1. Bonding: General Concepts AP Chemistry Unit 8 Author: BobCatChemistry

2. Types of Chemical Bonds

3. Ionic Bonds Ionic Bonds are formed when an atom that loses electrons relatively easily reacts with an atom that has a high attraction for electrons. Ionic Compounds results when a metal bonds with a nonmetal.

4. Bond Energy Bond energy is the energy required to break a bond. The energy of interaction between a pair of ions can be calculated using Coulomb’s law r = the distance between the ions in nm. Q 1 and Q 2 are the numerical ion charges. Q 1 and Q 2 are the numerical ion charges. E is in joules

5. Bond Energy When the calculated energy between ions is negative, that indicates an attractive force. A positive energy is a repulsive energy. The distance where the energy is minimal is called the bond length.

6. Covalent Bonds Covalent bonds form between molecules in which electrons are shared by nuclei. The bonding electrons are typically positioned between the two positively charged nuclei.

7. Polar Covalent Bonds Polar covalent bonds are an intermediate case in which the electrons are not completely transferred but form unequal sharing. A δ - or δ + is used to show a fractional or partial charge on a molecule with unequal sharing. This is called a dipole.

8. Electronegativity

9. Electronegativity Electronegativity is the ability of an atom in a molecule to attract shared electrons to itself. (electron love) Relative electronegativities are determined by comparing the measured bond energy with the “expected” bond energy. Measured in Paulings. After Linus Pauling the American scientist who won the Nobel Prizes for both chemistry and peace.

10. Electronegativity Expected H-X bond energy=

11. Electronegativity Electronegativity values generally increase going left to right across the periodic table and decrease going top to bottom.

12. Electronegativity and Bond type

13. Bond Polarity and Dipole

14. Dipoles and Dipole Moments A molecule that has a center of positive charge and a center of negative charge is said to be dipolar or to have a dipole moment. An arrow is used to show this dipole moment by pointing to the negative charge and the tail at the positive charge.

15. Dipoles and Dipole Moments Electrostatic potential diagram shows variation in charge. Red is the most electron rich region and blue is the most electron poor region.

16. Dipoles and Dipole Moments

19. Dipole moments are when opposing bond polarities don’t cancel out.

20. Dipoles and Dipole Moments

21. Example Problems For each of the following molecules, show the direction of the bond polarities and indicate which ones have a dipole moment: HCl, Cl 2, SO 3, CH 4, H 2 S

22. HCl

23. Cl 2

24. SO 3

25. CH 4

26. H2SH2SH2SH2S

27. Ions: Electron Configurations and Sizes

28. Electron Configurations of Compounds When two nonmetals react to form a covalent bond, they share electrons in a way that completes the valence electron configurations of both atoms. That is, both nonmetals attain noble gas electron configurations.

29. Electron Configurations of Compounds When a nonmetal and a representative-group metal react to form a binary ionic compounds, the ions form so that the valence electron configuration of the nonmetal achieves the electron configuration of the next noble gas atom and the valence orbitals of the metal are emptied. In this way both ions achieve noble gas electron configurations.

30. Predicting Ionic Formulas To predict the formula of the ionic compound, we simply recognize that the chemical compounds are always electrically neutral. They have the same quantities of positive and negative charges.

31. Sizes of Ions Size of an ion generally follows the same trend as atomic radius. The big exception to this trend is where the metals become nonmetals and the ions switch charge.

32. Sizes of Ions A positive ion is formed by removing one or more electrons from a neutral atom, the resulting cation is smaller than the neutral atom. Less electrons allow for less repulsions and the ion gets smaller.

33. Sizes of Ions An addition of electrons to a neutral atom produces an anion that is significantly larger than the neutral atom. An addition of an electron causes additional repulsions around the atom and therefore its size increases.

34. Energy Effects in Binary Ionic Compounds

35. Lattice Energy Lattice energy is the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid. The lattice energy is often defined as the energy released when an ionic solid forms from its ions. Lattice energy has a negative sign to show that the energy is released.

