Chemistry 103 1st Edition Exam 3 Study Guide Chapters 7 10 Chapter 7 The Quantum Mechanical Model of the Atom Quantum Numbers o n principal quantum number energy and size of the orbital o l angular momentum quantum number shape of the orbital o ml magnetic quantum number orientation of the orbital Chapter 8 Periodic Properties of the Elements Spin quantum number ms Pauli Exclusion Principle no two electrons in an atom can have the same four quantum numbers Electron configurations o need to know how many electrons o ground state electron configuration lowest energy configuration o order of increasing energy 1s 2s 2p 3s 3p 4s 3d 4p etc o Hund s Rule electrons will stay unpaired as long as possible o orbital diagrams putting electrons in boxes o difference between core and valence electrons o For condensed electron configurations Previous Noble Gas s p d f orbitals that come after that noble gas Magnetic Character o paramagnetic at least 1 unpaired electron o diamagnetic no unpaired electrons Periodic properties o elements in the same group have the same valence electron configuration Effective Nuclear Charge Zeff Z S o amount of charge in the nucleus that the valence electrons can actually see o increases as you go from L to R across periodic table no real trend from bottom to top Atomic Radius o decreases as you go from L to R across periodic table b c of Zeff o decreases as you go from bottom to top b c of decreasing n Ionic Radii o cations are smaller anions are larger than parent atoms Ionization Energy IE o energy needed for X X eo increases as you go from L to R across periodic table o increases as you go from bottom to top o IE s are very large for core electrons Electron Affinities EA o energy needed for X e Xo increases decreasing energy as you go from L to R across periodic table o increases as you go from bottom to top Metallic Character ability to form a cation conductive ductile malleable shiny etc o decreases as you go from L to R across periodic table o decreases as you go from bottom to top Chapter 9 Chemical Bonding I Lewis Theory 2 types of bonds o ionic e transferred metals and nonmetals o covalent e shared nonmetals Electronegativity the ability of an atom to attract electrons to itself in a bond o increases as you go from L to R across periodic table o increases as you go from bottom to top o a polar bond occurs between two atoms with different electronegativities can draw the dipole moment over the bond Lewis Dot Structures molecules o Steps for Lewis Structures 1 Add up the valence electrons from all the atoms in the molecule 2 Write the symbols for the atoms to show which atoms are attached to which and connect them with a single bond The more electronegative atoms generally sit on the outer parts of the molecule with a less electronegative atom in the center Hydrogens are NEVER the center of a molecule 3 Complete the octets of the atoms bonded to the central atom remember that hydrogen only gets two electrons 4 If there are not enough electrons to give the central atom an octet try forming multiple bonds Move unshared electron pairs on the outer atoms into bonds 5 If and ONLY if you still have electrons left over place them on the central atom even if this gives the central atom more than an octet the central atom must be a Row 3 element or heavier o formal charges help predict the best Lewis dot structures Look at how many electrons the element originally brought into the molecule Ex Carbon brings 4 When counting electrons for formal charges count lone pairs and one electron for each bond Ex If carbon makes a double bond and has two sets of lone pairs it would have six electrons in the molecule Original amount final amount Formal charge of the atom in the molecule Exceptions to the octet rule o molecules with an odd number of electrons o incomplete octets o expanded octets Chapter 10 Chemical Bonding II Molecular Shapes and Valence Bond Theory Valence Shell Electron Pair Repulsion VSEPR Theory the most likely structure will be one where the electron pairs are as far apart as possible o Use the handout for pictures of the geometries o Two electron groups linear electron geometry bond angles 180 o Three electron groups trigonal planar electron geometry bond angles 120 depending on the presence of lone pairs multiple bonds molecular geometries trigonal planar or bent o Four electron groups tetrahedral electron geometry bond angles 109 5 depending on the presence of lone pairs multiple bonds molecular geometries tetrahedral trigonal pyramidal or bent o Five electron groups trigonal bipyramidal electron geometry o Six electron groups octahedral electron geometry Molecular shape and polarity is the molecule polar or nonpolar o look at the polarity of the bonds and the arrangement of those bonds Valence Bond Theory explains how s and p orbitals mix hybridize to form compounds with non 90 bond angles o linear electron geometry sp hybrids o trigonal planar electron geometry sp2 hybrids o tetrahedral electron geometry sp3 hybrids o Count the number of sigma bonds and lone pairs in the atom this will help find the hybridization Sigma vs Pi bonds o end to end overlap usually formed by hybridized orbitals All single bonds are sigma bonds The first bond in a double or triple bond is also a sigma bond o edge to edge overlap formed by non hybridized orbitals The second bond in a double bond is a pi bond The second AND third bonds in a triple bond are also pi bonds