Bonding and molecular geometries

1. Bonding and molecular geometries chapters 8 and 9

2. types of bonds • Ionic • Formed by the Coulombic attraction between oppositely charged ions. • Covalent • Formed by the sharing of valence electrons between atoms. • Metallic • Formed by electrons that are free to move through a metallic substance.

3. octet rule • Atoms will gain, lose or share electrons to acquire a full valence shell. • What is energetically special about filled valence shells? • Usually 8 electrons • Exceptions to the octet rule: • H and He have a filled valence shell with 2 electrons. • Atoms with few valence electrons (Be, B, Al) can have less than an octet. • Atoms with unoccupied d orbitals (P and larger) can have up to 12.

4. Ionic bonding • Na(s) + ½ Cl2 (g) NaCl(s) ΔH = -410.9 kJ/mol • The ionization energy of sodium is 495.8 kJ/mol • The electron affinity of chlorine is -349 kJ/mol. • If it requires more energy to remove an electron from sodium than is gained by adding an electron to chlorine, why is this process exothermic? • The fact that ionic solids form crystalline lattices may have something to do with this.

5. Lattice energy • in NaCl, each sodium ion is surrounded by 6 chloride ions, and vice versa. • The lattice energy is defined as how much energy it takes to break a crystalline lattice into gaseous ions. • Since forming NaCl(s) is the opposite process, it makes sense that it should have the opposite sign. NaCl(s) Na+(g) + Cl-(g)ΔHlattice = 788 kJ/mol

6. more about lattice energy • Lattice energies follow Coulomb’s Law • Larger charges have higher lattice energies • Smaller volume ions have higher lattice energies. • For ionic compounds to dissolve, the interaction between the solvent and ions must be stronger than the lattice energy.

7. Covalent bonds • Electrons are shared between atoms due to the overlap of orbitals. • Direct overlap = sigma bond (σ) • Side by side overlap = pi bond (π) • Double bond = sharing of 4 electrons (1 sigma, 1 pi). • Triple bond = 6 electrons (1 sigma, 2 pi).

8. bond Length • Consider two hydrogen atoms in close proximity. We have three Coulombic forces: • The attraction for each electron to the neighboring nucleus. • The repulsion of electrons. • The repulsion of nuclei. • There will be a point when the energy of the system is at a minimum. The distance between the two nuclei is the length of the bond. • The energy of the system is the bond energy.

9. General relationships • Triple bonds are stronger (have higher bond energy) than double, which are stronger than single. • Stronger bonds have shorter lengths, generally. • C=C is shorter than C-C

10. Lewis structures • Shows connectivity and location of electrons • Rules • Count the number of valence electrons in the molecule. • Make a skeleton, placing the least electronegative element in the center (never H, always C). • Use molecular symmetry to help you. • Place single bonds between all adjacent atoms. • Add lone pairs around outside. • If you run out of electrons, make double/triple bonds. • Check formal charge. Best structures have least formal charges and – charge on most electronegative atom • Formal charge = valence electrons – unshared electrons - # of bonds.

11. Examples formaldehyde ammonium ion. Notice how all ions are placed in brackets with the charge.

12. Resonance • Many molecules can have multiple correct Lewis Structures. These exhibit resonance. • Resonance contributors must have the same skeleton. The only difference is the arrangement of lone pairs and multiple bonds. • The true structure is a hybrid of all contributors.

13. resonance There are no single or double bonds on the oxygens. Each oxygen has a “bond order” of 1.5, somewhere between a single and double bond. This has been proven by experimentation!

14. Polarity • Bonds formed between atoms with the same electronegativity are said to be nonpolar, meaning that there is no dipole moment. • Any bond formed by elements of differing electronegativities will be polar. • There is no “dividing” line between polar, nonpolar and ionic. Bond types exist on a spectrum. • Generally, metal-nonmetal bonds are ionic • Homonuclear bonds are nonpolar • Heteronuclear nonmetallic atoms form polar covalent bonds.

15. Polarity

16. Valence bond theory

17. VSEPR • Valence Shell Electron Pair Repulsion • Electron pairs want to be as far apart as possible. • All based on positions around the central atom. • Electron Domains are regions where we would expect to find electron pairs. • Can be bonding electrons or lone pairs. • Double/triple bonds still only count as one domain.

18. molecular geometry

19. Molecular Shapes • To determine the shape of a specific molecule, we must count the number of bonding domains and nonbonding domains. • Given in the form ABxEy, where x is the number of bonding domains and y is the number of lone pairs on the center.

20. Shapes

21. effect of lone pairs on bond angles • CH4, NH3 and H2O all have tetrahedral geometries, but different bond angles. • Lone pairs exert greater repulsive force because they are typically closer to the center atom.

22. Molecular Polarity • A molecule is polar if all of the bond dipoles fail to “cancel out” • Think vector addition! • Commonly caused by the presence of lone pairs, or the lack of symmetry.

23. molecule polarity

24. orbital hybridization • Look at the orbital diagram for carbon. How many bonds do you expect it to form? • We would expect it to form 2 bonds because it has two unpaired electrons to share. Why does it form 4 bonds?

25. Hybridization • If we combine the wave functions for carbon’s 2s and 2p orbitals, we can form hybrid orbitals that each have the same energy. • The number of hybridized orbitals formed are equal to the number of orbitals that were combined. In this case, 1 s orbital combines with 3 p orbitals to form 4 sp3 hybridized orbitals.

26. sp3 hybrid orbitals • sp3 hybridized orbitals form a tetrahedral geometry, with a bond angle of 109.5°. • sp3orbitals can only form sigma bonds with other molecules.

27. sp2 hybrid orbitals • sp2 hybridized orbitals form a trigonal planar geometry with bond angles of 120°. There is one p orbital that hasn’t hybridized. • sp2 hybridized atoms (that obey the octet rule) will have three sigma bonds (caused by the overlap of the hybrid orbitals) and one pi bond (caused by the side to side overlap of the unhybridized p orbital).

28. sp hybrid orbitals • sp hybridized orbitals will assume a linear geometry with a bond angle of 180°. There are two unhybridized p orbitals. • sp hybridized atoms (if they obey the octet rule) will form two sigma bonds (from the sp hybrid) and two pi bonds (from the unhybridized p orbitals). • This can be in the form of 2 double bonds, or 1 triple bond.

29. Molecular orbital theory

30. why MO theory? • Valence bond theory does a great job at showing us the geometry of molecules, their bonding, and their polarities. • It does not tell us anything about if a molecule will be para/diamagnetic, how it will absorb light, or anything about a molecule’s excited states. • Warning: You will never be asked any questions on the AP exam regarding MO Theory. I am including it here because it is taught in all college Gen Chem courses, and it is necessary to completely understand semiconductors.

31. molecular orbitals • If two orbitals overlap, their wave functions combine. This combination can be either constructive or destructive. • A constructive combination is called a bonding orbital. • A destructive combination is called an antibonding orbital.

32. Making molecular orbitals • When two orbitals combine, they will form one bonding and one antibonding orbital. • Given the fact that only one of the three p orbitals can directly overlap, p orbitals will form 1 sigma orbital and 2 pi orbitals. • We follow normal electron config rules to fill the boxes.

33. more homonucleardiatomics Notice how O2 has unpaired electrons. This makes it paramagnetic!