Chemistry Notes Chapter 8 Characteristics of Many Electron Atoms Schrodinger s equation was not applicable for any other element on the table but it was modified to be applicable to all other atoms Bohr s model was incapable of doing it Electron Configuration distribution of electrons in its orbitals controls properties Atomic orbitals are all hydrogen like except for magnitude in atomic number Must extend quantum mechanics model Intro of 4 th quantum number and limit number of electrons in one orbital to 2 a more complex set of subshell energy levels note energy levels 2S 2P Magnetic properties indicate 4th quantum number Ms associated with electrons Either Ms 1 2 or 1 2 Describes electron spin which is not a property of the orbital Spin quantum number Ms indicates direction of spin Pauli s exclusion principle no two electrons in an atom can have the same set of 4 quantum numbers First three describe location fourth describes state of the electron Atomic orbitals can have a max of 2 e and these must have opposite spins Down arrow 1 2 up arrow 1 2 spin paired Electrostatic effects can influence atom s energy states Coulomb s law Constant x product of charges divided by distance between the charges E k q1 x q2 d12 More far apart d is large makes energy of system smaller Less negative number Closer together larger more negative Q1 is nuclear charge More protons to electrons the more stable the orbital More protons more negative the energy Adding an electron destabilizes Number of electrons doubles energy lessens by 2x Two electrons easy to remove one One on one harder to remove Orbital more stable when one on one one proton one electron Effect of inner electrons on energy of an outer orbital shielding electron repulsion makes effect on outer electron easier to remove Stable more negative Easier to remove the outer electron less stable Effect of l value on orbital energy penetration effect put electron in 2s vs 2p 2p easier to remove that 2s because 2s is closer to nucleus 2s more stable Combo of shielding and penetration cause energy levels to be split into energy sublevel groups For a given n value shell the lower the l value and the lower the sublevel energy Sublevel groups in polyelectron atoms begin to overlap between the 3 and the 4 th shells rd Quantum Mechanical Model and the Periodic Table QM gives the theoretical basis for the periodic table Aufbau Principle repeating patterns arranged Electron configurations of elements recur in patterns that cause periodically varying element properties Starting at H add one electron per element to the lowest energy orbital available Orbital diagrams the superscript electrons in each orbital 2P Hund s rule we want to maximize the spin So we do parallels don t put two electrons in one orbital unless we have to or else they ll repel each other Electron configurations building up period 3 More negative more stable Energies 3s 3p 3d Electron configuration shorthand last noble gas and then its own config since that gas Pauli exclusion principle Outer electron configuration correlate closely with chemical behavior in the group s elements Group 1A 1 Noble gas ns 1 Group 7A 17 Noble gas ns2 np5 1 away from noble gas Fourth Period 4s sublevel has lower energy than 3d in K and Ca so 4s fills before 3d Cr Chromium is an exception 6 unpaired e not 4 Stability of filled electrons not Aufbau controls 4s1 3d5 Verrrrrry stable because all have same spin Molybdenum is also exception and silver Mn 4s2 3d5 Copper exception Unusual filling patterns of Cr Cu Mo Ag 4s1 3d10 show that half filled and filled subshells are especially stable The 4p subshell is then filled by the next six elements Some Features of Electron configurations Inner core are those in previous noble gas plus any completed transition series Outer electrons are those in the highest energy level n Only outer electrons participate chemical bonding Valence electrons are those that may be involved informing compounds In main group elements the valence electrons are the outer electrons Transition elements if d electrons are involved in bonding these are counted as valence electrons once d orbitals filled they are inner If not they are valence Ions of elements in groups 1a and 1a 6a and 7a are usually isoelectronic with the nearest noble gas Look back at labeled periodic table very clear 1a 1s1 2a 1s2 the middle transitions 3d10 halogens and nobles p Trends in three atomic properties Atomic radius ionization energy and electron affinity Metallic radius is one half the distance between nuclei of adjacent atoms in a Molecules atomic size covalent radius is one half the distance between crystal nuclei of identical atoms Variations in atomic size result from two influences atoms become larger as principle quantum no n increases Atoms also become smaller as effective nuclear charge Zeff increases Zeff Z inner electron shielding Down PT groups each step adds subshells of inner electrons shielding outer electrons from larger Z so Z changes a little Thus atomic radius increases going down a group Decreases going up Across a period electrons add to same outer level shielding each other poorly so Z increases Higher Z is smaller radius will be Atomic radius generally decreases from left to right in a period Transition metals vary every so slightly Cations are smaller than their parent atoms because it has more positive than negative charges meaning electron cloud is squeezed down Anions are larger than their parent atoms Positive charge can t hold the extra negative charge as well Ionic size increases down a group because the number of electron shells increases Across periods size decreases among the isoelectronic meaning same number of electrons different number of protons cations and less among isoelectronic anions The larger the difference in the number of electrons and protons the less the atom can pull in the electrons and the larger they will be Larger the positive charge the smaller the radius Trends in Ionization Energy IE is energy needed to remove one mole of e from 1 mole of gaseous atoms or ions Energy into system endothermic Always a positive number kJ IE1 energy needed to remove an electron from highest occupied sublevel of the gaseous atom 3a and 6a Poly Electron atoms can lose more than one e IE s of each electron are numbered in sequence Atom g Ion e Delta E IE1 0 IE2 is always larger than IE1 IE2 Ion Ion2 e delta E IE decreases down a group In a period GENERALLLY increases across a