Basic Concepts of Chemical Bonding

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1 of Chemical Mr. Matthew Totaro Legacy High School AP Chemistry

2 Chemical Bonds Three basic types of bonds Ionic Electrostatic attraction between ions. Covalent Sharing of electrons. Metallic Metal atoms bonded to several other atoms.

3 Types of Tro: Chemistry: A Molecular Approach, 2/e 3

4 Types of Bonds We can classify bonds based on the kinds of atoms that are bonded together Types of Atoms metals to nonmetals nonmetals to nonmetals metals to metals Tro: Chemistry: A Molecular Approach, 2/e Type of Bond Ionic Covalent Metallic 4 Bond Characteristic electrons transferred electrons shared electrons pooled

5 Lewis Theory One of the simplest bonding theories is called Lewis Theory Lewis Theory emphasizes valence electrons to explain bonding Using Lewis theory, we can draw models called Lewis structures aka Electron Dot Structures Lewis structures allow us to predict many properties of molecules such as molecular stability, shape, size, polarity G.N. Lewis ( ) Tro: Chemistry: A Molecular Approach, 2/e 5

6 Lewis Symbols G.N. Lewis pioneered the use of chemical symbols surrounded with dots to symbolize the valence electrons around an atom. When forming compounds, atoms tend to add or subtract electrons until they are surrounded by eight valence electrons (the octet rule).

7 Determining the Number of Valence Electrons in an Atom The column number on the Periodic Table will tell you how many valence electrons a main group atom has Transition Elements all have two valence electrons. Why? Tro: Chemistry: A Molecular Approach, 2/e 7

8 Lewis Structures of Atoms In a Lewis structure, we represent the valence electrons of main-group elements as dots surrounding the symbol for the element aka electron dot structures We use the symbol of element to represent nucleus and inner electrons And we use dots around the symbol to represent valence electrons pair first two dots for the s orbital electrons put one dot on each open side for first three p electrons then pair rest of dots for the remaining p electrons Tro: Chemistry: A Molecular Approach, 2/e 8

9 Practice Write the Lewis structure for arsenic Tro: Chemistry: A Molecular Approach, 2/e 9

10 Practice Write the Lewis structure for arsenic As is in column 5A, therefore it has five valence electrons. As Tro: Chemistry: A Molecular Approach, 2/e 10

11 Lewis Structures of Ions Cations have Lewis symbols without valence electrons lost in the cation formation Anions have Lewis symbols with eight valence electrons electrons gained in the formation of the anion Tro: Chemistry: A Molecular Approach, 2/e 11

12 Stable Electron Arrangements and Ion Charge Metals form cations by losing enough electrons to get the same electron configuration as the previous noble gas Nonmetals form anions by gaining enough electrons to get the same electron configuration as the next noble gas The noble gas electron configuration must be very stable Tro: Chemistry: A Molecular Approach, 2/e 12

13 Ionic

14 Energetics of Ionic As we saw in the last chapter, it takes 496 kj/mol to remove electrons from sodium.

15 Energetics of Ionic We get 349 kj/mol back by giving electrons to chlorine.

16 Energetics of Ionic But these numbers don t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!

17 Energetics of Ionic Bond Formation The ionization energy of the metal is endothermic Na(s) Na + (g) + 1 e H = +496 kj/mol The electron affinity of the nonmetal is exothermic ½Cl 2 (g) + 1 e Cl (g) H = 244 kj/mol Generally, the ionization energy of the metal is larger than the electron affinity of the nonmetal, therefore the formation of the ionic compound should be endothermic But the heat of formation of most ionic compounds is exothermic and generally large. Why? Na(s) + ½Cl 2 (g) NaCl(s) H f = 411 kj/mol Tro: Chemistry: A Molecular Approach, 2/e 17

18 Ionic & the Crystal Lattice The extra energy that is released comes from the formation of a structure in which every cation is surrounded by anions, and vice versa This structure is called a crystal lattice The crystal lattice is held together by the electrostatic attraction of the cations for all the surrounding anions The crystal lattice maximizes the attractions between cations and anions, leading to the most stable arrangement Tro: Chemistry: A Molecular Approach, 2/e 18

