“Loose Ends” of Bonding

1. “Loose Ends” of Bonding Chapters 8 and 9 Chemistry: Matter and Energy

2. Lattice Energy of Ionic Compounds • The lattice energy (bond energy) associated with an ionic compound is based on coulombs law. (Don’t worry about the mathematical statement of the law): When comparing two ionic compounds for bond strength, look at the ions’ charges that make up those compounds. The greater the ions’ charges, the stronger the bond. If the two compounds contain ions of the same charge, the smaller ion pair will have a stronger bond.

3. Examples of comparing compounds for relative lattice energy • CaS will have a higher lattice energy that KCl because the former has a 2+ ion bonded to a 2- ion, and the latter has a 1+ ion bonded to a 1- ion (The greater the ions’ charges, the stronger the bond.) • CaS will have a lower lattice energy that MgO, because even though they both have 2+ ions bonded to 2- ions, Ca2+ and S2- are larger than Mg2+ and O2-, respectively. (If the two compounds contain ions of the same charge, the smaller ion pair will have a stronger bond.)

4. Formal Charge • Formal charge is used to evaluate two Lewis dot structures that can both be correct. It is a calculated value for each atom in the structures. The formula is: FC= (# of valence e- in a free atom of the element)- (# of e- assigned to the atom in the structure) The structure will the lowest formal charges is the best structure!! (Also any negative formal charge should be on the most electronegative atom in the molecule.)

5. Example of determining formal charge Click on the link below and view the video: http://highered.mcgraw-hill.com/olcweb/cgi/pluginpop.cgi?it=swf::100%25::100%25::/sites/dl/free/0072512644/117354/05_Formal_Charge_Calculations.swf::Formal%20Charge%20Calculations

6. Resonance • Resonance occurs when more than one Lewis Structure can be drawn for a molecule or ion. The drawing above shows three resonance structures for carbonate ion (with nonzero formal charges written in). NOTE: The bond length of all three carbon – oxygen bonds is the same: an intermediate length between the length of a single bond and the length of a double bond

7. Bond dissociation energy (covalent bonds) • Bond length is inversely related to bond strength: longer bonds are weaker than shorter bonds. • Furthermore, bond length inversely related to the number of shared pairs between two atoms: single bonds are longer than double bonds, double bonds are longer than triple. • THEREFORE: single bonds are longer and weaker than double bonds, double bonds are longer and weaker than triple bonds

8. Polarity of Molecules • Polar molecules result when there is an uneven distribution of the electron density in the molecule. This results in the molecule having a partial positive charge (δ+) on one end and a partial negative charge (δ) on the other end. Another way of saying this is that the molecule has a “net dipole moment”.

9. More on Polarity • Whether a molecule is polar or nonpolar is dependent on the shape of the molecule. Molecules that are symmetrical are nonpolar and molecules that are asymmetrical are polar. • Ex: PCl3 is a trigonal pyramidal molecule with an unshared pair of electrons on the P. This is an asymmetrical shape so PCl3 is polar • Ex: PCl5 is a trigonalbipyramidal molecule with no unshared pairs of electrons on the central P. This is a symmetrical shape so PCl5 is nonpolar.

10. Hybridization • Hybridization is a theory that explains how atomic orbital “hybridize” to accommodate for bonding. To determine the hybridization on a molecule, draw the Lewis dot structure and count the total number of electron domains on the atom in question. The number of domains tells you how many atomic orbitals were involved in forming the hybrids. The only atomic orbitals that can be involved are s, p, and d.

11. Determining the hybridization • In CH4, there are 4 electron domains on carbon. Thus, there are 4 atomic orbitals that hybridize – one s and three p, so we say the carbon is sp3 hybridized. • In C2H4, there are 3 electron domains on each carbon. Thus, there are 3 atomic orbitals that hybridize – one s and two p, so we say each carbon is sp2 hybridized

12. Sigma (σ) and Pi (π) bonds • Molecular orbital theory is another theory that explains what happens to atomic orbitals when atoms bond. In this theory, overlap of s orbitals create sigma bonds and overlap of p orbitals create one sigma and two pi bonds. • Single bonds are all σ bonds • Double bonds contain one σ bond and one π bond • Triple bonds contain one σ bond and two π bonds

13. Bond order • Bond order is a quantitative way of expressing bond strength. The formula for calculating bond order is: BO = # of shared pairs on the central atom ÷ # of atoms bonded on the central atom EX: CO2 has a central carbon with two double bonds to oxygens (O=C=O). So there are 4 shared pairs and two bonded atoms so the bond order is 4 ÷ 2, which is 2

14. Bond order and bond strength • Molecules with a higher bond order contain stronger bonds • Carbon monoxide has carbon triple bonded to oxygen. There are three shared pairs and one bonded atom, so the bond order is 3 ÷ 1, which is equal to 3. Thus the bonding in CO is stronger than the bonding in CO2.