CHAPTER 10 THEORIES OF MOLECULAR BONDING I The Structure of the Molecular Bond While an atom can exist alone it usually chooses not to o Forms bonds with other atoms In the microscopic domain of quantum mechanics molecules form from atoms atoms form bonds because the formation of a molecule provides a means energy of the combined wavefunctions o A wavefunction that spreads out decreases in energy o Decrease in energy more stable configuration of electrons and protons Results in a stable union Coulombic attractions very important in determining the distribution of the wavefunction o Atom with greater ability to attract its valence electron s capable of extracting electron density from the atom that less strongly binds its valence electrons Leads to a shift in wavefunction Atoms combine when given the chance to create larger structures Rotations and vibrations in a molecule degrees of freedom Example H2 o Attraction between each electron and the two nuclei is offset by proton of one H repelling the proton of the other H o Net result small difference between strong attraction e to p and strong repulsion p to p e to e is attraction Attraction reduces potential energy and energy of system decreases Point of least amount of potential energy bond length Energy is released in the formation of a bond Electron sharing covalent bond o Electron distribution about each atom delocalizes within the structure of the molecule new molecular bond is equally distributed between two nuclei Polar covalent bond greater electron density centered on the atom with more tightly bound electrons Ionic bond delocalization of the combined wavefunction toward the atom with lower energy orbitals formed by simple Coulombic attraction between the cation that has donated electrons into the anion that has extracted electrons Ability of an atom to draw electrons to itself in a chemical bond electronegativity o Provides a useful way to estimate the degree to which electron density is delocalized in a chemical bond o Use Pauling electronegativity scale o Difference in EN between two species important in the formation of a chemical bond Larger the difference more polar the bond will be Types of Chemical Bonds 3 types ionic bonding covalent bonding metallic bonding Ionic o Results when there is a large difference in the electronegativity of two atoms o Metals bond to non metals Covalent o Non metal bonds with another non metal o EN difference is modest o Shared electron bond o Depending on EN difference can be polar covalent When there is a larger difference but not large enough to be ionic Metallic Bonding o Metal bonds with another metal o Bond in an ordered lattice o Have low electronegativity and low ionization energy lose electrons easily o Electron sea o Exhibit far less variety in chemical behavior than chemical bonds that explicitly bind electrons within a molecular structure Representation of Valence Electrons in a Chemical Bond Valence electrons ones that are engaged in the formation of a chemical bond GN Lewis developed a way of o 1 Representing valence electrons in an atom o 2 Representing those same electrons in the chemical bond formed o Only involves valence electrons Lewis Structure for Ionic Bonds Lattice Energy and the Formation of Ionic Crystals Energy trade off occurs when ionically bonded molecules form crystal structures 1 Reduction in potential energy 2 Energy expands to form cation and anion followed by release of energy resulting from the Coulomb attraction of the cation anion pairs that form the lattice crystal structure Lewis Structures and Covalent Bonding Central elements of theory of covalent bonding using Lewis s theory 1 Valence electrons play vital role in chemical bonding 2 Electrons are either transferred or shared 3 4 When electrons are shared a covalent bond results with small differences in Ionic bonds result from transfer of electrons electronegativity 5 Lewis theory is very effective in main group elements s and p better to use other theories for d and f orbitals Lewis Structures for Covalent Bonds Position and ordering of the nuclei in a bonding structure is controlled by the electron distribution that leads to the lowest possible potential energy of the molecular bonding structure Lewis Structures for Single Covalent Bonds Diatomics Complete respective closed shell configurations Lewis Structures and Bonding Character Idea in Lewis theory the attraction between two covalently bonded atoms is due to the sharing of one or more electron pairs o Shared electron idea implies that each bond links a specific pair of atoms Interactions within molecular structures intramolecular forces is much stronger than interactions between molecules intermolecular forces Lewis Structures for Polyatomic Molecules General Rules 1 Draw the correct skeletal structure for the molecule 2 Count total number of valence electron 3 Place e within the skeletal structure giving octets to the atoms and duets to hydrogen atoms Distribute lone pairs of electrons 4 Count the electrons around each atom and the total number available to make sure they match and satisfy the octet rule Doesn t always work can have resonance structures o Hybrid of two structures delocalized electron pairs o Electrons distribute themselves within the structure of a molecule such as to minimize the potential energy of the ensemble of all the electrons and protons in the molecule o Resonance structures same positioning of the atoms only difference location of single double or triple bond within the same structure o Use formal charges to find which structure is the correct one one that the molecule is most likely to be found in Method of Formal Charge Method assumes that the bonds are fully covalent with no delocalization of electron density toward a more electronegative member of a bonding pair Formula Rules for formal charge 1 Sum of formal charges in Lewis structure must equal 0 in a neutral molecule and must equal the net charge on a cation or anion 2 A smaller formal charge on an individual atom is favored over larger formal charges represent lower potential energy 3 Negative formal charges should appear on the most electronegative atoms 4 Structures having formal charges of the same sign on adjoining atoms are not favored Limitation to the Lewis Theory It is of use primarily for period 1 3 elements on the periodic table Free radicals have an odd number of valence electrons unpaired electron in the valence shell There are exceptions