2/16/2018 Ch. 4 - Chemical Bonding: Understanding Climate Change Chapter Outline . 4.1 Types of Chemical Bonds . 4.2 Naming Compounds and Writing Formulas . 4.3 Lewis Structures . 4.4 Electronegativity, Unequal Sharing, and Polar Bonds . 4.5 Vibrating Bonds and the Greenhouse Effect . 4.6 Resonance . 4.7 Formal Charge: Choosing among Lewis Structures . 4.8 Exceptions to the Octet Rule . 4.8 The Lengths and Strengths of Covalent Bonds 1 2/16/2018 Types of Chemical Bonds Ionic Covalent Metallic Chemical Bonds . Types of chemical bonds: • Ionic Bond – electrons transferred: chemical bond resulting from the electrostatic attraction of a cation for an anion. • Covalent Bond – chemical bond that results from a sharing of outermost electrons. • Metallic Bond – chemical bond consisting of the nuclei of metal atoms surrounded by a “sea” of shared electrons. 2 2/16/2018 푄 푥푄 Ionic Bonds 퐸 ∝ 1 2 푒푙 푑 Sample Exercise 4.1: Calculating the electrostatic Potential Energy of an Ionic Bond What is the electrostatic potential energy (Eel) of the ionic bond between a potassium ion and a chloride ion? 푄 푥푄 퐸 = 2.31푥10−19퐽 ∙ 푛푚 1 2 푒푙 푑 3 2/16/2018 Lattice Energy (U) • The energy released when one mole of an ionic compound forms from its free ions in the gas phase. 푄 푥푄 퐸 ∝ 1 2 푒푙 푑 • As charges increase, so does the lattice energy • As the distance between ions decreases, the lattice energy increases Covalent Bonds H2 4 2/16/2018 Metallic Bonds Cu = [Ar]3d104s1 3d10 4s1 Bond Type Summary 5 2/16/2018 Chapter Outline . 4.1 Types of Chemical Bonds . 4.2 Naming Compounds and Writing Formulas (Lab) . 4.3 Lewis Structures . 4.4 Electronegativity, Unequal Sharing, and Polar Bonds . 4.5 Vibrating Bonds and the Greenhouse Effect . 4.6 Resonance . 4.7 Formal Charge: Choosing among Lewis Structures . 4.8 Exceptions to the Octet Rule . 4.8 The Lengths and Strengths of Covalent Bonds 1 1 Sec. 4.3: Lewis Bonding Theory: Lewis Symbols for Atoms • • . Element symbol = nucleus + core electrons • • • O• . Valence electrons are drawn as dots around the symbol . Up to 4 valence electrons are placed around the symbol one at a time; additional electrons are paired up . The result is up to 4 pairs of electrons = octet . NOTE: hydrogen can not have an octet. When forming bonds with other atoms, it can have a • H maximum of 2 electrons in its valence shell 6 2/16/2018 Lewis Symbols and the Periodic Table Group e- configuration # of valence e- Lewis Dot Symbol 1A ns1 1 2A ns2 2 3A ns2np1 3 4A ns2np2 4 5A ns2np3 5 6A ns2np4 6 7A ns2np5 7 Lewis Symbols and the Periodic Table Unpaired dots = bonding capacity. Main Group Elements: Members of same family have same number of valence electrons, and similar bonding capacities. 7 2/16/2018 Lewis Structures of Ionic Compounds - Na Na+ + e [Ne]3s1 [Ne] + Cl - - - Na + e + Cl Cl - Na+ Cl [Ne]3s23p5 [Ne]3s23p6 = [Ar] Sample Exercise 4.8 (Modified) Draw the Lewis symbols of the monatomic ions formed by calcium and oxygen. Then draw the Lewis structure of calcium oxide (CaO). 8 2/16/2018 Lewis Structures of Molecular (Covalent) Compounds A covalent bond is a chemical bond in which two or more electrons are shared by two nonmetals, resulting in an octet for both atoms. For example - Lewis structure of F2 Lewis structure of H2O Guidelines for Drawing Lewis Structures (updated later on with the concept of “formal charge”) 1. Hydrogen is always a terminal atom because it can form only one bond. 2. The CENTRAL ATOM usually has the lowest electron affinity (or electronegativity as defined later) 3. Arrange the atoms geometrically and symmetrically. e.g. CHCl3 9 2/16/2018 Guidelines for Drawing Lewis Structures (updated later on with the concept of “formal charge”) 4. Sum up the total number of valence electrons (use the group number), and calculate the number of pairs. 5. Connect the atoms together so that each atom has an octet (except H). You may have to form multiple bonds. Multiple Bonds – sharing more than one pair of electrons Double bond – two atoms share two pairs of electrons CO2 H2CO Triple bond – two atoms share three pairs of electrons N2 C2H2 10 2/16/2018 Lewis Structures of Charged Species ClO- + NO2 Electronegativity, Unequal Sharing, and Polar Bonds . Electronegativity (): • Ability of an atom to attract bonding electrons. • Periodic trend similar to ionization energy. Electro- negativities 11 2/16/2018 Ionization Energies and Electronegativies EN increases EN decreases across a row. down a column. Polar Covalent Bonds • Unequal sharing of electrons in a covalent bond resulting in an uneven distribution of charge. • Results from differences in electronegativity. • Dipole Moment = polarity indicated by arrow pointing to more “partially negative” end, with a “partially positive” charge on the opposite size 훿+ 훿− H Cl 12 2/16/2018 Polar Covalent Bonds - Unequal sharing of electrons resulting in an uneven distribution of charge. +1 훿+ 훿− -1 Difference Bond Type 0 < 0.4 Covalent Cl2 0.4 < EN < 2 Polar Covalent HCl 2 Ionic NaCl Electronegativity Trends . As seen previously, electronegativity increases moving up to the right in the periodic table. (Noble gases not included.) . Bond polarity increases as ΔEN increases. ΔEN = 1.9 0.9 0.7 0.4 13 2/16/2018 Sample Exercise 4.12 Rank, in order of increasing polarity, the bonds formed between - O and C Cl and Ca N and S O and Si Are any of these bonds considered ionic? Chapter Outline . 