Covalent Bonding and Molecular Compounds

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1 Smith Covalent Bonding and Molecular Compounds 1

2 Covalent Bonding & Molecular Geometry Examine these two forms of the same compound, ibuprofen. 2

3 Covalent Bonding & Molecular Geometry 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

4 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. 4

5 Chemical Bonds 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) 5

6 1 Chlorine monoxide is A ClO 2 B C D E ClO OCl O 2 Cl I don't know how to answer this. 6

7 Chemical Bonds Ionic Bonding Ionic bonds occur when the difference in electronegativity between two atoms is more than 1.7. 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. 7

8 Ionic v. Covalent Bonding 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 results in individual molecules; each with its own unique shape. These shapes help determine the physical and chemical properties of everything around us! click here for an animation about ionic and covalent bonding 8

9 2 Which pair of atoms will form a covalent bond? A B C D Li and Ne K and Br C and O Na and Cl E I don't know how to answer this. 9

10 3 Which pair of atoms will form a covalent bond? A B C D E Li and I Na and Cl K and Fl H and O I don't know how to answer this. 10

11 Molecular Compounds Covalent compounds are formed between two nonmetals. When atoms are bonded covalently, the atoms are held together by sharing electrons. Such a compound is called a molecular compound which is also known as a molecule. In covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases. Both atoms used the shared electrons to reach that goal. 11

12 Naming Binary Molecular Compounds Use prefixes to indicate the number the atoms All end in "ide" Examples NO 2 nitrogen dioxide P 2 O 5 diphosphorous pentoxide ( penta oxide >pentoxide) 12

13 Naming Binary Molecular Compounds 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 CO carbon monoxide CO 2 carbon dioxide 13

14 4 Chlorine monoxide is A ClO 2 B C D E ClO OCl O 2 Cl I don't know how to answer this. 14

15 5 Dinitrogen tetroxide is A NO 2 B N 2 O 4 C NO 3 D N 4 O 2 E I don't know how to answer this. 15

16 6 H 2 O is A Hydrogen monoxide B C D E Dihydrogen monoxide Hydrogen oxide Hydrogen dioxide I don't know how to answer this. 16

17 Per 1 2 covalent bonding.notebook 7 SO3 is A sulfate B sulfur oxide C sulfur trioxide D sulfite E I don't know how to answer this. 17

18 8 MgO is A B C D E monomagnesium monoxide magnesium monoxide monomagnesium oxide magnesium oxide I don't know how to answer this. 18

19 9 P 4 O 10 is A B C D E Phosphorous pentoxide Tetraphosphorous decoxide Phosphorous oxide Phosphate I don't know how to answer this. 19

20 Lewis Structures 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. 20

21 10 How many valence electrons does nitrogen have? A 2 B 3 C 4 D 5 E 7 F I don't know how to answer this. 21

22 11 The Lewis structure for nitrogen is N True False 22

23 The Octet Rule Recall that atoms tend towards having the electron configuration of a noble gas.for most atoms, that means having 8 valence electrons. The Octet Rule also applies to molecular compounds. 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 Octet Rule H needs 2e Be needs 2e B needs 2e 23

24 How do electron dot structures represent shared electrons? Two atoms held together by sharing a pair of electrons are joined by a single covalent bond. H + H H H Shared pair of electrons Hydrogen atom Hydrogen atom Hydrogen molecule 1s H H 1s Hydrogen molecule 24

25 How do electron dot structures represent shared electrons? An electron dot structure such as H:H represents the shared pair of electrons of the covalent bond by two dots. H + H H H Shared pair of electrons Hydrogen atom Hydrogen atom Hydrogen molecule 1s H H 1s Hydrogen molecule 25

26 Structural Formulas 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. H H Hydrogen molecule H Shared pair of electrons H 26

27 12 How many electrons are shared by two atoms to create a single covalent bond? A 2 B 1 27

28 Single Covalent Bonds The halogens form single covalent bonds in their diatomic molecules. Fluorine is one example. F + F > F F OR F F Fluorine atom Fluorine atom Fluorine molecule 1s 2s 2p Fluorine molecule 1s 2s 2p 28

29 Lewis Structure of H 2 O 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. 2 H + O > O H or O H Hydrogen atoms Oxygen atom H H Water molecule 1s 2s 2p O 1s H 1s H Water molecule 29

