Chapter 16 Covalent Bonding

1. Chapter 16Covalent Bonding • The Nature of Covalent Bonding • Bonding Theories • Polar Bonds and Molecules

2. Chapter 16.1 The Nature of Covalent Bonding • Single Covalent Bonds • Double and Triple Covalent Bonds • Coordinate Covalent Bonds • Bond Dissociation Energies • Resonance • Exceptions to the Octet Rule

3. Single Covalent Bonds • A bond in which two atoms share a pair of electrons (usually non-metals) • The pair of electrons is represented by a dash • Structural formulas – chemical formulas that show the arrangement of atoms in molecules or poly atomic ions

4. Single Covalent Bonds • Hydrogen • H2 • H H • H-H

5. Single Covalent Bonds • Water • Ammonia • Methane

6. Double and Triple Covalent Bonds • Double Covalent Bonds – involve two shared pairs of electrons • Triple Covalent Bonds – involve three shared pairs of electrons

7. Double and Triple Covalent Bonds • Nitrogen • N2 • N N • N N

8. Coordinate Covalent Bond • One atom contributes both bonding electrons • Carbon Monoxide

9. Bond Dissociation Energies • The total energy required to break the bond between two covalently bonded atoms • H-H + 435kJ = H + H • C-C + 347kJ = C + C

10. Bond Dissociation Energies

11. Resonance • Structures that have two or more different electron dot structures that have the same number of electron pairs for the same molecule or ion • Ozone • NO3- • http://www.nku.edu/~russellk/tutorial/reson/NO3.gif • http://www.nku.edu/~russellk/tutorial/reson/CO3.gif

12. Exceptions to the Octet Rule • Occurs when the total number of valence electrons is an odd number. • NO2 (Two resonance structures) • BF3 • PCl3 and PCl5 • SF6

13. Diamagnetic • All electrons paired • The spinning of electrons creates magnetic fields. • The paired electrons spin in opposite directions, therefore their fields cancel each other out • Show a weak attraction to an external magnetic field

14. Paramagnetic • Contain one or more unpaired electrons • Creates a magnetic field • Show a strong attraction to an external magnetic field

15. Oxygen – An Exception • O=O • Actually a mix of • O=O and O-O

16. Ch 16.2 Bonding Theories • Molecular Orbitals • VSEPR Theory • Hybrid Orbitals

17. Molecular Orbitals • An orbital resulting from the overlapping of orbitals from two atoms when they bond • Bonding orbital – molecular orbital with an energy lower than the atomic orbital • Antibonding orbital – molecular orbital with an energy higher than the atomic orbital

18. Molecular Orbitals • Sigma Bond – a molecular orbital that is symmetrical along the axis connecting two atomic nuclei • Two s orbitals directly overlap

19. Molecular Orbitals • Pi bonds - two parallel 'p' orbitals in close proximity can overlap sideways (laterally) • A pi bond can only form after a sigma bond has already formed

20. VSEPR Theory • Electron pairs repel, so molecules adjust shape to keep the pairs as far apart as possible • Unbonded electrons repel bonded electrons more, causing the bonded electrons to be closer together • http://www.mikeblaber.org/oldwine/chm1045/notes/Geometry/VSEPR/Geom02.htm

21. Ch 16.3 Polar Bonds and Molecules • Bond Polarity • Polar Molecules • Attractions Between Molecules • Intermolecular Attractions and Molecular Properties

22. Bond Polarity • Nonpolar Covalent Bond – equal sharing of electrons, each atom pulls equally (same atoms) • Polar Covalent Bond – unequal sharing of electrons (different atoms)

23. Bond Polarity • The more electronegative atom will have a greater pull and acquire a slightly negative charge. • HCl • H2O

24. Bond Polarity

25. Polar Molecules • Dipole – a molecule with two poles (different charges) • Polarity depends on shape • Atoms that lie in the same axis will cancel each other out

26. Attractions Between Molecules • van der Waals forces – weak attractions between molecules – Two Types • Dispersion Forces – Weak attraction due to movement of molecules, increase with more electrons (diatomic halogens) • Dipole Interactions – polar molecules attracted to each other (water molecules)