1. Chapter 5 Compounds and Their Bonds
Covalent Bonds
Naming and Writing Formulas of Covalent Compounds
Bond Polarity

2. Ionic Bond H2, A Covalent Molecule hydrogen molecule
In hydrogen, two hydrogen atoms share their electrons to form a covalent bond.
Each hydrogen atom acquires a stable outer shell of two (2) electrons like helium (He).
H H H : H = HH = H2
hydrogen molecule

3. A covalent bond between two hydrogen atoms is shown in this picture.
Fig 5.1 A covalent bond is the result of attractive and repulsive forces between atoms.

4. 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.

5. Bond length: The optimum distance between nuclei involved in a covalent bond. If the atoms are too far apart, the attractive forces are small and no bond exists. If the atoms are too close, the repulsive interaction between the nuclei is so strong that it pushes the atoms apart, Fig 5.2.

6. When two chlorine atoms approach each other, the unpaired 3p electrons are shared by both atoms in a covalent bond. Each chlorine atom in the Cl2 molecule now have 6 electrons in its own valence shell and sharing two giving each valence shell octet.

7. Diatomic Elements
As elements, the following share electrons to form diatomic, covalent molecules.

8. In addition to H2 and Cl2, five other elements always exist as diatomic molecule.

9. Learning Check
What is the name of each of the following diatomic molecules?
H2 hydrogen
N2 nitrogen
Cl2 _______________
O2 _______________
I2 _______________

10. Solution
What are the names of each of the following diatomic molecules?
H2 hydrogen
N2 nitrogen
Cl2 chlorine
O2 oxygen
I2 iodine

11. Covalent Bonds in NH3 H Bonding pairs
The compound NH3 consists of a N atom and three H atoms.
 
 N  and 3 H  
By sharing electrons to form NH3, the electron dot structure is written as
H Bonding pairs
H : N : H
  Lone pair of electrons

12. Number of Covalent Bonds
Often, the number of covalent bonds formed by a nonmetal is equal to the number of electrons needed to complete the octet.

13. Dot Structures and Models of Some Covalent Compounds

14. Multiple Bonds
Sharing one pair of electrons is a single bond. X : X or X–X
In multiple bonds, two pairs of electrons are shared to form a double bond or three pairs of electrons are shared in a triple bond. X : : X or X =X X : : : X or X ≡ X

15. Multiple Bonds in N2
In nitrogen, octets are achieved by sharing three pairs of electrons.
When three pairs of electrons are shared, the multiple bond is called a triple bond.
octets
       
 N   N   N:::N
 
triple bond

16. 5.4 Coordinate Covalent Bonds
Coordinate Covalent Bond: The covalent bond that forms when both electrons are donated by the same atom.

17. Fig 5.7 Electronegativities and the periodic table

18. 5.6 Drawing Lewis Structure
Draw skeletal structure of compound showing what atoms are bonded to each other. Put the unique element ( or least electronegative atom) in the center.
Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.
Complete an octet for all atoms except hydrogen
If structure contains too many electrons, form double and triple bonds on central atom as needed.
UNKNOWN_PARAMETER_VALUE.mp K   Download  

19. 5 + (3 x 7) = 26 valence electrons
Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F, put N in center
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms and complete
octets on N and F atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons F N

20. 4 + (3 x 6) + 2 = 24 valence electrons
Write the Lewis structure of the carbonate ion (CO32-).
Step 1 – C is less electronegative than O, put C in center
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)
-2 charge – 2e-
4 + (3 x 6) + 2 = 24 valence electrons
Step 3 – Draw single bonds between C and O atoms and complete
octet on C and O atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
Step 5 - Too many electrons, form double bond and re-check # of e-
2 single bonds (2x2) = 4
1 double bond = 4
8 lone pairs (8x2) = 16
Total = 24 O C

21. ( ) - - Two possible skeletal structures of formaldehyde (CH2O) H C O
An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.
formal charge on an atom in a Lewis structure =
total number of valence electrons in the free atom -
total number of nonbonding electrons - 1 2
total number of bonding electrons ( )
The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.

