Covalent Bonding

Chapter 8 Covalent Bonding Gilbert N. Lewis 1875-1946. Originated idea of shared electrons in chemical bonds. First produced heavy water. Coined the term “photon.”

Cl gives electron to Most bonding discussed so far has been ionic bonding, i.e., bonding where electrons are transferred from one element to another forming ions. For example, in NaCl: Na Which then become: Cl- Na+ AND: NaCl is an ionic compound

Note octet + Cl Cl Na Na atom atom ion ion Transfer is of valence or outer electron or electrons, so that an octet of electrons is achieved. Na is: 1s22s22p63s1(note: 3s1 is a valence e-) 1s22s22p6 Note octet. The Na+ion is: Cl is: 1s22s22p63s23p5 (the 3s23p5 are valence e’s) The Cl-ion is: 1s22s22p63s23p6 (We normally “hide” the octet surrounding cations)

Na Mg O Ne Each Lewis dot represents a valence electron The number of valence electrons is the same as Group number. Na - Group 1- one valence electron Mg - Group 2 - two valence electrons O - Group 6 - six valence electrons Ne - Group 8 - eight valence electrons

Octet Rule Atoms tend to gain, lose or share electrons until each is surrounded by 8 valence electrons. This is an octet, or a configuration of s2p6 All inert gases (Group VIII) have s2p6 configuration. Which of following have inert gas configuration?: K K+ Mg Mg+ Mg2+ O O2- N3- Cl-

Octet Rule Note that the most stable form of atoms or ions from previous slide have an inert gas configuration K+ 1s22s22p63s23p6 = [Ar] 1s22s22p6 Mg2+ = [Ne] =[Ne] 1s22s22p6 O2- N3- 1s22s22p6 =[Ne] Mg2+ and O2- and N3- are isoelectronic

Octet Rule Element electron configurations can be represented as inert gas core plus valence e’s Na = [Ne]3s1 Al = [Ne]3s23p1 Cl = [Ne]3s23p5 Ar = [Ne]3s23p6 (= [Ar]) K = [Ar]4s1

. . . . . . x x x x . . Cl x x . . x x x Cl Cl Cl x . . . . x x x x . . . x x x x Cl Cl . . . x x Covalent Bonding is the sharing of electrons to achieve octet It usually involves non-metallic elements and hydrogen For example, Cl2 is: + Can also use a line to represent a pair of electrons in a bond

H . x . . C H H x x . x H H H C H H Exceptions to octet rule: H, Li, Be, B H can bond with only one other atom. It can never be “sandwiched” is o.k. X-H is not X-H-X Methane, CH4 is: Or:

OH- SO42- CH3OH The Rules 1. Add up all the valence electrons 2. With charged species (ions): a) if negative, add electrons equal to ion charge b) if positive, remove electrons equal to ion charge 3. Distribute around atoms so that octet rule is obeyed (note exceptions for H, Li, etc..) Examples

Bonding and Nonbonding electron pairs . . H O H . . . . . . Consider H2O non bonding pair bonding pair bonding pair non bonding pair In water: there are two bonding pairs and.... two non-bonding pairs of electrons

Multiple Bonds Each oxygen has only 7 (not 8) electrons O O O O + or O O O O A doublebond (two pairs of electrons) is formed or N N N N N N + A triplebond (three pairs of electrons) is formed

Polar Bonds 0 0 +1 -1 0 0 H H Cl Cl Na Cl Covalent Bonds Ionic Bond (sharing of e’s) (transfer of e’s) + < +1 = partial positive charge + - 0 0 - > -1 = partial negative charge +1 -1 H Cl How do we know which atom is + and which atom is - ? Polar Covalent Bond Electrons are shared, but not equally

Electronegativity Electronegativity is a measure of the tendency of an atom in a molecule to attract bonding electrons to itself. Atoms with large (positive) Ionization Energies and very negative Electron Affinities are more electronegative. Linus Pauling developed an arbitrary scale of electronegativities () with values ranging from: F: =4.0 (most electronegative) to Cs: =0.7 (least electronegative)