36. Lattice Energy Example Estimate the enthalpy of lithium fluoride and the changes of energy and lattice energy during formation: 1.Break down LiF into its standard state elements (use formation reaction): Li (s) + ½F 2(g)  LiF (s) Li + (g) + F - (g)  LiF (s)

37. Lattice Energy Example Li (s) + ½F 2(g)  LiF (s) Li + (g) + F - (g)  LiF (s) 2.Use sublimation and evaporation reactions to get reactants into gas form (since lattice energy depends on gaseous state). Find the enthalpies to these reactions: Li (s)  Li (g) 161 kJ/mol Li (g) + ½F 2(g)  LiF (s) Li (g) + ½F 2(g)  LiF (s)

38. Lattice Energy Example Li (g) + ½F 2(g)  LiF (s) Li + (g) + F - (g)  LiF (s) 3.Ionize cation to form ions for bonding. Use Ionization energy for the enthalpy of the reaction. Li (g)  Li + (g) + e - Ionization energy: 520 kJ/mol Li + (g) + ½F 2(g)  LiF (s)

39. Lattice Energy Example Li + (g) + ½F 2(g)  LiF (s) Li + (g) + F - (g)  LiF (s) 4.Dissociate diatomic gas to individual atoms: ½F 2(g)  F (g) ½ Bond dissociation energy of F-F = 154 kJ/ 2 = 77 kJ/mol Li + (g) + F (g)  LiF (s)

40. Lattice Energy Example Li + (g) + F (g)  LiF (s) Li + (g) + F - (g)  LiF (s) 5.Electron addition to fluorine is the electron affinity of fluorine: F (g) + e -  F - (g) -328 kJ/mol Li + (g) + F - (g)  LiF (s)

41. Lattice Energy Example Li + (g) + F - (g)  LiF (s) Li + (g) + F - (g)  LiF (s) 6.Formation of solid lithium fluoride from the gaseous ions corresponds to its lattice energy: Li + (g) + F - (g)  LiF (s) -1047 kJ/mol

42. Lattice Energy Example The sum of these five processes yields the overall reaction and the sum of the individual energy changes gives the overall energy change and the enthalpy of formation: Li (s)  Li (g) Li (g)  Li + (g) + e - ½F 2(g)  F (g) F (g) + e -  F - (g) Li + (g) + F - (g)  LiF (s) 161 kJ 520 kJ 77 kJ -328 kJ -1047 kJ Total = -617 kJ/mol

43. Lattice Energy

44. Lattice energy can be calculated with at form of Coulomb’s law: Q is the charges on the ions and r is the shortest distance between the centers of the cations and anions. k is a constant that depends on the structure of the solid and the electron configurations of the ions.

45. Partial Ionic Character of Covalent Bonds

46. Bond Character Calculations of ionic character: Even compounds with the maximum possible electronegativity differences are not 100% ionic in the gas phase. Therefore the operational definition of ionic is any compound that conducts an electric current when melted will be classified as ionic.

47. Bond Character

48. The Covalent Chemical Bond

49. Chemical Bond Model A chemical bond can be viewed as forces that cause a group of atoms to behave as a unit. Bonds result from the tendency of a system to seek its lowest possible energy. Individual bonds act relatively independent.

50. Example It takes 1652 kJ of energy required to break the bonds in 1 mole of methane. 1652 kJ of energy is released when 1 mole of methane is formed from gaseous atoms. Therefore, 1 mole of methane in gas phase has 1652 kJ lower energy than the total of the individual atoms. One mole of methane is held together with 1652 kJ of energy. Each of the four C-H bonds contains 413 kJ of energy.

51. Example CH 3 Cl contains 1578 kJ of energy: 1 mol of C-Cl bonds + 3 mol (C-H bonds)=1578 kJ C-Cl bond energy + 3 (413 kJ/mol) = 1578 kJ C-Cl bond energy = 339 kJ/mol

52. Properties of Models A model doesn’t equal reality; they are used to explain incomplete understanding of how nature works. Models are often oversimplified and are sometimes wrong. Models over time tend to get over complicated due to “repairs”.

