1. Chemical Bonding and Lewis Structures

2. Chemical Bonding Chemical Bonds are the forces that hold atoms together. Atoms form bonds in order to attain a minimal energy state. Bond formation is an exothermic process (just as bond breaking is endothermic) The type and strength of bond that forms between reacting particles dictates the physical and chemical characteristics of the molecule or ion in question. Bond energy is the energy needed to break a bond and is an indication of bond strength.

3. The Relationship Between Electronegativity and Bond Type

4. Ionic Bonds Ionic bonds occur between ions due to electrostatic attraction between positive cations and negative anions. Form between atoms with large differences in electronegativity (>1.7), usually between a metal and a nonmetal. Relative strength can be determined using Coulomb’s Law E = k(q 1 )(q 2 ) r 2 K= coulomb’s constant; q1=charge on one ion; q2= charge on other ion; r=distance between ions

5. Ionic Bond, continued The strength of an ionic bond is directly proportional to the magnitude of the charges involved and inversely proportional to the square of the distance between them. Ionic solids tend to have high melting points Often soluble in polar solvents, such as water.

6. Ionic Solid UNIT CELL Crystal lattice – at the lattice sites in the unit cell are positive and negative ions held together by Coulombic attraction force between positive and negative ions. CONDUCTIVITY Solid has no conductivity due to zero empty valence orbitals. The ions in the crystal are isoelectronic with Group 18. Molten and aqueous solutions will conduct due to mobile ions.

7. Covalent Bonds Nonpolar covalent bonds form between atoms of nonmetals with nearly identical electronegativities while polar covalent bonds for between nonmetals with dissimilar electronegativities (0.4-1.7). Covalent bonds within molecules are strong, but the binding forces between molecules are relatively weak. Molecular solids usually have low melting points. Usually soluble in nonpolar solvents (carbon tetrachloride)

8. Which of the following bonds would be the least polar yet still be considered polar covalent? Mg-O C-O O-O Si-O N-O React 5

9. Metallic Bonding Occurs in metallic solids Metal atoms usually have large, positively charged nuclei and few valence electrons. Nuclei are positioned in a regular geometric array (lattice) by electrostatic repulsion. Valence electrons are attracted equally by all nuclei. Leads to the “sea of electrons” model with nuclei bobbing in a “sea of electrons” Useful for explaining the physical characteristics of metals (ie, conductivity) Metals have a wide range of melting points.

10. Metallic Solid CRYSTAL LATTICE At lattice sites in unit cell are positive ions held together by mobile valence electrons traveling through empty valence orbitals. POSITIVE CHARGE DENSITY Smaller ionic radius = higher melting point CONDUCTIVITY Solid: excellent due to empty valence orbitals Liquid: good due to mobile ions.

11. Intermolecular Forces The group of weaker attractive forces between atoms or molecules. NOT BONDS!!!!! Van der Waals forces London dispersion forces-all atoms and molecules; caused by temporary dipoles created as electrons move about the nucleus. Strength depends on the number of electrons moving, so molecules with larger masses (more electrons) have greater London forces. Dipole-dipole force- an attraction between opposite polar ends of adjacent molecules. Hydrogen bonding- occurs when hydrogen bonds with a very electronegative anoin (F, O, N) resulting in a very polar molecule. Strongly positive and negative ends of the molecule have stronger interactions than either london or dipole-dipole forces.

12. The Effect of an Electric Field on Hydrogen Fluoride Molecules

13. Lewis Structure Lewis Structures – shows how the valence electrons are arranged among the atoms of a molecule There are rules for Lewis Structures that are based on the formation of a stable compound Atoms want to achieve a noble gas configuration

14. Octet & Duet Rules Octet Rule – atoms want to have 8 valence electrons Duet Rule – H is the exception. It wants to be like He & is stable with only 2 valence electrons

15. Steps for drawing Lewis Structures 1. Sketch a simple structure with a central atom and all attached atoms 2. Add up all of the valence electrons for each individual atom If you are drawing a Lewis structure for a negative ion add that many electrons to create the charge If you are drawing a Lewis structure for a positive ion subtract that many electrons to create the charge

16. 1. Subtract 2 electrons for each bond drawn 2. Complete the octet on the central atom & subtract those electrons 3. Complete the octet on the surrounding atoms & subtract those electrons 4. Get your final number If 0  you are done! If +  add that many electrons to the central atom If -  need to form multiple bonds to take away that many electrons Steps for drawing Lewis Structures

17. Examples CCl 4 Sketch a simple structure with a central atom and all attached atoms Cl │ Cl – C – Cl │ Cl

18. Examples Add up all of the valence electrons for each individual atom 4 + 4(7) = 32 Subtract 2 electrons for each bond drawn 32-8 = 24 Complete the octet on the central atom & subtract those electrons Done

19. Examples Complete the octet on the surrounding atoms & subtract those electrons 24 – 24 = 0 Final number = 0…DONE! Final structure is… __ │Cl │ __ │ __ │Cl – C – Cl │ │ │Cl │

20. Examples HF

21. Examples NH 3

22. Examples NO +

23. Exceptions to the octet rule Sometimes the central atom violates the octet rule and has more or less than 8 valence electrons Keep using the same rules to draw Lewis Structures

24. Examples SF 4 ICl 3 XeF 4 ICl 4 -

25. Resonance When more than one Lewis Structure can be written for a particular molecule Resonance structure – all possible Lewis structures that could be formed The actual structure is the average of all of the structures You MUST show all structures!

26. Examples SO 3 NO 2 - NO 3 -

27. Covalent Network Crystal CRYSTAL LATTICE At the lattice site in the unit cell there are atoms held together by strong covalent bonds. Si, SiO 2, C (diamond), and C (graphite) This crystal has the highest melting point.

28. SUMMARY Covalent network solids have the highest melting points because of the strongest forces holding the crystal together (covalent bonds). Metallic solids generally have the next highest melting points, but the larger the ionic radius, the lower the melting point. Ionic crystals next. Molecular crystals last. Consider IMF to rank these. The stronger the IMF, the higher the melting point.