Covalent Bonding. Slide 1 / 186. Slide 2 / 186. Slide 3 / 186. Table of Contents: Covalent Bonding. Covalent versus Ionic Bonds

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1 Slide 1 / 186 Slide 2 / 186 Covalent Bonding Note: Students and classrooms with ipads should download the free "Lewis Dots" App and can use that on all the slides where Lewis Dot drawings are to be done. Table of Contents: Covalent Bonding Slide 3 / 186 Click on the topic to go to that section Covalent versus Ionic Bonds Properties of Ionic and Covalent Materials Naming Binary Molecular Compounds Lewis Structures Resonance Structures VSEPR Theory Molecular Geometry Polarity

2 Slide 4 / 186 Covalent versus Ionic Bonds Return to Table of Contents Covalent Bonding & Molecular Geometry Slide 5 / 186 Examine these two forms of the same compound, ibuprofen. Covalent Bonding & Molecular Geometry Slide 6 / 186 This form of ibuprofen is about 100x more effective at alleviating pain than the other form. This form of ibuprofen has virtually no anti-inflammatory effect. Even though they consist of the exact same number and kinds of atoms, these two molecules have very different chemical properties.

3 Slide 7 / 186 Covalent Bonding & Molecular Geometry Take a look around you. The chemistry of everything you see, hear, feel, touch and taste is a result of not only what it's made of but also how it's put together.(remember this for next year in biology!) In this unit, we will explore what causes molecules to have various shapes. Later, we will then examine how molecular geometry affects different chemical properties. Chemical Bonds Slide 8 / 186 Chemical bonds hold atoms together to create chemical compounds. There are three basic types of bonds: Ionic - The electrostatic attraction between ions Covalent - The sharing of electrons between atoms Metallic - Each metal atom bonds to other metals atoms within a "sea" of electrons (covered in a later unit) Chemical Bonds Slide 9 / 186 ow ionic or covalent a bond is depends on the difference in electronegativity. The smaller the difference, the more likely electrons are "shared" and the bond is considered covalent, the greater the difference, the more likely electrons have been transferred and the atoms are ionized resulting in an ionic bond. Li Be B C N F Electronegativity Bond Li-F Be-F B-F C-F N-F - F-F Electronegativity Increasing Covalent Character

4 Chemical Bonds Slide 10 / 186 We can make a few simplifications... Ionic Bonding Ionic bonds occur when the difference in electronegativity between two atoms is more than 1.7. Na ---- F electronegativity = 3 Covalent Bonding If the difference of electronegativity is less than 1.7, neither atom takes electrons from the other; they share electrons. This type of bonding typically takes place between two non-metals or between two metals Cl electronegativity = 1.1 Ionic v. Covalent Bonding Slide 11 / 186 In the case of ionic bonding, a 3-D lattice of ions is the result... not individual molecules. The chemical formula for an ionic compound is just the ratio of each type of ion in the lattice, not a particular number of ions in a molecule. In contrast, covalent bonding can result in individual molecules or 3-D lattices depending on the elements involved. The bonding and the shapes of these molecules help determine the physical and chemical properties of everything around us! click here for an animation about ionic and covalent bonding Slide 12 / 186

5 Slide 13 / 186 Slide 14 / 186 Slide 15 / 186

6 Slide 16 / 186 Properties of Ionic and Covalent Materials Return to Table of Contents Properties of Ionic Compounds Boiling and Melting Points Slide 17 / 186 Since the attractions between the ions span a short distance, these forces are quite strong resulting in high melting points and boiling points! Na+ -- Cl- it takes a lot of energy to break an ionic lattice! Compound Melting Point (C) NaCl 801 Mg 2852 Properties of Ionic Compounds Conductivity Slide 18 / 186 Since ionic compounds consist of ions, when these ions are free to move, the substance can conduct electricity. To move, they must be in the liquid or molten state. NaCl (s) Lattice is strong, no conductivity Molten NaCl(l) Lattice is broken, ions are free to move and conduct

7 Properties of Metallic Substances Melting and Boiling Points Slide 19 / 186 Metallic compounds are held together by non-directional covalent bonds in which some electrons are shared but are loosely held and free to roam. The covalent bonds between the metal atoms are strong! This gives rise to high melting and boiling points! strong metallic covalent bonds Metallic Lattice Metal Cu Fe Melting Point 1085 C 1585 C REAL WRLD APPLICATIN Slide 20 / 186 In order to obtain pure metals, the ancients had to melt the metal (metallic substance) out of the rock (an ionic compound). Why do you think the bronze age (copper mixed with tin) came before the iron age? Copper has a lower melting point so it could be obtained in furnaces at lower temperatures. Move for Furnaces answer hot enough to extract iron would come later. Properties of Metallic Compounds Conductivity Slide 21 / 186 Since the electrons in metals are free to roam somewhat, metals are good conductors of electricity! Silver is the most conductive metal and is roughly 5-10 times more conductive than steel (mostly iron).

