Fundamentals of Chemical Bonding Valence Bond Theory n

Fundamentals of Chemical Bonding

Valence Bond Theory n n Chemical bonds involve the electrons A bond results if a more stable electron configuration results The s block and p block elements (often termed the representative elements) will form bonds such that there are eight electrons surrounding each atom (the octet rule) The exceptions are H, Li and Be, which tend to follow a “duet” rule.

Bonds are the forces that hold atoms together strongly n Remember that all bonding is electrical in nature n Covalent bonds involve a concentration of electrons between the two atoms n Result from overlap of two (or more) atomic orbitals n

Bond length and bond strength n n Atoms forming bonding interactions will more closer together until the cores begin to touch The bond length is the distance between the two atoms where the molecule is most stable The two atoms vibrate (I. e. the bond gets longer and shorter) For H 2, the bond length is 74 pm, and the bond strength is the energy that it takes to split the molecule into individual atoms (7. 22 10 -19 J)

Approaches to bonding n n Bonding can be treated by a localized or a delocalized method. Both have value. Generally u each electron is assigned to a specific orbital u the Pauli exclusion principle holds u the aufbau principle holds u the valence orbitals are all that are needed to describe the bonding

Localized bonding Electrons are either in bonding orbitals or nonbonding orbitals n Electrons in bonding orbitals are often unequally shared, resulting in a polar bond n

Electronegativity n n the ability of an atom to attract electrons in a bond to itself Differences in electronegativities of atoms that are bonded results in a partial transfer of electron charge to the more electronegative atom. The bond is therefore a polar covalent bond The polar bond has a negative end a positive end (a so-called dipole)

The octet rule n Atoms can obtain the octet by: u metals can lose one to three electrons to form a cation with the electron configuration of the previous noble gas u nonmetals can gain one to three electrons to form an anion with the electron configuration of the next noble gas u atoms can share electrons

Formation of ions Atoms can lose electrons to form ions u Na ([Ne]2 s 1) Na+ ([Ne]) + en Atoms can gain electrons to form ions u Cl ([Ne]2 s 22 p 5 + e- Cl ([Ne]2 s 22 p 6 n Lewis dot structures of the atoms can be very helpful here n

Forming ionic compounds Reaction of Li with O n Reaction of Na with N n Note that ionic compounds exist as solids in a three-dimensional array called a lattice n The electrostatic attractions between the ions in the lattice are the ionic bonds n

The Covalent Bond Covalent bonds result from electron sharing between two atoms n We use Lewis dot structures to show the order and arrangement of the atoms in a molecule and all of the valence electrons n

Lewis dot structures 1. 2. 3. Count the number of valence electrons for the atoms in the molecule For ions, add one electron for each negative charge or subtract one electron for each positive charge Place the most electropositive atom in the center (the inner atom). Array the more electronegative atoms around this atom as outer atoms

Lewis Dot structures 4. 5. 6. 7. Connect the outer atoms to the inner atoms with single bonds Subtract two electrons from the total number of valence electrons for each bond Array the remaining electrons around the outer atoms in pairs to complete their octet Check the octets of all atoms

Lewis dot structures 7. 8. 9. Use lone pairs on the outer atoms to form multiple bonds to the inner atom if needed to complete the octet Calculate formal charges on each of the atoms Arrange the electrons and atoms to minimize formal charges

Formal Charge n n n Difference between the number of valence electrons in the free atom, and the number of electrons located on the atom in the molecule The electrons in bonds are divided between the two atoms. The nonbonding electrons are assigned to the atom they are localized on Formal charge = (valence electrons of free atom) (valence electrons assigned in Lewis structure)

Formal Charge Formal charge on a given atom = # of valence electrons - 0. 5 (# of bonding electrons) - # of nonbonding electrons n From Pauling’s electroneutrality principle, atoms arrange themselves so that they minimize their formal charge (note that the book gives a narrower definition) n Consider N 2 O and SCNn