30 Lewis Structures of NH 3 In the ammonia molecule, NH 3, each atom attains a noblegas configuration by sharing electrons. This molecule has one unshared pair of electrons. H H 3 H + N > N H or N H Hydrogen atom Nitrogen atom H Ammonia molecule H 1s 2s 2p N Ammonia molecule 1s H 1s H 1s 30

31 Drawing Lewis Structures First, find the total number of valence electrons in the polyatomic ion or molecule. If it is an anion, add an electron for each negative charge. If it is a cation, subtract an electron for each positive charge. The P atom has 5 valence electrons. A Cl atom has 7, and there are three of them. The total number of valence electrons is: 31

32 Drawing Lewis Structures The central atom is the least electronegative element (excluding hydrogen). Connect the other atoms to it by single bonds. 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. The single bonds are shown as single lines. 32

33 Drawing Lewis Structures 1. Count each single bond as a pair (two) of electrons. 2. Add electons to the outer atoms to give each one 8 (a full shell), or just 2 electrons for hydrogen. 3. Do the same for the central atom. 4. Check: Does each atom have a full outer shell (8 except, 2 for hydrogen)? Have you used up all the valence electrons? Have you used too many electrons? 33

34 Drawing Lewis Structures NH 3 First, find the total number of valence electrons in the polyatomic ion or molecule. If it is an anion, add an electron for each negative charge. If it is a cation, subtract an electron for each positive charge. The N atom has 5 valence electrons and each of the three H atoms has 1 so the total number of valence electrons is, 5 + 3(1) = 8 34

35 Drawing Lewis Structures The central atom is the least electronegative element (excluding hydrogen because it can only have one bond). Connect the other atoms to it by single bonds. H can never be the central atom so N must be The H atoms will surround the N atom. The single bonds are shown as single lines. H NH 3 N H H 35

36 Drawing Lewis Structures 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 H). Next, do the same for the central atom. Check: Does each atom have a full outer shell? Have you used up all the valence electrons you started with? Have you used too many electrons? H N H H Each H already has two electrons, so that's done. But we have to add electrons to N to make 8. H N H H 36

37 F H N Cl C P Si O S B Se Xe I Slide for Answer H H C H H CH 4 Draw a Lewis Structure H H C H H Check to make sure that each atom has a full outer shell. 37

38 Draw a Lewis Structure F H N Cl C P Si O S B Se Xe I Slide for Answer F N F F NF 3 F N F Check to make sure that each atom has a full outer shell. F 38

39 Slide for Answer F H N Cl C Si P Si O S B Se Xe I SiF 4 Draw a Lewis Structure F F Si F Check to make sure that each atom has a full outer shell. F 39

40 Lewis Structures for ions If you are drawing the Lewis Structure for an ION... A negative ion has extra electrons, add the charge of the ion to your valence electron count. ClO 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. NH 4 + has 1(5) + 4(1) 1 = 8 electrons 40

41 F H N Cl C P Si O S B Se Xe I O O P O O PO 4 3 Draw a Lewis Structure Slide for OAnswer O P O O Check to make sure that each atom has a full outer shell. 41

42 Lewis Structures Draw the Lewis dot structure for the sulfate ion, SO

43 Lewis Structures Draw the Lewis dot structure for the hydronium ion, H 3 O + 43

44 F H N Cl C P Si O S B Se Xe I O C O CO 2 Draw a Lewis Structure Slide for Answer O C O We ran out of electrons, but carbon does not have an octet yet! Now What? 44

45 Double and Triple Covalent Bonds 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. 45

46 Double and Triple Covalent Bonds Carbon Dioxide, CO 2 1. Determine the # of valence electrons. 1 (4) + 2 (6) = 16 e 2. Form single bonds. O C O This leaves 12 electrons, 6 pairs 3. Place lone pairs on oxygen atoms to give each 8. O C O 46

47 Carbon Dioxide, CO 2 4. Check: We had 16 electrons to work with; how many have we used? O C O 5. There are too many electrons in our drawing. We must form DOUBLE BONDS between C and O. 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. O C O O C O 47

48 Covalent Bond Length 48

49 Covalent Bond Energy 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. 49

50 Covalent Bond Energies 50

51 Covalent Bonds Comparison Type of Bond Electrons shared Bond Strength Bond Length 2 weak long 4 intermediate intermediate 6 strong short 51