4.1 Types of Chemical Bonds . 4.2 Naming Compounds and Writing Formulas . 4.3 Lewis Structures . 4.4 Electronegativity, Unequal Sharing, and Polar Bonds . 4.5 Vibrating Bonds and the Greenhouse Effect . 4.6 Resonance . 4.7 Formal Charge: Choosing among Lewis Structures . 4.8 Exceptions to the Octet Rule . 4.8 The Lengths and Strengths of Covalent Bonds 2 8 14 2/16/2018 The Concept of “Resonance” Resonance structures (or hybrids) = some molecules have two or more plausible Lewis structures. The actual structure is “intermediate” or an “average” of all the possibilities. e.g. O3 Resonance Structures: Ozone . There are two plausible Lewis structures for ozone. The actual structure for O3 is the “average” of the two resonance structures. 15 2/16/2018 Sample Exercise 4.13: Drawing Resonance Structures Draw all the resonance structures of the nitrate ion. Resonance in Organic Compounds Benzene = C6H6 and it forms a ring 3 2 16 2/16/2018 Chapter Outline . 4.1 Types of Chemical Bonds . 4.2 Naming Compounds and Writing Formulas . 4.3 Lewis Structures . 4.4 Electronegativity, Unequal Sharing, and Polar Bonds . 4.5 Vibrating Bonds and the Greenhouse Effect . 4.6 Resonance . 4.7 Formal Charge: Choosing among Lewis Structures . 4.8 Exceptions to the Octet Rule . 4.8 The Lengths and Strengths of Covalent Bonds 3 3 Formal Charge and Lewis Structures . Three possible Lewis structures. Which one is best? . Formal charge (FC): FC = # valence electrons – - [ (# lone pair e ) + ½ (# shared)] 3 4 17 2/16/2018 Calculating Formal Charges FC = # valence electrons – [ (# lone pair e-) + ½ (# shared)] Choosing the Best Lewis Structure . Most Stable Resonance Structures: • Formal charges equal or close to zero. • Negative formal charges on the more electronegative element. 18 2/16/2018 Practice: Formal Charge What is the most likely Lewis structure for CO2? Summary: Formal Charge and Lewis Structures A way to decide which Lewis Structure, out of several possibilities, that’s the most stable and therefore the most likely. 1. For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. 2. Lewis structures with large formal charges are less plausible than those with small formal charges. 3. Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. 19 2/16/2018 Chapter Outline . 4.1 Types of Chemical Bonds . 4.2 Naming Compounds and Writing Formulas . 4.3 Lewis Structures . 4.4 Electronegativity, Unequal Sharing, and Polar Bonds . 4.5 Vibrating Bonds and the Greenhouse Effect . 4.6 Resonance . 4.7 Formal Charge: Choosing among Lewis Structures . 4.8 Exceptions to the Octet Rule . 4.8 The Lengths and Strengths of Covalent Bonds 3 9 Exceptions to the Octet Rule 1. The Incomplete Octet (Central atom in Group 3A) e.g. BF3 20 2/16/2018 Exceptions to the Octet Rule 2. Odd-Electron Molecules – “free radicals” (one unpaired electron = high reactivity) NO NO2 Exceptions to the Octet Rule 3. Expanded Octet – central atom has greater than 8 valence electrons surrounding it. Occurs only with elements in row 3 and higher because they have available d-orbitals A. Molecular (Covalent) Molecules SF6 21 2/16/2018 Exceptions to the Octet Rule B. Polyatomic Ions 3- 2- PO4 SO4 Chapter Outline . 4.1 Types of Chemical Bonds . 4.2 Naming Compounds and Writing Formulas . 4.3 Lewis Structures . 4.4 Electronegativity, Unequal Sharing, and Polar Bonds . 4.5 Vibrating Bonds and the Greenhouse Effect . 4.6 Resonance . 4.7 Formal Charge: Choosing among Lewis Structures . 4.8 Exceptions to the Octet Rule . 4.8 The Lengths and Strengths of Covalent Bonds 4 4 22 2/16/2018 The Lengths and Strengths of Covalent Bonds: Bond Length vs Bond Order Bond Bond Bond Bond Length Energy Order (pm) (kJ/mol) C-C 1 154 348 C=C 2 134 614 CC 3 120 839 C-N 1 143 293 C=C 2 138 615 CN 3 116 891 Resonance Structures have an intermediate bond order Bond order = 1.5 Bond order = 2 Bond order = 1 23 2/16/2018 Table of Average Covalent Bond Lengths and Bond Energies – can be used to estimate Hrxn Bond energy = the average energy required to break a particular type of bond.