Electronegativity

C-O N-C C-H Li-F Least Polar Most Polar (Ionic) Electronegativity and Bond Polarity (1) In a bond between two atoms, the atom with the higher electronegativity () is partially negative (-). (2) The larger the difference in electronegativities (), the more polar the bond. Which of the following bonds are the: (a) most polar, and (b) least polar. In each case, indicate the positive and negative ends of the bond. Atom  F 4.0 O 3.5 N 3.0 C 2.5 H 2.1 Li 1.0 + - - + - + + - =2.5-2.1 =0.4 =3.0-2.5 =0.5 =3.5-2.5 =1.0 =4.0-1.0 =3.0

+ - + - H-Cl H-F =3.0-2.1 =0.9 =4.0-2.1 =1.9 Alternative Notation H-Cl H-F Polar Bonds and Dipole Moments In physics, a positive and negative charge separated by a distance has a “dipole moment” () Since a polar diatomic molecule has a separated positive and negative charge, it has a dipole moment. Because HF has a more polar bond than HCl, it has a larger dipole moment.

Lewis Structures of Polyatomic Species (Dot structures, once again) 1. Count number of valence electrons. Remember to consider charge of polyatomic ions. 2. Arrange atoms and connect with single bonds. For species of the form XYn, the X atom is central. 3. Complete octets on outer atoms (other than H). 4. Put remaining electrons on central atom. Even if this creates greater than an octet. 5. If central atom has less than octet, create multiple bonds.

H Cl C Cl H Lewis Structures of Polyatomic Species CH2Cl2 1. # val e- = 1x4 + 2x1 + 2x7 = 20 e- 2. Arrange atoms and put in single bonds. Single bonds have used up 8 e-, leaving 20-8 = 12 e-. 3. Complete octets on Cl atoms. This uses remaining 12 e-. 4. No e- remaining to put on central atom. 5. Carbon has an octet of 8 e-. No need for multiple bonds.

F P F F Lewis Structures of Polyatomic Species PF3 1. # val e- = 1x5 + 3x7 = 26 e- 2. Arrange atoms and put in single bonds. Single bonds have used up 6 e-, leaving 26-6 = 20 e- remaining. 3. Complete octets on F atoms. This uses 18 e-, leaving 2 e- remaining. 4. Put remaining 2 e- on central atom. 5. Phosphorus has an octet of 8 e-. No need for multiple bonds.

N C S N C S N S C O O N Which of the following is correct structure for NCS- ? OR OR Don’t they all have correct number of valence e’s? Don’t they all obey the octet rule? Which is the preferred structure? Use the concept of FORMAL CHARGE. Formal Charge (FC) is charge atom would have if all atoms had the same electronegativity, i.e., if all bonding e’s were shared equally by all atoms. FC = # valence e’s - # assigned e’s O C O 6-5= +1 4-5= -1 6-6= 0 6-6= 0 5-5= 0

N C S N C S N S C Which of the following is correct structure for NCS- ? # val e- 5 4 6 5 4 6 5 4 6 # assigned e- 5 4 7 7 4 5 6 4 6 FC 0 0 -1 -2 0 1 -1 0 0 Formal Charge = # val e- - # assigned e- # val e- = # valence e- in isolated atom # assigned e- = # owned e- + (1/2) # shared e- (1)Lewis structure with smaller (closer to zero) formal charges is more realistic structure. Rules: 2) If any atoms must have negative formal charges, they should be the more electronegative atoms.

O C O O C O Formal Charge OR # val e- : 6 4 6 # val e- :6 4 6 # assigned e- :5 4 7 # assigned e- :6 4 6 Formal Charge:+1 0 -1 Formal Charge:0 0 0 Formal Charge = # val e- - # assigned e- # val e- = # valence e- in isolated atom # assigned e- = # owned e- + (1/2) # shared e- Rules: (1) Lewis structure with smaller formal charges is more realistic structure. (2) If any atoms must have negative formal charges, they should be the more electronegative atoms.

