Chapter 6. Bonding 6.1 Types of Chemical Bonds 6.2 Electronegativity 6.3 Bond Polarity and Dipole Moments 6.4 Ions: Electron Configurations and Sizes 6.5 Formation of Binary Ionic Compounds 6.6 Partial Ionic Character of Covalent Bonds 6.7 The Covalent Chemical Bond: A Model 6.8 Covalent Bond Energies and Chemical Reactions 6.9 The Localized Electron Bonding Model 6.10 Lewis Structure 6.11 Resonance 6.12 Exceptions to the Octet Rule 6.13 Molecular Structure: The VSEPR Model

What is a Chemical Bond? (1) Bonding is the force of attraction that holds atoms together in an element (N2) or compound (CO2 or NaCl). The distances between bonded atoms are less than those between non-bonded atoms. The forces between bonded atoms are greater than those between non-bonded atoms. The principal types of bonding are ionic, covalent, and metallic.

What is a Chemical Bond? (2) A chemical bond links two atoms or groups of atoms when the forces acting between them are sufficient to lead to the formation of an aggregate (a molecule) with sufficient stability to make it convenient for the chemist to consider it as an independent "molecular species” Paraphrased from Linus Pauling (1967).

Types of Chemical Bonds (1) Ionic Bonds Ionic substances are formed when an atom that loses electrons easily reacts with an atom that gains electrons easily. Na · → Na+ + e- Loss of an electron : Cl : . . - e- + Cl → Gain of an electron : Cl : . . - Combination to form the compound NaCl NaCl Na+ + → 4/19/2017

Types of Chemical Bonds (2) For ionic bonds, the energy of interaction between a pair of ions can be calculated by using Coulomb's law. The energy depends only on distance.

Types of Chemical Bonds (3) Covalent Bond is the sharing of pairs of electrons between atoms. Covalent bonding does not require atoms be the same elements but that they be of comparable electronegativity. Covalent bonds give an the angular relation between the atoms (in polyatomic molecules, does not apply to molecules like H2). H· ·H → H:H or H―H 4/19/2017

Electronegativity (EN) the tendency of an atom in a molecule to attract shared electrons. Shared electrons are closer to atoms with greater electronegativity. Trends in EN In a group (column): EN decreases w/ increasing Z (# of protons) In a period (row): EN increases w/ increasing Z Zumdahl Chapter 13 4/19/2017

The effect of an electric field on hydrogen fluoride molecules. Polar covalent bonds Dipole Moments Bonded atoms share electrons unequally, whenever they differ in Electronegativity HF: The F atom carries negative charge and the H atom positive charge of equal magnitude. The molecules align in an electric field. Polar molecules posses a dipole moment, μ

Covalent/Polar/Ionic Percent Ionic Character Covalent e.g., H2, Cl2 N2 Polar Covalent e.g., HF, H2O Ionic Bond e.g., LiF, NaCl

Ionic vs. Covalent Bonding

Electronegativity The Pauling Electronegativity values as updated by A.L. Allred in 1961.

The Process of Bond Formation (1). Two atoms with (i) unfilled electron shells or (ii) opposite charge, are separated by a great distance. The initial potential energy is zero. The atoms do not sense each other. Via some random process the atoms approach each other. There is an attractive force between the atoms. The PE goes negative. The distance between the atoms reaches the ideal bond distance. The PE reaches a minimum, attraction equals the repulsion. There is no net force. The distance continues to decrease Repulsion goes up sharply. PE goes positive. The atoms are forced back to the ideal bond distance.

The Process of Bond Formation (2).

The Process of Bond Formation (3) When two atoms (or molecules) approach each other, the result depends on the atom (or molecule) type. Cl + Cl both have unfilled shells. A bond can form. K+ + Cl- have opposing charges. A bond can form. Ar + Ar both have filled shells. A bond cannot form.

Core Electrons vs Valence Electrons Electrons can be divided into: Valence electrons (e- in unfilled shells, outermost electrons). Valence electrons participate in bonding. Core electrons (e- in a filled shells). Core electrons do not participate in bonding. Valence electrons in molecules are usually distributed in such a way that each main-group element is surrounded by eight electrons (an octet of electrons). Hydrogen is surrounded by two valence electrons.

