1. Chemical Bonding Chapters 8 and 9

2. Chemical Bonds What is a bond? What is a bond? A force that holds atoms together A force that holds atoms together We will look at it in terms of energy We will look at it in terms of energy Bond Energy is the NRG required to break a bond Bond Energy is the NRG required to break a bond Why are compounds formed? Why are compounds formed? Bonds give the system the lowest NRG Bonds give the system the lowest NRG

3. Bond Energy is the energy required to break a bond. REMEMBER: REMEMBER: It always takes energy to break a chemical bond. It always takes energy to break a chemical bond. AND, AND, To form a bond, requires a lowering of energy. To form a bond, requires a lowering of energy.

4. Types of Bonds Ionic Bonds – electrostatic forces that exist between ions of opposite charge Ionic Bonds – electrostatic forces that exist between ions of opposite charge Covalent Bonds – sharing of electrons between two atoms Covalent Bonds – sharing of electrons between two atoms Metallic Bonds – found typically in transition metals; each atom is bonded to several neighboring atoms; bonding electrons are relatively free to move throughout the structure of the metal. Metallic Bonds – found typically in transition metals; each atom is bonded to several neighboring atoms; bonding electrons are relatively free to move throughout the structure of the metal.

5. Ionic Bonding An atom with a low ionization energy reacts with an atom with high electron affinity. An atom with a low ionization energy reacts with an atom with high electron affinity. Typically a metal with a nonmetal Typically a metal with a nonmetal The electron transfers atoms. The electron transfers atoms. The ions each achieve a Noble Gas electron configuration = low energy state. The ions each achieve a Noble Gas electron configuration = low energy state. Opposite charges hold ions together. Opposite charges hold ions together.

6. Crystal Lattice A repeating + and  crystal lattice results. A repeating + and  crystal lattice results. NaCl, sodium chloride

7. Ionic Compounds Makes a solid crystal. Makes a solid crystal. Ions align themselves to maximize attractions between opposite charges, Ions align themselves to maximize attractions between opposite charges, and to minimize repulsion between like ions. and to minimize repulsion between like ions. Chemical formula is actually the empirical formula, called the “formula unit” Chemical formula is actually the empirical formula, called the “formula unit”ClNaClNaClNaClNaClNa NaClNaClNaClNaClNaCl = “NaCl” ClNaClNaClNaClNaClNa

8. Features of Ionic Compounds Brittle, hard, crystalline solids at room temperature Brittle, hard, crystalline solids at room temperature High melting points High melting points Metals bonded to non-metals Metals bonded to non-metals Elements from opposites sides of the Periodic Table Elements from opposites sides of the Periodic Table Dissolve in water to form ions Dissolve in water to form ions Conduct an electric current in water Conduct an electric current in water

9. Coulomb’s Law Expresses the NRG of interaction between a pair of ions. Expresses the NRG of interaction between a pair of ions. E= 2.31 x 10 -19 J · nm (Q 1 Q 2 ) / r E = energy of interaction between a pair of ions (in Joules) E = energy of interaction between a pair of ions (in Joules) r = distance (in nm) between ion centers r = distance (in nm) between ion centers Q 1 and Q 2 = charges of the ions Q 1 and Q 2 = charges of the ions Opposite charges means (–E) Opposite charges means (–E) Endo or Exo? What does that mean about NRG in the system? Endo or Exo? What does that mean about NRG in the system?

10. Size of Ions Ion size increases down a group. Ion size increases down a group. Cations are smaller than the atoms they came from. Cations are smaller than the atoms they came from. Anions are larger. Anions are larger. across a row they get smaller, and then suddenly larger. across a row they get smaller, and then suddenly larger.

