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Electrons in Atoms Chapter 5 Chem X “Indeed, to be useful, a model must be wrong, in some respects – else it would be the thing itself.” Models are never fixed, but change to explain new observations. The Nuclear Atom and Unanswered Questions Rutherford’s 1. 2. In model lacked details about… how e- occupy space around the nucleus Why the negatively charged electrons are not pulled into the atom’s positive nucleus early 1900’s…scientists observed that certain elements emitted visible light when heated in a flame The Nuclear Atom and Unanswered Questions In early 1900’s…scientists observed that certain elements emitted visible light when heated in a flame Analysis showed that an element’s chemical behavior is related to the arrangement of electrons in its atoms. More later…. We must first discuss light. Wave Nature of Light Radiation – a form of energy that exhibits wavelike behavior as it travels through space. Electromagnetic Travels at the speed of light through a vacuum. Wave Nature of Light Wave Nature of Light Electromagnetic wavelength – frequency Amplitude - Radiation – Wave Nature of Light spectrum – all the colors of the visible spectrum Continuous Wave Nature of Light Ex. 5-1 Calculating Wavelength of an EM Wave – Microwaves are used to transmit information. What is the wavelength of a microwave having a frequency of 2.44 x 109 Hz? 8.72 x 10-2 m Particle Nature of Light The 1. 2. wave model of light cannot explain… Why heated objects emit only certain frequencies of llight at a given temperature Or why some metals emit electrons when colored light of a specific frequency shines on them The Particle Nature of Light The quantum concept – waves don’t explain the emission of different wavelengths of light at different temperatures Max Plank (1858-1947) – concluded matter can gain or lose energy only in small, specific amounts called quanta Quantum – the minimum amount of energy that can be gained or lost by an atom. The Particle Nature of Light Prior experience had lead scientists to believe that energy could be absorbed and emitted in continual varying quantities, with no minimum limit to the amount. Actually, the water’s temperature changes by infinitesimal steps as its molecules absorb quanta of energy E = hn Particle Nature of Light effect – electrons, called photoelectrons, are emitted from a metal’s surface. Metals only ejected electrons above a certain minimum frequency. Photoelectric Particle Nature of Light Photon – a particle of electromagnetic radiation with no mass that carries a quantum of energy Einstein calculated Ephoton = hn KEY: Einstein was able to explain the photoelectric affect by…giving electromagnetic radiation particlelike properties. Atomic Emission Spectra – the set of frequencies of the electromagnetic waves emitted by atoms of the element. It is always several individual lines of color, NOT a continuous range of colors. Each element’s atomic emission spectrum is unique. Continuous Spectrum (top) Various elements’ emission spectra Atomic Emission Spectra The fact that only certain colors appear in an elements atomic emission spectrum means that… only certain specific frequencies of light are emitted. i.e. Only photons having certain specific energies are emitted. 5.2 Quantum Theory and the Atom Bohr Model of the Atom Answered the question why…atomic emission spectra are made up of only certain frequencies Correctly predicted… the frequencies of the lines in hydrogen’s atomic emission spectrum Energy States of hydrogen Proposed that… Ground State – lowest allowable energy state of an atom Excited state – when an atom gain energy e- jumps to higher E level Ionized state – (in class) Bohr’s Model of the Atom He proposed that the atom has only certain allowable energy states. ground state – lowest possible excited state – atom gain energy ionized state - atom gains enough E to lose an e- Bohr suggested that.. The single electron in a hydrogen atom moves around the nucleus in only certain allowed circular orbits. The smaller the orbit, the lower the atom’s energy state. Bohr’s Model of the Atom In order to go from one energy level to another, electrons can only gain/lose energy in certain amounts. This is the idea of “quantized” energy. Absorption Emission An explanation of hydrogen’s line spectrum n=1 ΔE = Efinal - Einitial = Ephoton = hn See next slide for bigger figures… An explanation of hydrogen’s line spectrum An explanation of hydrogen’s line spectrum n=1 n=5 n=4 n=3 ΔE = Efinal - Einitial = Ephoton = hn n=2 n=1 Atomic Emission Spectra Atomic Emission - Compound Flame Color LiCl NaCl KCl CaCl2 SrCl2 CuCl2 Electrons as Waves Louis De Broglie ( 1892-1987 ) - 1924 l = h__ mv Key: electrons have wave properties. Quantum Mechanical Model of the Atom Heisenberg Uncertainty Principle – it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time The Schrodinger wave equation (1926) Quantum mechanical model of the atom – model where electrons are treated as waves - limits the electron’s energy to certain values - Each solution is a wave function and is related to the probability of finding the electron within a particular volume of space around the nucleus Atomic orbital – 3 dimensional region around the nucleus that describes the electron’s probable location. (See fig. 5-13 – e- density) Joke… Heisenberg, Schrodinger and Ohm are in a car. They get pulled over. Heisenberg is driving and the cop asks him "Do you know how fast you were going?" "No, but I know exactly where I am" Heisenberg replies. The cop says "You were doing 55 in a 35." Heisenberg throws up his hands and shouts "Great! Now I'm lost!" The cop thinks this is suspicious and orders him to pop open the trunk. He checks it out and says "Do you know you have a dead cat back here?" "We do now!" shouts Schrodinger. The cop moves to arrest them. Ohm resists. Hydrogen’s Atomic Orbitals Orbital – contain 90% of e’s total probability Principal quantum numbers ( n ) Principal energy levels – specified by n Energy sublevels – as n increases, the number of sublevels increases Hydrogen’s Atomic Orbitals See fig. 5-14 (Stadium seating) Hydrogen’s Atomic Orbitals Principal quantum number (n) 1 2 3 4 Sublevels (types of orbitals) present Total number of Number of orbitals related orbitals related to principal to sublevel energy level (n2) 5.3 Electron Configurations Ground State Electron Configuration: configuration – the arrangement of electrons in an atom Electron 5.3 Electron Configurations The 1. 