Formal Charges n n n The oxidation state formalism is an exaggeration, that is that both electrons in the bond belong to the more electronegative atom The formal charge formalism goes the other way, and assumes that the electrons are equally shared From this, we can calculate a formal charge as a tool for assessing the stability of a given Lewis structure

Resonance structures n n n In compounds with multiple bonds, sometimes you can draw structures which vary only by placement of the double bonds The structures that vary only by the placement of the double bonds are called resonance structures, and are an approximation of the true structure Actually, the molecule is a superposition of the resonance structures, trying to show delocalization works

Molecules 3 D In the Novell delivered programs is an icon for Molecules 3 D-Pro (not Chem 3 D) n In this program, there is a Lewis dot structure icon for C on the toolbar n This is Coach Lewis, which will step you through the process of Lewis dot structures n

Polarity of bonds n n Bonds that involve atoms of differing electronegativities have a concentration of negative charge at the more electronegative atom, and a deficiency of charge at the less electronegative atom This unequal distribution of negative charge creates a dipole, where one end of the bond is slightly negative and the other is slightly positive

Polar molecules result from an arrangement of polar bonds such that the entire molecule has a dipole n The polar bonds can be arranged to cancel the polarity (CO 2) n The best way to predict this in complex cases is vector algebra, but it is important to learn to recognize this in clear cut structures n

Oxidation numbers n Please refer to Supplementary information on the web site to review this material

Octet Expansion n Elements in the second period only have four valence orbitals, which limits the total number of bonds that can be formed to a central atom to four Elements in the third and higher periods have accessible d-orbitals available for bonding, allowing more than four bonds to the central atom Consider BF 5

VSEPR Electron pairs will repel each other, and will govern the structure of the molecule, all other things being equal n Electron pairs will arrange themselves to be as far apart as possible n Note that molecular geometry is described by the bonded atoms and does not include the lone pairs n

Electron pair repulsion n n Degree of repulsion depends on the electron pair types; in order of decreasing repulsion u lone pair-lone pair u lone pair-bonding pair u bonding pair-bonding pair We will also treat all of the electrons that bond together two atoms as one electron group regardless of whether the bond is single, double or triple

Parent structures n n n We will consider three parent structures (there are more, but will be covered in a later class) The central or inner atom is designated A, outer atoms are designated X, and lone pairs are designated E The parent structures are based on the number of electron pairs that surround the central atom

Hybridization The shapes of molecules cannot be reconciled to the directionality of the bonds in most molecules n This is treated by mixing the atomic orbitals of the central atom to form new, hybrid orbitals n For s and p block elements, there are three hybrids; sp, sp 2, sp 3 n

Parent structures Parent structure Name/hybid Bond angles AX 4 tetrahedron/sp 3 109. 5° AX 3 trigonal planar/sp 2 120° AX 2 linear/sp 180°

Parent structure derivatives Each parent structure can give rise to a family of derivatives, simply by replacing bonding pairs with lone pairs n For AX 4, there are two derivatives u AX 3 E - the trigonal pyramid (NH 3) u AX 2 E 2 - bent (H 2 O) n Replacement of the bonding pairs with the lone pair(s) compresses the bond angles n

Coordination and steric number Coordination number is number of atoms bound to the central atom n Steric number is the number of pairs of electrons around the central atom (the sum of the coordination number and the number of lone pairs of the central atom) n Steric number defines the parent structure n

Higher order structures Octet expansion is treated by forming hybrids with d orbitals n dsp 3 yields trigonal bipyramidal n d 2 sp 3 yields octahedral n Bond angles are either 120 or 90 degrees n See Tables 8 -4 and 8 -5 n

Confirmation of structure Molecules with an asymmetric distribution of electron density have a dipole moment n Polar bonds are necessary for dipole moments, but symmetrically shaped molecules can have zero dipole moments due to cancellation of the polar bonds n