Chapter 10 Chemical Bonding Bonding Theories Bonding the

Chapter 10 Chemical Bonding

Bonding Theories • Bonding-- the way atoms attach to make molecules. • Bonding can help us: 1. Predict the shapes of molecules and their properties as well. 2. Design and build molecules with particular sets of chemical and physical properties. 2

Lewis Theory • Uses dots to represent valence electrons. • Arranges bonding between atoms to attain certain sets of stable valence electron arrangements. 3

Lewis Theory • Lewis theory emphasizes the importance of valence electrons. • Uses symbol of element to represent nucleus and inner core electrons. • Uses dots around the symbol to represent valence electrons. Li • Be • • • B • • C • • • N • • • O: • • • : F: • • • : Ne: • • 4

Example—Write the Lewis Symbol for Arsenic, • As is in column 5 A, therefore it has 5 valence electrons. 5

Lewis Bonding Theory • Atoms bond because it results in a more stable electron configuration. • Atoms bond together by either transferring or sharing electrons. • Usually this results in all atoms obtaining an outer shell with 8 electrons. Ø Octet rule. Ø some exceptions to this rule—Li and Be 6

Ionic Bonds • Bonds formed between metals and nonmetals Recall • metals lose electrons--- Cations • Nonmetals gain electrons--- Anions • Therefore electrons must be transferred from atom to another for bonding to occur 7

Example 10. 3—Using Lewis Theory to Predict Chemical Formulas of Ionic Compounds Predict the formula of the compound that forms between calcium and chlorine. ∙∙ ∙∙ ∙∙ Ca ∙ Cl ∙∙ ∙∙ ∙∙ Transfer all the valence electrons from the metal to the nonmetal. ∙ Cl ∙∙ ∙ ∙ Ca Draw the Lewis dot symbols of the elements. Ca 2+ Ca. Cl 2 8

Covalent Bonds • Bond form between two nonmetals. • The type of compound formed are known as molecular or covalent compounds • Atoms share pairs of electrons to attain octets. • Types covalent bonds ØSingle ØDouble ØTriple 9

Single Covalent Bonds • Two atoms share one pair of electrons. H • • H • • H O H • • F F • • F • • • • • O • • • F • • • F 10

Double Covalent Bond • Two atoms sharing two pairs of electrons. Ø 4 electrons. • Shorter and stronger than single bond. O • • O • • • • • O • • O 11

Triple Covalent Bond • Two atoms sharing 3 pairs of electrons. Ø 6 electrons. • • N • • • N • • N • • • • Shorter and stronger than single or double bond. • • • N N 12

Trends in Bond Length and Energy Bond C-C C=C C C C-N C=N C N C-O C=O Length (pm) 154 134 120 147 128 116 143 123 Energy (k. J/mol) 346 602 835 305 615 887 358 799

Bonding and Lone Pair Electrons • Bonding pairs--Electrons that are shared by atoms • Lone pairs-- Electrons that are not shared by atoms but belong to a particular atom ØAlso known as nonbonding pairs. Bonding pairs • • O • • S • • O • • Lone pairs 14

Lewis Structures of Molecules • Some common bonding patterns. Ø C = 4 bonds & 0 lone pairs. Ø 4 bonds = 4 single, or 2 double, or single + triple, or 2 single + double. Ø N = 3 bonds & 1 lone pair. Ø O = 2 bonds & 2 lone pairs. Ø H and halogen = 1 bond. Ø Be = 2 bonds & 0 lone pairs. Ø B = 3 bonds & 0 lone pairs. B C N O F 15

Writing Lewis Structures for Covalent Molecules Example: Write the Lewis Structure of CO 2 16

Lewis Structure CO 2 Write skeletal structure. ØMost metallic atom central. ØH terminal. 17

Information: Given: CO 2 Find: Lewis structure Solution : skeletal → electron distribution → Lewis Example: Write the Lewis structure of CO 2. • Apply the solution map. Ø Count and distribute the valence electrons. ØCount valence electrons. 1 A 2 A 8 A 3 A 4 A 5 A 6 A 7 A C O C=4 O=2∙ 6 Total CO 2 = 16 18

Example: Write the Lewis structure of CO 2. Information: Given: CO 2 Find: Lewis structure Solution Map: formula → skeletal → count distribute valence electrons Ø Count and distribute the valence electrons. C=4 O=2∙ 6 Total CO 2 = 16 Start = 16 e. Used = 4 e. Left = 12 e 19

Example: Write the Lewis structure of CO 2. Information: Given: CO 2 Find: Lewis structure Solution Map: formula → skeletal → electron distribution → Lewis Ø Count and distribute the valence electrons. ØComplete octets by distributing to the outside atoms first. C=4 O=2∙ 6 Total CO 2 = 16 Start = 16 e. Used = 4 e. Left = 12 e. Start = 12 e. Used = 12 e. Left = 0 e 20

