Exam 3 Study Guide November 11 2014 1 Chapter 6 Electronic Structure and the Periodic Table a Introduction i Electron configurations show the number of electrons in each energy level ii Orbital diagrams show the arrangement of electrons within orbitals b Light photon energies and atomic spectra i The Wave Nature of Light Wavelength and Frequency troughs which sink below it 1 Light can travel as a wave successive crests which rise above the midline and a Wavelength the distance between 2 consecutive crests or troughs b Frequency the number of wave cycles that pass a given point in a unit of time v c Equation i Measured in Hz hertz represents 1 cycle per second i v c 1 C is the speed of light 2 998x10 8 m s 2 should be expressed in meters 3 v should be expressed in Hz 2 3 The Electromagnetic Spectrum a b The color of gases and liquids are due to the selective absorption of certain components of visible light the wavelength reflected is the color shown Wavelength nm Color transmitted Color absorbed 400 400 450 Ultraviolet Violet Colorless Red orange yellow 450 500 500 550 550 580 580 650 650 700 700 Blue Green Yellow Orange Red Infrared Purple Blue green Colorless ii The particle nature of light photon energies 1 Photons light can be considered photons which are a stream of particles whose energy E is given by the following equation a E hv hc b H 6 626x10 34 Js iii Atomic spectra 2 Joule unit to express energy this is a relatively small unit so we use kilojoules 1J 10 3 J 3 Energy is inversely proportional to wavelength 1 Isaac Newton showed that visible white light from the sun can be broken down into its various color components by a prism a The spectra is continuous contains essentially all wavelengths between 400 and 700 nm 2 Since photons are produced when an electron moves from 1 energy level to another the electronic levels in an atom must be quantized limited to particular values c The Hydrogen Atom i Bohr model electron s orbit definite energies RH n2 1 Assumed that a hydrogen atom consisted of a central photon about which an electron moves in a circular orbit 2 He related the electrostatic force of attraction of the proton for the election to the centrifugal force due to the circular motion of the electron a Could express the energy of the atom in terms of the radius of the 3 He then assumed that the electron in the hydrogen atom can have only certain a En b En is the energy of the electron c R H is a constant 2 180x10 18J d nis an integer called the principal quantum number e Zero electron and nucleus are infinitely separated f Lowest energy state ground state e as close to nucleus as possible g Highest energy state excited state e further from nucleus i Further away from nucleus absorb E ii Move to nucleus release E 4 Excited electrons will give off energy to return to ground state and emit light 5 Light absorption formulas a E Ehi Elo i ii If positive give off E If negative take in E b E Rh 1 2 1 2 nlo nhi ii Quantum Mechanical Model 1 Heisenberg s uncertainty principal it s impossible to know to the precise location of the e at any time 2 Quantum numbers n l ml ms a First Quantum Number n i N period row on the periodic table ii A higher n means higher energy for the e b Second Quantum Number l shape of sublevel s p d f i Integral number from 0 to n 1 1 0 s group 1 2 2 1 p group 13 18 3 2 d transition metal a n 1 a n 2 a n 3 4 3 f lanthanide and actinide series ml specific orbital area of orbital that hold Integral values from l to l Indicates the specific orbital that the electron is in a n 4 c Third Quantum Number 2e i ii iii S 0 iv P 1 1 0 1 v D 2 2 1 1 2 vi F 3 3 2 1 0 1 2 3 d Fourth Quantum Number ms spin of the electron 1 2 or 1 2 i ii Pauli exclusion principle in an atom no 2e can have the same 4 quantum numbers d Atomic Orbitals shape and size i S sublevels are spherical 1 As n increases the radius of the orbital becomes bigger a Ex An electron in 2s orbital is more likely to be found far out from the ii P sublevels are dumbbells nucleus than is a 1s e 1 As n increases the radius of the orbital becomes bigger e Electron configurations in atoms i Introduction 1 Electron configuration simple way to describe the arrangement of electrons in an atom ii Electron configuration from sublevel energies 1 Electrons enter the available sublevels in order of increasing sublevel energy iii 1 Abbreviated electron configuration a Use the closest ideal gas and then continue with electron configuration iv Filling of sublevels and the periodic table 1 The elements in Group 1 and 2 are filling s sublevels 2 The elements in group 13 18 fill p sublevels 3 The transition metals fill the d sublevel 4 The lanthanides and actinides are filling f sublevels a Lanthanides first row Z 57 70 b Actinides second row Z 89 102 f Orbital diagrams of atoms i Intro 1 Orbital diagrams used to show electron distribution among orbitals 2 The arrows in the boxes represent the spin of the electron 3 Hund s Rule when several orbitals of equal energy are available as in a given sublevel electrons enter singly with parallel spins a b In all filled orbitals the two electrons have opposite spins In accordance with Hund s rule within a given sublevel there are as many half filled orbitals as possible 4 Paramagnetic if there are unpaired electrons present the solid will be attracted into the field this substance is paramagnetic 5 Diamagnetic if the atoms in a solid contain only paired electrons it is slightly repelled by the field this substance is diamagnetic g Electron arrangements in monoatomic ions i Ions with noble gas structure 1 Ex The 3 atoms preceding Neon N O F and the three elements following Na Mg Al all form ions with the neon configuration of neon a The nonmetal atoms N O F do this by gaining electrons and forming b The metal atoms Na Mg Al do this by losing electrons and forming anions cations c All of these atoms are isoelectric the have the same electron configuration i N3 O2 F Ne Na Mg2 Al3 ii Transition metal cations 1 Transition metals to the right of the scandium subgroup don t form ions with noble gas configuration 2 They can form cations with 1 2 3 charges following this rule a Electrons are removed from the sublevel of highest n you can predict correctly that when transition metal atoms form positive …