1. Covalent Bonds and Lewis Structures
Covalent bond: A bonds formed by sharing electrons between atoms. Usually nonmetals
Molecule: A group of atoms held together by covalent bonds.
The nonmetals near the middle of the periodic table reach an electron octet by sharing an appropriate number of electrons.
General, Organic, and Biological Chemistry; Prentice 2003, Ch. 5

2. Spherical 1S orbital of two individual hydrogen atoms bends together and overlap to give an egg shaped region in the hydrogen molecule. The shared pair of electrons in a covalent bond is often represented as a line between atoms.
*Remember, hydrogen doesn’t get a full octet,
it only fills the first shell s orbital

3. A water molecule results when two hydrogen atoms and one oxygen atom are covalently bonded in a way shown in the following picture:
How many electrons are in the
valence shell of each atom in water?

4. Let’s go back to the electron configurations:
Hydrogen atom: 1 proton, 1 electron s1
Oxygen atom: 8 protons, 8 neutrons, 8 electrons. 1s2 2s2 2p4
Which electronic orbitals are involved in the bonds?
Hydrogen:
Oxygen: 1s 2p

5. Fig 5.4 For P, S, Cl, and other elements in the third period and below, the number of covalent bonds may vary, as indicated by the numbers shown in parentheses.

6. Numbers of Covalent Bonds
- single bond– sharing of two electrons
- longest bond length
- double bond- sharing of four electrons
- middle bond length
- triple bond- sharing of six electrons
- shortest bond length
- strongest covalent bond
- The electrons in the valence shell that are not shared are called lone pairs.

7. Drawing Lewis Structures
To draw Lewis structure, you need to know the connections among atoms. Also, common bonding patterns, shown below, for C, N, O, X, and H simplifies writing Lewis structure.
*O is not drawn correctly here. How should it be drawn?

8. Method #1- for drawing structures (
Method #1- for drawing structures (*note- you actually need to know how to determine Formal Charge to get the best representation, but that typically isn’t covered in Chem I course)
Determine the number of valence electrons in the molecule.
Determine the arrangement of atoms in the molecule (least electronegative –center)
Connect the atoms with single bonds
Give each peripheral atom an octet by adding lone pairs
Place all remaining electrons on central atom by adding lone pairs
If the central atom still doesn’t have an octet, take a lone pair from a peripheral atom to form a multiple bond

9. 4 for C and 6 for O (twice) = 16 electrons
LEWIS DOT STRUCTURES
1. Arrange the symbols such that the least electronegative element is in the center and the other elements are surrounding the central atom.
O C O
2. Give each of the elements their appropriate number of valence electrons (dots). Remember the number of valence electrons for a representative element is the same as the group number.
: O : C : O :
3. Keep track of the total numbr of valence electrons for the compound by adding the valence electrons from each atom. If the compound is an ion then add electrons (dots) for each negative charge or subtract electrons (dots) for each positive charge.
4 for C and 6 for O (twice) = 16 electrons

10. LEWIS DOT STRUCTURES
4. Now move the dots around so that you have 8 dots (the octet rule) around each element (do not forget the exceptions) while at the same time keeping the dots in pairs. Electrons, at this point, exist as pairs (the buddy system).
5. EXCEPTIONS TO THE OCTET RULE: Group I, II, and III need only 2, 4, and 6 electrons, respectively, around that atom (if they are covalently bonding).
6. If there are too few pairs to give each atom eight electrons, change the single bonds between two atoms to either double or triple bonds by moving the unbonded pairs of electrons next to a bonding pair.
: O : : C : : O :

11. LEWIS DOT STRUCTURES
7. Once the octet rule has been satisfied for each atom in the molecule then you may replace each pair of dots between two atoms with a dash.
: O = C = O :
8. Now check your structure by
a) count the total number of electrons to make sure you did not lose or gain electrons during the process.
b) Use FORMAL CHARGE (FC) calculations as a guideline to the correct structure. A zero formal charge is usually a good indication of a stable structure.
FC (X) = # of valence electrons - (1/2 bonding electrons + nonbonding electrons)
For our example: FC(C) = 4 - (1/ ) = 0
FC(O) = 6 - (1/ ) = 0

12. The Basics: Drawing Lewis Structures
Step 1: Calculate the total number of valence electrons in the molecule or ion
Step 2: Determine the central atom(s) of the molecule or ion – usually it’s the least electronegative atom.
Step 3: Draw a tentative diagram for the molecule or ion.
Rules
A hydrogen atom always forms one bond. Hydrogen is always a terminal atom in a Lewis diagram – an atom that is bonded to only one other atom.
A carbon atom normally forms four bonds
When several carbon atoms appear in the same molecule, they are often bonded to each other.

