1 Chapter 8 Covalent bonding

2 Covalent Bonding  A metal and a nonmetal transfer electrons –An ionic bond  Two metals just mix and don’t react –An alloy  What do two nonmetals do? –Neither one will give away an electron –So they share their valence electrons –This is a covalent bond

3 Covalent bonding  Makes molecules –Specific atoms joined by sharing electrons  Two kinds of molecules:  Molecular compound –Sharing by different elements  Diatomic molecules –Two of the same atom –O 2 N 2

4 Diatomic elements  There are 8 elements that always form molecules  H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2, and At 2  Oxygen by itself means O 2  The –ogens and the –ines  pattern on the periodic table

5 1 and 7

6 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  C 6 H 12 O 6  Formula doesn’t tell you about how atoms are arranged

7 Polar Bonds  When the atoms in a bond are the same, the electrons are shared equally.  This is a nonpolar covalent bond.  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?

8 Electronegativity  A measure of how strongly the atoms attract electrons in a bond.  The bigger the electronegativity difference the more polar the bond.  Use table 12-3 Pg. 285  Covalent nonpolar  Covalent moderately polar  Covalent polar  >2.0 Ionic

9 Chapter 9: Chemical Bonds9 Electronegativity Electronegativity (EN) is a measure of the ability of an atom to attract bonding electrons to itself EOS The greater the electronegativity of an atom in a molecule, the more strongly it attracts the electrons in a covalent bond

10 Chapter 9: Chemical Bonds10 Pauling’s Electronegativities EOS Electronegativity Illustrated

11 Chapter 9: Chemical Bonds11 Electronegativity Difference and Bond Type Two identical atoms have the same electronegativity and share a bonding electron pair equally. This is called a nonpolar covalent bond Example: chlorine gas EOS All homonuclear diatomic molecules have nonpolar covalent bonds: H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2

12 Chapter 9: Chemical Bonds12 Electronegativity Difference and Bond Type In covalent bonds between atoms with somewhat larger electronegativity differences, electron pairs are shared unequally. This is called a polar covalent bond Example: hydrogen chloride gas, HCl EOS The electrons are drawn closer to the atom of higher electronegativity, Cl

13 Covalent Bonding  Electrons are shared by atoms.  These are two extremes.  In between are polar covalent bonds.  The electrons are not shared evenly.  One end is slightly positive, the other negative.  Indicated using small delta 

14 How to show 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  

15 H - F ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ --

16 © 2009, Prentice-Hall, Inc. Polar Covalent Bonds  When two atoms share electrons unequally, a bond dipole results.  The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:  = Qr  It is measured in debyes (D).

17 Chapter 9: Chemical Bonds17 Electronegativity Difference and Bond Type With still larger differences in electronegativity, electrons may be completely transferred from metal to nonmetal atoms to form ionic bonds Example: sodium chloride, NaCl EOS

18 Chapter 9: Chemical Bonds18 Electronegativity Differences EOS

19 Chapter 9: Chemical Bonds19 Electron Distributions and Covalent Bonds Symmetric distribution EOS Asymmetric Distribution

20 Electronegativity difference Bond Type Zero Intermediate Large Covalent Polar Covalent Ionic Covalent Character decreases Ionic Character increases

21 How does H 2 form?  The nuclei repel ++

22 How does H 2 form? ++  The nuclei repel  But they are attracted to electrons  They share the electrons

23 Chapter 9: Chemical Bonds23 The Lewis Theory of Chemical Bonding: An Overview Electrons, particularly valence electrons, play a fundamental role in chemical bonding EOS In losing, gaining, or sharing electrons to form chemical bonds, atoms tend to acquire the electron configurations of noble gases

24 Chapter 9: Chemical Bonds24 Lewis Symbols Valence electrons are shown by dots around the element symbol EOS Use rules of electron configurations when forming dot structures … e.g., electrons remain unpaired if possible

25 Covalent bonds  Nonmetals hold onto their valence electrons.  They can’t give away electrons to bond.  Still need noble gas configuration.  Get it by sharing valence electrons with each other.  By sharing both atoms get to count the electrons toward noble gas configuration.