8 REAL WRLD APPLICATIN Slide 22 / 186 Copper is often used in electrical cable rather than silver even though it is roughly 10% less conductive than silver. Why? Copper currently trades Move for roughly for answer 3 dollars an ounce while silver trades for about 30 dollars a month. It's about the money!!!! 5 Which of the following would NT conduct electricity in the solid state? Slide 23 / 186 A Al B Al 2 3 C NaCl D Both A and B E Both B and C Properties of Covalent Network Substances Melting Point and Boiling Point Slide 24 / 186 Like ionic and metallic substances, covalent network solids are giant molecules arranged in 3-D crystalline shapes. ere, the atoms involved tend to semi-metals like Silicon or Germanium or elemental carbon. Since the bonds are covalent, they are quite strong! This gives rise to high melting and boiling points! Glass (75% Si 2) Diamond (pure C) Melts at 1500 C Melts at 3500 C

9 Properties of Covalent Network Substances Conductivity Slide 25 / 186 Since these substances have higher electronegativities, they keep good tabs on their electrons thereby preventing the electrons from moving. As a result they are largely non-conductive. Diamond and graphite are both allotropes or different versions of carbon and vary somewhat in their conductivity. Diamond (C) non-conductive Graphite (C) a little conductive REAL WRLD APPLICATIN Slide 26 / 186 Diamond is notorious for being ARD! This is true for lots of covalent network crystals. Can you think of some applications where hardness is important? Body Armor Drill Bits B 4C (boron carbide) polycrystalline diamond slide for answers 6 Which of the following would be classified as a covalent network solid? Slide 27 / 186 A NaCl B F C C 2 D Ge 2 3 E Fe

10 Molecular Compounds Slide 28 / 186 When atoms are bonded covalently, the atoms are held together by sharing electrons. This occurs between non-metals such as C,,S,,P,N, etc. Unlike in all of the other substances, the atoms form small individual molecules that then interact with each other and their environment. These are called molecular compounds. P Cl Cl Cl = C = In covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases. Both atoms use the shared electrons to reach that goal. Click here to view interactive website Properties of Molecular Substances Melting and Boiling Points Slide 29 / 186 Since these substances contain lots of small molecules, the bonds holding these small molecules together are fundamentally different from the covalent bonds found inside the molecule. They cover a much larger distance and are quite weak giving rise to LW melting and boiling points! weak inter-molecular forces between molecules Properties of Molecular Substances Conductivity Slide 30 / 186 Molecular compounds contain electronegative non-metals and do not lose their electrons easily so they are non-conductive. As a result they are excellent INSULATRS! Rubber: (C 5 9) 250

11 Summary of Substances Slide 31 / 186 Ionic Metallic Cov. Network Molecular metals and nonmetals metals semi-metals and pure carbon non-metals Na 2 Fe C(diamond) C 4 igh MP igh MP igh MP Low MP conduct as liquid conduct in all states non-conductive non-conductive Brittle Malleable Brittle Brittle 7 Which of the following would have the lowest melting point? Slide 32 / 186 A N 2 B C(graphite) C C(diamond) D E W LiF 8 Which of the following will not conduct electricity in any state? Slide 33 / 186 A Cu B NaF C Fe D C 2 E All of these will conduct

12 9 Which of the following consists of small individual molecules? Slide 34 / 186 A C(diamond) B Si 2 C Cu 2 D Na E S 3 10 Which of the following substances has both ionic and covalent bonding within the crystal? Slide 35 / 186 A Cu B CuC 3 C LiCl D Ba E BaF 2 Slide 36 / 186 Naming Binary Molecular Compounds Return to Table of Contents

13 Naming Binary Molecular Compounds Slide 37 / 186 Use prefixes to indicate the number the atoms. All end in "ide" Examples N2 nitrogen dioxide P25 diphosphorous pentoxide ( penta-oxide-->pentoxide) Naming Binary Molecular Compounds Slide 38 / 186 Look on your reference sheets for the prefixes. The atom with the lower electronegativity is usually written first. If there is only one of the first atom, the mono- is left off. Examples C carbon monoxide C2 carbon dioxide 11 Chlorine monoxide is Slide 39 / 186 A Cl 2 B C D Cl Cl 2Cl

14 12 Dinitrogen tetroxide is Slide 40 / 186 A N 2 B N 2 4 C N - 3 D N is A ydrogen monoxide Slide 41 / 186 B C D Dihydrogen monoxide ydrogen oxide ydrogen dioxide 14 S 3 is Slide 42 / 186 A B C D sulfate sulfur oxide sulfur trioxide sulfite

15 15 Mg is Slide 43 / 186 A B C D monomagnesium monoxide magnesium monoxide monomagnesium oxide magnesium oxide 16 P 4 10 is A Phosphorous pentoxide B Tetraphosphorous decoxide C Phosphorous oxide D Phosphate Slide 44 / 186 Slide 45 / 186 Lewis Structures Return to Table of Contents

