1. 1 Covalent bonding Molecules & Structures

2. 2 What do you already know about Covalent bonding?

3. 3  What kind of bonding occurs when a metal and a nonmetal transfer electrons? –Ionic bonding  What is made when two metals just mix and don’t react? –An alloy  What do two nonmetals or some metalloids with nonmetals form when they bond together? –Covalent bond Bonding Review

4. 4 Covalent bonding makes molecules  Molecules are –Specific atoms, usually nonmetals, joined by sharing electrons  Two major kinds of molecules:  Molecular compound –Sharing e - between different elements –Example: CH 4  Diatomic molecules –Sharing e - between two of the same atom –These atoms occur naturally as compounds b/c they are more stable that way –Examples:

5. 5 Diatomic elements  There are 7 elements that always form molecules  H 2, N 2, O 2, F 2, Cl 2, Br 2 and I 2  1 + 6 pattern on the periodic table

6. 6 1 and 6

7. 7 Properties of Molecular Compounds  Tend to have low melting and boiling points  Have a molecular formula which shows type and number of atoms in a molecule –Not necessarily the lowest ratio of elements  Ex: C 6 H 12 O 6 or H 2 O  The molecular formula doesn’t tell you how bonded atoms are arranged

8. 8 How does H 2 form?  The nuclei repel ++

9. 9 How does H 2 form? ++  The nuclei repel  But they are attracted to electrons  They share the electrons  So the bond forms when the attractive forces balance the repulsive forces

10. 10 How do we show bonding?  Lewis structures –Use electron-dot diagrams to show how electrons are arranged in molecules –Ex: H 2  These are called structural formulas –They show what atoms are bonded together and the type(s) of bond(s).

11. 11 Covalent bonds  Nonmetals hold onto their valence electrons.  Need noble gas configuration to be stable –Usually with 8 valence electrons  Get it by sharing valence electrons with each other.

12. 12 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons l Both end with full orbitals FF

13. 13 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons l Both end with full orbitals FF 8 Valence electrons

14. 14 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons l Both end with full orbitals FF 8 Valence electrons

15. 15 Single Covalent Bonds:  A sharing of two valence electrons.  Ex: H 2 Double Covalent Bonds:  A sharing of four valence electrons  Ex: O 2 Triple Covalent Bonds:  A sharing of six valence electrons  Ex: N 2 As the number of shared electrons pairs increase, bond length decreases The shorter the bond length, the stronger the bond

16. 16 Drawing Structural Formulas  We use Lewis Structures to show this:Lewis Structures 1. Predict the location of atoms -Terminal atoms will be an end atom on the structure because they can only form one bond -Ex: H & F ALWAYS -The central atom is usually the one that is less electronegative 2. Make sure you have the correct number of valence e- for all of your atoms –For polyatomic ions, add one e - for each negative charge & subtract one e - for each positive charge 3. Start bonding by creating single bonds between the central atom(s) and each of the terminal atoms

17. 17 4. Make sure to share electrons between atoms as needed to get an octet on EACH atom. –SOME EXCEPTIONS: H wants only 2 e -, Be is OK with only 4 e-, and B is OK with only 6 e - –Remember C, N, O, & S can form double or triple bonds with the same element or another element if cannot get an octet with only single bonds 5. Redraw with lines for each shared pair of electrons (covalent bonds). Remember to enclose a polyatomic ion in brackets and indicate the overall charge on the ion.

18. 18 Some things to think about:  Too few e- for octets? Consider a double or triple bond.  Too many electrons? Can an atom have an expanded octet? Any atom in rows 3-7 of the periodic table can have an expanded octet. –Why? B/c of empty d-orbitals –The atom with the expanded octet is usually the central atom.

19. 19

20. 20 Examples:  PH 3 H2SH2S  CCl 4  SiO 2  NO 3 -

21. 21 Practice  Draw Lewis structures for the following:  PCl 3  CH 2 O  C 3 H 6  SO 4 2-

22. 22 Practice on Molecules with More than One Central Atom  C 2 H 2  CH 3 COOH  H 2 O 2

23. 23 Dealing with Exceptions to the Octet Rule

24. 24 Resonance Structures  When more than one correct structure can be written for a molecule or ion. –Usually happens with molecules or polyatomic ions that have both double and single bonds.  They only differ in the position of the electron pairs NOT the atom positions.  EX: NO 2 -  Which one is the true structure?  Does it go back and forth?  Double bonds are shorter than single  In NO 2 - all the bonds are the same length  The actual molecule behaves as if it only has one structure.

