Chapter 9 Lewis Theory of Chemical Bonding Lewis Bonding Theory Emphasizes valence electrons to explain bonding Lewis structures - Electron Dot Structures Lewis structures allow us to predict many properties of molecules - molecular stability, shape, size, polarity Why Do Atoms Bond? A chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms. To calculate this potential energy, you need to consider the following interactions: nucleus–to–nucleus repulsions electron–to–electron repulsions nucleus–to–electron attractions Types of Bonds Types of Atoms Type of Bond metals to nonmetals Ionic nonmetals to nonmetals Covalent metals to metals Metallic Bond Characteristic electrons transferred electrons shared electrons pooled Types of Bonding Ionic Bonds Metal atoms lose an electrons and become cations. Nonmetal atoms gain electrons and become anions. The oppositely charged ions are then attracted to each other, resulting in an ionic bond. Covalent Bonds Nonmetal atoms have relatively high ionization energies, so it is difficult to remove electrons from them. When nonmetals bond together, it is better in terms of potential energy for the atoms to share valence electrons. The potential energy is lowest when the electrons are between the nuclei. Atoms held together because shared electrons are attracted to both nuclei. Metallic Bonds The relatively low ionization energy of metals allows them to lose electrons easily. Metal atoms release their valence electrons to be shared as a pool by all the atoms/ions in the metal. An organization of metal cation islands in a sea of electrons. Bonding results from attraction of cation for the delocalized electrons. Valence Electrons & Bonding Because valence electrons are held most loosely, and Because chemical bonding involves the transfer or sharing of electrons between two or more atoms, Valence electrons are the most important in bonding. 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. IA IIA IIIA IVA VA VIA VIIA VIIIA Li Be B C N O F Ne 1e- 2e- 3e- 4e- 5e- 6e- 7e- 8e- Lewis Structures of Atoms We represent the valence electrons of main-group elements as dots surrounding the symbol for the element. IA H VIIIA IIA III A IVA VA VIA VIIA He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar Practice – Write the Lewis structure for Arsenic •• • As • • A lithium ion A fluoride ion Stable Electron Arrangements and Ion Charge Main-group ions and the noble gas configurati 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. Lewis Bonding Theory ⇒ Octet Rule When atoms bond, they tend to gain, lose, or share electrons to result in eight valence electrons ns2np6 (noble gas configuration) Exceptions H, Li, Be, B attain an electron configuration like He He = two valence electrons (a duet) Li loses its one valence electron H may share or gain one electron It commonly loses its one electron to become H+ Be loses two electrons to become Be2+ It commonly shares its two electrons in covalent bonds, resulting in four valence electrons B loses three electrons to become B3+ It commonly shares its three electrons in covalent bonds, resulting in six valence electrons Expanded octets for elements in Period 3 or below Lewis Theory and Ionic Bonding 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. Sodium Chloride Formation Lewis Theory Predictions for Ionic Bonding Lewis theory predicts the number of electrons a metal atom should lose or a nonmetal atom should gain. 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. 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. Li2O Use Lewis theory to predict the chemical formula of calcium chloride · Cl ·· · Cl ·· ·· ·· ·· ·· · · Ca Ca2+ CaCl2 Use Lewis symbols to predict the formula of an ionic compound made from reacting a metal, M, that has two valence electrons with a nonmetal, X, that has five valence electrons 3M2+ M3X2 2X3Sr3N2 Energetics of Ionic Bond Formation The ionization energy of the metal is endothermic: Na(s) → Na+(g) + 1 e ─ " ΔH° = +496 kJ/mol The electron affinity of the nonmetal is exothermic: ½Cl2(g) + 1 e ─ → Cl─(g)" ΔH° = −244 kJ/mol Therefore the formation of the ionic compound should be endothermic. Na(s) + ½Cl2(g) → NaCl(s)" ΔH°f = + “something” But the heat of formation of most ionic compounds is exothermic and generally large. Na(s) + ½Cl2(g) → NaCl(s)" ΔH°f = −411 kJ/mol Why? Ionic Bonding & the Crystal Lattice The extra energy that is released comes from the formation of a structure in which every cation is surrounded by anions. This structure is called a crystal lattice. The crystal lattice is held together by electrostatic attractions. The crystal lattice maximizes these attractions between cations and anions, leading to the most stable arrangement. Crystal Lattice Electrostatic attraction is nondirectional!! There is no direct anion–cation pair “bond” Therefore, there is no ionic molecule. The chemical formula for an ionic compound is an empirical formula, simply giving the ratio of ions based on charge balance. 