26 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

27 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

28 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

29 Lewis Structure  Shows how the valence electrons are arranged.  One dot for each valence electron.  A stable compound has all its atoms with a noble gas configuration.  Hydrogen follows the duet rule.  The rest follow the octet rule.  Bonding pair is the one between the symbols.

30 Chapter 9: Chemical Bonds30 Multiple Covalent Bonds The covalent bond in which one pair of electrons is shared is called a single bond e.g., H : Clor H—Cl Double bonds have two shared pairs of electrons EOS Triple bonds have three shared pairs of electrons

31 Rules  Sum the valence electrons.  Use a pair to form a bond between each pair of atoms.  Arrange the rest to fulfill the octet rule (except for H and the duet). H2OH2O  A line can be used instead of a pair.

32 A useful equation  (happy-have) / 2 = bonds  CO 2 C is central atom  POCl 3 P is central atom  SO 4 2- S is central atom  SO 3 2- S is central atom  PO 4 3- P is central atom  SCl 2 S is central atom

33 Exceptions to the octet  BH 3  Be and B often do not achieve octet  Have less than an octet, for electron deficient molecules.  SF 6  Third row and larger elements can exceed the octet  Use 3d orbitals?  I 3 -

34 Exceptions to the octet  When we must exceed the octet, extra electrons go on central atom.  (Happy – have)/2 won’t work  ClF 3  XeO 3  ICl 4 -  BeCl 2

35 Chapter 9: Chemical Bonds35 Writing Lewis Structures Hydrogen atoms are terminal atoms (bonded to only one other atom) EOS The central atom of a structure usually has the lowest electronegativity and the terminal atoms (except H) generally have higher electronegativities

36 Single Covalent Bond  A sharing of two valence electrons.  Only nonmetals and Hydrogen.  Different from an ionic bond because they actually form molecules.  Two specific atoms are joined.  In an ionic solid you can’t tell which atom the electrons moved from or to.

37 How to show how they formed  It’s like a jigsaw puzzle.  I have to tell you what the final formula is.  You put the pieces together to end up with the right formula.  For example- show how water is formed with covalent bonds.

38 Water H O Each hydrogen has 1 valence electron and wants 1 more The oxygen has 6 valence electrons and wants 2 more They share to make each other “happy”

39 Water  Put the pieces together  The first hydrogen is happy  The oxygen still wants one more H O

40 Water  The second hydrogen attaches  Every atom has full energy levels H O H

41 Multiple Bonds  Sometimes atoms share more than one pair of valence electrons.  A double bond is when atoms share two pair (4) of electrons.  A triple bond is when atoms share three pair (6) of electrons.

42 Carbon dioxide  CO 2 - Carbon is central atom ( I have to tell you)  Carbon has 4 valence electrons  Wants 4 more  Oxygen has 6 valence electrons  Wants 2 more O C

43 Carbon dioxide  Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short O C

44 Carbon dioxide l Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short O C O

45 Carbon dioxide l The only solution is to share more O C O

46 Carbon dioxide l The only solution is to share more O C O

47 Carbon dioxide l The only solution is to share more O CO

48 Carbon dioxide l The only solution is to share more O CO

49 Carbon dioxide l The only solution is to share more O CO

50 Carbon dioxide l The only solution is to share more O CO

51 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond O CO

52 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond O CO 8 valence electrons

53 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond O CO 8 valence electrons

54 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond O CO 8 valence electrons

55 How to draw them  To figure out if you need multiple bonds  Add up all the valence electrons.  Count up the total number of electrons to make all atoms happy.  Subtract.  Divide by 2  Tells you how many bonds - draw them.  Fill in the rest of the valence electrons to fill atoms up.