16 Lewis Structures Slide 46 / 186 Lewis structures are diagrams that show valence electrons as dots. Lewis structures are also known as Lewis dot or electron dot diagrams. Note that no electrons are paired until after the fourth one. 17 ow many valence electrons does nitrogen have? Slide 47 / 186 A 2 B 3 C 4 D 5 E 7 18 The Lewis structure for nitrogen is N Slide 48 / 186 True False

17 The ctet Rule Recall that atoms tend towards having the electron configuration of a noble gas.for most atoms, that means having 8 valence electrons. The ctet Rule also applies to molecular compounds. Slide 49 / 186 In covalent bonding, an atom will share electrons in an effort to obtain eight electrons around it (except hydrogen which will attempt to obtain 2 valence electrons). A pair of valence electrons that is not shared between atoms is called an unshared pair, also known as a lone pair or a nonbonding pair. Exceptions to the ctet Rule needs 2e- Be needs 4e- B needs 6e- ow do electron dot structures represent shared electrons? Slide 50 / 186 An electron dot structure such as : represents the shared pair of electrons of the covalent bond by two dots. + Shared pair of electrons ydrogen atom ydrogen atom ydrogen molecule 1s 1s ydrogen molecule Structural Formulas Slide 51 / 186 A structural formula represents the covalent bonds by dashes and shows the arrangement of covalently bonded atoms. As in the example below, one shared pair of electrons is represented by one dash. ydrogen molecule Shared pair of electrons

18 19 ow many electrons are shared by two atoms to create a single covalent bond? Slide 52 / 186 A 2 B 1 Single Covalent Bonds Slide 53 / 186 The halogens form single covalent bonds in their diatomic molecules. Fluorine is one example. F + F # # # F F R F F Fluorine atom Fluorine atom Fluorine molecule 1s 2s 2p Fluorine molecule 1s 2s 2p Lewis Structure of 2 Slide 54 / 186 In a water molecule, each hydrogen and oxygen atom attains a noble-gas configuration by sharing electrons. The water molecule has two unshared, or lone, pairs of electrons > or ydrogen atoms xygen atom Water molecule 1s 2s 2p 1s 1s Water molecule

19 Lewis Structures of N 3 Slide 55 / 186 In the ammonia molecule, N 3, each atom attains a noblegas configuration by sharing electrons. This molecule has one unshared pair of electrons. 3 + N --> N or N ydrogen atom Nitrogen atom Ammonia molecule 1s 2s 2p N Ammonia molecule 1s 1s 1s Slide 56 / 186 Drawing Lewis Structures Slide 57 / The central atom is the least electronegative element (excluding hydrogen). P has an electronegativity of 2.1 and Cl has an electronegativity of 3.0 P will be the central atom. The Cl atoms will surround the P atom. Cl Cl P Cl The single bonds are shown as single lines. 3. Connect the other atoms to it by single bonds.

20 Drawing Lewis Structures Slide 58 / Count each single bond as a pair (two) of electrons. 5. Add electons to the outer atoms to give each one 8 (a full shell), or just 2 electrons for hydrogen. 6. Do the same for the central atom. 7. Check: Does each atom have a full outer shell (8 except, 2 for hydrogen)? ave you used up all the valence electrons? ave you used too many electrons? Drawing Lewis Structures Slide 59 / Find the total number of valence electrons in the polyatomic ion or molecule. N 3 The N atom has 5 valence electrons and each of the three atoms has 1 so the total number of valence electrons is, 5 + 3(1) = 8 Drawing Lewis Structures Slide 60 / The central atom is the least electronegative element (excluding hydrogen because it can only have one bond). 3. Connect the other atoms to it by single bonds. can never be the central atom so N must be The atoms will surround the N atom. The single bonds are shown as single lines. N 3 N

21 Drawing Lewis Structures Slide 61 / Count each single bond as a pair (two) electrons. Now add electons to the outer atoms to give each one a full shell (2 in the case of ). 5. Next, do the same for the central atom. 6. Check: Does each atom have a full outer shell? 7. ave you used up all the valence electrons you started with? ave you used too many electrons? N Each already has two electrons, so that's done. But we have to add electrons to N to make 8. N 20 ow many total valence electrons does 2 have? Slide 62 / 186 A 8 B 10 C D Which element in 2 is the least electronegative? Slide 63 / 186 A B