2. 3. 4. aufbau principle – All orbitals related to an energy sublevel are of equal energy In a multi-electron atom, the energy sublevels within a principal energy level have different energies In order of increasing energy, the sequence of energy sublevels within a principal energy level is s, p, d, and f Orbitals related to energy sublevels within one principal energy level can overlap orbitals related to energy sublevels within another principal level. 5.3 Electron Configurations Pauli Exclusion Principle– A maximum of two electrons may occupy a single atomic orbital, but only if the electrons have opposite spins. 5.3 Electron Configurations Hund’s Rule– single electrons with the same - spin must occupy each equal energy orbital before additional electrons with opposite spins can occupy the same orbitals. Quantum Mechanical Model of the Atom Heisenberg Uncertainty Principle – Schrodinger Wave Equation – waves the Electrons are treated as ________, equation can tell you the coordinates most probable to find an electron of a given energy atomic orbital – Joke… Heisenberg, Schrodinger and Ohm are in a car. They get pulled over. Heisenberg is driving and the cop asks him "Do you know how fast you were going?" "No, but I know exactly where I am" Heisenberg replies. The cop says "You were doing 55 in a 35." Heisenberg throws up his hands and shouts "Great! Now I'm lost!" The cop thinks this is suspicious and orders him to pop open the trunk. He checks it out and says "Do you know you have a dead cat back here?" "We do now!" shouts Schrodinger. The cop moves to arrest them. Ohm resists. Quantum Mechanical Model of the Atom Energy Level 1 2 3 4 # electrons # orbitals needed Types & #’s of orbitals Orbitals s-orbitals p-orbitals d-orbitals f-orbitals Quantum Mechanical Model of the Atom ___ 4s ___ ___ ___ 4p ___ ___ ___ ___ ___ 4d ___ ___ ___ ___ ___ ___ ___ ___ ___ 3s 3p 3d ___ ___ ___ ___ 2s 2p ___ 1s ___ ___ ___ ___ ___ ___ ___ 4f Quantum Mechanical Model of the Atom ___ ___ ___ 6p ___ 5s 5p 4d ___ ___ ___ ___ ___ 3d 4s ___ ___ ___ ___ 3p 3s ___ ___ ___ ___ 2p 1s ___ ___ ___ ___ ___ ___ ___ ___ ___ ___ 4f ___ ___ ___ 4p 2s ___ ___ ___ ___ ___ ___ ___ 5d ___ 6s ___ ___ ___ ___ ___ Link to Notebook Periodic table Electron Configuration bau principle – an electron will occupy the _____________________ that can receive it. auf exclusion principle – (no two electrons can have the same set of four quantum numbers) i.e. If two electrons are in the same orbital, they must have _______________. Pauli Electron Configuration Hund’s Rule – the lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals. i.e. as electrons fill into degenerate orbitals one must go into each orbital will ___________ spins before you _______________________. Electron Configuration Noble gas Short-hand - a shortened econfiguration that starts with the previous noble gas Orbital Notation - uses lines to denote each orbital with electrons written as arrows (the orbital types 4s etc. are sometime written underneath the lines) Electron Configuration Hydrogen Scandium Helium Zinc Lithium Barium Carbon Thallium Argon Radon Calcium Electron Configuration Hydrogen 1s1 Helium 1s2 Lithium 1s22s1 Carbon 1s22s22p2 Argon 1s22s22p6 3s23p6 Calcium 1s22s22p6 3s23p64s2 Electron Configuration Scandium 1s22s22p6 3s23p64s23d1 Zinc 1s22s22p6 3s23p64s23d10 Barium 1s22s22p6 3s23p64s23d104p65s24d105p66s2 Thallium 1s22s22p6 3s23p64s23d104p65s24d105p66s2 4f145d106p1 Radon 1s22s22p6 3s23p64s23d104p65s24d105p66s2 4f145d106p6 Electron Configuration Electron configuration - shows how many electrons are in each type of orbital Fluorine 1s22s22p5 Sulfur 1s22s22p63s23p4 Germanium 1s22s22p63s23p64s23d104p2 Radon 1s22s22p63s23p64s23d104p65s24d105p6 6s24f145d106p6 Electron Configuration Noble gas Short-hand - a shortened econfiguration that starts with the previous noble gas Fluorine [He]2s22p5 Sulfur [Ne] 3s23p4 Germanium [Ar] 4s23d104p2 Radon [Xe] 6s24f145d106p6 Electron Configuration Orbital Notation - uses lines to denote each orbital with electrons written as arrows (can use long or shorthand) (the orbital types 4s etc. are sometime written underneath the lines) [He]2s2p [Ne] 3s3p [Ar] 4s3d 4p . [Xe] 6s4f 5d 6p Radon [Xe] 6s24f145d106p It’s usually only necessary to write orbital notation for the highest energy orbital (as done with radon). Fluorine Sulfur Germanium Radon Electron Configuration Review Write the electron configuration for each of the following: 1. 1. 2. 3. 4. Argon Carbon Osmium Barium Electron Configuration Review 1. What is the electron configuration for Xenon? 2. What is the orbital notation for Xenon? 3. What’s Xenon’s highest energy level? What’ Xenon’s highest energy orbital? 4. Electron Configuration Review 1. How many valence electrons do each of the following elements have? 1. 2. 3. 4. 5. Chlorine Calcium Aluminum Carbon Tin