Example: Write the Lewis structure of CO 2. Information: Given: CO 2 Find: Lewis structure Solution Map: formula → skeletal → electron distribution → Lewis Ø Count and distribute the valence electrons. ØComplete octets. – If not enough electrons to complete octet of central atom, bring in pairs of electrons from attached atom to make multiple bonds. Start = 12 e. Used = 12 e. Left = 0 e 21

Writing Lewis Structures for Polyatomic Ions • Procedure is the same, the only difference is in counting the valence electrons. • For polyatomic cations, take away one electron from the total for each positive charge. • For polyatomic anions, add one electron to the total for each negative charge. 22

Example NO 3 ─ 1. Write skeletal structure. Ø N is central because it is the most metallic. 2. Count valence electrons. N=5 O 3 = 3∙ 6 = 18 (-) = 1 Total = 24 e 23

Example NO 3 ─, Continued 3. Attach atoms with pairs of electrons and subtract from the total. N=5 O 3 = 3∙ 6 = 18 (-) = 1 Total = 24 e- Electrons Start 24 Used 6 Left 18 24

Example NO 3 ─, Continued 3. Complete octets, outside-in. Ø Keep going until all atoms have an octet or you run out of electrons. N=5 O 3 = 3∙ 6 = 18 (-) = 1 Total = 24 e- Electrons Start 24 Used 6 Left 18 Electrons Start 18 Used 18 Left 0 25

Example NO 3 ─, Continued 5. If central atom does not have octet, bring in electron pairs from outside atoms to share. Ø Follow common bonding patterns if possible. 26

Examples—Lewis Structures • NCl. O 18 e- • H 3 BO 3 24 e- • NO 2 -1 18 e- • H 3 PO 4 32 e- • SO 3 -2 26 e- • P 2 H 4 14 e 27

Exceptions to the Octet Rule • H and Li, lose one electron to form cation. Ø Li now has electron configuration like He. Ø H can also share or gain one electron to have configuration like He. • Be shares two electrons to form two single bonds. • B shares three electrons to form three single bonds. • Expanded octets for elements in Period 3 or below. Ø Using empty valence d orbitals. • Some molecules have odd numbers of electrons. Ø NO 28

Resonance • We can often draw more than one valid Lewis structure for a molecule or ion. • In other words, no one Lewis structure can adequately describe the actual structure of the molecule. • The actual molecule will have some characteristics of all the valid Lewis structures we can draw. 29

Resonance, Continued • Lewis structures often do not accurately represent the electron distribution in a molecule. • Real molecule is a hybrid of all possible Lewis structures. • Resonance stabilizes the molecule. 30

Drawing Resonance Structures 1. Draw first Lewis structure that maximizes octets. 2. Move electron pairs from outside atoms to share with central atoms. 31

Example—Draw Lewis Resonance Structures for CNO− (C Is Central with N and O Attached) C=4 N=5 O=6 (-) = 1 Total = 16 e- 32

Molecular Geometry • Molecules are three-dimensional objects. • We often describe the shape of a molecule with terms that relate to geometric figures. • The geometric figures also have characteristic angles that we call bond angles. 33

Electron Groups—what are they? • Each lone pair of electrons constitutes one electron group on a central atom. • Each bond constitutes one electron group on a central atom. Ø regardless of whether it is single, double, or triple • Electron groups determine geometry 34

Linear Geometry • When there are two electron groups the geometry/shape of the molecule will be linear. (true only in cases that the electron groups are bonds) • The bond angle is 180°. 35

Trigonal Planar Geometry • When there are three electron groups around the central atom, the geometry/shape the molecule will be trigonal planar. (true only in cases that the electron groups are bonds) • The bond angle is 120°. 36

Tetrahedral Geometry • When there are four electron groups around the central atom, the geometry/shape of the molecule will be tetrahedral. (true only in cases that the electron groups are bonds) • The bond angle is 109. 5°. 37

The Effect of Lone Pairs • Lone pair groups “occupy more space” on the central atom. • relative sizes of repulsive force interactions are: Lone Pair – Lone Pair > Lone Pair – Bonding Pair > Bonding Pair – Bonding Pair • These repulsions affect bond angles. ØMakes them to be smaller! 38

Derivative Shapes • If the electron groups are all bonds, the molecule’s shape will be one of basic molecular geometries— linear, trigonal planar tetrahedral. • Molecules with lone pairs or different kinds of surrounding atoms will have distorted bond angles and different bond lengths, but the shape will be a derivative of one of the basic shapes. 39

Derivative of Trigonal Geometry • Three electron groups around the central atom, and one of them is a lone pair, the resulting molecule is called a trigonal bent shape. • The bond angle is <120°. 40