22 22 Which of the following is the correct Lewis Structure for 2? Slide 64 / 186 A B C D 1. Find the total number of valence electrons: 2. Central atom is the least electronegative: 3. Connect the other atoms to it by single bonds. 4. Count each single bond as a pair of electrons. 5. Add electrons to the outer atoms to give each one 8 (except only gets 2). 6. Add electrons to the central atom to give it Check to make sure all valence electrons are used. 23 Which of the following is the correct Lewis Structure for C 2 6? A B C C C C Slide 65 / 186 C C C D C C Lewis Structures for ions Slide 66 / 186 If you are drawing the Lewis Structure for an IN... A negative ion has extra electrons, add the charge of the ion to your valence electron count. Cl 2 - has 1(7) + 2(6) + 1 = 20 electrons A positive ion is missing electrons, subtract the charge of the ion to your valence electron count. N 4 + has 1(5) + 4(1) -1 = 8 electrons

23 24 ow many valence electrons does C 3 2- have? Slide 67 / 186 A B C D ow many valence electrons does 3 + have? A B C D Slide 68 / 186 Formal Charge Slide 69 / 186 The "Formal Charge" method tells us how the electrons are distributed within a molecule. For example, depending on how the electrons are shared, some atoms may have more electrons than others resulting in a semi-charged state for that atom. Formal Charge = # of valence electrons - # of electrons atom possesses within the lewis structure. P FC for P: 5-4= +1 (count each bond as one) FC for each : 6-7= -1 (count each bond as one) Note: The charges must add to the charge of the molecule. So for P P atom x +1 = atoms x -1 = = -3

24 Formal Charge The best Lewis structure will have the formal charge = 0 on each atom. owever, if the molecule carries a charge, the more electronegative atoms should carry a charge as they have the greater attraction for electrons! Slide 70 / 186 The oxygen is more electronegative so it makes sense that it carries the negative charge. Each bond is counted as one in a formal charge calculation as each atom forming part of the bond contributes just one electron to that bond. [ - ] -1 FC on = 6-7 = -1 FC on = 1-1 = 0 Formal Charge Slide 71 / 186 Example: Below are two possible lewis structure for the phosphate ion, P Which Lewis structure is considered to more closely represent the actual molecule based on formal charge calculations? P P Structure 1 Structure 2 Structure 2 is superior as all formal charges = 0 whereas in structure 1, the P carries a +1 charge and each oxygen carries a -1 charge slide for answer 26 Which of the following would be the formal charge on the N in the ammonium ion? Slide 72 / 186 A +1 B 0 C -1 D -2 E -3

25 27 In which of the following molecules would N carry a non-zero formal charge? A CN B N 3 C N 3- D N 2- E N 4+ Slide 73 / 186 Lewis Structures Slide 74 / 186 Draw the Lewis dot structure for the sulfate ion, S4 2-, and find the formal charge on each atom. FC on S = 6-4 = +2 FC on = 6-7 = (+2) + 4(-1) = -2 slide for answer Lewis Structures Slide 75 / 186 Draw the Lewis dot structure for the hydronium ion, 3 + and find the formal charge on each atom. FC on = 6-5 = +1 FC on = 1-1 = 0 * note how in this case the more electronegative atom () is carrying a + charge relative to. This demonstrates the theory is imperfect.

26 F N Cl C P Si S B Se Xe I C 2 Draw a Lewis Structure Slide for Answer C We ran out of electrons, but carbon does not have an octet yet! Now What? Slide 76 / 186 Double and Triple Covalent Bonds Slide 77 / 186 Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two pairs or three pairs of electrons. A bond that involves two shared pairs of electrons is a double covalent bond. A bond formed by sharing three pairs of electrons is a triple covalent bond. Double and Triple Covalent Bonds Slide 78 / 186 Carbon Dioxide, C2 1. Determine the # of valence electrons. 1 (4) + 2 (6) = 16 e - 2. Form Single Bonds C This leaves 12 electrons, 6 pairs 3. Place lone pairs on oxygen atoms to give each 8. C

27 Slide 79 / 186 Carbon Dioxide, C 2 4. Check: We had 16 electrons to work with; how many have we used? C 5. There are too many electrons in our drawing. We must form DUBLE BNDS between C and. Instead of sharing only 1 pair, a double bond shares 2 pairs. So one pair is taken away from each atom and replaced with another bond. C C Covalent Bond Length Slide 80 / 186 Covalent Bond Energy Slide 81 / 186 Bond Type Bond Energy Bond Type Bond Energy C C 348 kj N N 163 kj C C 614 kj N N 418 kj C C 839 kj N N 941 kj It requires more energy to break double and triple bonds compared to single bonds. Triple bonds are the strongest of the three.