52 13 As the number of bonds between a pair of atoms increases, the distance between the atoms: A B C D E increases decreases remains unchanged varies, depending on the atoms I don't know how to answer this 52

53 14 As the number of bonds between a pair of atoms increases, the strength of the bond between the atoms: A B C D E increases decreases remains unchanged varies, depending on the atoms I don't know how to answer this 53

54 15 As the number of bonds between a pair of atoms increases, the energy of the bond between the atoms: A B C D E increases decreases remains unchanged varies, depending on the atoms I don't know how to answer this 54

55 16 How many electrons are shared by two atoms to create a single bond? 55

56 17 How many electrons are shared by two atoms to create a double bond? 56

57 18 How many electrons are shared by two atoms to create a triple bond? 57

58 Writing Lewis Structures If you run out of electrons before the central atom has an octet form multiple bonds until it does. 58

59 Bonding of O 2 Oxygen molecule O + O > O O or O O Oxygen atom Oxygen atom Oxygen molecule O 1s 2s 2p Oxygen molecule O 1s 2s 2p 59

60 F H N Cl C P Si O S B Se Xe I CO Draw a Lewis Structure Slide for Answer Carbon has the lower electronegativity, so we will consider it the "central" atom... C O 60

61 Coordinate Covalent Bonds 61

62 Coordinate Covalent Bonds In carbon monoxide, oxygen has a stable configuration but the carbon does not. C + O > C O Carbon atom Oxygen atom Carbon monoxide 1s 2s 2p C O 1s 2s 2p Carbon monoxide molecule 62

63 Coordinate Covalent Bonds 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. 63

64 F H N Cl C P Si O S B Se Xe I F 2 Draw a Lewis Structure Slide for Answer F F 64

65 Diatomic Molecules 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 NO LONGER HAPPENS! Remember: HONClBrIF Hydrogen Nitrogen Oxygen Fluorine Chlorine Bromine Iodine 65

66 F H N Cl C P Si O S B Se Xe I O 3 Draw a Lewis Structure Slide for Answer O O O 66

67 Exceptions to the Octet Rule 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) 67

68 Exception 1: Odd Number of Electrons Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons. NO is an example: 68

69 Exception 2: Fewer Than Eight Electrons 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 Memorize these exceptions 69

70 Exception 3: Expanded Octet The only way PCl 5 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. The d orbitals in these atoms participate in bonding, allowing the expanded octet. 70

71 Exception 3: Expanded Octet How many electrons do these central atoms have around them? 71

72 Exceptions to the Octet Rule Draw the Lewis dot structure for phosphorous pentachloride, PCl 5 : 72

73 Exceptions to the Octet Rule Draw the Lewis dot structure for the xenon tetrafluoride, XeF 4. 73

74 Exceptions to the Octet Rule Draw the Lewis dot structure for boron trifluoride, BF 3 : 74

75 Exceptions to the Octet Rule Draw the Lewis dot structure for the iodine tricholoride, ICl 3. 75

76 Polarity of Bonds 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. 76

77 Polarity of Bonds Identical atoms will have an electronegativity difference of ZERO. As a result, the bond is NONPOLAR. 77

78 Bonds and Electronegativity 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) 78

79 Polarity of Bonds Therefore, the fluorine end of the molecule has more electron density than the hydrogen end. H 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. 79

80 Bond Dipoles and Electronegativity When two atoms share electrons unequally, a bond dipole results. 80

81 Polarity of Bonds Bond lengths, Electronegativity, Differences and Dipole Moments of the Hydrogen Halides Compound Bond Electronegativity Dipole length (A0) Difference Moment (D) HF HCl HBr HI

82 [*] 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 CO 2 : The polar bond is shown as a dipole, the arrow points to the more negative atom. Dipoles add as vectors. 82

83 [*] Polarity of Molecules 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 NONPOLAR. Many molecules with lone pairs of electrons will be POLAR. 83

84 [*] Polarity of Molecules These are some examples of polar & nonpolar molecules. Slide 110(?), for Answer polar 431, polar Slide for Answer Slide 440, nonpolar for Answer Slide 330, for nonpolar Answer Slide 440, for polar Answer 84

85 [*] 19 Which of these are polar molecules? A B C D E a, b a, b, c a, c a, c, d c, e 85

86 86

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