Ch 9 Models of Chemical Bonding.pdf

Ch. 9 – Models of Chemical Bonding Chemistry The Molecular Nature of Matter and Change Ninth Edition Martin S. Silberberg and Patricia G. Amateis ©McGraw-Hill Education. All rights reserved. Authorized only for instructor use in the classroom. No reproduction or further distribution permitted withou...

Ch. 9 – Models of Chemical Bonding Chemistry The Molecular Nature of Matter and Change Ninth Edition Martin S. Silberberg and Patricia G. Amateis ©McGraw-Hill Education. All rights reserved. Authorized only for instructor use in the classroom. No reproduction or further distribution permitted without the prior written consent of McGraw-Hill Education. Models of Chemical Bonding 9.1 Atomic Properties and Chemical Bonds. 9.2 The Ionic Bonding Model (no Haber-Born cycle, no lattice energy) 9.3 The Covalent Bonding Model. 9.4 Bond Energy and Chemical Change. 9.5 Between the Extremes: Electronegativity and Bond Polarity. © McGraw Hill Why Do Atoms Bond? Chemical bonds between atoms form because the bonds lower the potential energy of the valence electrons. 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, the following interactions must be considered in setting up quantum-mechanical calculations: – Nucleus-electron attractions – Nucleus-nucleus repulsions – Electron-electron repulsions © 2017 Pearson Education, Inc. Types of Chemical Bonding Ionic bonding involves the transfer of electrons and is usually observed when a metal bonds to a nonmetal. Covalent bonding involves the sharing of electrons and is usually observed when a nonmetal or metalloid bonds to a nonmetal. Metallic bonding involves electron pooling and occurs when a metal bonds to another metal. ©McGraw-Hill Education. Three Models of Chemical Bonding Figure 9.2 ©McGraw-Hill Education. Gradations in Bond Type Among Period 3 and Group 4A Elements Figure 9.3. Gradations in bond type among Period 3 (black type) and Group 4A (red type) elements. ©McGraw-Hill Education. Lewis Model (1916, before Quantum Mech.) 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. – Also known as electron-dot structures Lewis structures allow us to predict bond types in molecules. Valence electrons are held most loosely. Chemical bonding involves the transfer or sharing of electrons between two or more atoms. In the Lewis model, a chemical bond is the sharing or transfer of valence electrons to attain stable electron configurations for the bonding atoms. G.N. Lewis – “The Atom and the Molecule” (J. Am. Chem. Soc., 38:762 – 1916) © 2017 Pearson Education, Inc. Gilbert N. Lewis 1875-1946 American chemist Lewis Electron-Dot Symbols In a Lewis electron-dot symbol of an element: – The element symbol represents the nucleus AND inner electrons – Dots around the symbol represent the valence electrons. To draw the Lewis symbol for any main-group element: Note the A-group number, which gives the number of valence electrons. Place one dot at a time on each of the four sides of the element symbol. Keep adding dots, pairing them, until all are used up. Example: Nitrogen, N, is in Group 5A and therefore has 5 valence electrons. ©McGraw-Hill Education. Lewis Dot Symbols for Some Elements Figure 9.4 ©McGraw-Hill Education. Lewis Symbols and Bonding For a metal, the total number of dots in the Lewis symbol is the number of electrons the atom may lose to form a cation. For a nonmetal, the number of unpaired dots equals: – Either the number of electrons the atom gains to form an anion – or the number it shares to form covalent bonds. The octet rule states that when atoms bond, they lose, gain, or share electrons to attain a filled outer level of 8 electrons (or 2, a duet, for H and Li). ©McGraw-Hill Education. 9.2 The Ionic Bonding Model An ionic bond is formed when a metal transfers electrons to a nonmetal to form ions, which attract each other to give a solid compound. The total number of electrons lost by the metal atom(s) equals the total number of electrons gained by the nonmetal atom(s). ©McGraw-Hill Education. Sample Problem 9.1 – Problem, Plan and Solution Depicting Ion Formation PROBLEM: Use partial orbital diagrams and Lewis symbols to depict the formation of Na+ and O2– ions from the atoms, and determine the formula of the compound formed. PLAN: Draw orbital diagrams and Lewis symbols for Na and O atoms. To attain filled outer levels, Na loses one electron and O gains two. Two Na atoms are needed for each O atom so that the number of electrons lost equals the number of electrons gained. SOLUTION: ©McGraw-Hill Education. 9.3 Covalent Bonding Covalent compounds make up the overwhelming majority of compounds in Chemistry, from diatomic molecules to macromolecules. Covalent bonding means sharing of valence electrons between participating atoms. What does that mean? What happens to the atomic orbitals in the process of forming a covalent bond? ©McGraw-Hill Education. Covalent Bond Formation in H2 1) Hydrogen atoms are far apart (no interaction). 2) Atoms are closer to each other. The electron from one atom experiences the pull from the nucleus of the other atom. 3) Bottom of the potential well: the optimal distance between atoms called the equilibrium bond length. At this distance, the attraction forces between electrons and nuclei exactly compensate the repulsion forces between nuclei. 