hati in Fall 2015. Since its upload, it has received 12 views. For similar materials see General Chemistry I in Chemistry at University of Wisconsin - Eau Claire.

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Date Created: 11/06/15

L IGHT AND E LECTRON M OTION Variables for This Chapter and the Motion of Electrons R = Rydberg constant 1.0974 x 10 m -1 c = speed of light 2.99792458 x 10 m/s -34 h = Planck’s constant 6.6260693 x 10 Js λ wavelength nm ν frequency unit/s n = integer of electron energy level 1, 2, 3,.∞+ E= energy Joules (J) ∆ change l =Orbital Angular Momentum Quantum Number 0, 1, 2,.. n-1 m =Magnetic Quantum Number ± l l Increasing frequency and Energy Increasing wavelength Max Planck – discovered that oscillations had only certain frequencies that could occur. All frequencies were whole number multiples of one base frequency. - Therefore the radiation given off can only have certain Energies E=nhν n- must be positive (energy level of electron) o Quantized- only certain whole number energies ≥ 1 can exist h – proportionality constant (Planck’s Constant) -34 o 6.6260693 x 10 Js ℎ???? ???? = ℎν = ???? Wave-particle duality – electromagnetic radiation can behave as a particle and a wave but not be observed as both simultaneously. Photon= light particle When electrons move from higher level (energy) to a lower level (energy), light is emitted - Similar to Potential energy, o Higher levels (orbits)(energy) correspond to greater potential to create light as the drop levels o The potential energy in this case is converted into light energy rather than traditional kinetic energy - Energy is put in in order to move the electron to a higher level and let out when the electron moves to a lower level - Energy levels increase as n increases n=1 n=2 1 1 1 = ????[ 2 − 2] ????ℎ???????? ???? > 2 n=3 λ 2 ???? 7 -1 n=4 R = Rydberg constant = 1.0974 x 10 m Bohr’s Model of the Atom o Electrons move in circular orbits around the nucleus o Only certain orbits are allowed (quantized orbits) o n refers to different allowed orbits o as n increases, the radius and energy of the orbit increase o electrons are permitted or orbit only at a specific energy called the “allowed energy state” because energy is quantized o energy is lowest when n=1 and increases with increasing n value Balmer Equation – can be used to find the wavelength of emitted light when only the energy level of the electron is known. 1 1 1 = ????[ − ] ????ℎ???????? ???? > 2 ???? 22 ????2 R = Rydberg constant = 1.0974 x 10 m-1 n = energy level of the electron λ = wavelength To find the total energy of an electron in the n energy level Rydberg Plank’s constant constant Potential energy Speed of Rhc light of electron in the E =n − 2 nth level ???? Principle quantum number (energy level) n > 1 are excited states of electrons 7 -1 R = Rydberg constant = 1.0974 x 10 m n = Principal quantum number corresponding to the energy level of the electron λ = wavelength -34 h = 6.6260693 x 10 Js c = 2.99792458 x 10 m/s How an electron becomes excited Energy Energy emitted absorbed Ground state Excited state Ground state - Electron energy when in orbit is negative because electrons in orbit have less energy that when they are free - n=1 corresponds to the ground state for an atom because it is the lowest possible energy level - n>1 is as excited state - as distance from the nucleus increases so does the Energy (becomes less negative) ∆ E when the electron changes energy levels is shown by ???? − ???? = ∆???? = [− ????ℎ???? ] − [− ????ℎ???? ] ???????????????????????? ???????????????????????????????? ???? ???????????????????? ???? ???????????????????????????? 7 -1 R = Rydberg constant = 1.0974 x 10 m n = Principal quantum number corresponding to the energy level of the electron h = 6.6260693 x 10-34Js 8 c = 2.99792458 x 10 m/s E = energy level + (positive) value is absorbtion - (negative) value is emission ????ℎ???? ????ℎ???? ???? ???????????????????????? ????????????????????????????????????