28 Covalent Bond Energies Slide 82 / 186 Slide 83 / 186 Covalent Bonds Comparison Type of Bond Electrons shared Bond Strength Bond Length 2 weak long 4 intermediate intermediate 6 strong short 28 As the number of bonds between a pair of atoms increases, the distance between the atoms: Slide 84 / 186 A increases B decreases C remains unchanged D varies, depending on the atoms

29 29 As the number of bonds between a pair of atoms increases, the strength of the bond between the atoms: Slide 85 / 186 A increases B decreases C remains unchanged D varies, depending on the atoms 30 As the number of bonds between a pair of atoms increases, the energy of the bond between the atoms: Slide 86 / 186 A increases B decreases C remains unchanged D varies, depending on the atoms 31 ow many electrons are shared by two atoms to create a single bond? Slide 87 / 186

30 32 ow many electrons are shared by two atoms to create a double bond? Slide 88 / ow many electrons are shared by two atoms to create a triple bond? Slide 89 / Using Lewis structure drawings, determine which molecule below would have the shortest bond length between atoms? Slide 90 / 186 A 2 B F 2 C Cl 2 D C E I 2

31 35 Which of the following molecules would have the longest C- bond length? Use Lewis structures. Slide 91 / 186 A C B C 2 C 2C D C 3 E The lengths are all the same Slide 92 / 186 Writing Lewis Structures If you run out of electrons before the central atom has an octet form multiple bonds until it does. Slide 93 / 186 Bonding of 2 xygen molecule + --> or xygen atom xygen atom xygen molecule 1s 2s 2p xygen molecule 1s 2s 2p

32 F N Cl C P Si S B Se Xe I C Draw a Lewis Structure Slide for Answer C Carbon has the lower electronegativity, so we will consider it the "central" atom... Slide 94 / 186 Coordinate Covalent Bonds Slide 95 / 186 Coordinate Covalent Bonds Slide 96 / 186 In carbon monoxide, oxygen has a stable configuration but the carbon does not. C + # # # C Carbon atom xygen atom Carbon monoxide 1s 2s 2p C 1s 2s 2p Carbon monoxide molecule

33 Coordinate Covalent Bonds Slide 97 / 186 A coordinate covalent bond is a covalent bond in which one atom contributes both bonding electrons. In a structural formula, you can show coordinate covalent bonds as arrows that point from the atom donating the pair of electrons to the atom receiving them. In a coordinate covalent bond, the shared electron pair comes from one of the bonding atoms. Carbon has 4 valence electrons, oxygen has 6. Slide 98 / 186 F N Cl C P Si S B Se Xe I F 2 Draw a Lewis Structure Slide for Answer F F Diatomic Molecules Slide 99 / 186 A molecule is a neutral group of atoms joined together by covalent bonds. Air contains oxygen molecules. A diatomic molecule is a molecule consisting of two atoms. Certain elements do not exist as single atoms; they always appear as pairs. When atoms turn into ions, this N LNGER APPENS! Remember: NClBrIF ydrogen Nitrogen xygen Fluorine Chlorine Bromine Iodine 2 N N N 2 2

34 36 n the periodic table below, mark which elements exist as diatomic molecules. Note the pattern. Slide 100 / 186 Exceptions to the ctet Rule Slide 101 / 186 There are three types of ions or molecules that do not follow the octet rule: #1 Ions or molecules with an odd number of electrons #2 Ions or molecules with less than an octet #3 Ions or molecules with more than eight valence electrons (an expanded octet) Exception 1: dd Number of Electrons Slide 102 / 186 Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons. N is an example:

35 Exception 2: Fewer Than Eight Electrons Slide 103 / 186 Beryllium (Be) - this metal is shown to form molecular compounds, rather than ionic compounds as expected; only needs 4 electrons to be stable Boron (B) - only needs 6 electrons to be stable Be B Memorize these exceptions Exception 3: Expanded ctet Slide 104 / 186 The only way PCl5 exists is if phosphorus has 10 electrons around it. This is called an expanded octet. Atoms on the third energy level or higher are allowed to expand their octet to 10 or 12 electrons. These atoms are larger and can accommodate more electrons. Exception 3: Expanded ctet Slide 105 / 186 ow many electrons do these central atoms have around them?

36 Exceptions to the ctet Rule Slide 106 / 186 Draw the Lewis dot structure for sulfur hexaflouride, SF 6: Move for answer Exceptions to the ctet Rule Slide 107 / 186 Draw the Lewis dot structure for the xenon tetrafluoride, XeF4. Move for answer Exceptions to the ctet Rule Slide 108 / 186 Draw the Lewis dot structure for boron trifluoride, BF3: Move for answer

37 Exceptions to the ctet Rule Slide 109 / 186 Draw the Lewis dot structure for the iodine tricholoride, ICl3. Cl - I - Cl Move for answer Cl [*] 37 Which of the following need fewer than 8 valence electrons to be stable? A B C D E Boron and Beryllium Boron and elium Boron, Beryllium, and ydrogen Boron, Beryllium, ydrogen and elium Boron, Beryllium, ydrogen, elium and xygen Slide 110 / The correct lewis structure for BeCl 2 is Slide 111 / 186 Cl - Be - Cl True False

38 39 Elements in the first two rows of the periodic table cannot have expanded octets because their atoms do not have enough space. Slide 112 / 186 True False Slide 113 / 186 Resonance Structures Return to Table of Contents F N Cl C P Si S B Se Xe I 3 Draw a Lewis Structure and use that to determine the VSEPR number Slide for Answer For the central oxygen: Electron domains = 3 Bonding domains = 2 Unpaired electrons = 1 Its VSEPR number is Slide 114 / 186