25. 25 Let’s Draw Some Resonance Structures:  SO 2  SO 3 2-  O 3

26. 26 Coordinate Covalent Bond  When one atom donates both electrons in a covalent bond.  Ex: Carbon monoxide OC

27. 27 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Ex: Carbon monoxide OC

28. 28 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Ex: Carbon monoxide OC OC

29. 29 Another Example:  The ammonium ion (NH 4 + ), which is formed from the combination of ammonia and a H + ion.

30. 30 Molecules with odd numbers of electrons  Molecules when valence e - are counted cannot form octets around each atom  Examples: NO 2, ClO 2 and NO

31. 31 Don’t forget Expanded Octets  The central atoms contain more than 8 valence electrons  Ex: XeF 4  Ex: SF 6

32. 32 Bond Dissociation Energy  The energy required to break a bond  C - H + 393 kJ  C + H  You ALWAYS have to add energy to break bonds (it will always be a positive number)  Double bonds have larger bond dissociation energies than single  Triple bonds are even larger –Examples of bond dissociation energy between different types of carbon-carbon bonds –C-C 347 kJ –C=C 657 kJ –C≡C 908 kJ

33. 33 Bond Dissociation Energy  The larger the bond energy, the harder it is to break the bond  Large bond energies make chemicals less reactive  In chemical reactions, bonds in the reactants are broken and new bonds are formed to make products.  The total energy change of a chemical reaction is determined from the energy of the bonds broken and formed. –This is where Endothermic and Exothermic reactions come from

34. 34 Let’s Do An Example:  First balance the equation for the combustion of methane: CH 4 + O 2 → CO 2 + H 2 O  Draw the Lewis structures of the reactants and products

35. 35  Refer to the structures and add up the energy released from forming all the bonds in the products. (For example: if 2 moles of water are formed, that’s 4 O-H bonds)  Add up the bond energies for the moles of the reactants.  Subtract the energy used to break all the bonds in the reactants from the energy used to make the bonds in the products. This is how much energy released in the combustion reaction.  Is the reaction endothermic or exothermic?

36. 36 Bond Polarity

37. 37 Electronegativity and Polarity  Covalent bonds are the sharing of electrons between atoms.  The amount of sharing can change depending on how strongly an atom holds onto its electrons.  We use the periodic table and values of electronegativity to determine how strongly an atom will pull electrons in a bond.  The electronegativity of the atoms was assigned by Linus Pauling when he studied the bonding abilities of atoms in molecules.

38. 38  When the atoms in a bond are the same, the electrons are shared equally.  Also when there is little difference in electronegativity, the electrons are essentially shared equally  These are considered nonpolar covalent bonds.  When two different atoms are connected, the electrons may not be shared equally.  This is a polar covalent bond.  How do we measure how strong the atoms pull on electrons?

39. 39 Electronegativity  A measure of how strongly the atoms attract electrons in a bond.  The bigger the electronegativity difference the more polar the bond.  Use Figure 9-15 Pg. 263 in text to get electronegativities of atoms.  We use general guidelines to determine if a bond is polar, nonpolar or ionic. –Chemical bonds between different elements are never completely ionic or covalent

40. 40  Use the following differences in electronegativities of bonded atoms as general guidelines for bond polarity:  0.0 - 0.4 nonpolar covalent bond  0.5 - 1.7 polar covalent bond  >1.7 Ionic

41. 41 How to show when a bond is polar  Isn’t a whole charge just a partial charge   means a partially positive   means a partially negative  The Cl pulls harder on the electrons  The electrons spend more time near the Cl H Cl  

42. 42 For each pair of elements, calculate the electronegativity difference and label the bond type (polar covalent, nonpolar covalent, or ionic).  H, Cl  H, S  S, Cl  Na, F  Cl, Br  Al, Br

43. 43 Examples with compounds  Let’s determine if the bonds in the following compounds are polar, nonpolar, or ionic. You will need to show your calculations!  HCl  CH 4  HSF

44. 44 Molecular Shapes Some Theories

45. 45 VSEPR Theory  Valence Shell Electron Pair Repulsion.  Predicts three dimensional geometry of molecules based on the number of pairs of valence electrons both bonded and unbonded.  Name tells you the theory.  Valence shell - outside electrons.  Electron Pair repulsion - electron pairs try to get as far away as possible.  Can determine both the angles of bonds  AND the shape of molecules

46. 46 VSEPR  The bond angle is the angle formed by two terminal atoms  Both shared pair e - and lone pair e - repel each other –Lone pair e - repel more than shared e -