22. ( ) - -1 +1 H C O = 1 2 = 4 - 2 - ½ x 6 = -1 = 6 - 2 - ½ x 6 = +1
C : 4 e-
O : 6 e-
2H :2x1 e-
12 e-
2 single bonds (2x2) = 4
1 double bond = 4
2 lone pairs (2x2) = 4
Total = 12 H C O
formal charge on an atom in a Lewis structure = 1 2
total number of bonding electrons ( )
total number of valence electrons in the free atom -
total number of nonbonding electrons
formal charge on C
= 4 -
2 -
½ x 6 = -1
formal charge on O
= 6 -
2 -
½ x 6 = +1

23. ( ) - H C O = 1 2 = 4 - 0 - ½ x 8 = 0 = 6 - 4 - ½ x 4 = 0 C – 4 e-H C O
C – 4 e-
O – 6 e-
2H – 2x1 e-
12 e-
2 single bonds (2x2) = 4
1 double bond = 4
2 lone pairs (2x2) = 4
Total = 12
formal charge on an atom in a Lewis structure = 1 2
total number of bonding electrons ( )
total number of valence electrons in the free atom -
total number of nonbonding electrons
formal charge on C
= 4 -
0 -
½ x 8 = 0
formal charge on O
= 6 -
4 -
½ x 4 = 0

24. Formal Charge and Lewis Structures
For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present.
Lewis structures with large formal charges are less plausible than those with small formal charges.
Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms.
Which is the most likely Lewis structure for CH2O? H C O -1 +1 H C O

25. 5.7 Shape of Molecules VSEPR
The shape of a molecule is predicted from the geometry of the electron pairs around the central atom.
In the valence-shell electron-pair repulsion theory (VSEPR), the electron pairs are arranged as far apart as possible to give the least amount of repulsion of the negatively charged electrons.

26. Two Electron Pairs
In a molecule of BeCl2, there are two bonding pairs around the central atom Be. (Be is an exception to the octet rule.)
The arrangement of two electron pairs to minimize their repulsion is 180° or opposite each other.
The shape of the molecule is linear.

27. Two Electron Pairs with Double Bonds
The electron-dot structure for CO2 consists of two double bonds to the central atom C.
Because the electrons in a double bond are held together, a double bond is counted as a single unit.
Repulsion is minimized when the double bonds are placed opposite each other at 180° to give a linear shape.

28. Three Electron Pairs
In BF3, there are 3 electron pairs around the central atom B. (B is an exception to the octet rule.)
Repulsion is minimized by placing three electron pairs in a plane at angles of 120°, which is a trigonal planar arrangement.
The shape with three bonded atoms is trigonal planar.

29. Two Bonding Pairs and A Nonbonding Pair
In SO2, there are 3 electron units around the central atom S.
Two electron units are bonded to atoms and one electron pair is a nonbonding pair.
Repulsion is minimized by placing three electron pairs in a plane at angles of 120°, which is trigonal planar.
The shape with two bonded atoms is bent.

30. Four Electron Pairs
In CH4, there are 4 electron pairs around the central atom C.
Repulsion is minimized by placing four electron pairs at angles of 109°, which is a tetrahedral arrangement.
The shape with four bonded atoms is called tetrahedral.

31. Three Bonding Atoms and One Nonbonding Pair
In NH3, there are 4 electron pairs around the N.
Three pairs are bonded to atoms and one is a nonbonding pair.
Repulsion is minimized by placing four electron pairs at angles of 109°, which is a tetrahedral arrangement.
The shape with three bonded atoms is pyramidal.

32. Two Bonding Atoms and Two Lone Pairs
In H2O, there are 4 electron pairs around O.
Two pairs are bonded to atoms and two are nonbonding pairs.
Repulsion is minimized by placing four electron pairs at angles of 109° called a tetrahedral arrangement.
The shape with two bonded atoms is called bent.

33. Some Steps Using VSEPR to Predict Shape
Draw the electron dot structure.
Count the charged clouds around the central atom.
Arrange the charged clouds to minimize repulsion.
Determine the shape using the number of bonded atoms in the electron arrangement.