56 Examples  NH 3  N - has 5 valence electrons wants 8  H - has 1 valence electrons wants 2  NH 3 has 5+3(1) = 8  NH 3 wants 8+3(2) = 14  (14-8)/2= 3 bonds  4 atoms with 3 bonds N H

57 NHH H Examples  Draw in the bonds  All 8 electrons are accounted for  Everything is full

58 Examples  HCN C is central atom  N - has 5 valence electrons wants 8  C - has 4 valence electrons wants 8  H - has 1 valence electrons wants 2  HCN has = 10  HCN wants = 18  (18-10)/2= 4 bonds  3 atoms with 4 bonds -will require multiple bonds - not to H

59 HCN  Put in single bonds  Need 2 more bonds  Must go between C and N NHC

60 HCN l Put in single bonds l Need 2 more bonds l Must go between C and N l Uses 8 electrons - 2 more to add NHC

61 HCN l Put in single bonds l Need 2 more bonds l Must go between C and N l Uses 8 electrons - 2 more to add l Must go on N to fill octet NHC

62 Where do bonds go?  Think of how many electrons they are away from noble gas.  H should form 1 bond- always  O should form 2 bonds – if possible  N should form 3 bonds – if possible  C should form 4 bonds– if possible

63 Practice  Draw electron dot diagrams for the following.  PCl 3  H 2 O 2  CH 2 O  C 3 H 6

64 Another way of indicating bonds  Often use a line to indicate a bond  Called a structural formula  Each line is 2 valence electrons HHO = HHO

65 Structural Examples H CN C O H H  C has 8 electrons because each line is 2 electrons  Ditto for N  Ditto for C here  Ditto for O

66 Coordinate Covalent Bond  When one atom donates both electrons in a covalent bond.  Carbon monoxide  CO OC

67 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide l CO OC

68 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide l CO OC OC

69 How do we know if  Have to draw the diagram and see what happens.  Often happens with polyatomic ions  If an element has the wrong number of bonds

70 Polyatomic ions  Groups of atoms held by covalent bonds, with a charge  Can’t build directly, use (happy-have)/2  Have number will be different  Surround with [ ], and write charge  NH 4 2+  ClO 2 1-

71 Resonance  When more than one dot diagram with the same connections is possible.  Choice for double bond  NO 2 -  Which one is it?  Does it go back and forth?  Double bonds are shorter than single  In NO 2 - all the bonds are the same length

72 Formal Charge  For molecules and polyatomic ions that exceed the octet there are several different structures.  Use charges on atoms to help decide which.  Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work.

73 Formal Charge  The difference between the number of valence electrons on the free atom and that assigned in the molecule or ion.  We count half the electrons in each bond as “belonging” to the atom.  SO 4 -2  Molecules try to achieve as low a formal charge as possible.  Negative formal charges should be on electronegative elements.

74 Chapter 9: Chemical Bonds74 Formal Charge EOS Formal charge is the difference between the number of valence electrons in a free (uncombined) atom and the number of electrons assigned to that atom when bonded to other atoms in a Lewis structure

75 Chapter 9: Chemical Bonds75 Formal Charge Usually, the most plausible Lewis structure is one with no formal charges When formal charges are required, they should be as small as possible and negative formal charges should appear on the most electronegative atoms EOS Adjacent atoms in a structure should not carry formal charges of the same sign

76 Examples  XeO 3  NO 4 3-  SO 2 Cl 2

77 Resonance  It is a mixture of both, like a mule.  CO 3 2-

78 Bond Dissociation Energy  The energy required to break a bond  C - H kJ C + H  Double bonds have larger bond dissociation energies than single  Triple even larger –C-C 347 kJ –C=C 657 kJ –C≡C 908 kJ

79 Bond Dissociation Energy  The larger the bond energy, the harder it is to break  Large bond energies make chemicals less reactive.