47. 47 # of e- domains Shared e- pairs Lone pair e- Molecular shape Bond Angle Example 220 Linear 180 o BeH 2 330 Trigonal planar 120 o BH 3 440 Tetrahedral 109.5 o CH 4 431 Trigonal pyramidal 107.3 o NH 3 3 or 421 or 2 Bent 104.5 o H2OH2O 550 Trigonal bipyramidal 90 o / 120 o PF 5 660 Octahedral 90 o SF 6 http://intro.chem.okstate.edu/1314F00/Lecture/Chapter10/VSEPR.html

48. 48 Examples of how we get the molecular shapes  Single bonds fill all atoms.  There are 4 pairs of electrons pushing away.  The furthest they can get away is 109.5º. CHH H H

49. 49 4 atoms bonded  Basic shape is tetrahedral.  A pyramid with a triangular base.  Same basic shape for everything with 4 pairs. C HH H H 109.5º

50. 50 3 bonded - 1 lone pair NHH H N HH H <109.5º  Still basic tetrahedral but you can’t see the electron pair.  Shape is called trigonal pyramidal.

51. 51 2 bonded - 2 lone pair OH H O H H <109.5º  Still basic tetrahedral but you can’t see the 2 lone pair.  Shape is called bent.

52. 52 3 atoms no lone pair C H H O  The farthest you can the electron pair apart is 120º.  Shape is flat and called trigonal planar.  Will require 1 double bond C H HO 120º

53. 53 2 atoms no lone pair  With three atoms the farthest they can get apart is 180º.  Shape called linear.  Will require 2 double bonds or one triple bond C O O 180º

54. 54 Try to determine the shapes of these molecules:  SiH 4  PF 3  HBr

55. 55 How do we get the shapes of these?  C 2 H 2  CH 3 COOH

56. 56 Molecular Orbitals (MO)  The overlap of atomic orbitals from separate atoms makes molecular orbitals  Each molecular orbital has room for two electrons  Two types of MO –Sigma ( σ ) between atoms –Pi ( π ) above and below atoms

57. 57 Sigma bonding orbitals  From s orbitals on separate atoms ++ s orbital +++ Sigma bonding molecular orbital

58. 58 Sigma bonding orbitals  From p orbitals on separate atoms p orbital Sigma bonding molecular orbital  

59. 59 Pi bonding orbitals  P orbitals on separate atoms      Pi bonding molecular orbital

60. 60 Sigma & Pi Bonds  Sigma bonds (  occur from overlap of orbitals between the atoms  Pi bond (  bond) occur between p orbitals. above and below atoms  All single bonds are  bonds  Double bond is 1   and 1  bond  Triple bond is 1   and 2  bonds

61. 61 Hybrid Orbitals Combines bonding with geometry

62. 62 Hybridization  The mixing of different atomic orbitals to form the same type of hybrid orbitals.  All the hybrid orbitals that form are identical.  Each hybrid orbital contains one electron that it can share with another atom.  The number of atomic orbitals mixed to form the hybrid orbital equals the total number of pairs of electrons (double and triple bonds get treated as though they are one pair of electrons)  Lone pairs on the central atom also occupy hybrid orbitals.

63. 63 Types of Hybridization  sp 3 -1 s and 3 p orbitals mix to form 4 sp 3 orbitals.  EX: CH 4, NH 3, H 2 O  sp 2 -1 s and 2 p orbitals mix to form 3 sp 2 orbitals leaving 1 regular p orbital.  EX: BH 3, AlCl 3, C 2 H 4  sp -1 s and 1 p orbitals mix to form 2 sp orbitals leaving 2 regular p orbitals.  Ex: BeCl 2, CO 2, O 2, N 2

64. 64 sp 3 hybridizaiton

65. 65

66. 66 sp 2 hybridization

67. 67

68. 68 Where is the p orbital?  Perpendicular  The overlap of orbitals makes a sigma bond (  bond)

69. 69 CC H H H H

70. 70 sp hybridization  when two things come off  one s and one p hybridize  linear

71. 71 sp hybridization  end up with two lobes 180º apart.  p orbitals are at right angles  makes room for two  bonds and two sigma bonds.  a triple bond or two double bonds

72. 72 CO 2  C can make two  and two   O can make one  and one  COO

73. 73 N2N2

74. 74 N2N2

75. 75 sp 3 d  1 s, 3 p, and 1 d orbitals mix together making 5 sp 3 d hybrid orbitals  Ex: PCl 5

76. 76 sp 3 d 2  1 s, 3 p, and 2 d orbitals mix together making 6 sp 3 d 2 hybrid orbitals  Ex: SF 6

77. 77 Molecular Polarity How to show if the entire molecule is polar or not.

78. 78 Molecular Polarity  Molecules are either nonpolar or polar, depending on the location and nature of the covalent bonds.