53. Properties of Models Remember that simple models often require restrictive assumptions. Best way to use models is to understand their strengths and weaknesses. We often learn more when models are incorrect than when they are right. Cu and Cr.

54. Covalent Bond Energies and Chemical Reactions

55. Bond Energies Bond energy averages are used for individual bond dissociation energies to give approximate energies in a particular bond. Bond energies vary due to several reasons: multiple bonds, 4 C-H bonds in methane different elements in the molecule, C-H bond in C 2 H 6 or C-H bond in HCCl 3

56. Bond Energy Example CH 4(g)  CH 3(g) + H (g) CH 3(g)  CH 2(g) + H (g) CH 2(g)  CH (g) + H (g) CH (g)  C (g) + H (g) 435 kJ 453 kJ 425 kJ 339 kJ Total 1652 kJ Average 413 kJ

57. Bond Energy Example HCBr 3 HCCl 3 HCF 3 C 2 H 6 380 kJ 430 kJ 410 kJ

58. Average Bond Energies

59. Bond Energy A relationship also exists between the number of shared electron pairs. single bond – 2 electrons double bond – 4 electrons triple bond – 6 electrons

60. Bond Energy Bond energy values can be used to calculate approximate energies for reactions. Energy associated with bond breaking have positive signs Endothermic process Energy associated with forming bonds releases energy and is negative. Exothermic process

61. Bond Energy A relationship exists between the number of shared electron pairs and the bond length. As the number of electrons shared goes up the bond length shortens.As the number of electrons shared goes up the bond length shortens.

62. Bond Energy

63. ΔH = sum of the energies required to break old bonds (positive signs) plus the sum of the energies released in the formation of new bonds (negative signs). D represents bond energies per mole and always has positive signD represents bond energies per mole and always has positive sign n is number of molesn is number of moles

64. Bond Energy Example H 2(g) + F 2(g)  2HF (g) 1 H-H bond, F-F bond and 2 H-F bonds ΔH = D H-H + D F-F – 2D H-F ΔH= (1mol x 432 kJ/mol) + (1mol x 154 kJ/mol) – (2mol x 565 kJ/mol) ΔH = -544 kJ

65. The Localized Electron Bonding Model

66. Localized Electron Model The localized electron model assumes that a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. Electrons are assumed to be localized on a particular atom individually or in the space between atoms.

67. Localized Electron Model Pairs of electrons that are localized on an atom are called lone pairs. Pairs of electrons that are found in the space between the atoms are called bonding pairs

68. Localized Electron Model Three parts of the LE Model: 1.Description of the valence electron arrangement in the molecule using Lewis structures. 2.Prediction of the geometry of the molecule using VSEPR model 3.Description of the type of atomic orbitals used by the atoms to share electrons or hold lone pairs.

69. Lewis Structures

70. The Lewis structure of a molecule show how the valence electrons are arranged among the atoms in the molecule. Named after G. N. Lewis Rules are based on observations of thousands of molecules. Most important requirement for the formation of a stable compound is that the atoms achieve noble gas electron configurations.

71. Lewis Structures Only the valence electrons are included. The duet rule: diatomic molecules can find stability in the sharing of two electrons. The octet rule: since eight electrons are required to fill these orbitals, these elements typically are surrounded by eight electrons.

72. Lewis Structure Steps 1.Sum the valence electrons from all the atoms. Total valence electrons. 2.Use a pair of electrons to form a bond between each pair of bound atoms. 3.Arrange the remaining electrons to satisfy the duet rule for hydrogen and the octet rule for the others. a)Terminal atoms first. b)Check for happiness

73. Examples HF N 2 NH 3 CH 4 CF 4 NO +

74. Exceptions to the Octet Rule

75. Incomplete: An odd number of electrons are available for bonding. One lone electron is left unpaired. Suboctet: Less than 4 pairs of electrons are assigned to the central atom Suboctets tend to form coordinate covalentbonds BH 3 + NH 3

76. Exceptions to the Octet Rule Extended: The central atom has more than 4 pairs of electrons. At the third energy level and higher, atoms may have empty d orbitals that can be used for bonding.