4) At shorter distances, the electrostatic repulsion between the 2 nuclei increases the potential energy. ©McGraw-Hill Education. Potential energy curve (red) between 2 Hydrogen atoms. Distribution of Electron Density in H2 Figure 9.13 A. Top view electron density contour map is high around and between the nuclei (blue shading). Notice the elliptical symmetry of the electron density around nuclei located at the 2 foci of the ellipse. Electron density doubles with each concentric curve. B. Side view of electron density contour map. The highest regions of electron density are shown as peaks. ©McGraw-Hill Education. Bonding Pairs and Lone Pairs Atoms share electrons to achieve a full outer level of electrons. The shared electrons are called a shared pair or bonding pair. The shared pair is represented as a pair of dots or a line: An outer-level electron pair that is not involved in bonding is called a lone pair, or unshared pair. ©McGraw-Hill Education. Single Covalent Bonds When two atoms share one pair of electrons, it is called a single covalent bond. One atom may use more than one single bond to fulfill its octet. Hydrogen can only achieve a duet. It is a terminal atom, cannot be in the middle of a chain of atoms. © 2017 Pearson Education, Inc. Double Covalent Bond When two atoms share two pairs of electrons the result is called a double covalent bond. – Four electrons are shared. © 2017 Pearson Education, Inc. Triple Covalent Bond When two atoms share three pairs of electrons the result is called a triple covalent bond. – Six electrons are shared. : 𝐍 ≡ 𝐍: © 2017 Pearson Education, Inc. Properties of a Covalent Bond The bond order is the number of electron pairs being shared by a given pair of atoms. A single bond consists of one bonding pair (two electrons) and has a bond order of 1. A double bond consists of two bonding pairs (four electrons) and has a bond order of 2. A triple bond consists of three bonding pairs (six electrons) and has a bond order of 3. The bond energy (BE) is the energy needed to overcome the attraction between the nuclei and the shared electrons. The stronger the bond the higher the bond energy. The bond length is the distance between the nuclei of the bonded atoms. © McGraw Hill Trends in Bond Order, Energy, and Length For a given pair of atoms, a higher bond order results in a shorter bond length and higher bond energy. For a given pair of atoms, a shorter bond is a stronger bond. Bond length increases down a group in the periodic table and decreases across the period. Bond energy shows the opposite trend. © McGraw Hill Bond Length and Covalent Radius Figure 9.14 © McGraw Hill Single vs Double vs Triple Average Bond Lengths Bond length is the equilibrium distance between nuclei. Multiple bonds are shorter than single bonds. ©McGraw-Hill Education. Bond Average Bond Length (pm) C─O 143 C═O 123 C≡O 113 C─C 154 C═C 134 C≡C 121 N─N 146 N═N 122 N≡N 110 Average Bond Energies (kJ/mol) and Lengths (pm) © McGraw Hill Trends in Bond Lengths In general, the more electrons two atoms share, the shorter the covalent bond. – Must be comparing bonds between like atoms – C≡C (120 pm) < C═C (134 pm) < C—C (154 pm) – C≡N (116 pm) < C═N (128 pm) < C—N (147 pm) Generally, bond length decreases from left to right across period, just like the atomic radii. – C—C (154 pm) > C—N (147 pm) > C—O (143 pm) Generally, bond length increases down the column. – F—F (144 pm) > Cl—Cl (198 pm) > Br—Br (228 pm) In general, as bonds get longer, they also get weaker (see the Bond Energies in the table). © 2017 Pearson Education, Inc. Sample Problem 9.3 Problem and Plan Comparing Bond Length and Bond Strength PROBLEM: Without referring to Table 9.2, rank the bonds in each set in order of decreasing bond length and decreasing bond strength: (a) S─F, S─Br, S─Cl (b) C═O, C─O, C≡O PLAN: (a) S is singly bonded to three different halogen atoms, so the bond order is the same. Bond length increases and bond strength decreases as the halogen’s atomic radius increases. (b) The same two atoms are bonded, but the bond orders differ. In this case, bond strength increases and bond length decreases as bond order increases. ©McGraw-Hill Education. Sample Problem 9.3 - Solution SOLUTION: (a) Atomic size increases going down a group, so F < Cl < Br. Bond length: S–Br > S–Cl > S–F Bond strength: S–F > S–Cl > S–Br (b) By ranking the bond orders, we get Bond length: C–O > C=O > CΞO Bond strength: CΞO > C=O > C–O ©McGraw-Hill Education. 9.4 Bond Energy (Bond Enthalpy) The bond energy (BE) is the energy needed to overcome the attraction between the nuclei and the shared electrons. The stronger the bond the higher the bond energy. Bond Energy (also known as the bond dissociation energy, D) is the standard enthalpy change for breaking a bond in 1 mol of gaseous molecules (to 1 mol of each gaseous individual atoms or radicals): 𝑘𝐽 𝐶𝑂 𝑔 → 𝐶 𝑔 + 𝑂 𝑔 𝐷 𝐶 ≡ 𝑂 = 1070 𝑚𝑜𝑙 𝑘𝐽 𝐻2 𝑂 𝑔 → 𝐻 𝑔 + 𝑂𝐻 𝑔 𝐷 𝑂 − 𝐻 = 467 𝑚𝑜𝑙 As bond dissociation is always an endothermic process, the bond energy is always a positive number. ©McGraw-Hill Education. Bond Energies and DH°rxn The heat released or absorbed during a chemical change is due to differences between the bond energies of reactants and products: 1. First, the heat is absorbed by reactants to break the existing bonds (ΔH > 0). 2. Second, heat is released by products in forming new bonds (∆H