Chemical Bonding and Lewis Structures

1. Chemical Bonding and Lewis Structures 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(q1)(q2) • r2 • 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. React 5 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

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 • Sketch a simple structure with a central atom and all attached atoms • 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. Steps for drawing Lewis Structures • Subtract 2 electrons for each bond drawn • Complete the octet on the central atom & subtract those electrons • Complete the octet on the surrounding atoms & subtract those electrons • 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

17. Examples • CCl4 • 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 • NH3

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 • SF4 • ICl3 • XeF4 • ICl4-

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 • SO3 • NO2- • NO3-

27. Covalent Network Crystal • CRYSTAL LATTICE • At the lattice site in the unit cell there are atoms held together by strong covalent bonds. • Si, SiO2, 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.