Ch. 10: Chemical Bonding

1. Ch. 10: Chemical Bonding Dr. Namphol Sinkaset Chem 152: Introduction to General Chemistry

2. I. Chapter Outline • Introduction • Lewis Dot Structures • Ionic Compounds • Covalent Compounds • Shapes of Molecules • Polar Bonds and Polar Molecules

3. I. Introduction • Bonding theories allow prediction of how atoms form compounds and the shapes they will take. • The 3-D shape of a molecule determines many of its physical properties. • With the advent of faster computers, simulations allow screening of drug candidates.

4. I. HIV-protease Inhibitors

5. I. Lewis Theory • Lewis theory is simple to apply, but amazingly powerful. • It can be used to go from a formula of a compound to its 3-D structure. • Lewis theory centers on how valence electrons are used by atoms to form compounds.

6. II. Lewis Dot Structures • Valence e-’s are the most important e-’s in bonding. • Lewis dot structures are a way to depict the valence e-’s of atoms. • Lewis dot symbols have two parts: • element symbol: represents nucleus and core e- • dots around symbol: represent valence e-’s

7. II. Origin of Dot Structures • Oxygen has 6 valence electrons, so it’s dot structure will have 6 dots.

8. II. Lewis Dot Structures • The number of valence e- is given by the element’s group number!!

9. II. The Central Idea • Lewis realized that noble gases all had 8 valence e-’s (except He). • He reasoned that having 8 valence e-’s (or 2 for H and He) leads to stability, or being unreactive. • These special configurations of valence e-’s are known as an octet or a duet.

10. II. The Octet Rule • In Lewis theory, atoms bond in order to obtain eight valence electrons. • The octet rule states that in chemical bonding, atoms transfer or share electrons to obtain outer shells with eight electrons. • The octet rule is a main-group concept, generally applying to all except H and Li.

11. II. Sample Problem • Draw Lewis dot structures for Ca, Ge, Se, and Br.

12. III. Transfer of Electrons • When metals bond with nonmetals, the metal transfers electrons to the nonmetal to form an ionic compound. • The positive charge of the metal cation and the negative charge of the nonmetal anion holds the ionic compound together. • We can show this with Lewis theory.

13. III. Formation of Potassium Chloride

14. III. Formation of Sodium Sulfide

15. III. Sample Problem • Use Lewis dot structures to depict the compound that forms between: • magnesium and nitrogen • aluminum and bromine

16. IV. Sharing Electrons • When nonmetals bond with other nonmetals, a covalent compound is formed. • Electrons are shared in covalent compounds in order to achieve an octet. • Electrons that appear in the space between atoms count towards the octet of both atoms.

17. IV. Lewis Structure for Water

18. IV. Octet/Duets Satisfied

19. IV. Bonding vs. Lone Pairs • Electrons shared between atoms are bonding pair electrons. • Electrons that are only on one atom are lone pair or nonbonding electrons.

20. IV. Lines as Bonds • Generally, bonding pair electrons are represented by lines. • Note that a line equals two electrons.

21. IV. Why Some Elements Exist as Diatomics • If two hydrogens combine, they can satisfy their duet. • If two chlorines combine, they can satisfy their octet.

22. IV. Diatomics

23. IV. Higher Order Bonds • In some compounds, atoms need to share more than one electron pair to reach an octet. • Double bond – atoms share 4 electrons • Triple bond – atoms share 6 electrons • However, higher order bonds are a last resort used by atoms to reach an octet!

24. IV. Diatomic Oxygen • Oxygen has 6 valence electrons. • Sharing one pair satisfies the octet of one oxygen atom, but not the other. • Since there are no more electrons that can be used, higher order bonds must be made.