13. Draw the Lewis Structure for the following molecules.
1. H2O
Oxygen has 6 valence electrons & Hydrogen has 1 valence electron for a total of 8 electrons.
. .
H : O : H or
2. CO
Oxygen has 6 valence electrons & Carbon has 4 valence electrons for a total of 10 electrons.
: C : : : O :

14. Draw the Lewis Structure for the following molecules.
3. BH3
Boron has 3 valence electrons & Hydrogen has 1 valence electron for a total of 6 electrons
H : B : H
. . H
4. NH3
Nitrogen has 5 valence electrons & Hydrogen has 1 valence electron for a total of 8 electrons
H : N : H

15. p. 362

16. CW- Draw CCl4

17. p. 374

18. CW- Draw PBr3

19. p. 362

20. Draw OCl2 **Exception here. Put the most eneg one in center

21. p. 375

22. Draw NH4+ *Cation **Coordinate covalent bond- one atom donates both shared e-

23. Cations
p. 363

24. Do:
OH-

25. Solution:

26. Classwork:
Do “WS Covalent Lewis Dot Structures”. You can check your answers on youtube or on kentchemistry.com.
Interactive Practice in class today: Must do a MINIMUM of 5 structures to get credit. I will be watching!

27. The next few slides have….
Some more “challenging” Lewis structures. I will let you know if we are doing them or not.

28. Draw SO3 2- *anion; name it!

29. Anions
p. 366

30. Draw SO2 *Resonant structure. Draw both

31. Resonance
p. 366

32. Practice:
Predict the most stable structure: ONC- or OCN- or NOC-

33. LEWIS DOT STRUCTURES
Predict the most stable structure: ONC- or OCN- or NOC-
:O::N::C: or :O::C::N: or :N::O::C:
1) Total electrons is:
6 e- for O + 5 e- for N + 4 e- for C + 1 e- for negative charge = 16 e- total. All structures fulfill the octet rule.
2) FC (X) = # of valence electrons - (1/2 bonding electrons + nonbonding electrons)
structure#1: structure #3:
FC(C) = 4 - (1/ ) = -2 FC(C) = 4 - (1/ ) = -2
FC(O) = 6 - (1/ ) = 0 FC(O) = 6 - (1/ ) = +2
FC(N) = 5 -(1/ ) = +1 FC(N) = 5 -(1/ ) = -1
structure#2:
FC(C) = 4 - (1/ ) = 0
FC(O) = 6 - (1/ ) = 0
FC(N) = 5 -(1/ ) = -1
structure #2 has the combination with the lowest formal charge. It also has the negative formal charge on one of the more electronegative atoms. Calculate the formal charge for the most stable structure:
:O:C:::N:
. .
(-1, 0, 0)

34. Practice problems
Draw: ClO2-
SiH4
AsH3

35. Practice Problems ClO2- SiH4 .. .. .. : O : Cl : O : .. .. .. AsH3 ..
: O : Cl : O : H
H : Si : H ..
H : As : H H

36. Group Study Problems 1. Draw the Lewis structure for the following
a) H2S b) PH3 c) CH2O
d) NO2- e) H2CO3 f) CBr4
g) CH2FCl h) C2H2 I) O3
2. Calculate the formal charge for “c”, “d”, “f”, and “I”.

37. So, how do you draw these structures?
* Practice any we didn’t do in class at home
Br2
CH4
HCl
SiF4
PF3
H2CO
C2H6
C2H4
CO2 CO
H2O
OH-
NH3
HCN
NH4 + (What is a coordinate covalent bond?
NO3- (*use the ‘rules” for this one. Also,
this one displays resonance. What is that?)
CO3 2-
SO2

38. Shape of Molecules
Molecular shapes can be predicted by noting how many bonds and electron pairs surround individual atoms and applying what is called the valence-shell electron-pair repulsion (VSEPR) model. The basic idea of VSEPR model is that the negatively charged clouds of electron in bonds and lone pair repel each other therefore tends to keep apart as far as possible causing molecules to assume specific shape.

