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# CHM 113 Week 8 Notes Bio 113

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This 5 page Class Notes was uploaded by Andrew Notetaker on Saturday October 8, 2016. The Class Notes belongs to Bio 113 at Arizona State University taught by Udo Savalli in Spring 2015. Since its upload, it has received 3 views. For similar materials see Dinosaurs in Science at Arizona State University.

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Date Created: 10/08/16

Chapter 6: Electronic Structure of Atoms Monday, October 3, 2016 2:56 PM 6.1 The Wave Nature of Light Visible light is a only a small part of a broad spectrum of energy known as electromagnetic radiation. Electromagneticradiation, including visible light, travels through space at a velocity of 3.0 x 10 m/s (speed of light) All types of radiation has wavelike properties. The distance between corresponding points on adjacent waves is the wavelength ( ) The number of waves passing a given point per unit of time is the frequency (v). For waves traveling at the same velocity,the speed of light c is shorter than the wave length, the greater the frequency. (inverse relationship) C = V 3.0 x 10 = 700 nm (nm 10 m/nm) x v Shorter wavelengths/higher frequencies = higher energy Longer wavelengths/lowerfrequencies = lower energy Quantization of Energy Blackbody Radiation Light emitted from heated metal atoms is not continuous it is emitted in discrete amounts (quanta). This emitted energy is related to a wavelength of the emitted energy by: E = hν = hc/λ h is Planck’s constant, 6.63 x 1034J•s Photoelectric Effect When light is shined on metals,particles are ejected by the surface of the material at a minimum frequency. The frequency of light must be above some minimum energy (wavelength). Light is made up of photons Energy exist not as a continuum but at discrete small steps or levels. Dual nature of electromagnetic radiation EMR has properties of waves and particles Line Spectra & Bohr Model White light is composedof all wavelengths of visible light. Diffraction produces a continuous spectrum. Diffraction of light emitted from pure elements results in a line spectrum, not a continuous spectrum. Not only do pure elements emit light (Emission)at discrete, specific wavelengths (monochromatic)but they absorb electromagneticradiation at those same specific wavelengths (Absorption). Emission spectra is when an electron drops from a higher energy level to a lower energy level, emitting a wavelength. Absorption spectrum is when light passes through a cold gas and atoms in the gas absorb the frequencies, showing an absence of light in the spectrum. Rydberg Equation: CHM 113 Lecture Page 1 1913, Niels Bohr proposed the model of the hydrogen atom. 1. One electron circles the nucleus in orbits of certain radii. 2. An electron permitted in orbit has a specific allowed energy 3. Energy is emitted or absorbed as the electron moves from one energy state to another. The energy of each level was calculated using 2 -18 2 E = (-hcRH) = (1/n )= (-2.18 x 10 J) (1/n ) h= Planck's constant C =speed of light R = Rydberg's constant H N is called the PrincipleQuantum number (1,2,3..) Balmer series n =f-Visible Lyman series n =f-Ultraviolet Paschen Series n = 3-Infrared f Limitationsof the Bohr Model- 1. Cant explain why a negatively charged electrondoes not fall into the nucleus 2. Cant explain the spectra of multi-electron atoms 3. Electrons have wavelike properties, which prohibits describing electrons as circling in orbits Electrons exist only at certain discrete energy levels Energy is involved in electron transitions between energy levels 2 de Broglie used Einstein’s equation: E = mc Combining : m = E/c 2 and E = h c / Gives: m = h / c and = h / mc Where mv for any object with mass is called momentum What is the wavelength of an electron that travels at 6.0 x 10 m/s and has a mass of 9.11 x 10 -31kg? h = 6.63 x 10 -34J•s (J = kg • m / s )2 -34 2 -31 6 = (6.63 x 10 kg • m / s) / (9.11 x 10 kg)(6.0 x 10 m/s) − = 1.21 x 10 10m = 0.121 nm (x-rays) Heisenberg claimed that because of the wave nature of electronsit was impossible to know the precise position and velocity of an electron in an atom. He also claimed that because of the nature of electrons it was impossible to know the position and velocity of an electron at the same time. Δx • mΔv ≥ h/4p Where Δx represents the uncertainty of position and Δv represents the uncertainty of velocity. Bohrs: electrons exist in discrete energy levels De Broglie's: all matter has wave properties Heisenberg: the position/velocitycan never be determined simultaneously Schrodinger proposed the mathematicallybased model of the atom. This gave a set of equations that determinethe probability of finding the electron at any point in space. The wave equations are designated with a lowercaseGreek psi ( CHM 113 Lecture Page 2 Schrodinger proposed the mathematicallybased model of the atom. This gave a set of equations that determine the probability of finding the electron at any point in space. The wave equations are designated with a lowercaseGreek psi ( The square of a wave equation gives a probability density map where an electron has a likelihood of being an any given time. Solving the wave equations gives a set of wave functions or orbitals. Each orbital is described by a set of 3 quantum numbers n, l and m . 