Sizes of Ions Cations (+): smaller than parent atom Anions (-): larger than parent atom Isoelectronic series: O2- F- Na+ Mg2+ Al3+ As Z, the nuclear charge, increases the atomic radii decreases Ionic Radii In picometers

Sizes of ions

Ions: Predicting Forumlae of Salts K: [Ar]4s1 Na+: [Ar] Can lose 1 electron Ca: [Ar]4s2 Ca+2: [Ar] Can lose 2 electrons O-2: [Ne] O: [He] 2s22p4 Can gain 2 electrons

Dot Structures: closed shell elements 2 Xe 2, 8, 8, 18, 18 Ne 2 8 Ar 2 8 Rn 2, 8, 8, 18, 18, 32 Kr 2 8 18

Dot Structures: valence electrons and the periodic table

Dot Structures: steps in converting a molecular formula into a Lewis Dot structure Draw the individual atoms/ions - with their valence electrons. H has 1, Carbon has 4, N has 5, O has 6, etc. Determine the total number of valence electrons. Account for the net charge. Step 1 Molecular formula Make the molecule. Place the atom with lowest EN in center (CO2, SO4 ). Step 2 Atom placement Step 3 Make the bonds. Add 2 e- to each bond Follow the octet rule. Add remaining e- (double, triple bonds, account for net charge) Make single bonds Step 4 Impose octet rule to C, N, O, F, Check the formal and total charges, and octet rule. Step 4 Lewis structure

Lewis Dot Structure NF3 Example

Lewis Dot Structure: NF3

Dot Structures: NF3 Molecular formula count valence e- Atom placement Add in valence e- done

Dot Structures Write Lewis Structures for Molecules with One Central Atom PROBLEM: Write a Lewis structure for CCl2F2, one of the compounds responsible for the depletion of stratospheric ozone.

Dot Structures The Lewis Structures for Molecules with One Central Atom PROBLEM: Write a Lewis structure for CCl2F2, one of the compounds responsible for the depletion of stratospheric ozone. SOLUTION: C Cl F :

Dot Structures The Lewis Structure for Molecules with More than One Central Atom PROBLEM: Write the Lewis structure for methanol (molecular formula CH4O, i.e., CH3OH), an important industrial alcohol that is used as a gasoline alternative in car engines. If you drink it you get drunk and then go blind and die. Hydrogen can have only one bond. C and O are bonded. H fills in the rest of the bonds.

Dot Structures H : H C O H : H The Lewis Structure for Molecules with More than One Central Atom PROBLEM: Write the Lewis structure for methanol (molecular formula CH4O, i.e., CH3OH), an important industrial alcohol that is used as a gasoline alternative in car engines. If you drink it you get drunk and then go blind and die. Hydrogen can have only one bond. C and O are bonded. H fills in the rest of the bonds. There are 4(1) + 4 + 6 = 14 valence e-. C forms 4 bonds, O forms 2. Each H forms 1 bond. O has 2 paisr of nonbonding e-. H : H C O H : H

Dot Structures Write Lewis structures for the following: Writing Lewis Structures for Molecules with Multiple Bonds. PROBLEM: Write Lewis structures for the following: (a) Ethylene (C2H4), the most important reactant in the manufacture of polymers. (b) Nitrogen (N2), the most abundant atmospheric gas. Try O2, H2.

Dot Structures The Lewis Structures for Molecules with double or triple bond. PROBLEM: Write Lewis structures for: (a) Ethylene (C2H4), the most important reactant in the manufacture of polymers (b) Nitrogen (N2), the most abundant atmospheric gas If a central atom does not have 8e-, an octet, then a pair of e- can be moved from an adjacent atom to form a multiple bond. PLAN: Ethylene SOLUTION: (a) There are 2(4) + 4(1) = 12 valence e-. H can can form only one bond. C H C H :

Dot Structures The Lewis Structures for Molecules with double or triple bond. PROBLEM: Write Lewis structures for: (a) Ethylene (C2H4), the most important reactant in the manufacture of polymers (b) Nitrogen (N2), the most abundant atmospheric gas If a central atom does not have 8e-, an octet, then a pair of e- can be moved from an adjacent atom to form a multiple bond. PLAN: N2 SOLUTION: (b) N2 has 2(5) = 10 valence e-. Therefore a triple bond is required to make the octet around each N. N : . N : . N :

Write the Lewis Structure for O3 (Ozone) Dot Structures Write the Lewis Structure for O3 (Ozone) Zumdahl Chapter 13 4/19/2017

Resonance Delocalized Electron-Pair Bonding O3 can be drawn in 2 ways - Neither structure is acurate. Reality is hybrid of the two. Resonance structures have the same relative atom placement but a difference in the locations of bonding and nonbonding electron pairs. is used to indicate that resonance occurs.