39. There are three step to applying the VSEPR model:
Step 1: Draw a Lewis structure of the molecule, and identify the atom whose geometry is of interest.
Step 2: Count the number of electron charge clouds surrounding the atom of interest.
Step 3: Predict molecular shape by assuming that the charge clouds orient in space so that they are as far away from one another as possible.

40. VSEPR Model
The shape depends on the number of charged clouds surrounding the atom
© by Prentice-Hall, Inc. A Pearson Company

41. Bond angles to know: Linear-180
Trigonal planar-120 (If all e- are in bonded pairs like BF3. If one pair is lone….and thus “bent”, like SO2 , then it’s actually 119)
Tetrahedral-109.5
Pyramidal-107 (lone pair pushes bonded pairs closer together)
Bent (water) (2 lone pairs)
Practice:

42. Polar Covalent Bonds and Electronegativity
Electrons in a covalent bond occupy the region between the bonded atoms.
If the atoms are identical, as in H2 and Cl2, electrons are attracted equally to both atoms and are shared equally.
If the atoms are not identical, however, as in HCl, the bonding electrons may be attracted more strongly by one atom than by the other and may thus are shared unequally. Such bonds are known as polar covalent bonds.

44. In HCl, electron spend more time near the chlorine than the hydrogen
In HCl, electron spend more time near the chlorine than the hydrogen. Although the molecule is overall neutral, the chlorine is more negative than the hydrogen, resulting in partial charges on the atoms.
Partial charges are represented by placing d- on the more negative atom and d+ on the more positive atom.
Ability of an atom to attract electrons is called the atom’s electronegativity.
Fluorine, the most electronegative element, assigned a value of 4, and less electronegative atoms assigned lower values
Electroneg differences > considered ionic; < covalent

45. Fig 5.7 Electronegativities and the periodic table

46. Polar Molecules
Entire molecule can be polar if electrons are attracted more strongly to one part of the molecule than to another.
Molecule’s polarity is due to the sum of all individual bond polarities (sometimes called dipolar forces or dipoles) and lone-pair contribution in the molecule.

47. Molecular polarity is represented by an arrow pointing at the negative end and is crossed at the positive end to resemble a positive sign.

48. Molecular polarity depends on the shape of the molecule as well as the presence of polar covalent bonds and lone-pairs.

49. Naming Binary Molecular Compounds
When two different elements combines together they form binary compound.
The formulas of binary compounds are usually written with the less electronegative element first. Thus, metals are always written before non-metals. Prefix such as mono, di, tri, tetra etc, are used to indicate number of atoms of each element.

50. A few examples of binary compounds are
given below:

51. The following two steps guide is helpful in naming binary compounds:
Step 1: Name the first element in the formula, using a prefix if needed to indicate the number of atoms.
Step 2: Name the second element in the formula, using an –ide ending as for anions, along with a prefix if needed to indicate the number of atoms.

52. Characteristic of Molecular Compounds
Molecules are neutral as a result there is no strong electrical attractions between the molecules to hold them together. However, there are several weaker forces exist between molecules, known as intermolecular forces (van der Waals forces).
When intermolecular forces are very weak, molecules are weakly attracted to one another and that the substance is gas at ordinary temperature.

53. If the intermolecular forces are somewhat
If the intermolecular forces are somewhat stronger, the molecules are pulled together into a liquid.
If the forces are stronger, the substance becomes
a molecular solid.
Hydrogen bond-specific type of intermolecular force involving hydrogen (that is cov. bonded to a highly electronegative atom) attracted to an unshared electron pair of another molecule. Ex-water
Melting points and boiling points of molecular solids are lower than those of ionic solids.
Most molecular compounds are insoluble in water.
Molecular compounds do not conduct electricity when melted because they have no charged particles.

54. Remember
Electronegativity increases across periods and decreases down groups
Ionization Energy increases across periods and decreases down groups
Atomic radius decreases across periods and increases down groups

55. Linear
Bent
Trigonal planar
Tetrahedral