19 Crystal Lattice Electrostatic attraction is nondirectional!! no direct anion cation pair Therefore, there is no ionic molecule the chemical formula is an empirical formula, simply giving the ratio of ions based on charge balance Tro: Chemistry: A Molecular Approach, 2/e 19

20 Lattice Energy The extra stability that accompanies the formation of the crystal lattice is measured as the lattice energy The lattice energy is the energy released when the solid crystal forms from separate ions in the gas state always exothermic hard to measure directly, but can be calculated from knowledge of other processes Lattice energy depends directly on size of charges and inversely on distance between ions Tro: Chemistry: A Molecular Approach, 2/e 20

21 Lattice Energy Lattice energy, then, increases with the charge on the ions. It also increases with decreasing size of ions.

22 Energetics of Ionic By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.

23 Trends in Lattice Energy Ion Size The force of attraction between charged particles is inversely proportional to the distance between them Larger ions mean the center of positive charge (nucleus of the cation) is farther away from the negative charge (electrons of the anion) larger ion = weaker attraction weaker attraction = smaller lattice energy Tro: Chemistry: A Molecular Approach, 2/e 23

24 Lattice Energy vs. Ion Size Tro: Chemistry: A Molecular Approach, 2/e Pearson Education, Inc.

25 Trends in Lattice Energy Ion Charge The force of attraction between oppositely charged particles is directly proportional to the product of the charges Larger charge means the ions are more strongly attracted larger charge = stronger attraction stronger attraction = larger lattice energy Of the two factors, ion charge is generally more important Lattice Energy = 910 kj/mol Lattice Energy = 3414 kj/mol Tro: Chemistry: A Molecular Approach, 2/e 25

26 Example: Order the following ionic compounds in order of increasing magnitude of lattice energy: CaO, KBr, KCl, SrO First examine the ion charges and order by sum of the charges Ca 2+ & O 2-, K + & Br, K + & Cl, Sr 2+ & O 2 (KBr, KCl) < (CaO, SrO) Then examine the ion sizes of each group and order by radius; larger < smaller KBr < KCl < (CaO, SrO < CaO SrO) (KBr, KCl) same cation, Br > Cl (same Group) (CaO, SrO) same anion, Sr 2+ > Ca 2+ (same Group) Tro: Chemistry: A Molecular Approach, 2/e 26

27 Practice Order the following ionic compounds in order of increasing magnitude of lattice energy: MgS, NaBr, LiBr, SrS First examine the ion charges and order by sum of the charges Mg 2+ & S 2-, Na + & Br, Li + & Br, Sr 2+ & S 2 (NaBr, LiBr) < (MgS, SrS) Then examine the ion sizes of each group and order by radius; larger < smaller NaBr < LiBr < (MgS, SrS < MgS SrS) (NaBr, LiBr) same anion, Na + > Li + (same Group) (MgS, SrS) same anion, Sr 2+ > Mg 2+ (same Group) Tro: Chemistry: A Molecular Approach, 2/e 27

28 Ionic Model vs. Reality Lewis theory implies that the attractions between ions are strong Lewis theory predicts ionic compounds should have high melting points and boiling points because breaking down the crystal should require a lot of energy the stronger the attraction (larger the lattice energy), the higher the melting point Ionic compounds have high melting points and boiling points MP generally > 300 C all ionic compounds are solids at room temperature Tro: Chemistry: A Molecular Approach, 2/e 28

29 Practice Which ionic compound below has the highest melting point? KBr (734 ºC) CaCl 2 (772 ºC) MgF 2 (1261 ºC) Tro: Chemistry: A Molecular Approach, 2/e 29