39 [*] Resonance Slide 115 / 186 Consider the Lewis structure we would draw for ozone, 3: We would expect the double bond to have a shorter bond length than the single bond. owever, the true, observed structure of ozone shows that both - bonds are the same length. ow can this be? [*] Slide 116 / 186 Resonance ne Lewis structure cannot accurately depict a molecule like ozone. Therefore, we use multiple structures, called resonance structures, to describe the molecule. zone has two resonance structures. [*] Resonance The actual ozone molecule is a synthesis of these two resonance structures. The bond length for both outer oxygen atoms falls somewhere between the single and double bond length. Slide 117 / 186 Resonance structure Resonance structure zone molecule

40 [*] Resonance Slide 118 / 186 The nitrate ion, N3 1- also requires resonance structures to explain its covalent bonding. There are three resonance structures for the nitrate ion: [*] Resonance Structures Slide 119 / 186 Draw the Lewis dot structure for S 3: move for answer [*] 40 ow many resonance structures can be drawn for the carbonate ion, C3 2-? A 1 B 2 C 3 D 4 E 5 Slide 120 / 186

41 [*] Benzene The benzene molecule is a regular hexagon of carbon atoms with a hydrogen atom bonded to each one. There are two resonance structures for benzene. Slide 121 / 186 Benzene, C66, is obtained from the distillation of fossil fuels. More than 4 billion pounds of benzene is produced annually in the United States. Because benzene is a carcinogen, its use is closely regulated. [*] Slide 122 / 186 Localized v. Delocalized electrons In truth, the shared pairs of electrons do not always remain between adjacent C atoms. They are not localized. Instead, the electrons are said to be delocalized, meaning that they they can move around the 6-carbon ring. # # # # or Benzene is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring... we will talk more about this at the end of the year when we study organic chemistry. Slide 123 / 186 VSEPR Theory Return to Table of Contents

42 VSEPR Theory Slide 124 / 186 Valence Shell Electron Pair Repulsion According to VSEPR theory, the molecules will adopt a shape/geometry so as to reduce the repulsion between the bonded electrons. Click here to view a PhET simulation VSEPR Numbers Slide 125 / 186 The VSEPR number of a molecule is a three digit number that can be used to determine a molecule's shape. ere's how you find it: 1. Draw the Lewis structure for the molecule. Locate the central atom, if applicable. 2. The first digit of the VSEPR number is the total number of electron-domains around the central atom. Electron domains are either shared pairs of electrons or lone pairs of electrons Multiple bonds (i.e. double or triple bonds) count as only NE electron domain. VSEPR Numbers (cont) Slide 126 / The second digit of the VSEPR number is the total number of bonding-domains around the central atom. Bonding domains are single, double or triple bonds. 4. The third digit of the VSEPR number is the total number of lone pairs around the central atom. Each pair of electrons that are not involved in bonds counts as one lone pair. 5. Check your work - the first digit is equal to the sum of the second and third.

43 41 ow many electron domains does C 4 have? Slide 127 / 186 A 1 B 2 C 3 D 4 E 5 42 ow many electron domains does 2 have? Slide 128 / 186 A 1 B 2 C 3 D 4 E 5 43 ow many electron domains does C 2 have? Slide 129 / 186 A 1 B 2 C 3 D 4 E 5 C

44 F N Cl C P Si S B Se Xe I C 4 Draw a Lewis Structure and use that to determine the VSEPR number C Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 4 Bonding domains = 4 Lone pairs of electrons = 0 Slide 130 / 186 Its VSEPR number is F N Cl C P Si S B Se Xe I NF 3 Draw a Lewis Structure and use that to determine the VSEPR number Slide for Answer F N F F Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 4 Bonding domains = 3 Lone pairs of electrons = 1 Slide 131 / 186 Its VSEPR number is F N Cl C P Si S B Se Xe I SiF 4 Draw a Lewis Structure and use that to determine the VSEPR number Slide for F Answer F Si F Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. F Electron domains = 4 Bonding domains = 4 Lone pairs of electrons = 0 Slide 132 / 186 Its VSEPR number is 4 4 0

45 F N Cl C P Si S B Se Xe I P 4 3- Draw a Lewis Structure and use that to determine the VSEPR number Slide for Answer P Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 4 Bonding domains = 4 Lone pairs of electrons = 0 Slide 133 / 186 Its VSEPR number is F N Cl C P Si S B Se Xe I IF 5 Draw a Lewis Structure and use that to determine the VSEPR number Slide for F Answer F F I F F Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 6 Bonding domains = 5 Lone pairs of electrons = 1 Slide 134 / 186 Its VSEPR number is Slide 135 / 186 Molecular Geometry Return to Table of Contents