25. IV. Double Bond Formation

26. IV. Triple Bonds • Sometimes six electrons need to be shared by two atoms. • An example is diatomic nitrogen.

27. IV. Single, Double, Triple Bonds • Higher order bonds mean more stability and shorter internuclear distances.

28. IV. Steps for Drawing Lewis Structures • Determine total # of valence e-. (Cation, subtract e-’s for charge; anion, add e-’s for charge). • Place atom w/ lower Group # (lower electronegativity) as the central atom. • Attach other atoms to central atom with single bonds. • Fill octet of outer atoms. (Why?) • Count # of e- used so far. Place remaining e- on central atom in pairs. • If necessary, form higher order bonds to satisfy octet rule of central atom. • Allow expanded octet for central atoms from Period 3 or lower.

29. IV. Sample Problem • Draw correct Lewis structures for NF3, CO2, SeCl2, CCl4, and H2CO.

30. IV. Exceptions to the Octet Rule • Lewis theory is too simple to cover all bonding possibilities. • Some exceptions exist to the octet rule: • e- deficient atoms like Be and B, e.g. BeCl2 and BF3. • Compounds w/ odd # of e-’s: free radicals. Examples include NO and NO2. • Expanded valence – when d orbitals are used to accommodate more than an octet.

31. V. VSEPR Theory • From a correct Lewis structure, we can get to the 3-D shape using this theory. • VSEPR stands for valence shell electron pair repulsion. • The theory is based on the idea that electron groups (lone pairs, single bonds, or multiple bonds) repel each other.

32. V. Linear Geometry • CO2 has two electron groups. • The two double bonds try to get as far away from each other as possible.

33. V. Trigonal Planar • Formaldehyde has three electron groups around the central atom. • They form 120° angles to get away from each other.

34. V. Tetrahedral • Methane has four electron groups. • A bond angle of 109.5° keeps them furthest apart.

35. V. Tetrahedral-Based Shapes • Other shapes are based on tetrahedral with bonding groups being replaced by lone pair electrons. • The electron geometry is the shape based on all electron types (bonding and lone pair). • The molecular geometry is the shape based on just atoms.

36. V. Ammonia • The central atom has three bonds and one lone pair (4 electron groups). • EG = tetrahedral • When drawing MG, lone pairs are left off!

37. V. Trigonal Pyramidal • Ammonia has a trigonal pyramidal molecular geometry.

38. V. Water • The central atom has two bonds and two lone pairs (4 electron groups). • EG = tetrahedral • Again, lone pairs omitted when drawing MG!

39. V. Bent • Water has a bent molecular geometry.

40. V. Geometry Summary

41. V. Drawing w/ Perspective • We use the conventions below to depict a 3-D object on a 2-D surface.

42. V. Practice Problem • Draw the molecular shapes for ClO2-, BF3, and NF3. Indicate the name of the molecular and electronic geometries for each as well.

43. VI. Sharing Electrons • It is not reasonable to assume that all atoms will share electrons equally. • Some atoms will pull electrons closer to them. • Electronegativity is the ability of an element to pull electrons in a covalent bond closer.

44. VI. Effect of Electronegativity • O is more electronegative than H, so in an O-H bond, the bonding electrons are more likely found around the O.

45. VI. Dipole Moment • The unequal sharing leads to a partial charge separation in the bond called a dipole moment. • Covalent bonds with a dipole moment are polar covalent bonds. • The greater the difference in electronegativity, the greater the polarity.

46. VI. Electronegativity Values

47. VI. Bonding Type • Electronegativity difference can be used to determine the type of bond.

48. VI. Sample Problem • Calculate the difference in electronegativity for the following pairs of atoms and determine of the bond between them is pure covalent, polar covalent, or ionic. • I and I • Cs and Br • P and O

49. VI. Polar Molecules • Just because a molecule has polar bonds doesn’t mean that the molecule is polar overall. • The degree of polarity in each bond and the orientation of those polar bonds determines whether the molecule is polar.

50. VI. Polarity Vectors • We can use vectors to represent dipole moments and analyze the resultant.