Chemistry 4th Edition McMurry/Fay
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1 7 Chapter Covalent Bonding Chemistry 4th Edition McMurry/Fay Dr. Paul Charlesworth Michigan Technological University The Covalent Bond 01 Covalent bonds are formed by sharing at least one pair of electrons. Chapter 07 Slide 2 The Covalent Bond 02 Every covalent bond has a characteristic length that leads to maximum stability. This is the bond length. Chapter 07 Slide 3
2 The Covalent Bond 03 Energy required to break a covalent bond in an isolated gaseous molecule is called the bond dissociation energy. Chapter 07 Slide 4 Polar Covalent Bonds 01 Bond polarity is due to electronegativity differences between atoms. Pauling Electronegativity: is expressed on a scale where F = 4.0 Chapter 07 Slide 5 Polar Covalent Bonds 02 Pauling Electronegativities Chapter 07 Slide 6
3 Polar Covalent Bonds 03 % Ionic Character: As a general rule for two atoms in a bond, we can calculate an electronegativity difference (?EN ):?EN = EN(Y) EN(X) for X Y bond. If?EN < 0.5 the bond is covalent. If?EN < 2.0 the bond is polar covalent. If?EN > 2.0 the bond is ionic. Chapter 07 Slide 7 Polar Covalent Bonds 04 Using electronegativity values, predict whether the following compounds are nonpolar covalent, polar covalent, or ionic: SiCl 4 CsBr FeBr 3 C 4 Cl CCl 4 N 3 2 O Chapter 07 Slide 8 Electron-Dot Structures 01 Using electron-dot (Lewis) structures, the valence electrons in an element are represented by dots. Valence electrons are those electrons with the highest principal quantum number (n). Chapter 07 Slide 9
4 Electron-Dot Structures 02 The electron-dot structures provide a simple, but useful, way of representing chemical reactions. Ionic: Covalent: Chapter 07 Slide 10 Electron-Dot Structures 03 Single Bonds: C C Double Bonds: C C C C Triple Bonds: C C C C Chapter 07 Slide 11 Drawing Lewis-Dot Structures 01 Rule 1: Count the total valence electrons. Rule 2: Draw the structure using single bonds. Rule 3: Distribute the remaining electron pairs around the peripheral atoms. Rule 4: Put remaining pairs on central atom. Rule 5: Share lone pairs between bonded atoms to create multiple bonds. Chapter 07 Slide 12
5 Drawing Lewis-Dot Structures 02 N 2 F Amino Fluoride: is the central atom. In this molecule, nitrogen Rule 1: Number of electrons = 5 + (2 x 1) + 7 = 14 = 7 pairs N F N F N F Rule 2 Rule 3 Rule 4 Chapter 07 Slide 13 Drawing Lewis-Dot Structures 03 Chapter 07 Slide 14 Drawing Lewis-Dot Structures 04 Polyatomic molecules with central atoms below the second row ten: In this compound there are 10 valence electrons on bromine; this is called an expanded octet. The extra pairs go into unfilled d orbitals. Chapter 07 Slide 15
6 Drawing Lewis-Dot Structures 05 Draw electron-dot structures for: C O 2 CO 2 N 2 4 C 5 N C 2 4 C 2 2 Cl 2 CO Draw electron-dot structures for: SF 4 SF 6 XeOF 4 XeF 5 + XeF 4 3 S + CO 3 Chapter 07 Slide 16 Resonance Structures 01 ow is the double bond formed in O 3? Move lone pair from this oxygen? O O O Or from this oxygen? O O O O O O The correct answer is that both are correct, but neither is correct by itself. or Chapter 07 Slide 17 Resonance Structures 02 When multiple structures can be drawn, the actual structure is an average of all possibilities. The average is called a resonance hybrid. A straight double-headed arrow indicates resonance. O O O O O O Chapter 07 Slide 18
7 Resonance Structures 03 The nitrate ion, NO 3, has three equivalent oxygen atoms, and its electronic structure is a resonance hybrid of three electron-dot structures. Draw them. Draw as many electron-dot resonance structures as possible for: SO 2, CO 3 2, CO 2, SO 4 2, PO 4 3. Chapter 07 Slide 19 Formal Charge 01 Formal Charge: Determines the best resonance structure. We determine formal charge and estimate the more accurate representation. Chapter 07 Slide 20 Formal Charge 02 Formal Charge= # of Valence e - # of bonding ē # of nonbonding - e 2 Calculate the formal charge and determine the most favorable of the following electron dot structures: SO 2 NO 3 NCO N 2 O O 3 CO 2 3 Chapter 07 Slide 21