46 Slide 136 / 186 VSEPR and molecule shape prediction According to VSEPR theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible. The shape of a molecule plays an important role in determining its chemical and physical properties. To determine a molecule's shape, i.e. its molecular geometry, we must first determine its electron-domain geometry. ow does VSEPR theory help predict the shapes of molecules? Slide 137 / 186 Recall: Electron domains are either shared pairs of electrons or lone pairs of electrons Bonding domains are single, double or triple bonds. Each pair of electrons that are not involved in bonds counts as one lone pair. To determine the electron-domain geometry, look at the first number and use the following chart... Electron Domain Geometry Slide 138 / 186

47 Electron-Domain Geometry (EDG) Slide 139 / 186 The EDG (2,3,4,5,or 6) gives us the general shape of the molecule, as shown here. owever, these domains do not have to be bonds. The molecular geometry tells us if there is a bond or lone pair of electrons present, thereby specializing the general shape. Let's take a closer look... Linear Electron-Domain Geometry Slide 140 / 186 Two atoms around a central one will form a linear shape with bond angles of 180 o Linear Linear Molecular Geometry Slide 141 / 186 There is only one molecular geometry for linear electron-domain: linear molecular geometry (220).

48 Trigonal Planar Electron-Domain Geometry Slide 142 / 186 Three atoms around a central one will form a trigonal planar shape with bond angles of 120 o trigonal planar Trigonal Planar Molecular Geometry Slide 143 / 186 There are two molecular geometries: Trigonal planar, if all the electron domains are bonding (330) Bent, if one of the domains is a nonbonding pair (321) Trigonal Planar Molecular Geometry Slide 144 / trigonal planar (330) bent (321) It is very important to note that unbonded pairs of electrons repel more strongly than bonded electrons thereby shrinking the bond angle between atoms

49 Tetrahedral Electron-Domain Geometry Slide 145 / 186 Four atoms around a central one will form a tetrahedral shape with bond angles of o tetrahedral Tetrahedral Molecular Geometry Slide 146 / 186 There are three molecular geometries: Tetrahedral, if all are bonding pairs (440) Trigonal pyramidal, if one is a nonbonding pair (431) Bent, if there are two nonbonding pairs (422) Tetrahedral Molecular Geometry Slide 147 / tetrahedral (440) trigonal pyramidal (431) bent (422) Again, note the decrease in bond angle as the number of high repelling unbonded pairs of electrons increase.

50 Trigonal Bipyramidal Electron-Domain Geometry Slide 148 / 186 Five atoms around a central one will form a trigonal bipyramidal shape with bond angles of 120 o and 90 o trigonal bipyramidal Trigonal Bipyramidal Molecular Geometry Slide 149 / 186 Trigonal bipyramidal Seesaw T-shaped Linear Trigonal Bipyramidal Molecular Geometry Slide 150 / 186 There are four molecular geometries for the trigonal bipyramidal electron domain geometry: Trigonal Bipyramidal (550) See-Saw (541) T-Shape (532) Linear (523)

51 ctahedral Electron-Domain Geometry Slide 151 / 186 Six atoms around a central one will form an octahedral shape with bond angles of 90 o octahedral ctahedral Molecular Geometry Slide 152 / 186 ctahedral Square Pyramidal Square Planar ctahedral Molecular Geometry Slide 153 / 186 There are only three molecular geometries for the octahedral electron domain geometry: ctahedral (660) Square Pyramidal (651) Square Planar (642)

52 VSEPR and molecular geometry Slide 154 / 186 Using VSEPR numbers, you can determine molecular geometry. VSEPR numbers are a set of 3 numbers. 1) the total number of electron domains 2) the number of bonding domains* 3) the number of unshared pairs of electrons (*Remember that multiple bonds count as NE domain) Electron-domain geometry has the same name as the first shape. VSEPR Numbers and Molecular Geometries Slide 155 / 186 Draw the Lewis structure for ammonia, N3. What are the VSEPR numbers for N3? 4,3,1 What is the electron-domain geometry of N3? What is the molecular shape of N3?# tetrahedral triangular pyramidal What is the N- bond angle in the molecule? 107 What is the formal charge on the N atom? 5-5 = 0 slide for answer VSEPR Numbers and Molecular Geometries Slide 156 / 186 Draw the Lewis structure for ClF3. What are the VSEPR numbers for ClF3? 5,3,2 What is the electron-domain geometry of ClF3? What is the molecular shape of ClF3? trigonal bipyramidal T What would be the Cl-F bond angle(s)? 180, 90 What would be the formal charge on Cl? 7-7 = 0 slide for answer

53 44 The methane molecule (C4) has which geometry? Slide 157 / 186 A linear B trigonal bipyramidal C trigonal planar D tetrahedral [*] Slide 158 / Give the VSEPR number for this molecule. 46 Give the VSEPR number for this molecule. Slide 159 / 186