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 1) Always exothermic 2) Can be calculated from knowledge of other processes Lattice energy depends directly on size of charges and inversely on distance between ions. Practice – Given the information below, determine the lattice energy of NaCl Na(s) → Na(g) +108 kJ ½ Cl2(g) → Cl(g) +½(244 kJ) Na(g) → Na+(g) +496 kJ Cl (g) → Cl−(g) −349 kJ Na(s) + ½ Cl2(g) → NaCl(s) −411 kJ ? Na+ (g) + Cl−(g) → NaCl(s) ΔH (NaCl lattice) Determining Lattice Energy The Born–Haber Cycle The Born–Haber Cycle is a hypothetical series of reactions that represents the formation of an ionic compound from its constituent elements. The reactions are chosen so that the change in enthalpy of each reaction is known except for the last one, which is the lattice energy. Naº (s) + ½ Cl2 (g) NaCl (s) Born–Haber Cycle for NaCl ΔH°f (metal atoms, g) separating atoms ΔH°f (nonmetal atoms, g) ΔH°f (salt) ΔH°f (cations, g) forming ions ΔH°f (anions, g) ΔH°(crystal lattice) forming lattice Born–Haber Cycle Use Hess’s Law to add up enthalpy changes of other reactions to determine the lattice energy. ΔH°f(salt) = ΔH°f(metal atoms, g) + ΔH°f(nonmetal atoms, g) + ΔH°f(cations, g) + ΔH°f(anions, g) + ΔH°(crystal lattice) ΔH°f(NaCl, s) = ΔH°f(Na atoms,g) ΔH°f(NaCl, s) = ΔH°f [Na(s)--->Na(g)] (Heat of vaporization) + ΔH°f(Cl atoms,g) + ΔH°f (Cl–Cl bond energy) + ΔH°f(Na+,g) + Na 1st Ionization Energy + ΔH°f(Cl−,g) + Cl Electron Affinity + ΔH°(NaCl lattice) + NaCl Lattice Energy Na(s) → Na(g) +108 kJ ½ Cl2(g) → Cl(g) +½(244 kJ) Na(g) → Na+(g) +496 kJ Cl (g) → Cl−(g) −349 kJ Na+ (g) + Cl−(g) → NaCl(s) ΔH (NaCl lattice) Na(s) + ½ Cl2(g) → NaCl(s) −411 kJ (measured in an experiment) NaCl Lattice Energy = ΔH°f(NaCl, s) − [ΔH°f(Na atoms,g) + ΔH°f(Cl–Cl bond energy) + Na 1st Ionization Energy + Cl Electron Affinity] NaCl Lattice Energy = (−411 kJ) − [(+108 kJ) + (+122 kJ) + (+496 kJ) + (−349 kJ) ] = −788 kJ Practice – Given the information below, determine the lattice energy of MgCl2 Mg(s) ➔ Mg(g) ΔH1°f = +147.1 kJ/mol ½ Cl2(g) ➔ Cl(g) ΔH2°f = +122 kJ/mol Mg(g) ➔ Mg+(g) ΔH3°f = +738 kJ/mol Mg+(g) ➔ Mg2+(g) ΔH4°f = +1450 kJ/mol Cl(g) ➔ Cl−(g) ΔH5°f = −349 kJ/mol Mg(s) + Cl2(g) ➔ MgCl2(s) ΔH6°f = −641 kJ/mol Practice – Given the information below, determine the lattice energy of MgCl2 Mg(s) ➔ Mg(g) 2{½ Cl2(g) ➔ Cl(g)} Mg(g) ➔ Mg+(g) Mg+(g) ➔ Mg2+(g) 2{Cl(g) ➔ Cl−(g)} ΔH1°f = +147.1 kJ/mol 2ΔH2°f = 2(+122 kJ/mol) ΔH3°f = +738 kJ/mol ΔH4°f = +1450 kJ/mol 2ΔH5°f = 2(−349 kJ/mol) − 2+ Mg (g) + 2 Cl (g) ➔ MgCl2(s) ΔH° lattice energy = ? kJ/mol Mg(s) + Cl2(g) ➔ MgCl2(s) ΔH6°f = −641 kJ/mol 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 Lattice Energy vs. Ion Size 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 Lattice Energies of Some Ionic Solids (kJ/mole) -- - → ++ + 23+ 2+ 2+ 2(M + X (2M X M2X) (M + X MX3) (M (M + 2X → MX2) MX) 3X → → MX) M2X3) Anions Cations F- Cl- Br- I- O2- Li+ 1036 853 807 757 2,925 Na+ 923 787 747 704 2,695 K+ 821 715 682 649 2,360 Be2+ 3,505 3,020 2,914 2,800 4,443 Mg2+ 2,957 2,524 2,440 2,327 3,791 Ca2+ 2,630 2,258 2,176 2,074 3,401 Al3+ 5,215 5,492 5,361 5,218 15,916 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 (KBr, KCl) < (CaO, SrO) Then examine the ion sizes of each group and order by radius larger < smaller KBr < KCl < SrO < CaO 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 (NaBr, LiBr) < (MgS, SrS) Then examine the ion sizes of each group and order by radius larger < smaller NaBr < LiBr < SrS < MgS Ionic Bonding-Model vs. Observations Lewis theory implies strong attractions between ions. Lewis theory predicts high melting points and boiling points for ionic compounds. 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. Ionic Compounds Melt Melting an ionic solid Ionic Bonding-Model vs. Observations Lewis theory implies that the positions of the ions in the crystal lattice are critical to the stability of the structure Lewis theory predicts that moving ions out of position should therefore be difficult, and ionic solids should be hard Ionic solids are relatively hard (compared to most molecular solids) Ionic Bonding-Model vs. Observations Lewis theory implies that if the ions are displaced from their position in the crystal lattice, that repulsive forces should occur This predicts the crystal will become unstable and break apart. Lewis theory predicts ionic solids will be brittle. Ionic solids are brittle. When struck they shatter. Ionic Bonding-Model vs. Observations To conduct electricity, a material must have charged particles that are able to flow through the material Lewis theory implies that, in the ionic solid, the ions are locked in position and cannot move around Lewis theory predicts that ionic solids should not conduct electricity Ionic solids do not conduct electricity Conductivity of NaCl in NaCl(s), the ions are stuck in position and not allowed to move to the charged rods Ionic Bonding-Model vs. Observations Lewis theory implies that, in the liquid state or when dissolved in water, the ions will have the ability to move around Lewis theory predicts that both a liquid ionic compound and an ionic compound dissolved in water should conduct electricity Ionic compounds conduct electricity in the liquid state or when dissolved in water Conductivity of NaCl in NaCl(aq), the ions are separated and allowed to move to the charged rods