11. Periodic Trends Across the period effective nuclear charge increases so they get smaller. Across the period effective nuclear charge increases so they get smaller. Energy level changes between anions and cations. Energy level changes between anions and cations. Li +1 Be +2 B +3 C +4 N -3 O -2 F -1

12. Size of Isoelectronic ions Positive ions have more protons so they are smaller. Positive ions have more protons so they are smaller. A stronger + charged nucleus pulls the electrons inward. A stronger + charged nucleus pulls the electrons inward. Al +3 Mg +2 Na +1 Ne F -1 O -2 N -3

13. Lattice Energy The energy release that occurs when separated gaseous ions are packed together to form an ionic solid The energy release that occurs when separated gaseous ions are packed together to form an ionic solid X +x (g) + Y -y (g)  X y Y x (s) + energy X +x (g) + Y -y (g)  X y Y x (s) + energy Lattice NRG (E el ) = k(Q 1 Q 2 )/r Lattice NRG (E el ) = k(Q 1 Q 2 )/r k = constant k = constant Q 1 and Q 2 = charges of ions Q 1 and Q 2 = charges of ions r = distance between ion centers r = distance between ion centers

14. Example: Which has the most exothermic lattice energy, NaCl or KCl? Which has the most exothermic lattice energy, NaCl or KCl? NaCl … why? NaCl … why? Since both have the same charges (+1 and -1), the distance between the charges needs to be considered. Since Na+ is smaller than K+, the distance between the centers of Na and Cl is less and therefore has a greater lattice NRG. Since both have the same charges (+1 and -1), the distance between the charges needs to be considered. Since Na+ is smaller than K+, the distance between the centers of Na and Cl is less and therefore has a greater lattice NRG.

15. Example 2 and 3 Arrange the following ionic compounds in order of increasing lattice energy: NaF, CsI, and CaO Arrange the following ionic compounds in order of increasing lattice energy: NaF, CsI, and CaO CsI < NaF < CaO CsI < NaF < CaO Which substance would you expect to have the greatest lattice energy, AgCl, CuO or CrN? Which substance would you expect to have the greatest lattice energy, AgCl, CuO or CrN? CrN CrN

16. Covalent Bonds Bond is a force which causes a group of atoms to behave as a single unit. Bond is a force which causes a group of atoms to behave as a single unit. Electrons are shared by atoms. Electrons are shared by atoms. Electron orbitals must overlap. Electron orbitals must overlap.

17. The Covalent Bond The electrons in each atom are attracted to the nucleus of the other. The electrons in each atom are attracted to the nucleus of the other. The electrons repel each other, The electrons repel each other, The nuclei repel each other. The nuclei repel each other. They reach a distance with the lowest possible energy. They reach a distance with the lowest possible energy. The distance between is the bond length. The distance between is the bond length.

18. 0 Energy Internuclear Distance

19. 0 Energy Internuclear Distance

20. 0 Energy Internuclear Distance

21. 0 Energy Internuclear Distance

22. 0 Energy Internuclear Distance Bond Length (They reach a distance apart with the lowest possible energy.)

23. 0 Energy Internuclear Distance Bond Energy

24. Features of Covalent Compounds (aka: “Molecular Compounds”) 1.Electrons are shared, (Single, Double or Triple Bonds are possible.) 2. Non-metals bonded to Non-metals. 3. Includes all diatomic molecules. 4. Relatively low melting/boiling points. 5. No repeating formula, particle is a single unit called a “molecule”.

25. How We Represent Covalent Compounds 1. Molecular Formulas 2. Structural Formulas H H - C - H H CH 4 Show the bonds between the atoms.

26. Covalent vs. Ionic Bonding Covalent is sharing, ionic is stealing. Covalent is sharing, ionic is stealing. Totally different from each other. Totally different from each other. Reality Check: There are many compounds which exhibit both traits! Reality Check: There are many compounds which exhibit both traits! These are called polar covalent bonds. These are called polar covalent bonds. The electrons are shared, but shared unequally. The electrons are shared, but shared unequally.