Derivatives of Tetrahedral Geometry • Four electron groups around the central atom, and one is a lone pair, ---trigonal pyramidal shape. • bond angle is <109. 5°. 41

Bond Angle Distortion from Lone Pairs 42

Derivatives of Tetrahedral Geometry • Four electron groups around the central atom, and two are lone pairs, --- tetrahedral bent shape. • bond angle is <109. 5°. 43

Bond Angle Distortion from Lone Pairs 44

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Predicting the Shapes Around Central Atoms Example: Predict the shape and bond angle of Cl. O 2− 1. Draw the Lewis structure. 2. Determine the number of electron groups around the central atom. 4 electron groups---ie 2 bonds and 2 lone pairs 46

Predicting the Shapes Around Central Atoms Example: Predict the shape of Cl. O 2− 2. Determine the number of electron groups around the central atom. 4 electron groups---ie 2 bonds and 2 lone pairs 3. Use electron groups to determine shape. Tetrahedral bent 4. From your shape, predict the angle. ~105 o or less than 109 o 47

Practice —Predict the Shape Around the Central Atom • H 3 BO 3 H 3 PO 4 NO 2 -1 SO 32− 48

Sketching a Molecule • Because molecules are three-dimensional objects, our drawings should indicate their three-dimensional nature • By convention: ØA filled wedge indicates that the attached atom is coming out of the paper toward you. ØA dashed wedge indicates that the attached atom is going behind the paper away from you. 49

Sketching a Molecule, Continued 50

Electronegativity • Electronegativity—The ability of an atom to pull electrons towards itself in a chemical bond. d+ H — F d. Periodic trends • Increases across the period (left to right). • Decreases down the group (top to bottom). 51

Electronegativity, Continued 2. 1 1. 0 1. 5 2. 0 2. 5 3. 0 3. 5 4. 0 0. 9 1. 2 1. 5 1. 8 2. 1 2. 5 3. 0 0. 8 1. 0 1. 3 1. 5 1. 6 1. 5 1. 8 1. 9 1. 6 1. 8 2. 0 2. 4 2. 8 0. 8 1. 0 1. 2 1. 4 1. 6 1. 8 1. 9 2. 2 1. 9 1. 7 1. 8 1. 9 2. 1 2. 5 0. 7 0. 9 1. 1 1. 3 1. 5 1. 7 1. 9 2. 2 2. 4 1. 9 1. 8 1. 9 2. 0 2. 2 0. 7 0. 9 1. 1 52

Electronegativity Determines the Type of bond Difference in Electronegativity Type of Bond 0. 0 -0. 3 pure covalent. 0. 4 -1. 9 polar covalent. 2. 0 and above ionic 53

Bond Types 3. 0 -3. 0 = 0. 0 4. 0 -2. 1 = 1. 9 Covalent Ionic Polar Pure 0 3. 0 -0. 9 = 2. 1 0. 4 2. 0 Electronegativity difference 4. 0 54

Bond Polarity • Bonding between unlike atoms results in unequal sharing of the electrons. ØOne atom pulls the electrons in the bond closer to its side. ØOne end of the bond has larger electron density than the other. • The result is bond polarity. ØThe end with the larger electron density gets a partial negative charge and the end that is electron deficient gets a partial positive charge. d+ H • • Cl d- 55

Dipole Moments • A dipole is a material with positively and negatively charged ends. • Polar bonds or molecules have one end slightly positive, d+, and the other slightly negative, d-. • Dipole moment, m, is a measure of the size of the polarity. 56

Polarity of Molecules • In order for a molecule to be polar it must: 1. Have polar bonds. 2. The bond dipole MUST not Cancel each other. The O—C bond is polar. The bonding electrons are pulled equally toward both O ends The net result is a nonpolar molecule. 57

Molecule Polarity, Continued The H—O bond is polar. Both sets of bonding electrons are pulled toward the O end of the molecule. The net result is a polar molecule. 58

Practice: Determine if the following molecules are polar 1) BF 3 2) HCl 59

Molecular Polarity Affects Solubility in Water • Polar molecules are attracted to other polar molecules. • Since water is a polar molecule, other polar molecules dissolve well in water. ØAnd ionic compounds as well. • Some molecules have both polar and nonpolar parts. 60

Recommended Study Problems Chapter 10 NB: Study problems are used to check the student’s understanding of the lecture material. Students are EXPECTED TO BE ABLE TO SOLVE ALL THE SUGGESTED STUDY PROBLEMS. If you encounter any problems, please talk to your professor or seek help at the HACC-Gettysburg learning center. Questions from text book Chapter 10 16, 17, 25, 29, 31, 33, 35, 39, 49, 51, 53, 57, 63, 65, 71, 73, 81, 85, 87, 89, 93 ANSWERS -The answers to the odd-numbered study problems are found at the back of your textbook 61