30 Lewis Theory Atoms bond because it results in a more stable electron configuration. more stable = lower potential energy Atoms bond together by either transferring or sharing electrons Usually this results in all atoms obtaining an outer shell with eight electrons octet rule there are some exceptions to this rule the key to remember is to try to get an electron configuration like a noble gas Tro: Chemistry: A Molecular Approach, 2/e 30

31 Octet Rule When atoms bond, they tend to gain, lose, or share electrons to result in eight valence electrons ns 2 np 6 noble gas configuration Many exceptions H, Li, Be, B attain an electron configuration like He He = two valence electrons, a duet Li loses its one valence electron H shares or gains one electron though it commonly loses its one electron to become H + Be loses two electrons to become Be 2+ though it commonly shares its two electrons in covalent bonds, resulting in four valence electrons B loses three electrons to become B 3+ though it commonly shares its three electrons in covalent bonds, resulting in six valence electrons expanded octets for elements in Period 3 or below using empty valence d orbitals Tro: Chemistry: A Molecular Approach, 2/e 31

32 Lewis Theory and Ionic Lewis symbols can be used to represent the transfer of electrons from metal atom to nonmetal atom, resulting in ions that are attracted to each other and therefore bond + Tro: Chemistry: A Molecular Approach, 2/e 32

33 Lewis Theory Predictions for Ionic Lewis theory predicts the number of electrons a metal atom should lose or a nonmetal atom should gain in order to attain a stable electron arrangement the octet rule This allows us to predict the formulas of ionic compounds that result It also allows us to predict the relative strengths of the resulting ionic bonds from Coulomb s Law Tro: Chemistry: A Molecular Approach, 2/e 33

34 Predicting Ionic Formulas Using Lewis Symbols Electrons are transferred until the metal loses all its valence electrons and the nonmetal has an octet Numbers of atoms are adjusted so the electron transfer comes out even Li 2 O Tro: Chemistry: A Molecular Approach, 2/e 34

35 Example: Using Lewis theory to predict chemical formulas of ionic compounds Predict the formula of the compound that forms between calcium and chlorine. Draw the Lewis dot symbols of the elements. Transfer all the valence electrons from the metal to the nonmetal, adding more of each atom as you go, until all electrons are lost from the metal atoms and all nonmetal atoms have eight electrons. Tro: Chemistry: A Molecular Approach, 2/e 35 Ca Ca Cl Cl Cl Ca 2+ CaCl 2

36 Practice Use Lewis symbols to predict the formula of an ionic compound made from reacting a metal, Mg, that has two valence electrons with a nonmetal, N, that has five valence electrons Mg 3 N 2 Tro: Chemistry: A Molecular Approach, 2/e 36

37 Lewis Theory of Covalent Lewis theory implies that another way atoms can achieve an octet of valence electrons is to share their valence electrons with other atoms The shared electrons would then count toward each atom s octet The sharing of valence electrons is called covalent bonding 37

38 Covalent In covalent bonds, atoms share electrons. There are several electrostatic interactions in these bonds: Attractions between electrons and nuclei, Repulsions between electrons, Repulsions between nuclei.

39 Covalent : and Lone Pair Electrons Electrons that are shared by atoms are called bonding pairs Electrons that are not shared by atoms but belong to a particular atom are called lone pairs aka nonbonding pairs pairs O S. O. 39 Lone pairs

40 F Single Covalent Bonds When two atoms share one pair of electrons it is called a single covalent bond 2 electrons One atom may use more than one single bond to fulfill its octet to different atoms H only duet F F F F F 40 H O H O H H

41 Double Covalent Bond When two atoms share two pairs of electrons the result is called a double covalent bond four electrons O O O O Tro: Chemistry: A Molecular Approach, 2/e 41

42 Triple Covalent Bond When two atoms share three pairs of electrons the result is called a triple covalent bond six electrons N N N N Tro: Chemistry: A Molecular Approach, 2/e 42

43 Polar Covalent Bonds Though atoms often form compounds by sharing electrons, the electrons are not always shared equally. Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

44 Electronegativity Electronegativity is the ability of atoms in a molecule to attract electrons to themselves. On the periodic chart, electronegativity increases as you go from left to right across a row. from the bottom to the top of a column.

