Science league topic 7: bond energy and resonance structure

1. Science league topic 7: bond energy and resonance structure

2. Bond Energies • chemical reactions involve breaking bonds in reactant molecules and making new bond to create the products • the amount of energy it takes to break one mole of a bond in a compound is called the bond energy (D) • Bond energy is measured in the gas state • homolytically – each atom gets ½ bonding electrons

3. Bond Lengths • the distance between the nuclei of bonded atoms is called the bond length • because the actual bond length depends on the other atoms around the bond we often use the average bond length • averaged for similar bonds from many compounds

4. Trends in Bond Energies and bond length • the more electrons two atoms share, the stronger the covalent bond • Energy: Triple bonds > double bonds > single bond • C≡C (837 kJ) > C=C (611 kJ) > C−C (347 kJ) • the shorter the covalent bond, the stronger the bond • F−F (237 kJ) > Cl−Cl (218 kJ) > Br−Br (193 kJ) • bonds get weaker down the column • Bond length: Triple bonds > double bonds > single bond

5. Using Bond Energies to Estimate DH°rxn • the actual bond energy depends on the surrounding atoms and other factors • we often use average bond energies to estimate the DHrxn • works best when all reactants and products in gas state • bond breaking is endothermic, DH = + • bond making is exothermic, DH = − DHrxn= ∑nr(D(reactants)) + ∑np (D(products))

7. Estimate the Enthalpy of the Following Reaction

8. Estimate the Enthalpy of the Following Reaction H2(g) + O2(g) ® H2O2(g) reaction involves breaking 1mol H-H and 1 mol O=O and making 2 mol H-O and 1 mol O-O bonds broken (energy cost) (+436 kJ) + (+498 kJ) = +934 kJ bonds made (energy release) 2(464 kJ) + (142 kJ) = -1070 DHrxn = (+934 kJ) + (-1070. kJ) = -136 kJ (Appendix DH°f = -136.3 kJ/mol)

9. Resonance structures

10. •• •• •• •• O S O •• •• O S O •• •• • • • • • • • • Resonance • when there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures • the actual molecule is a combination of the resonance forms – a resonance hybrid • it does not resonate between the two forms, though we often draw it that way • look for multiple bonds. Most oxyanions and oxyacidshave resonance structures.

11. Resonance • The bonds between Os have exactly the same length! • The resonance bond length is < single bond, but > double bond.

12. Rules of Resonance Structures • Resonance structures must have at least one double bond and at least two same surrounding atoms. • One of the same surrounding atoms should be connected to the double bond. • only electron positions can change • C, N, O, and F elements have a maximum of 8 electrons • bonding and nonbonding • Others don’t follow octet rule • Resonance structures of a molecule are the same. All the bonds which could form a double bond are the same with same bond energy.

13. -1 -1 +1 Drawing Resonance Structures -1 • draw first Lewis structure that maximizes octets • assign formal charges • move electron pairs from atoms with (-) formal charge toward atoms with (+) formal charge • Make the resonance structure: move in electrons to form the double bond from a different surrounding atom after you can move out electron pairs from one double bond. -1 +1

14. Exceptions to the Octet Rule • There are three types of ions or molecules that do not follow the octet rule: • Atoms in compounds with less than an octet. Boron and Hydrogen • Center atom in compounds with more than eight valence electrons (an expanded octet). All the nonmetals from period 3 and below. • The empty d orbitals in the outermost shell can be used to accommodate the e-

15. Fewer Than Eight Electrons • Consider BF3: • Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine. • Fluorine only forms single bond in the compounds. • Therefore, structures that put a double bond between boron and fluorine are much less important than the single-bond structure.

16. More Than Eight Electrons • The only way PCl5 can exist is if P has 10 electrons around it. • If Cl is negative charged in the compounds, it forms single bond like F. • It is allowed to expand the octet of atoms on the 3rd row or below. • Presumably dorbitals in these atoms participate in bonding.

17. More Than Eight Electrons Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

18. More Than Eight Electrons • This eliminates the charge on the phosphorus and the charge on one of the oxygens. • The lesson is: When the central atom is on the 3rd row or below and expanding its octet eliminates some formal charges, do so.

19. Something to remember • H and F can only form single bond. • When negatively charged, Cl, Br and I can only form single bond.

20. Type of covalent bonds

21. Types of Bonds • sigma (s) bond: the bonding points along the axis connecting the two bonding nuclei • s-to-sor p-to-p • pi (p) bond: the bonding is parallel to each other and perpendicular to the axis connecting the two bonding nuclei • between parallel porbitals • the interaction between parallel orbitals is not as strong as between orbitals that point at each other; therefore s bonds are stronger than p bonds

22. Tro, Chemistry: A Molecular Approach

23. How many s and p bonds in Formaldehyde?

24. Bond Rotation • the sigma (s)bond can rotate along the internuclear axis without breaking the bond. • the pi (p)bond can NOT rotate around the axis without the breaking of the pi bond.

25. Homework • Page 400: 61 b, c and d, 62 c and d, 71 a, 73, 77 • How many sigma and pi bonds in the CO2?