27. Polar Covalent Bonds The electrons are shared, but they are not shared evenly. The electrons are shared, but they are not shared evenly. One atom has the pair more often than the other. One atom has the pair more often than the other. A “polar” molecule results: One end is slightly positive, while the other is slightly negative. A “polar” molecule results: One end is slightly positive, while the other is slightly negative. A partial charge is called a “dipole”. A partial charge is called a “dipole”.

28. H - F ++ -- Example:

29. H - F ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ --

30. ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- + -

31. ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- - +

32. How Do We Know Which Atom Has the Electron Pair More Often? Answer: Electronegativity Values Answer: Electronegativity Values Electronegativity - The ability of an atom to attract shared electrons to itself. Electronegativity - The ability of an atom to attract shared electrons to itself. The more electronegative an atom, the more often it has the shared pair. The more electronegative an atom, the more often it has the shared pair. Greater electronegativity =  pole Greater electronegativity =  pole

33. Electronegativity E.N. values are assigned for almost every element (Figure 8.6, p. 285) E.N. values are assigned for almost every element (Figure 8.6, p. 285) Gives us relative electronegativities of all elements. Gives us relative electronegativities of all elements. Tends to increase left to right and decreases as you go down a group. Tends to increase left to right and decreases as you go down a group. Noble gases aren’t discussed. Noble gases aren’t discussed. Difference in electronegativity between atoms tells us how polar. Difference in electronegativity between atoms tells us how polar.

34. Helpful Number Line: Determine the Electronegativity Value difference between two atoms Determine the Electronegativity Value difference between two atoms 0 0.4 2.0 3.4 Non-polar covalent Polar covalent ionic

35. Electronegativity difference Bond Type Zero Intermediate Large Covalent Polar Covalent Ionic Covalent Character decreases Ionic Character increases

36. Polar Covalent Bond- How it is drawn H - F ++ --

37. Reminder on Writing Formulas and Nomenclature Ionic Compounds: Name the cation then name the anion Name the cation then name the anion When writing formulas, add subscripts to make sure that the charges balance When writing formulas, add subscripts to make sure that the charges balance Covalent Compounds: When naming, use prefixes for the subscripts, the 2 nd atom will end in –ide When naming, use prefixes for the subscripts, the 2 nd atom will end in –ide Write formulas, assign subscripts based on the prefixes Write formulas, assign subscripts based on the prefixes

38. Lewis Dot Structures are Models to represent both ionic and covalent What is a Model? Explains how nature operates. Explains how nature operates. Derived from observations. Derived from observations. It simplifies and categorizes the information. It simplifies and categorizes the information. A model must be sensible, but it has limitations. A model must be sensible, but it has limitations.

39. Properties of a Model A human invention, not a blown up picture of nature. A human invention, not a blown up picture of nature. Models can be wrong, because they are based on speculations and oversimplification. Models can be wrong, because they are based on speculations and oversimplification. You must understand the assumptions in the model, and look for weaknesses. You must understand the assumptions in the model, and look for weaknesses. We learn more when the model is wrong than when it is right. We learn more when the model is wrong than when it is right.

40. Lewis Structures Show how the valence electrons are arranged. Show how the valence electrons are arranged. One dot for each valence electron. One dot for each valence electron. A stable compound has all its atoms with a noble gas configuration. A stable compound has all its atoms with a noble gas configuration. Hydrogen follows the duet rule. Hydrogen follows the duet rule. The rest follow the octet rule. The rest follow the octet rule. Bonding pair is the one between the symbols. Bonding pair is the one between the symbols.

41. X 3 2 1 4 7 5 8 6 The X represents the symbol for the element. The dots are placed around the symbol in the order shown above.

42. Rules 1. Sum the # of ALL the valence electrons. 2. Determine the central atom. The least electronegative element is central (H never central, C nearly always central) 3. Write symbols for the atoms to show which atoms are attached and connect them with a single bond (a dash) 4. Complete the octets of the atoms attached to the central atom (except for H, follows a duet rule). 5. Place any leftover electrons on the central atom. Not enough electrons - consider a double or triple bond.