45 Electronegativity Scale Tro: Chemistry: A Molecular Approach, 2/e 45

46 Polar Covalent Bonds When two atoms share electrons unequally, a bond dipole results. The dipole moment, µ, produced by two equal but opposite charges separated by a distance, r, is calculated: µ = Qr It is measured in debyes (D).

47 Polar Covalent Bonds \ The greater the difference in electronegativity, the more polar is the bond.

48 HF EN 2.1 H F EN 4.0 δ+ H F δ Tro: Chemistry: A Molecular Approach, 2/e Pearson Education, Inc.

49 Electronegativity Difference and Bond Type If difference in electronegativity between bonded atoms is 0, the bond is pure covalent equal sharing If difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent If difference in electronegativity between bonded atoms is 0.5 to 1.6, the bond is polar covalent If difference in electronegativity between bonded atoms is larger than or equal to 1.6, the bond is ionic Percent Ionic Character 4% 51% 100% Electronegativity Difference 49

50 Bond Polarity EN Cl = = 0 Pure Covalent EN Cl = 3.0 EN H = = 0.9 Polar Covalent EN Cl = 3.0 EN Na = = 2.1 Ionic Tro: Chemistry: A Molecular Approach, 2/e 50

51 Water a Polar Molecule stream of water attracted to a charged glass rod stream of hexane not attracted to a charged glass rod 51

52 Tro: Chemistry: A Molecular Approach, 2/e Pearson Education, Inc.

53 Dipole Moments Tro: Chemistry: A Molecular Approach, 2/e 53

54 Percent Ionic Character The percent ionic character is the percentage of a bond s measured dipole moment compared to what it would be if the electrons were completely transferred The percent ionic character indicates the degree to which the electron is transferred Tro: Chemistry: A Molecular Approach, 2/e 54

55 Example: Determine whether an N O bond is ionic, covalent, or polar covalent Determine the electronegativity of each element N = 3.0; O = 3.5 Subtract the electronegativities, large minus small (3.5) (3.0) = 0.5 If the difference is 2.0 or larger, then the bond is ionic; otherwise it s covalent difference (0.5) is less than 2.0, therefore covalent If the difference is 0.5 to 1.9, then the bond is polar covalent; otherwise it s covalent difference (0.5) is 0.5 to 1.9, therefore polar covalent Tro: Chemistry: A Molecular Approach, 2/e 55

56 Lewis Structures Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

57 Writing Lewis Structures PCl 3 Keep track of the electrons: 5 + 3(7) = Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. If it is an anion, add one electron for each negative charge. If it is a cation, subtract one electron for each positive charge.

58 Writing Lewis Structures Keep track of the electrons: 2. The central atom is the least electronegative element that isn t hydrogen. Connect the outer atoms to it by single bonds = 20

59 Writing Lewis Structures 3. Fill the octets of the outer atoms. Keep track of the electrons: 26 6 = 20; = 2

60 Writing Lewis Structures 4. Fill the octet of the central atom. Keep track of the electrons: 26 6 = 20; = 2; 2 2 = 0

61 Writing Lewis Structures 5. If you run out of electrons before the central atom has an octet form multiple bonds until it does.

62 Writing Lewis Structures Then assign formal charges. For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms. Subtract that from the number of valence electrons for that atom: the difference is its formal charge.

63 Writing Lewis Structures The best Lewis structure is the one with the fewest charges. puts a negative charge on the most electronegative atom.

64 Practice Assign formal charges CO 2 H 3 PO 4 SeOF 2 SO 3 2 NO 2 P 2 H 4 64

65 Practice - Assign formal charges CO 2 all 0 SeOF 2 Se = +1 H 3 PO 4 P = +1 rest 0 SO 3 2 S = +1 NO 2 65 P 2 H 4 all 0

66 Exceptions to the Octet Rule There are three types of ions or molecules that do not follow the octet rule: ions or molecules with an odd number of electrons, ions or molecules with less than an octet, ions or molecules with more than eight valence electrons (an expanded octet).

