1. The Chemical Bond

2. Chemical Bonds  Are the forces that hold atoms together to form compounds  Bond energy – the amount of energy needed to break a bond and produce a neutral atom  Bond strength – the amount of force holding two atoms together Ex. 85.9 kcal/mole Ex. 85.9 kcal/mole

3. Types of Bonds  Ionic Bond – involves the transfer of electrons between atoms  Covalent Bond – involves that sharing of electrons between atoms Types of Covalent bonds Types of Covalent bonds Single C:CSingle C:C Double C::CDouble C::C Triple C:::CTriple C:::C  Coordinate Covalent Bond – one atoms donates a pair of electrons to be shared

4. Ionic Bonds  Involves electron transfer and ion formation  Cation – has a positive charge  Anion – has a negative charge  The cation is much smaller than the anion  In an ionic solid ions pack together in a way that is dependant of the size of the ions  Ions arrange in a way that there is local neutrality  High melting point crystalline solids

5. Electron Transfer / Ion Formation

6. Ionic Bond

7. Metallic Bonds  The force of attraction that holds metals together  It consists of the attraction of free-floating valence electrons for positively charged metal ions

8. Metals  The valence electrons of metal atoms can be modeled as a sea of electrons

9. Properties of Metals (due to the ‘sea’ of electrons)  Conduct electricity – flow of electrons  Malleable – hammered into sheets  Ductile – draw into wire

10. Ionic Crystal Structure Note: the smaller size of the cations relative to the anions. The packing of the ions determines crystal shape.

11. Covalent Bond  Interatomic forces are created by the sharing of electrons.  The atoms share their s and p electrons to attain the electron configuration of a noble gas  Atoms have small differences in electronegativity  Generally low melting points (many are liquids and gases at room temperature)

12. The Octet Rule  Atoms react by gaining or losing electrons so as to acquire the stable electron configuration of a noble gas. Usually eight valence electrons

13. Lewis Dot Structures

14. Bond Formation  A bond forms when two electron clouds overlap and occupy a common orbital (molecular orbital)

15. Overlap of ‘s’ and ‘p’ orbitals s, s overlap s, p overlap p, p overlap Sigma Bonds Pi Bonds p, p side by side overlap

16. Nonpolar Covalent Bonds When the electrons are equally shared the bond between the atoms is nonpolar. Note the even distribution of the electron cloud of the hydrogen molecule

17. Polar Covalent Bonds  Polar – an unsymmetrical distribution of electric charge due to the unequal sharing of electons  The electronegativity difference between the atoms determines the degree of polarity

18. What Type of Bond Is It? Electronegativity Difference Most Probably Type of Bond Example 0.0 – 0.4 Nonpolar Covalent H-H 0.4-1.0 Moderately Polar Covalent H-Cl 1.0-2.0 Very Polar Covalent H-F > 2.0 IonicNaCl

19. Classifying Compound  Molecular – held together with covalent bonds  Network Solid (Ionic) – held together with ionic bonds

20. Dipole Moment  The measure of the force exerted on a dipole ( a single bond)

21. Dipole  A molecule that has an uneven distribution of charge even though the molecule as a whole is electrically neutral

22. The Water Molecule A Polar Molecule There are two polar covalent bonds and the bent shape of the molecule causes the uneven distribution of charge resulting in a polar molecule.

23. The Methane Molecule The even distribution of the charge results in a nonpolar molecule

24. Non-polar Molecules The individual C-O bonds are polar in nature but the overall molecule is nonpolar due to the even or balanced distribution of charge.

25. Another Example The CF 4 molecule has 4 evenly distributed polar bonds resulting in no net dipole for the molecule. The result is a nonpolar molecule

26. Coordinate Covalent Bond  Also known as a “Dative Bond”  A covalent bond in which both electrons are donated by a single atom

27. Expressions of Chemical Formulas  Chemical formula H 2  Lewis Dot structure (dots represent valence electrons) H:H  Dash formula H-H (dash represents a pair of electrons –a bond)

28. Electron Dot Formulas

29. Let Us Practice Some Lewis Dot Structures!  Water H 2 O  Methane CH 4  Ammonia NH 3  Carbon Tetrachloride CCl 4

30. Let’s see how we did! Water Methane AmmoniaCarbon Tetrachloride

31. Some Common Molecular Shapes Linear Bent BentPyramidalTetrahedral

32. Bonding and Molecular Orbitals  Sigma Bonds Single bonds Single bonds Overlap of two s orbitalsOverlap of two s orbitals Overlap of an s and a p orbitalOverlap of an s and a p orbital  Pi Bonds Double or Triple bonds Double or Triple bonds Side by side interaction of two p orbitalsSide by side interaction of two p orbitals

33. Sigma Bonds (   When two atomic orbitals combine to form a molecular orbital along the internuclear axis

34. When two carbon atoms bond there is an overlap of atomic orbitals along the internuclear axis. When carbon bonds with the hydrogens there is an overlap of hydrogen’s ‘s’ orbitals with carbons atomic orbitals to produce 6 sigma bonds. Ethane C 2 H 6

35. Pi Bonds (   When two atomic orbitals combine to form a molecular orbital above and below the internuclear axis  Can result from the side by side interaction between two ‘p’ orbitals

36. The Carbon Carbon Double Bond C=C Consists on one  and one  bond

37. So why does carbon bond with 4 equal energy orbitals? Why does carbon form tetrahedral geometry? Answer: Hybrid Orbital Theory

39. Carbons 1-s and 3-p valence orbitals combine to result in 4 equal energy bonding orbitals

40. The four equal energy orbitals account for carbons tetrahedral geometry

41. sp 2 Hybridization in Boron Results in trigonal planar geometry

43. sp Hybridization in Beryllium Explains linear geometry

45. Molecular Shapes LinearBentPyramidal Trigonal Planar Tetrahedral

46. Energy Changes in Bond Formation Bonding and Antibonding Orbitals * Notice that the molecular orbital is lower energy then the atomic orbitals The energy levels in a hydrogen molecule can be represented in a diagram - showing how the two 1s atomic orbitals combine to form two molecular orbitals, one bonding (  ) and one antibonding (  * )