Chemical Bonding PDF

Chemistry 101 Chapter 9 CHEMICAL BONDING · Chemical bonds are strong attractive force that exists between the atoms of a substance Chemical Bonds are commonly class...

Chemistry 101 Chapter 9 CHEMICAL BONDING · Chemical bonds are strong attractive force that exists between the atoms of a substance Chemical Bonds are commonly classified into 3 types: 1. IONIC BONDING Ø Ionic bonds usually form between metals and nonmetals Ø Bonds form by transfer of electrons Ø Ex: NaCl, MgBr2, AlF3, etc 2. COVALENT BONDING Ø Covalent bonds usually form between atoms of nonmetals Ø Nonmetallic atoms may be the same or different Ø Bonds form by sharing of electrons Ø Ex: H2, Cl2, N2, HCl, H2O, NH3, CO2, CCl4 3. METALLIC BONDING Ø Metallic bonds form between metallic ions Ø Bonds are formed through “non­localized” sharing of electrons commonly referred to as “Sea of electrons” Ø Ex: Cu, Al, Na 1 Chemistry 101 Chapter 9 LEWIS ELECTRON – DOT SYMBOLS · When atoms bond, the bonding is usually achieved by the interaction of valence electrons · Recall: Valence electrons are the “s” and “p” electrons of the outermost shell · In Lewis Electron – Dot Symbols the dots represent valence electrons Examples: Na Ga 1s 2s2 2p6 3s1 2 1s 2s 2p6 3s2 3p6 3d10 4s24p1 2 2 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2 3p6 3d10 4s24p1 [Ne] 3s1 [Ar] 3d10 4s24p1 Lewis Electron­Dot Na Ga Symbol Note: Na Ga stands for stands for [Ne] [Ar] 3d10 Noble Gas Core Pseudo­Noble Gas Core 2 2 6 1s 2s 2p 1s2 2s2 2p6 3s2 3p6 3d10 [Ne] [Ar] 3d10 Conclusion: In Lewis Electron – Dot Symbols: Ø dots represent valence electrons Ø symbols of elements represent Noble Gas Cores or Pseudo – Noble Gas Cores 2 Chemistry 101 Chapter 9 LEWIS ­ ELECTRON DOT SYMBOLS FOR REPRESENTATIVE ELEMENTS IA IIA IIIA IVA VA VIA VIIA VIIIA Li Be B C N O F Ne [He]2s [He]2s [He]2s 2p [He]2s 2p [He]2s 2p [He]2s 2p [He]2s 2p [He]2s22p6 1 2 2 1 2 2 2 3 2 4 2 5 2 Li Be B C N O F Ne Na Mg Al Si P S Cl Ar [Ne]3s [Ne]3s [Ne]3s 3p [Ne]3s 3p [Ne]3s 3p [Ne] 3s 3p [Ne] 3s 3p [Ne]3s23p6 1 2 2 1 2 2 2 3 2 4 2 5 3 Na Mg Al Si P S Cl Ar ns1 ns2 ns2 np1 ns2 np2 ns2 np3 ns2 np4 ns2 np5 ns2 np6 X X X X X X X X Number of Dots = Number of Valence Electrons = Group Number 3 Chemistry 101 Chapter 9 IONIC BOND · Ionic bonds are commonly found in “SALTS” · Recall SALTS are made up of : CATIONS and ANIONS any positive ion any negative ion other than H+ other than OH- · Examples of Salts: NaCl KNO3 (NH4)2SO4 NH4C2H3O2 Properties of Salts: Ø crystalline solids at room temperature (RT) Ø have high Melting Points Ø are Strong Electrolytes in molten form and in aqueous solution FORMATION OF AN IONIC BOND Na + Cl Na+ Cl electron transfer 1 [Ne]3s [Ne] 3s23p5 [Ne] [Ne] 3s23p6 [Ar] Oppositely charged ions attract · The Ionic Bond is the electrostatic attraction between the oppositely charged ions. 4 Chemistry 101 Chapter 9 OTHER EXAMPLES OF FORMATION OF IONIC COMPOUNDS Mg + Cl Mg2+ 2 Cl MgCl2 [Ne]3s2 [Ne] 3s23p5 [Ne] [Ar] Ionic Formula Empirical Cl (Simplest) Formula [Ne] 3s23p5 2 Na + S 2 Na+ S 2 Na2S [Ne]3s1 [Ne] 3s23p4 [Ne] [Ar] 2 e- Ionic Empirical Formula (Simplest) Formula 5 Chemistry 101 Chapter 9 2 Al + 3 O 2 Al3+ 3 O 2 Al2O3 [Ne]3s2 [He]2s22p4 [Ne] [Ne] 6 electrons · Ionic Bond forms by the transfer of one or more valence electrons from a metallic atom to a nonmetallic atom transfer valence electrons Metals Nonmetals Metallic Nonmetallic Cation Anion (Electronic (Electronic Configuration Configuration of Noble Gas of Noble Gas preceding