∆???? = [− 2 ] − [− 2 ] ???????????????????????? ???????????????????????????????? ???? = ℎ???? ???? = ℎ???? ???? = ???? ???? ∆???? ℎ energy wavelength frequency DeBroglie – matter has wave properties. For large objects position and velocity can be determined because the wavelength is so small. In the quantum world the position and velocity can never be determined without error and NEVER both with high n=3 certainty. n=2 - A large mass OR velocity produces such a small wavelength that the vibrations are not noticeable with today’s measurement techniques - ℎ e n=1 ???? = ???????? λ – wavelength m – mass of object -34 h – 6.6260693 x 10 (Js) v – velocity (m/s) Heisenberg (1932 Nobel Prize) Uncertainty Principle – It is not possible to know both position and momentum of an electron with certainty. If the position is known with great certainty, then momentum is known with little certainty and if momentum is known with great certainty then position is known with little certainty. ℎ ∆???? ∗ ∆???? ≥ 4???? x- position p – momentum Erwin Schrödinger – discovered the property of electrons based on wave characteristics (ψ) also known as quantum mechanics (or wave mechanics) - electrons vibrate on wavelengths of standing waves and only certain frequencies or vibrations are allowed for these waves meaning that they are quantized - the wave function is described as (ψ) and has no physical meaning 2 - however ψ is the 2robability of finding as electron in a certain area of space proportional to ψ called the electron density - probability is highest close to the nucleus wave functions are called orbits Standing Waves ½ λ 1 λ 3/2 λ Bohr model of an atom used n to describe an orbital, while Schrödinger uses three quantum numbers n, l, and mlto describe the orbit. n – Principle Quantum Number – valued [1, ∞) and a whole number - greater n = greater E and size of orbital l – Orbital Angular Momentum Quantum Number - any value [0, 1, 2,.. n-1), has a corresponding letter l value 0 1 2 3 Subshell label s p d f m l Magnetic Quantum Number - related to the spatial orientation of orbitals within a subshell valued ± l Quantum Number interrelationships and Orbital Information Implied Shell Sub-shell Individual orbital Principal Angular Size of orbitals Quantum Momentum Magnetic Quantum Number and the Type of Orbitals in the increase Number Quantum Number Subshell Number s as n n= number of subshells value n l m l Number of Orbitals in Shell = n2 increase Number of Orbitals in Subshell 2l+1 1 0 0 One 1s orbital (one orbital of one type in the n=1 shell) One 2s orbital 2 0 0 Three 2p orbitals 1 -1, 0, +1 (four orbitals of two types in the n=2 shell) 0 0 One 3s orbital 3 1 -1, 0, +1 Three 3p orbitals 2 -2, -1, 0, +1, +2 Five 3d orbitals (nine orbitals of three types in the n=3 shell) One 4s orbital 0 0 Three 4p orbitals 1 -1, 0, +1 4 2 -2, -1, 0, +1, +2 Five 4d orbitals 3 -3, -2, -1, 0, +1, +2, +3, Seven 4f orbitals (16 orbitals of four types in the n=16 shell) Nodal Surface – surface on which the probability of finding an electron is zero l value = the number of nodal surfaces in a configuration - there is never an impenetrable surface within which the electron can be contained - the probability of finding the electron isn’t constant throughout the volume - “electron cloud” implies the electron is a particle however it is viewed as a wave in this instance S orbitals – are spheres containing the probability of finding the electron within the sphere - Size of the sphere is greater for greater n values P orbitals – are dumbbell shaped and pass through the nucleus though there is a nodal surface that also passes through the nucleus D orbital – roughly cone shaped but can vary F orbital – electron density lies in 8 regions of space separated by the planes electron clouds with corresponding orbitals Stern- Gerlach experiment- discovered that electrons spin at a quantized level m = + ½ m = - ½ s s α β Favored electron Less favored electron arrangement arrangement Pauli Exclusion Principle – no two electrons can have the same four Quantum numbers Ex. 4p , 4p , 4p , and each can hold two electrons per orbital x y z 4p x 4p y 4pz Examples of orbitals filled by electrons with atoms H He Li Be

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