43. A useful equation: ( happy - have )  2 = # bonds (what they want - what they have)  2 = # bonds (what they want - what they have)  2 = # bonds H 2 O (12 - 8)  2 = 2 bonds  H  O  H 

44. Practice Structures PCl 3 PCl 3 ICl ICl CH 4 CH 4 CH 2 Cl 2 CH 2 Cl 2 HCN HCN NO + NO + CO 2 CO 2 H 2 H 2 NH 3 NH 3 C 2 H 4 C 2 H 4 BrO 3 -1 BrO 3 -1 O 2 O 2 ClO 2 -1 ClO 2 -1 OH -1 OH -1 PO 4 -3 PO 4 -3

45. Partial Ionic Character There are probably no totally ionic bonds between individual atoms. There are probably no totally ionic bonds between individual atoms.

46. % Ionic Character Electronegativity difference 25% 50% 75%

47. How do we deal with it? If bonds can’t be ionic, what are ionic compounds? If bonds can’t be ionic, what are ionic compounds? An ionic compound will be defined as any substance that conducts electricity when melted. An ionic compound will be defined as any substance that conducts electricity when melted. Also use the generic term salt. Also use the generic term salt. As it turns out, most compounds fall somewhere between ionic and covalent. As it turns out, most compounds fall somewhere between ionic and covalent.

48. The bond is a human invention. The bond is a human invention. It is a method of explaining the energy change associated with forming molecules. It is a method of explaining the energy change associated with forming molecules. Bonds don’t exist in nature, but are useful. Bonds don’t exist in nature, but are useful. We have a model of a bond. We have a model of a bond.

49. Exceptions to the octet Less than an Octet Less than an Octet BH 3 BH 3 Be and B often do not achieve octet, but form highly reactive compounds Be and B often do not achieve octet, but form highly reactive compounds More than an Octet More than an Octet SF 6 and I 3 - SF 6 and I 3 - Third row and larger elements can exceed the octet. Third row and larger elements can exceed the octet. How? Use 3d orbitals How? Use 3d orbitals Odd Number of Electrons Odd Number of Electrons NO, NO 2, and ClO 2 NO, NO 2, and ClO 2

50. Exceptions to the octet When we must exceed the octet, extra electrons go on central atom. When we must exceed the octet, extra electrons go on central atom. ClF 3 XeO 3 ICl 4 - BeCl 2

51. Resonance Structures When more than one valid Lewis structure can be written for a particular molecule. When more than one valid Lewis structure can be written for a particular molecule. Actual structure is an average of the depicted resonance structures Actual structure is an average of the depicted resonance structures Drawn by writing the variant structures connected by a double-headed arrow Drawn by writing the variant structures connected by a double-headed arrow

52. What to do when more than one Lewis structure works? Use Formal Charge Assign formal charges on atoms to help decide which is best. Assign formal charges on atoms to help decide which is best. Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work. Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work. Molecules try to achieve as low a formal charge as possible. Molecules try to achieve as low a formal charge as possible. Negative formal charges should be on the most electronegative elements. Negative formal charges should be on the most electronegative elements.

53. Formal Charge Number of valence electrons on the free atom minus number of valence electrons assigned to the atom in the molecule Number of valence electrons on the free atom minus number of valence electrons assigned to the atom in the molecule Lone pair e- belong to atom in question Lone pair e- belong to atom in question Shared e- are divided equally between the sharing atoms Shared e- are divided equally between the sharing atoms The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge on that species The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge on that species If the charge on an ion is -2, the sum of the formal charges must be -2. If the charge on an ion is -2, the sum of the formal charges must be -2.

54. Resonance Examples NO 3 -1 NO 3 -1 SO 3 SO 3 HCO 2 -1 HCO 2 -1

55. Bond Lengths are Averages Have made a table of the averages of different types of bonds pg. 305 Have made a table of the averages of different types of bonds pg. 305 single bond one pair of electrons is shared. single bond one pair of electrons is shared. double bond two pair of electrons are shared. double bond two pair of electrons are shared. triple bond three pair of electrons are shared. triple bond three pair of electrons are shared. More bonds, shorter bond length. More bonds, shorter bond length.