8 Molecular Shapes: VSEPR 01 The approximate shape of molecules is given by Valence- Shell Electron-Pair Repulsion (VSEPR). Chapter 07 Slide 22 Molecular Shapes: VSEPR 02 Step 01: Count the total electron groups. Step 02: Arrange electron groups to maximize separation. Groups are collections of bond pairs between two atoms or a lone pair. Groups do not compete equally for space: Lone Pair > Triple Bond > Double Bond > Single Bond Chapter 07 Slide 23 Molecular Shapes: VSEPR 02 Two Electron Groups: Electron groups point in opposite directions. Chapter 07 Slide 24
9 Molecular Shapes: VSEPR 03 Three Electron Groups: Electron groups lie in the same plane and point to the corners of an equilateral triangle. Chapter 07 Slide 25 Molecular Shapes: VSEPR 06 Four Electron Groups: Electron groups point to the corners of a regular tetrahedron. Chapter 07 Slide 26 Molecular Shapes: VSEPR 09 Five Electron Groups: Electron groups point to the corners of a trigonal bipyramid. Chapter 07 Slide 27
10 Molecular Shapes: VSEPR 11 Six Electron Groups: Electron groups point to the corners of a regular octahedron. Chapter 07 Slide 28 Molecular Shapes: VSEPR 12 Electron Groups Lone Pairs Bonds Geometry Examples Linear BeCl Trigonal planar BF Bent SO Tetrahedral C Trigonal pyramidal N Bent 2O Trigonal bipyramidal PCl See-saw SF T-Shaped ClF linear I Octahedral SF Square pyramidal SbCl Square planar XeF4 Chapter 07 Slide 29 Molecular Shapes: VSEPR 13 Draw the Lewis electron-dot structure and predict the shapes of the following molecules or ions: O 3 3 O + XeF 2 PF 6 XeOF 4 Al 4 BF 4 SiCl 4 ICl 4 AlCl 3 Chapter 07 Slide 30
11 Valence Bond Theory Covalent bonds are formed by overlapping of atomic orbitals, each of which contains one electron of opposite spin. 2. Each of the bonded atoms maintains its own atomic orbitals, but the electron pair in the overlapping orbitals is shared by both atoms. 3. The greater the amount of orbital overlap, the stronger the bond. Chapter 07 Slide 31 Valence Bond Theory 02 Linus Pauling: Wave functions from s orbitals & p orbitals could be combined to form hybrid atomic orbitals. Chapter 07 Slide 32 Valence Bond Theory 03 sp hybrid: Chapter 07 Slide 33
12 Valence Bond Theory 04 sp 2 hybrid: Chapter 07 Slide 34 Valence Bond Theory 05 sp 2 hybrid (p bond): Chapter 07 Slide 35 Valence Bond Theory 06 sp 3 hybrid: Chapter 07 Slide 36
13 Valence Bond Theory 07 sp 3 d hybrid: Chapter 07 Slide 37 Valence Bond Theory 08 sp 3 d 2 hybrid: Chapter 07 Slide 38 Molecular Orbital Theory 01 The molecular orbital (MO) model provides a better explanation of chemical and physical properties than the valence bond (VB) model. Atomic Orbital: Probability of finding the electron within a given region of space in an atom. Molecular Orbital: Probability of finding the electron within a given region of space in a molecule. Chapter 07 Slide 39
14 Molecular Orbital Theory 02 Additive combination of orbitals (σ) is lower in energy than two isolated 1s orbitals and is called a bonding molecular orbital. Chapter 07 Slide 40 Molecular Orbital Theory 03 Subtractive combination of orbitals (σ ) is higher in energy than two isolated 1s orbitals and is called an antibonding molecular orbital. Chapter 07 Slide 41 Molecular Orbital Theory 04 Molecular Orbital Diagram for 2 : Chapter 07 Slide 42
15 Molecular Orbital Theory 05 Molecular Orbital Diagrams for 2 and e 2 : Chapter 07 Slide 43 Molecular Orbital Theory 06 Additive and subtractive combination of p orbitals leads to the formation of both sigma and pi orbitals. Chapter 07 Slide 44 Molecular Orbital Theory 07 Second-Row MO Energy Level Diagrams: Chapter 07 Slide 45
16 Molecular Orbital Theory 08 MO Diagrams Can Predict Magnetic Properties: Chapter 07 Slide 46 Molecular Orbital Theory 09 Bond Order is the number of electron pairs shared between atoms. Bond Order is obtained by subtracting the number of antibonding electrons from the number of bonding electrons and dividing by 2. Chapter 07 Slide 47 Molecular Orbital Theory 10 Construct an MO diagram for e 2 + ion. Determine the bond order and whether the ion is likely to be stable? The B 2 and C 2 molecules have MO diagrams similar to N 2. What MOs are occupied in B 2 and C 2, and what is the bond order in each? Would any of these be paramagnetic? Chapter 07 Slide 48
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