1 n-Principal Quantum number (Energy, Distance from nucleus) l- Azimuthal Quantum number (Shape) M 1 Magnetic quantum number (direction) Principal quantum number describes the energy level of the orbital and average distance of the electrons n are integers > or = 1. Angular momentum quantum number (l) Defines the shape, allowed values are integers ranging from 0 to n-1. s,p,d,f designations are called subshells Magnetic Quantum number (m ) 1 The three dimensional orientation of the orbital Allowed values of m are integers ranging from -l to l: 1 -l< or = m1< or = +l There can be up to 1s orbital 3p orbitals 5d orbitals 7f orbitals S orbitals have 1 orientation(sphere) P orbitals have 3 (x,y,z) D orbitals have 5 Orbitals with the same value of n form a shell Different orbital types within a shell are subshells. Within one shell there are n possible orbitals. Orbitals with the same value of n and l are in the same subshell. Each orbital can contain a max of two electrons. Orbitals on the same energy level have the same energy. The value of l for s orbitals is 0. They are spherical in shape, the radius increases with the value of n Radial Probability distribution graphs show electrons are at increased distance from the nucleus as n increases. Regions of low probability in n>1 s orbitals are called nodes. Nodes represent areas with 0 probability of finding electrons,these increase with n. The value of l for p orbitals is 1. These are dumb-bell shaped. Each orbital has two lobes with a node between (3p orbitals in a subshell) As n increases, p orbitals acquire nodes. The value of l for a d orbital is 2. Four of the five d orbitals have 4 lobes. The other resembles a p orbital with a doughnut around the center. First shell n=1, first row of periodic table has 1 subshell 1s- 1 orbital 2nd shell (n=2, second row) has 2 subshells (s & p) 2s-1 orbital 2p- set of 3 orbitals 3rd shell (n=3,third row) 3 subshells (s,p and d) 4th shell (n=4,4th row) has 4 subshells (s,p,d, and f) Electron Spin and the Pauli Exclusion Principle The 4th quantum number is the spin quantum number m s M =+1/2 or -1/2 s No two electronsin the same atom can have exactly the same energy, therefore no two electrons in the same atom can have identical sets of quantum numbers Electron Configurations CHM 113 Lecture Page 3 No two electronsin the same atom can have exactly the same energy, therefore no two electrons in the same atom can have identical sets of quantum numbers Electron Configurations Element Configuration H 1s1 He 1s2 Li 1s 2s1 Be 1s 2s2 B 1s 2s 2p 1 1 2 2 C 1s 2s 2p 1 2 3 N 1s 2s 2p 1 2 4 O 1s 2s 2p 1 2 5 F 1s 2s 2p Ne 1s 2s 2p 6 Diagonal rule shows how to fill the orbitals 1s 2s, 2p 3s, 3p, 3d 4s, 4p, 4d, 4f 5s, 5p, 5d, 5f, (5g) 6s, 6p, 6d, 6f, (6g, 6h) Valence Electrons Electrons in the outermostshell- electrons with highest principle quantum number, n Electrons will always be in s or p subshells Aufbau Principle Electrons fill into the lowest energy orbitals first Hund's Rule Electrons fill singly into different orbitals of the same sublevel with the same spin before pairing up. Pauli Exclusion Principle Two electrons can occupy each orbital, they must have opposite spin Rather than writing out all the shells, previous filled shell (noble gas) and add on: 2 Ca: [Ar] 4s Cl: [Ne] 3s 3p 5 2 2 Ti:[Ar] 4s 3d Electrons populate the 4s orbital before the 3d orbitals because of lower 4s energy. Max number of electronscan occupy each subshell= number of columns for each type of subshell s-block,2 columns, max of 2 electrons per s orbital D-block, 10 columns, max of 10 electrons per d orbital Some exceptions to the Aufbae order Expected electron configurations for Cr and Cu? Cr: Predicted 1s 2s 2p 3s 3p 4s 3d2 4 2 2 6 2 6 1 5 Actual: 1s 2s 2p 3s 3p 4s 3d Also true for Mo and W. 2 2 6 2 6 2 9 Cu: Predicted- 1s 2s 2p 3s 3p 4s 3d Actual: 1s 2s 2p 3s 3p 4s 3d1 10 Also true for Ag and Au Elements in the same group share physical and chemical properties because they have the same valence electron configurations Commonions form by gaining or losing valence electrons in order to have the same number of electrons as the closest noble gas. Cations: remove electrons in the reverse of the Aufbau order CHM 113 Lecture Page 4 Cations: remove electrons in the reverse of the Aufbau order 3+ Al Mg 2+ + Na These ions are isoelectronic. Anions: Add electrons in the Aufbau order F- 2- O N 3- These ions are also isoelectronic. Transition metal cations removeelectrons from the p/s orbitals first, then the (n-1)d orbitals. CHM 113 Lecture Page 5

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