the following the the metallic atom nonmetallic atom) 6 Chemistry 101 Chapter 9 ENERGY INVOLVED IN IONIC BONDING Na(s) + ½ Cl2 Na+Cl-(s) DH = ­ 411 kJ/mol NaCl (exothermic) · Where is this energy coming from? The reaction can be considered to take place in 5 steps: Step 1: Sublimation of Sodium (solid Na ® gaseous Na) Na (s) Na (g) DH1 = + 108 kJ/mol Na endothermic Step 2: Dissociation of Chlorine molecules (Cl 2molecules ® Cl atoms) ½ Cl2 Cl (g) DH2 = + 120 kJ/mol Cl endothermic Step 3: Ionization of Sodium (Na atoms ® Na+ ions) Na (g) Na+ (g) + e- DH3 = + 496 kJ/mol Na endothermic (First Ionization Energy of Na) Step 4: Formation of Cl- (Cl atoms ® Cl- ions) Cl (g) + e- Cl- (g) DH4 = –349 kJ/mol Na exothermic (Electron Affinity of Cl Step 5: Formation of NaCl(s) from ions Na+ (g) + Cl- (g) Na+Cl- (s) DH5 = –786 kJ/mol Na exothermic (Lattice Energy of NaCl = U) ­energy given off when ions form a crystalline structure Net Energy Released = + 108 kJ + 120 kJ + 496 kJ – 349 kJ – 786 kJ = – 411 kJ/mol 7 Chemistry 101 Chapter 9 Na+(g) + Cl(g) E. A. of Cl I. E. of Na (–349 kJ/mol) (+ 496 kJ/mol) Na+ (g) + Cl-(g) Na(g) + Cl(g) Dissociation of Cl2 (+ 120 kJ/mol) Na(g) + ½ Cl2(g) Sublimation of Na Lattice Energy (+ 108 kJ/mol) (–786 kJ/mol) Na(s) + ½ Cl2(g) DHf (NaCl) ­ 411 kJ/mol 8 Chemistry 101 Chapter 9 INTERPRETING THE ENERGETICS OF IONIC BONDING · The Net Ionic Reaction of the formation of Na+Cl-(s) is exothermic: Na(s) + ½ Cl2(g) Na+Cl-(s) DHf = ­ 411 kJ/mol Higher Energy Lower Energy Less stable More stable NOTE: 1. The system’s energy is lowered 2. The larger the energy given off: Ø the more stable the ionic compound Ø the easier the ionic compound forms CONCLUSIONS 1. Formation of an Ionic Compound is favored by any factor that increases the Net Energy Released: (a) Low Ionization Energy (characteristic of Metals) (b) High Electron Affinity (characteristic of Nonmetals) (c) High Lattice Energy. The Lattice Energy depends on: Ø respective charges of ions Ø respective sizes of ions 2. Most compounds formed between Metals and Nonmetals are Ionic Compounds Elements on the Elements on the lower left­hand side upper right­hand side of Periodic Table of Periodic Table (Metals with low I.E.) (Nonmetals with high E.A.) (Noble Gases are excluded) 9 Chemistry 101 Chapter 9 IONS OF MAIN GROUP ELEMENTS IA IIA IIIA IVA VA VIA VIIA Period 1 H– Period 2 Li+ covalent covalent covalent N3– O2– F– Period 3 Na+ Mg2+ Al3+ covalent P3– S2– Cl– Period 4 K+ Ca2+ Ga3+ covalent As3– Se2– Br– Period 5 Rb+ Sr2+ In3+ Sn2+ Sn4+ Sb3+ Te2– I– Period 6 Ca+ Ba2+ Tl3+ Tl+ Pb2+ Pb4+ Bi3+ covalent radioactive NOTE: 1. Group IA through Group IVA Ø Charge of Cation = Group Number Ø Atoms lose all their valence electrons (“s” and “p” electrons) Ø Electronic Configuration of cations = Preceding Noble Gas electronic Configuration 2. Group IIIA through Group VA Ø Charge of Cation = Group Number – 2 Atoms lose only their “p” valence electrons ( “s” electrons are kept) Electronic Configuration of cations = Preceding Pseudo Noble Configuration + ns2 Sn – 2e Sn2+ Similarly: Pb2+ [Kr] 5s 4d 5p2 2 10 [Kr] 5s2 4d10 Sb – 3e Sb3+ Similarly: Bi3+ [Kr] 5s2 4d10 5p3 [Kr] 5s2 4d10 3. Group VA through Group VIIA Ø Charge of Anion = Group Number ­ 8 Atoms gain electrons to achieve stable octet (or doublet in the case of H) Electronic Configuration of anions = Noble Gas electronic Configuration following the element. 