O O O O O O O O O O O O 2. 14 e- left 3. 2 e- left 4. No e- left 5. Make double bond “Average” Structure Resonance Structures Ozone (O3) 1. 18 val e- OR Resonance Structures

O O O O O O O O O The Structure of Ozone “Average” Structure Resonance Structure BO: 1.5 1.5 BO: 2 1 BL: 1.28  1.28  BL: 1.21  1.48  This is the actual measured ozone Bond Length Actual Structure 117o BO = Bond Order BL = Bond Length

C C C C C C Resonance Structures Examples NO2- SO3 NO3- H H H H H H

Exceptions to the Octet Rule Odd Number of Electrons Examples: NO: # val e- = 5 + 6 = 11 e- NO2: # val e- = 5 + 2x6 = 17 e- ClO2: # val e- = 7 + 2x6 = 19 e- Molecules with odd numbers of electrons are comparatively rare, and are often unstable

F F B B F F F Be F F Be F F F FC: +1 -1 0 0 Exceptions to the Octet Rule Some Boron and Beryllium Compounds BeF2: # val e- = 2 + 2x7 = 16 e- X FC: 0 0 0 FC: +1 -2 +1 BF3: # val e- = 3 + 3x7 = 24 e- FC: 0 0 0 X One of 3 resonance structures

Cl Cl 3. All gone 2. 30 e- left P Cl Cl Cl Exceptions to the Octet Rule Expanded Valence PCl5 2 4 1. 40 val e- 10 6 8 10 e- around P More than an octet Where do the extra electrons go??

3d o o o 3p x x o o 3s Expanded Valence PCl5 Phosphorus Z=15 1s22s22p63s23p3 Where can we sit? x x x 5 e- from P: o 5 e- from Cl’s: x

Expanded Valence Nitrogen Z=7 1s22s22p3 NCl5 ?? There are no empty 2d orbitals to accept the extra electrons. Expanded Valence Compounds (1) Central atom is 3rd. period or below (2) Outer atoms are usually halogens or oxygen

Expanded ValenceAdditional Examples I3- AsF6- SF4 XeF4

-C- .. O .. .. .. C= O N .. C Lewis Structures of Organic Molecules Guidelines: X-H (1) Hydrogen atoms have 1 bond (2) Carbon atoms have 4 bonds • Nitrogen atoms have 3 bonds and 1 e- lone pair (4) Oxygen atoms have 2 bonds and 2 e- lone pairs These guidelines result from the preference of the above atoms to have a formal charge of zero in organic molecules

Bond Enthalpy The Bond Enthalpy (aka Bond Dissociation Enthalpy) is H to break a bond in a gas phase molecule. H-H(g) --> H(g) + H(g) H = D(H-H) = 436 kJ/mol Cl-Cl(g) --> Cl(g) + Cl(g) H = D(Cl-Cl) = 242 kJ/mol H-F(g) --> H(g) + F(g) H = D(H-F) = 567 kJ/mol H(g) + F(g) --> H-F(g) ??? H = = - D(H-F) = -567 kJ/mol When a bond is formed, H is the negative of the Bond Enthalpy.

Bond Enthalpies and Bond Lengths – criteria • The greater the bond order (triple vs. double vs. single bonds), the stronger and shorter the bonds. • The closer the two bonding atoms, the stronger the bond (compare H-H vs. C-C). • The greater the difference in electronegativity, the shorter the bond (compare O-H vs. C-H).

Bond Order (BO), Bond Enthalpy (BE), Bond Length (BL) Bond BO BE BL C-C C=C CC H-H H-C H-O 1 2 3 1 1 1 348 kJ 614 839 436 413 460 1.54  1.34 1.20 0.74 1.09 0.96 N-N N=N NN C-O C=O CO 1 2 3 1 2 3 163 418 941 359 806 1083 1.47 1.24 1.10 1.43 1.20 1.13

Covalent Bonding