77. General Rules The second row elements C, N, O, and F always obey the octet rule The second row elements B and Be often have fewer than eight electrons around them in their compounds. They are electron deficient and very reactive. The second row elements never exceed the octet rule, since their valence orbitals can only hold 8.

78. General Rules Third-row and heavier elements often satisfy the octet rule but can exceed the octet rule by using their empty valence d orbitals. When writing the Lewis structure for a molecule, satisfy the octet rule for the atoms first. If electrons remain after the octet rule has been satisfied, then place them on the elements having available d orbitals

79. Resonance

80. Resonance Resonance is when more than on valid Lewis structure can be written for a particular molecule. The resulting electron structure of the molecule is given by the average of these resonance structures.

81. Resonance The concept of resonance is necessary because the localized electron model postulates that electrons are localized between a given pair of atoms. However, nature does not really operate this way. Electrons are really delocalized- they move around the entire molecule. The valence electrons in a resonance structure distribute themselves equally and produce equal bonds.

82. Formal Charge Some molecules or polyatomic ions can have several non-equivalent Lewis structures. Example: SO 4 2-Example: SO 4 2- Because of this we assign atomic charges to the molecules in order to find the right structure.

83. Formal Charge The formal charge of an atom in a molecule is the difference between the number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule Formal charge = (# of valence electrons on neutral ‘free atom’) – (# of valence electrons assigned to the atom in the molecule)

84. Formal Charge Assumptions on electron assignment: Lone pair electrons belong entirely to the atom in question. Shared electrons are divided equally between the two sharing atoms.

85. Formal Charge Example SO 4 2- : All single bonds Formal charge on each O is -1 Formal charge on S is 2

86. Formal Charge Example SO 4 2- : two double bonds, two single Formal charge on single bonded O is -1 Formal charge on double bonded O is 0 Formal charge on S is 0

87. Formal Charges 1.Atoms in molecules try to achieve formal charges as close to zero as possible. 2.Any negative formal charges are expected to reside on the most electronegative atoms. If nonequivalent Lewis structures exist for a species, those with formal charges closest to zero and with any negative formal charges on the most electronegative atoms are considered to best describe the bonding in the molecule or ion.

88. Molecular Structure: The VSEPR Model

89. VSEPR Valence shell electron repulsion model is useful in predicting the geometries of molecules formed from nonmetals. The structure around a given atom is determined principally by minimizing electron – pair repulsion.

90. VSEPR From the Lewis structure, count the electron pairs around the central atom. Lone pairs require more room than bonding pairs and tend to compress the angles between the bonding pairs. Multiple bonds should be counted as one effective pair. With a molecule with resonance, all structures should yield the same shape.

91. Linear 180° 3-D linear 3-D linear

92. Trigonal Planer 120° 3-D trigonal planar 3-D trigonal planar 3-D trigonal w/lone pair 3-D trigonal w/lone pair

93. Tetrahedral 109.5° 3-D tetrahedral 3-D tetrahedral

94. Trigonal Pyramidal 107° 3-D tetrahedral 1 lone pair / trigonal pyramidal 3-D tetrahedral 1 lone pair / trigonal pyramidal

95. Bent/V 104.5° 3-D tetrahedral 2 lone pair / bent 3-D tetrahedral 2 lone pair / bent

96. Tetrahedral Arrangements

97. Bipyramidal Arrangements trigonal bipyramidal bipyramidal 1 lone pair / see saw bipyramidal 2 lone pair / T shape bipyramidal 3 lone pair / linear

98. Octahedral Arrangements octahedral octahedral 1 lone pair / square pyramidal octahedral 2 lone pair / square planar

99. Molecules without a central atom The molecular structure of more complicated atoms can be predicted from the arrangement of pairs around the center atoms. A combination of shapes will result that allows for minimum repulsion throughout.

100. Molecules without a central atom

102. The End