56. How do they share electrons? Localized Electron Model Simple model, easily applied. Simple model, easily applied. A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. Three Parts: 1) Valence electrons using Lewis structures 2) Prediction of geometry using VSEPR 3) Description of the types of orbitals

57. VSEPR: the 3-D shape Lewis structures tell us how the atoms are connected to each other. Lewis structures tell us how the atoms are connected to each other. They don’t tell us anything about shape. They don’t tell us anything about shape. The shape of a molecule can greatly affect its properties. The shape of a molecule can greatly affect its properties. Valence Shell Electron Pair Repulsion Theory allows us to predict geometry Valence Shell Electron Pair Repulsion Theory allows us to predict geometry

58. VSEPR Molecules take a shape that puts electron pairs as far away from each other as possible. Molecules take a shape that puts electron pairs as far away from each other as possible. Have to draw the Lewis structure to determine electron pairs. Have to draw the Lewis structure to determine electron pairs. count # bonding pairs count # bonding pairs count # nonbonding (lone) pairs count # nonbonding (lone) pairs Lone pair take up more space. Lone pair take up more space. Multiple bonds count as one pair. Multiple bonds count as one pair.

59. VSEPR The number of pairs determines The number of pairs determines bond angles bond angles underlying structure underlying structure The number of atoms determines The number of atoms determines actual shape actual shape

60. VSEPR Electron pairs Bond Angles Underlying Shape 2180° Linear 3120° Trigonal Planar 4109.5° Tetrahedral 5 90° & 120° Trigonal Bipyramidal 6 90°Octagonal

61. Actual shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 220linear 330trigonal planar 321bent 440tetrahedral 431trigonal pyramidal 422bent

62. Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 550trigonal bipyrimidal 541See-saw 532T-shaped 523linear

63. Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 660Octahedral 651Square Pyramidal 642Square Planar 633T-shaped 621linear

64. How well does it work? Does an outstanding job for such a simple model. Does an outstanding job for such a simple model. Predictions are almost always accurate. Predictions are almost always accurate. Like all simple models, it has exceptions. Like all simple models, it has exceptions. We’ll spend some time in this class drawing structures and making models to better understand the different VSEPR models We’ll spend some time in this class drawing structures and making models to better understand the different VSEPR models

65. Hybridization The mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies. The mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies. Creates Hybrid Orbitals Creates Hybrid Orbitals Methane goes from 2s 2 2p 2 to 2sp 3 Methane goes from 2s 2 2p 2 to 2sp 3 Draw the sub orbitals according to Hund’s rule Draw the sub orbitals according to Hund’s rule

66. Types of Hybrid Orbitals Two orbital sets = sp Two orbital sets = sp Three orbital sets = sp 2 Three orbital sets = sp 2 Four orbital sets = sp 3 Four orbital sets = sp 3 Five orbital sets = sp 3 d Five orbital sets = sp 3 d Six orbital sets = sp 3 d 2 Six orbital sets = sp 3 d 2

67. Bond Types Sigma Bonds (  ) Sigma Bonds (  ) Bond in which the electron pair is shared in an area centered on a line running between the atoms Bond in which the electron pair is shared in an area centered on a line running between the atoms Lobes of bonding orbital point toward each other Lobes of bonding orbital point toward each other Pi Bonds (  ) Pi Bonds (  ) Electron pair above and below the  bond Electron pair above and below the  bond Created by overlapping of non-hybridized p orbitals Created by overlapping of non-hybridized p orbitals

68. Single bonds consist of one  bond Single bonds consist of one  bond Double bonds consist of one  bond and one  bond Double bonds consist of one  bond and one  bond Triple bonds consist of one  bond and 2  bonds Triple bonds consist of one  bond and 2  bonds