10 Chemistry 101 Chapter 9 CATIONS OF TRANSITION METALS · These ions are very often colored Loss of electrons Ionic Charge Examples The “ns” electrons are lost first (most common) 2+ Zn2+ Cd2+ 2+ In some cases in addition to the “s” electrons, some “d” and/or larger Fe2+ Fe3+ electrons from an inner shell (n­1) may be lost. than 2+ Unusual charges are due to irregular electronic 2+ Cu2+ Cu+ configurations associated with the stability of half­filled and/or 1+ Ag+ or filled subshells Fe2+ [Ar] -¯ - - - -. 4s 3d Fe0 [Ar] -¯ -¯ - - - - Fe3+ [Ar] - - - - -. 4s 3d 4s 3d Cu 2+ Cu+ Ag+ [Ar] 4s0 3d 9 [Ar] 4s 0 3d 10 [Kr] 5s 0 4d10 11 Chemistry 101 Chapter 9 IONIC RADII Ionic Radius Ø is the radius of the spherical region around the nucleus within which the electrons are most likely to be found. Ø is obtained from known distances between nuclei in ionic crystals (obtained by X­ray diffraction crystallography, a method that makes the positions of the nuclei visible) Determining the Radius of the I - (iodide) ion in LiI (lithium iodide) · Distance between the nuclei of two I - ions = 426 pm 426 pm - Radius of the I ion = ¾¾¾ = 213 pm 2 · From this information the radius of the Li+ ion can be determined from the experimentally determined distance between Li+ ion and the I - ion. d Radius of Li + = d (Li + ­ I - ) – 213 pm 12 Chemistry 101 Chapter 9 RADII OF CATIONS Radii of atoms > Radii of Cations Na0 Na+ 1s2 2s22p6 3s1 1s2 2s22p6 3 shells 2 shells isoelectronic with Ne · Isoelectronic = species having the same number of electrons and electron configuration. · Arrange the following species in order of increasing size: Fe 0 Fe 3+ Fe 2+ 26 p, 26 e 26 p, 23 e 26 p, 24 e [Ar] 4s2 3d 6 [Ar] 3d 5 [Ar] 3d 6 4 shells 3 shells 3 shells weaker stronger interelectron interelectron repulsions repulsions smaller ion larger ion Fe 3+ < Fe 2+ < Fe 0 smallest largest Conclusion: · For cations of the same element (same number of protons): The larger the positive ionic charge, the smaller the cation · Reason: Larger positive ionic charge means fewer electrons 13 Chemistry 101 Chapter 9 Size of Cations in The Periodic Table (Main Group Elements) IA IIA IIIA 2 Li+ (isoelectronic with He) 3 Na+ Mg 2+ Al 3+ (isoelectronic with Ne) 11 protons 12 protons 13 protons 4 K+ Ca 2+ Ga3+ (isoelectronic with Ar) 19 protons 20 protons 5 Rb + Sr2+ In3+ (isoelectronic with Kr) 37 protons 38 protons 6 Cs + Ba2+ Tl 3+ (isoelectronic with Xe) 55 protons 56 protons 14 Chemistry 101 Chapter 9 RADII OF ANIONS Radii of Atoms < Radii of Anions Cl 0 Cl - 1s2 2s22p6 3s23p5 1s2 2s22p6 3s23p6 isoelectronic with Ne 3 shells 3 shells 17 p , 17 e 17p, 18 e Nuclear Charge= + 17 Nuclear Charge = + 17 Holds 17 electrons strongly Holds 18 electrons less strongly · Arrange the following species in order of increasing size: O0 O 2- O- 8 p, 8 e 8 p, 10 e 8 p, 9 e [He] 2s2 2p 4 [He] 2s2 2p 6 [He] 2s2 2p5 2 shells 2 shells 2 shells some strongest stronger interelectron interelectron interelectron repulsions repulsions repulsions larger ion smaller anion O0 < O- < O 2- smallest largest Conclusion: · For anions of the same element (same number of protons): The larger the negative ionic charge, the larger the anion · Reason: Larger negative ionic charge means more electrons. 15 Chemistry 101 Chapter 9 Size of Anions in the Periodic Table VA VIA VIIA Noble Gases 2 N3- O2- F- (isoelectronic with Ne) 7 protons 8 protons 9 protons 3 P3- S2- Cl - (isoelectronic with Ar) 15 protons 16 protons 17 protons 4 As 3 - Se 2 - Br - (isoelectronic with Kr) 33 protons 34 protons 35 protons Te 2 - I- 5 52 protons 53 protons (isoelectronic with Xe) 16