79. 79 Nonpolar Molecules  There is symmetry with regard to the distribution of electrons.  Determine the shape!  If there is an electronegative atom on one part of the molecule and one that “balances” it on another part, then the molecule is nonpolar. If not, it is a polar molecule  Ex: CH 4 and CCl 4 and CH 4 Cl

80. 80 Polar Molecules  Molecules with a partially positive end and a partially negative end  Symmetry can not cancel out the effects of the polar bonds. (There is no “balancing” of electronegative atoms on another part of the molecule) Must determine shape first. Examples: H 2 Oand NF 3

81. For each molecule, draw the Lewis structure, predict the shape and bond angle, and identify as polar or nonpolar. Br 2 HCN C 2 H 2 NH 4 + H 2 S PF 3 CH 2 O MgO

82. 82 ““ Symmetrical” shapes are those without lone pair on central atom –T–Tetrahedral –T–Trigonal planar –L–Linear TT he molecule will be nonpolar if all the atoms are the same or have low differences in electronegativities SS hapes with lone pair on central atom are not symmetrical CC an be polar even with the same atom bonded to the central atom

83. 83 Is it a polar or nonpolar molecule?  HF H2OH2O  NH 3  CBr 4  CO 2  CH 3 Cl

84. 84 Properties of Molecules  Most have LOW melting & boiling points  tend to be gases and liquids at room temperature  Ex: CO 2, NH 3, H 2 O  Polar and Nonpolar molecules have a little bit different properties due to the partial charges.

85. 85 H - F ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- + -

86. 86 Properties of Solid Molecules  Two kinds of crystals: –Molecular solids – molecules held together by attractive forces Ex: BI 3, Dry Ice, sugar –Network solids- atoms held together by bonds One big molecule (diamond, graphite) High melting & boiling points, brittle, extremely hard

87. 87 Graphite Diamond

88. 88 Intermolecular Forces What holds molecules to each other

89. 89 Intermolecular Forces  They are what make solid and liquid molecular compounds possible.  The weakest are called van der Waal’s forces - there are two kinds –Dispersion forces –Dipole Interactions

90. 90 Dispersion Force  Depends only on the number of electrons in the molecule  Bigger molecules more electrons  More electrons stronger forces F 2 is a gas Br 2 is a liquid I 2 is a solid

91. 91 Dispersion force HH HH HH HH ++ -- HH HH ++ -- ++ 

92. 92 Dipole interactions  Occur when polar molecules are attracted to each other.  Slightly stronger than dispersion forces.  Opposites attract but not completely hooked like in ionic solids.

93. 93 Dipole interactions  Occur when polar molecules are attracted to each other.  Slightly stronger than dispersion forces.  Opposites attract but not completely hooked like in ionic solids. HFHF  HFHF 

94. 94 + - + - + - + - + - + - + - + - + - + -

95. 95 Hydrogen bonding  Are the attractive force caused by hydrogen bonded to F, O, or N.  F, O, and N are very electronegative so it is a very strong dipole.  They are small, so molecules can get close together  The hydrogen partially share with the lone pair in the molecule next to it.  The strongest of the intermolecular forces.

96. 96 Hydrogen Bonding H H O ++ -- ++ H H O ++ -- ++

97. 97 Hydrogen bonding H H O H H O H H O H H O H H O H H O H H O

98. 98 Video lesson  Water, a polar molecule, on YouTube: https://www.youtube.com/watch?v=iOOvX 0jmhJ4 https://www.youtube.com/watch?v=iOOvX 0jmhJ4

99. 99 Review Ionic and Covalent Compounds  Practice Quiz and Graphics: http://www.elmhurst.edu/~chm/vchemboo k/145Areview.html http://www.elmhurst.edu/~chm/vchemboo k/145Areview.html

100. 100 Internet resources  Molecular polarity: http://www.elmhurst.edu/~chm/vchembook/210polarity.html http://www.elmhurst.edu/~chm/vchembook/210polarity.html  Polar covalent compounds: http://www.elmhurst.edu/~chm/vchembook/152Apolar.html http://www.elmhurst.edu/~chm/vchembook/152Apolar.html  Nonpolar covalent compounds: http://www.elmhurst.edu/~chm/vchembook/150Anpcovalent. html http://www.elmhurst.edu/~chm/vchembook/150Anpcovalent. html  Ionic compounds: http://www.elmhurst.edu/~chm/vchembook/143Aioniccpds.ht ml http://www.elmhurst.edu/~chm/vchembook/143Aioniccpds.ht ml  Compare Ionic, Polar, and Nonpolar Bonds: http://www.elmhurst.edu/~chm/vchembook/153Acompare.ht ml http://www.elmhurst.edu/~chm/vchembook/153Acompare.ht ml