Chapter 9: Models of Chemical Bonding
Chapter 9: Models of Chemical Bonding CH 121
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Date Created: 12/01/15
Chapter 9 Lecture Notes Forms of chemical bonds 0 There are 2 extreme forms of connecting or bonding atoms 0 Ionic complete transfer of 1 or more electrons from on atom to another 0 Covalent some valence electrons shared between atoms 0 Most other bonds are somewhere in between Ionic bonds 0 Essentially completely electron transfer from an element of low ionization energy metal to an element of high af nity for electrons nonmetal o Ionic compounds exist primarily between metals at the left of the periodic table groups 1A and 2A and transition metals and nonmetals at the right Energetics of ionic bonding O O Electrostatic attraction between the reactants By accounting for all 3 energies ionization energy electron af nity and lattice energy we can get a good idea of the energetics involved in such a process Lattice energy 0 O O O O The energy required to completely separate a mole of a solid ionic compound into its gaseous ions The energy associated with electrostatic interactions is governed by coulomb39s law Ee k QlQZId Increases with the charge on the ions Also increases with decreasing size of ions Covalent bonding Bond arises from the mutual attraction of 2 nuclei for the same electrons Electron sharing Several electrostatic interactions in these bonds Attractions between electrons and nuclei Repulsions between electrons Repulsions between nuclei 0 Bond is a balance of attractive and repulsive forces Polar covalent bonds 0 Although atoms often form compounds by sharing electrons the electrons are not always shared equally o Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does 0 Therefore the uorine end of the molecule has more electron density than the hydrogen end 0 When 2 atoms share electrons unequally a bond dipole results 0 Dipole moment produced by 2 equal but opposite charges separate by a distance r is calculated 0 o It is measured in debyes D 0 Greater the difference in electronegativity the more polar is the bond Electronegativity 0 Ability of atoms in a molecule to attract electrons to itself 0 On the period chart electronegativity increases from left to right and from bottom to top Electron distribution in molecules 0 Electron distribution is depicted with lewis electron dot structure 0 Valence electrons are distributed as shared or bond pairs and unshared or lone pairs Bond and lone pairs 0 Valence electrons are distributed as shared or bond pairs and unshared or lone pairs Bond formation 0 A bond can result from a quotheadtoheadquot overlap of atomic orbitals on neighboring atoms Note that each atom contributes a single unpaired electron Valence electrons o Electrons are divided between core and valence electrons Rules of the game 0 Number of valence electrons of a main group atoms group number Except for H and sometimes atoms of 3rd and higher periods 0 This is because atoms like to have a noble gas con guration Octet Rule Lewis structures 0 Representations of molecules showing all electrons bonding and nonbonding Writing lewis structures 0 Find the sum of valence electrons of all atoms in the polyatomic ion or molecule If it is an anions add one electron for each negative charge If it is a cation subtract one electron for each positive charge 0 The central atom is the least electronegative element that isn39t hydrogen Connect the outer atoms to it by single bonds 0 Fill the octets of the outer atoms 0 Fill octet of the central atoms o If you run out of electrons before the central atom has an octet form multiple bonds until it does 0 Assign formal charges 0 For each atom count the electrons in lone pairs and half the electrons it shares with the other atoms Subtract that from the number of valence electrons for that atom the difference is its formal charge 0 The best lewis structure Is the one with the fewest charges Puts a negative charge on the most electronegative atom Resonance 0 One lewis structure can39t accurately depict a molecule such as ozone 0 We use multiple structures resonance structures to describe the molecule 0 They aren39t localized but are rather delocalized Building a dot structure 0 Decides the central atom never hydrogen 0 Count valence electrons 0 Form a single bond between the central atom and each surrounding atom o The remaining electrons form lone pars to complete octet as needed Variations of the octet rule 0 3 types of ions or molecules that don39t follow the octet rule lons or molecules with an odd number of electrons lons or molecules with less than an octet lf lling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atoms don39t ll the octet of the central atom lons or molecules with more than eight valence electrons expanded octet When the central atom is on the 3rd row or below and expanding its octet eliminates some formal charges Be B N Cl P S Covalent charge strength 0 Most simply the strength of a bond is measured by determining how much energy is required to break the bond 0 This is the bond enthalpy designated D Average bond enthalpies 0 There are tables that list the average bond enthalpies for many different types of bonds They are AVERAGE bond enthalpies not absolute ones 0 Average bond enthalpies are positive because bond breaking is an endothermic process Enthalpies of reaction 0 Yet another way to estimate delta H for a reaction is to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed o In other words delta H rxn sum of bond enthalpies broken sum of bond enthalpies of bonds formed