54 47 Give the VSEPR number for this molecule. Slide 160 / 186 F Xe F [*] 48 Which compound below contains an atom that is surrounded by more than an octet of electrons? Slide 161 / 186 A PF 5 B C 4 C NBr 3 D F 2 49 Which of the following molecules would have a bent shape? Slide 162 / 186 A S 2 B S 3 C C 4 D C 2 2 E F

55 50 Which of the following molecules would have a degree bond angle between atoms? Slide 163 / 186 A 2S B CF 3Cl C C 2 D PCl 3 E N The molecular shape and geometry of the nitrate ion (N 3- ) would be: Slide 164 / 186 A bent B linear C trigonal planar D trigonal bipyramidal E tetrahedral [*] YBRIDIZATIN TERY Slide 165 / 186 According to carbon's orbital diagram, it should only be able to form two bonds... 1s 2s 2p But we know carbon forms 4 bonds, not 2!!!

56 [*] YBRIDIZATIN TERY Slide 166 / 186 Scientists propose that the outermost s and p orbitals are actually combined to create 4 "hybrid" orbitals of equal energy. Carbon 1s sp 3 hybrid orbitals This explained how carbon could form 4 bonds [*] YBRIDIZATIN TERY Slide 167 / 186 To predict the hybridization involved in a compound, simply look at the first VSEPR numbers, this tells you how many electron domains(orbitals) need to be hybridized. For example: = 4 electron domains sp 3 Carbon requires 4 hybrid orbital so it hybridizes it's outermost "s" orbital and all three of the "p" orbitals to give 4 sp 3 hybrids. [*] YBRIDIZATIN TERY Slide 168 / 186 Example: Find the hybridization of the N atom in N 3? VSEPR # = 4 so the hybridization is sp 3

57 [*] YBRIDIZATIN TERY Slide 169 / 186 Example: What is the hybridization of C in C 2? VSEPR # = 2 nly the s orbital and 1 p orbital are needed to be hybridized so the hydridization is sp Note: The other 2 p orbitals not involved in hybridization are used to form the double bonds (called Pi bonds) [*] 52 Which of the following would require sp 2 hybridization? Slide 170 / 186 A BF 3 B 2 C PCl 3 D F 2 E N 2 [*] 53 What would be the hybridization found on in F 2? Slide 171 / 186 A sp B sp 2 C sp 3 D s 2 p 3 E s 3 p 3

58 Slide 172 / 186 Polarity Return to Table of Contents Polarity of Bonds Slide 173 / 186 Though atoms often form compounds by sharing electrons, the electrons are not always shared equally. In a covalent bond, one atom has a greater ability to pull the shared pair toward it. Polarity of Bonds Slide 174 / 186 Identical atoms will have an electronegativity difference of ZER. As a result, the bond is NNPLAR.

59 Bonds and Electronegativity Slide 175 / 186 Bond Type Electronegativity Difference Non-Polar Covalent very small or zero Polar Covalent about 0.2 to 1.6 Ionic above 1.7 (between metal & non-metal) Polarity of Bonds Slide 176 / 186 Therefore, the fluorine end of the molecule has more electron density than the hydrogen end. F We use the symbol to designate a dipole (2 poles). The "+" end is on the more positive end of the molecule and the arrow points towards the more negative end. Slide 177 / 186

60 Polarity of Bonds Slide 178 / 186 Bond lengths, Electronegativity, Differences and Dipole Moments of the ydrogen alides Compound Bond Electronegativity Dipole length (A0) Difference Moment (D) F Cl Br I [*] Polarity of Molecules But just because a molecule possesses polar bonds does not mean the molecule as a whole will be polar. For instance, in the case of C2: Slide 179 / 186 The polar bond is shown as a dipole, the arrow points to the more negative atom. Dipoles add as vectors. [*] Polarity of Molecules Slide 180 / 186 By adding the individual bond dipoles, one can determine the overall dipole moment for the molecule. For a molecule to be polar, it must a) contain one or more dipoles AND b) have these polar bonds arranged asymmetrically In other words, if all the dipoles are symmetrical, they will cancel each other out and the molecule will be NNPLAR. Many molecules with lone pairs of electrons will be PLAR.

61 [*] Polarity of Molecules These are some examples of polar & nonpolar molecules. What are their VSEPR numbers? Slide 181 / (?), polar Slide for Answer Slide 431, polar for Answer Slide 440, nonpolar for Answer Slide 330, nonpolar for Answer Slide 440, polar for Answer [*] 54 Which of these are polar molecules? Slide 182 / 186 A a, b B a, b, c C a, c D a, c, d E c, e 55 Sulfur trioxide (S 3) is polar. Slide 183 / 186 True False

62 56 ydrogen sulfide gas ( 2S) is non-polar. Slide 184 / 186 True False 57 Which of the following contains polar bonds but is a non-polar molecule? Slide 185 / 186 A C 4 B CS 2 C 2S D CF 4 E All of these are polar Slide 186 / 186