67 Odd Number of Electrons Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

68 Fewer Than Eight Electrons Consider BF 3 : Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine. This would not be an accurate picture of the distribution of electrons in BF 3.

69 Fewer Than Eight Electrons Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

70 Fewer Than Eight Electrons The lesson is: If filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don t fill the octet of the central atom.

71 More Than Eight Electrons The only way PCl 5 can exist is if phosphorus has 10 electrons around it. It is allowed to expand the octet of atoms on the third row or below. Presumably d orbitals in these atoms participate in bonding.

72 More Than Eight Electrons Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

73 More Than Eight Electrons This eliminates the charge on the phosphorus and the charge on one of the oxygens. The lesson is: When the central atom is on the third row or below and expanding its octet eliminates some formal charges, do so.

74 Resonance This is the Lewis structure we would draw for ozone, O 3.

75 Resonance But this is at odds with the true, observed structure of ozone, in which both O O bonds are the same length. both outer oxygens have a charge of 1/2.

76 Resonance One Lewis structure cannot accurately depict a molecule like ozone. We use multiple structures, resonance structures, to describe the molecule.

77 Resonance Just as green is a synthesis of blue and yellow ozone is a synthesis of these two resonance structures.

78 Resonance In truth, the electrons that form the second C O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. They are not localized; they are delocalized.

79 Resonance The organic compound benzene, C 6 H 6, has two resonance structures. It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

80 Rules of Resonance Structures Resonance structures must have the same connectivity only electron positions can change Resonance structures must have the same number of electrons Second row elements have a maximum of eight electrons bonding and nonbonding third row can have expanded octet Formal charges must total same 80

81 Drawing Resonance Structures NO 3-81

82 Covalent Bond Strength Most simply, the strength of a bond is measured by determining how much energy is required to break the bond. This is the bond enthalpy. The bond enthalpy for a Cl Cl bond, D(Cl Cl), is measured to be 242 kj/mol.

83 Average Bond Enthalpies Table 8.4 lists the average bond enthalpies for many different types of bonds. Average bond enthalpies are positive, because bond breaking is an endothermic process.

84 Average Bond Enthalpies Note: These are average bond enthalpies, not absolute bond enthalpies; the C H bonds in methane, CH 4, will be a bit different than the C H bond in chloroform, CHCl 3.

85 Enthalpies of Reaction Yet another way to estimate H for a reaction is to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed. In other words, H rxn = Σ(bond enthalpies of bonds broken) Σ(bond enthalpies of bonds formed)

86 Enthalpies of Reaction CH 4 (g) + Cl 2 (g) CH 3 Cl(g) + HCl(g) In this example, one C H bond and one Cl Cl bond are broken; one C Cl and one H Cl bond are formed.

87 Enthalpies of Reaction So, H = [Σ(C H) + Σ(Cl Cl)] [Σ(C Cl) + Σ(H Cl)] = [(413 kj) + (242 kj)] [(328 kj) + (431 kj)] = (655 kj) (759 kj) = 104 kj

88 Practice Estimate the enthalpy of the following reaction H 2 (g) + O 2 (g) H 2 O 2 (g) Reaction involves breaking 1 mol H H and 1 mol O=O and making 2 mol H O and 1 mol O O bonds broken (energy cost) (+436 kj) + (+498 kj) = +934 kj bonds made (energy release) 2( 464 kj) + ( 142 kj) = kj H rxn = (+934 kj) + ( kj) = 136 kj (Appendix H f = kj/mol) Tro: Chemistry: A Molecular Approach, 2/e 88

89 Bond Lengths The distance between the nuclei of bonded atoms is called the bond length Because the actual bond length depends on the other atoms around the bond we often use the average bond length averaged for similar bonds from many compounds 89

90 Bond Enthalpy and Bond Length We can also measure an average bond length for different bond types. As the number of bonds between two atoms increases, the bond length decreases.

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