Atomic Theory I Chap 5 Wave Nature of Light Atomic Theory Learning Targets 5.5 Know the history of the development of the atomic theory 5.12 Apply quantum mechanics Light and Elements • Since ancient times that elements emitted their own specific colors when vaporized • That light can be separated into a line spectra by the use of a prism. What is light? • Electromagnetic radiation (a.k.a. light) is a form of energy that behaves like a wave. • Waves have 3 properties: – Wavelength = Distance between consecutive peaks or troughs – Frequency = How many peaks pass a given point per time – Speed = How fast a given peak moves through space (3x108 m/s) How are the wavelength & energy related? • Shorter wavelength is higher frequency which means greater energy Let’s look at visible light • Purple light is of slightly higher energy than red light (far left versus far right of the spectrum) • Purple light has a higher frequency and a shorter wavelength than red light. Light behaves like a wave… …but it also showed particle properties • Debate on light properties raged until the 1900’s because light: – behaves like a wave AND – behaves like a particle Photoelectric Effect • Light shined on metal could eject electrons – Only certain wavelengths/frequencies ejected electrons • Demonstrated Particle Nature of Light! OK so why are Light Emission Spectra Unique for each Atom Quantum Theory • Planck suggested since only certain wavelengths of light are given off that the energy of an element is restricted to “pieces” of particular sizes called a quantum (plural is quanta) Max Planck & Neils Bohr Einstein said that light are packets (particles) of a specific amount of energy called photons. How can you explain the spectrum of an element such as hydrogen? • How can something with a positive nucleus and negative electrons moving around it only give off energy in small discrete areas • We need some help ……along came Niels Bohr Bohr’s Contribution • Bohr proposed a planetary model of the atom: electrons orbit the nucleus in distinct energy levels (n = 1, 2, 3, ...). • The lowest energy level is called the ground state (n = 1) • When an electron absorbs energy it is promoted to an excited state (n = 2,3,4…) Contributions restated • Energy increases as shell (level) number increases. • Light can be absorbed to make an electron jump to a higher level (excited). • Light is emitted when an electron falls to a lower energy level • Bohr predicted all the lines seen for the spectrum of Hydrogen. • Bohr’s model worked for hydrogen, but he was unable to explain the spectra of atoms with more than one electron. • However, his ideas led to the current model of the atom we use today. A New Approach to the Atom • Quantum-Mechanical Model: – Energy of electrons is quantized (Planck) – De Broglie reasoned: if light had a dual nature matter could also (wave & particle properties). – Heisenberg Uncertainty Principle • It is impossible to know the exact position as well as the momentum of an electron at any given instant Schrodinger • developed equation for location of electron Atomic orbital • Exists around the nucleus • Location where you have a chance of finding the electron • Orbitals have different shapes Quantum Numbers • Quantum numbers describe the position of an electron. 1. Principle quantum number (n): • Positive numbers starting with 1 (n = 1,2,3,….) • n=1 is lowest energy level (ground state) • Designates the energy level (higher n, farther from nucleus, higher energy) Quantum Numbers 2. Sublevel tell shape of orbital Principle energy level divided into sublevels Designated by letters (s,p,d,f) Letter s p d f Shape spherical dumbbell complex complex # in a level 1 3 5 7 Increasing energy • • f-orbitals How do levels and sublevels fit? • # of sublevels in a main level equals the principle quantum number – level 1 = has 1 sublevel = s – level 2 = has 2 sublevels = s & p – level 3 = has 3 sublevels = s, p, & d – level 4 = has 4 sublevels = s, p, d, & f – level 5 = has 5 sublevels = s, p, d, f & g How do electrons fit into sublevels? • Pauli Exclusion Principle: an individual orbital can hold up to 2 electrons Sublevel # of orbitals Maximum number of electrons in sublevel s 1 2 p 3 6 d 5 10 f 7 14 Electron Configuration • A distribution of atom’s electrons among levels, sublevels, and orbital's •Aufbau Principle – electrons added to lowest energy orbital's available until all electrons of atom have been used Electron Configuration • Beryllium (4 electrons) Orbital Type # of Electrons »1s2 2s2 Energy Level (n) Electron Configurations • Start in 1st energy level • Fill electrons into orbitals and levels starting with lowest and going to highest. • Note: starting in period 3, 4s fills before 3d, use diagonal diagram to get order correct Examples H 1s 1 Mg 1s2 2s2 2p6 3s2 V 1s2 2s2 2p6 3s2 3p64s2 3d3 Electron filling clues come from the periodic table? If you look at the table there are 7 rows: Row (periods) = 7 levels there are 4 major areas (blocks): Blocks = 4 sublevels there are 2 columns in the first area: s sublevel(1 orbital can hold 2 electrons) There are 6 columns in the next main block; p sublevel has 3 orbitals w/ 2 e- per orbital Electron filling clues come from the periodic table? the transition area has 10 columns: d sublevel(5 orbitals w/ 2 e- per orbital) 14 columns in the inner transition; f sublevel has 7 orbitals w/ 2 e- per orbital d (n-1) f (n-2) back • Example – Germanium (atomic # 32) Row # 11 22 33 44 55 66 77 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2 Short Hand • Valence electrons – The electrons in the highest energy level – they are involved in all reactions • Electron configuration short hand only has valence electrons written with core electrons represented by the Nobel gas from the row before. Electron Configuration Shorthand • Longhand Configuration Cl 17e 1s2 2s2 2p6 3s2 3p5 Core Electrons Valence Electrons • Shorthand Configuration (noble gas configuration) 2 5 Cl 17e [Ne] 3s 3p Shorthand Configuration • Example - Germanium 1 2 3 4 5 6 7 [Ar] 4s2 3d10 4p2 Try the noble gas shortcut: 53I 30Zn [Kr] 5s24d105p5 [Ar] 4s23d10 Electron Configs of Ions • Cations lose the valence electrons first (outer most) • Na 1s2 2s2 2p63s1 – loses 3s1 becomes Na+1 • Ca 1s2 2s2 2p63s2 3p6 4s2 – loses 4s2 to be Ca+2 Electron Configs of Ions (cont) • Anions(-) gain electrons to fill outer level • F 1s2 2s2 2p5 – gains e- F 1s2 2s2 2p6 - • O 1s2 2s2 2p4 – gains e- O 1s2 2s2 2p6 -2 Orbital Filling (Box) Diagrams • • • • Show distribution of electrons in orbitals. Electrons are represented by arrows Orbitals are represented by boxes Hund’s Rule – Within a sublevel, place one e- per orbital before pairing them. (“Empty Bus Seat Rule”) WRONG RIGHT # of Orbitals • • • • - s sublevel = 2 e = 1 orbital (1 box) p sublevel= 6 e = 3 orbitals (3 boxes) d sublevel=10 e =5 orbitals (5 boxes) f sublevel=14 e =7 orbitals (7 boxes) Other Electron Notations • Electron Configuration 2 2s2 2p4 O 1s 8 • Orbital Diagram O 8e- 1s 2s 2p What will be the correct filling? Example • Na 1s2 2s2 2p6 3s1 Example 47Ag = [Kr] [Kr] 5s2 4d9 Example 47Ag = [Kr] [Kr] 5s2 4d9 Electron Dot Structure • An atom’s valence electrons are represented by an electron dot structure. Na = 1s22s22p63s1 (1valence electron) Na Mg = 1s22s22p63s2 (2valence electrons) Mg the number of valence electrons is the same as the column number in periodic table (except helium) All the elements within a Group have the same number of valence electrons, therefore they will have the same number of dots in their Lewis Structures.
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