Covalent Bonding 10/29/2013

Transcription

1 Bond Energies or Bond Dissociation Energies Tables 8.4 and 8.5 on page 72 gives a list of the energy required to dissociate or break bonds. This value is used to determine whether covalent bonds will form or to tell if reactions are favorable (exothermic). This is the only time that = reactants - products = Σ (bond broken) - Σ ( bonds formed) Where Σ is the sum of, let s try one Br from its elements 2(g) + Br 2(g) 2 Br (g) = Σ (- + Br-Br) - Σ (2 (-Br) = ( ) (2(6)) = kj =exothermic Covalent Bonding Percent Ionic Character- a comparison of the measured dipole moment of the molecule and the calculated dipole moment from the ions. This is measured for the covalent molecules and will never reach 100% - due to the fact that no individual bonds are ionic. In simpler terms, in a covalent bond no element can completely take another atoms valence electrons away. % ionic = measured dipole moment of X-Y x 100 character calculated dipole moment of X + Y - Localized Electron Bonding Model (LE) There are two different theories that are used to predict accurate models for molecules, there is this one LE and the Molecular rbital Theory (M). (chapter 9) We will use the LE method and not cover the M theory, it is not part of this course-however, you may want to look over pages The LE assumes that a molecule is composed of atoms that are bound together by sharing pairs of electrons using those that are free to bond (valence electrons). 1

2 Using the LE Theory There are three parts to this theory 1. Description of valence electron arrangement using Lewis structures. 2. Prediction of the geometry of the molecule using the VSEPR model.. Description of the type of atomic orbitals used by the atoms to share electrons or hold lone pairs. The electrons involved in the bond(s) are BP or bonding pairs. The electrons around the atoms that remain unshared are called the LP or lone pairs Modeling Covalent Molecules To begin we need a molecule, let s use ammonia (N ) Start by getting a total of all the valence electrons (include in this process any Formal Charge) Formal Charge = the charge on the entire molecule as in a polyatomic ion For ammonia there is no formal charge so let s total the valence electrons N = 5 + s = So our total is 8 electrons or 4 pairs. Lewis Dot Structures Gilbert Lewis devised a way to show the bonding that occurs in molecules by using dots. Each dot represents a valence electron and when a pair is shared by two atoms it can be replaced by a stick that represents the bonding electrons. There are 4 bond sites on each atom; top, bottom; right and left- except only has one and B has only three arranged symmetrically around it (triangular) 2

3 When Drawing Electron Dot (Lewis Dot) Diagrams 1. Each dot represents a valence electron 2. Sticks are used to show a shared pair. Each element wants an CTET 4. Unpaired electrons move to become paired 5. Unshared pairs can move to form a double or triple bond-coordinate covalent bond 6. Unshared pairs are always shown 7. Bonds are arranged to show shape as far apart as possible 8. An unshared pair takes up space and must be shown Back to ammonia, N Select your central atom (most bond sites) and attach everything to it N This used pairs leaving one more pair to use or a LP on nitrogen (hydrogen is full) BP + 1 LP = 4 pairs N Now we have a complete octet! Another one- C 2 Total the valence electrons = C =4 + 2 = 6 = 16 or 8 pairs-start by connecting to the central atom. C We have 6 pairs left, if we count the number of pairs, that could occupy the bond sites we need 8 ( on each and 2 on C) since we don t have enough, try a double bond. C now we have an octet on C and have 4 pairs left over so put them on (2 pairs each) C Move electrons around symmetrically. ctets Complete!

4 Structural Diagrams Let s show both models we made using the structural formulas Show LPs N = C = TRY P 4 - P = = ( more) = 2 r 16 pairs, put P in the middle (most bond sites) and connect all the oxygens to it P P 4 - This uses 4 pairs leaving us 12 more Use them for completing each s octet - P - - Always label the charge! Sulfite ion, S 2- Step 1. Central atom = S Step 2. Count valence electrons S = 6 x = x 6 = 18 Negative charge = 2 TTAL = 26 e- or 1 pairs Step. Form bonds 10 pairs of electrons are now left. S 4

5 Sulfite ion, S 2- Remaining pairs become lone pairs, first on outside atoms and then on central atom. Additional electrons-due to charge - are added where needed. 2- S Each atom is surrounded by an octet of electrons. 5