Formal Charge Selecting the Best Resonance Structure An atom “owns” all of its nonbonding electrons and half of its bonding electrons. Formal charge is the charge an atom would have if the bonding electrons were shared equally. Formal charge of atom = # valence e- - (# unshared electrons + 1/2 # shared electrons) For OC # valence e- = 6 # nonbonding e- = 6 # bonding e- = 2 X 1/2 = 1 Formal charge = -1 For OA # valence e- = 6 # nonbonding e- = 4 # bonding e- = 4 X 1/2 = 2 Formal charge = 0 For OB # valence e- = 6 # nonbonding e- = 2 # bonding e- = 6 X 1/2 = 3 Formal charge = +1

Resonance SAMPLE PROBLEM 10.4 Writing Resonance Structures Write resonance structures for the nitrate ion, NO3-. PROBLEM:

Resonance SAMPLE PROBLEM 10.4 Writing Resonance Structures Write resonance structures for the nitrate ion, NO3-. PROBLEM: Use 5+(3x6)+1=24 valence e-. After you write out the dot structure see if other structures can be drawn in which the electrons can be delocalized over more than two atoms. PLAN: SOLUTION: Nitrate has 1(5) + 3(6) + 1 = 24 valence e- Something is wrong. N does not have an octet; shift a pair of e- to form a double bond

Resonance None of these are the ‘true’ structure. The true structure is a average of all three, and cannot be represented by a simple dot structure. 37

Resonance and Formal Charge Problem: Draw all the resonance structures of NCO-. Determine the relative weights of the resonance structures.

Resonance and Formal Charge Problem: Draw all the resonance structures of NCO-. A B C formal charges -2 +1 -1 -1 The true structure is a weighted average of Forms A, B and C. Form A: Formal charge of -2 on N and +1 on O. O is the most electro-negative atom in the molecule. Low weight. Form B: Formal charge on the N is more negative than that of O. O is the most electro-negative atom in the molecule. Low weight. Form C: Formal charge on O is -1. High weight.

Resonance and Formal Charge Three criteria for estimating the weights of resonance structures: Smaller formal charges (absolute value) give higher weight than larger charges. Opposing charges on adjacent atoms gives a low weight. A negative formal charge is on an electronegative atom (O and sometimes N and S) gives a high weight.

Resonance and Weighted Averages EXAMPLE: NCO- has 3 possible resonance forms - A B C None of these are the ‘true’ structure. The true structure is a weighted average of all three. The weight of C is greater than the weights of A and B. 41

Carbon Monoxide

Dot Structures - Octet Expansion & Contraction PROBLEM: Write Lewis structures for (a) H3PO4 (pick the most heavily weighted structure); (b) BFCl2. PLAN: Draw the Lewis structures for the molecule. Note that these dare exceptiosn to the octet rule. P is a Period-3 element and can have an expanded valence shell. SOLUTION: (a) H3PO4 has two resonance forms and formal charges indicate the more important form. -1 +1 (b) BFCl2 has only 1 Lewis structure. more weight lower formal charges

VSEPR - Valence Shell Electron Pair Repulsion Theory Each electron pair (bonded and non-bonded) around a central atom is located as far away as possible from the others, to minimize electrostatic repulsions. But a double bond counts as one bond. A triple bond counts as one bond. These repulsions maximize the space allowed for each electron pair. The result is five electron-group arrangements of minimum energy seen in a large majority of molecules and polyatomic ions. Linear, trigonal planer, tetrahedral, trigonal bipyramidal, octahedral.

Electron-pair repulsions and the five basic molecular shapes. linear trigonal planar tetrahedral trigonal bipyramidal octahedral

The single molecular shape of the linear electron-group arrangement. Examples: CS2, HCN, BeF2

Figure 10.4 The two molecular shapes of the trigonal planar electron-group arrangement. Class Shape Examples: SO2, O3, PbCl2, SnBr2 Examples: SO3, BF3, NO3-, CO32-

Factors Affecting Actual Bond Angles Bond angles are consistent with theoretical angles when the atoms attached to the central atom are the same and when all electrons are bonding electrons of the same order. 1200 Effect of Double Bonds larger EN 1200 ideal greater electron density

Factors Affecting Actual Bond Angles Bond angles are consistent with theoretical angles when the atoms attached to the central atom are the same and when all electrons are bonding electrons of the same order. 1200 1220 1160 real Effect of Double Bonds larger EN 1200 ideal greater electron density Effect of Nonbonding(Lone) Pairs Lone pairs take more space than bonding pairs.. 950

Factors Affecting Actual Bond Angles Bond angles are consistent with theoretical angles when the atoms attached to the central atom are the same and when all electrons are bonding electrons of the same order. 1200 1220 1160 real Effect of Double Bonds larger EN 1200 ideal greater electron density

Factors Affecting Actual Bond Angles Bond angles are consistent with theoretical angles when the atoms attached to the central atom are the same and when all electrons are bonding electrons of the same order. 1200 1220 1160 real Effect of Double Bonds larger EN 1200 ideal greater electron density Effect of Nonbonding(Lone) Pairs Lone pairs repel bonding pairs more strongly than bonding pairs repel each other. 950

The three molecular shapes of the tetrahedral electron-group arrangement. Figure 10.5 Examples: CH4, SiCl4, SO42-, ClO4- NH3 PF3 ClO3 H3O+ H2O OF2 SCl2

Lewis structures and molecular shapes. Figure 10.6 Lewis structures and molecular shapes.

Figure 10.7 The four molecular shapes of the trigonal bipyramidal electron-group arrangement. PF5 AsF5 SOF4 SF4 XeO2F2 IF4+ IO2F2- XeF2 I3- IF2- ClF3 BrF3

Figure 10.8 The three molecular shapes of the octahedral electron-group arrangement. SF6 IOF5 BrF5 TeF5- XeOF4 XeF4 ICl4-

A summary of common molecular shapes with two to six electron groups.

SAMPLE PROBLEM Predicting Molecular Shapes with Five or Six Electron Groups PROBLEM: Determine the molecular shape and predict the bond angles (relative to the ideal bond angles) of (a) SbF5 and (b) BrF5.

SAMPLE PROBLEM 10.7 Predicting Molecular Shapes with Five or Six Electron Groups PROBLEM: Determine the molecular shape and predict the bond angles (relative to the ideal bond angles) of (a) SbF5 and (b) BrF5. SOLUTION: (a) SbF5 - 40 valence e-; all electrons around central atom will be in bonding pairs; shape is AX5 - trigonal bipyramidal. (b) BrF5 - 42 valence e-; 5 bonding pairs and 1 nonbonding pair on central atom. Shape is AX5E, square pyramidal.

SAMPLE PROBLEM Predicting Molecular Shapes with More Than One Central Atom PROBLEM: Determine the shape around each of the central atoms in acetone, (CH3)2C=O. PLAN: Find the shape of one atom at a time after writing the Lewis structure. SOLUTION: tetrahedral tetrahedral trigonal planar >1200 <1200

The tetrahedral centers of ethane and ethanol. Figure 10.11 The tetrahedral centers of ethane and ethanol. ethanol CH3CH2OH ethane CH3CH3

The orientation of polar molecules in an electric field. Figure 10.12 The orientation of polar molecules in an electric field. Electric field ON Electric field OFF

Dipole moment = μ = QR Zumdahl Chapter 13 4/19/2017

Non-Polar covalent bonding Zumdahl Chapter 13 4/19/2017

SAMPLE PROBLEM 10.9 Predicting the Polarity of Molecules PROBLEM: From electronegativity (EN) values (button) and their periodic trends, predict whether each of the following molecules is polar and show the direction of bond dipoles and the overall molecular dipole when applicable: (a) Ammonia, NH3 (b) Boron trifluoride, BF3 (c) Carbonyl sulfide, COS (atom sequence SCO)

SAMPLE PROBLEM Predicting the Polarity of Molecules continued (b) BF3 has 24 valence e- and all electrons around the B will be involved in bonds. The shape is AX3, trigonal planar. F (EN 4.0) is more electronegative than B (EN 2.0) and all of the dipoles will be directed from B to F. Because all are at the same angle and of the same magnitude, the molecule is nonpolar. 1200 (c) COS is linear. C and S have the same EN (2.0) but the C=O bond is quite polar(DEN) so the molecule is polar overall.

SAMPLE PROBLEM Predicting the Polarity of Molecules continued (b) BF3 has 24 valence e- and all electrons around the B will be involved in bonds. The shape is AX3, trigonal planar. F (EN 4.0) is more electronegative than B (EN 2.0) and all of the dipoles will be directed from B to F. Because all are at the same angle and of the same magnitude, the molecule is nonpolar. 1200 (c) COS is linear. C and S have the same EN (2.0) but the C=O bond is quite polar(DEN) so the molecule is polar overall.

Carbonate, CO3-2 Thiocyanate ion, NCS-1 N2O, laughing gas

SN = 5 (bonded atoms) + 0 (lone pairs) 5: Trigonal bipyramidal The VSEPR Theory SN = 5 (bonded atoms) + 0 (lone pairs) 5: Trigonal bipyramidal Zumdahl Chapter 13 4/19/2017

SN = 6 (bonded atoms) + 0 (lone pairs) 6: Octahedral The VSEPR Theory SN = 6 (bonded atoms) + 0 (lone pairs) 6: Octahedral Zumdahl Chapter 13 4/19/2017

Pyramidal The VSEPR Theory SN = 3 (bonded atoms) + 1 (lone pairs) Zumdahl Chapter 13 4/19/2017

Bent SN = 2 (bonded atoms) + 2 (lone pairs) Zumdahl Chapter 13 4/19/2017

Zumdahl Chapter 13 4/19/2017

Zumdahl Chapter 13 4/19/2017

Zumdahl Chapter 13 4/19/2017

Zumdahl Chapter 13 4/19/2017