34. Summary of Electron Arrangements and Shapes
Number of atoms bonded
to the central atom

35. The shape depends on the number of charged clouds surrounding the atom as summarized in Table 5.1

36. Learning Check
Use VSEPR theory to determine the shape of the following molecules or ions.
1) tetrahedral ) pyramidal 3) bent
A. PF3
B. H2S
C. CCl4
D. PO43-

37. Solution
Use VSEPR theory to determine the shape of the following molecules or ions.
1) tetrahedral 2) pyramidal 3) bent
A. PF3 2) pyramidal
B. H2S 3) bent
C. CCl4 1) tetrahedral
D. PO43- 1) tetrahedral

38. Comparing Nonpolar and Polar Covalent Bonds

39. Electronegativity
Electronegativity is the attraction of an atom for shared electrons.
The nonmetals have high electronegativity values with fluorine as the highest.
The metals have low electronegativity values.

40. Fig 5.7 Electronegativities and the periodic table

42. Nonpolar Covalent Bonds
The atoms in a nonpolar covalent bond have electronegativity differences of 0.4 or less.
Examples: Atoms Electronegativity Type of Difference Bond N-N = Nonpolar covalent Cl-Br = Nonpolar covalent H-Si = Nonpolar covalent

43. Polar Covalent Bonds
The atoms in a polar covalent bond have electronegativity differences of 0.5 to 1.9.
Examples: Atoms Electronegativity Type of Difference Bond O-Cl = 0.5 Polar covalent Cl-C = 0.5 Polar covalent O-S = 1.0 Polar covalent

44. Comparing Nonpolar and Polar Covalent Bonds

45. Polar, Nonpolar and Ionic Bond

46. Ionic Bonds
The atoms in an ionic bond have electronegativity differences of 2.0 or more.
Examples: Atoms Electronegativity Type of Difference Bond Cl-K – 0.8 = 2.2 Ionic N-Na 3.0 – = 2.1 Ionic

47. Predicting Bond Type

48. Learning Check Identify the type of bond between the following as
1) nonpolar covalent
2) polar covalent
3) ionic
A. K-N
B. N-O
C. Cl-Cl

49. Solution A. K-N 3) ionic B. N-O 2) polar covalent C. Cl-Cl
1) nonpolar covalent

50. 5.9 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 and lone-pair contribution in the molecule.

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

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

53. Would a linear water molecule be Polar?
Why is water not linear? O .. H  

55. Learning Check Identify each of the following molecules as
1) polar or 2) nonpolar. Explain.
A. PBr3
B. HBr
C. Br2
D. SiBr4

56. Solution Identify each of the following molecules as
1) polar or 2) nonpolar. Explain.
A. PBr3 1) polar; pyramidal
B. HBr 1) polar; polar bond
C. Br2 2) nonpolar, nonpolar bond
D. SiBr4 2) nonpolar; dipoles cancel

57. Naming Covalent Compounds
In the name of a covalent compound, the first nonmetal is named followed by the name of the second nonmetal ending in –ide.
Prefixes indicate the number of atoms of each element.

58. Learning Check Complete the name of each covalent compound:
CO carbon ______oxide
CO2 carbon _______________
PCl3 phosphorus ___________
CCl4 carbon _______________
N2O ______________________

59. Solution Complete the name of each covalent compound:
CO carbon monoxide
CO2 carbon dioxide
PCl3 phosphorus trichloride
CCl4 carbon tetrachloride
N2O dinitrogen monoxide

60. Formulas and Names of Some Covalent Compounds

61. Learning Check Select the correct name for each compound.
A. SiCl4 1) silicon chloride
2) tetrasilicon chloride
3) silicon tetrachloride
B. P2O5 1) phosphorus oxide
2) phosphorus pentoxide
3) diphosphorus pentoxide
C. Cl2O7 1) dichlorine heptoxide
2) dichlorine oxide
3) chlorine heptoxide

62. Solution Select the correct name for each compound.
A. SiCl4 3) silicon tetrachloride
B. P2O5 3) diphosphorus pentoxide
C. Cl2O7 1) dichlorine heptoxide

63. Chapter Summary
Covalent bond: Bond formed by sharing of electrons between the atoms.
Molecule: A group of atoms held together by covalent bonds.
Coordinate covalent bond: Bond formed when a filled orbital containing lone pair of electrons on one atom overlaps a vacant orbital on another atom.
Molecular formula: Formula that shows the numbers and kinds of atoms in a molecule.
Lewis structure: shows how atoms are connected in a molecule.

64. Chapter Summary Contd.
Molecules have specific shapes that depend on the number of electron charge clouds surrounding the various atoms (VSEPR model)
Bonds between atoms are polar covalent if the bonding electrons are not shared equally between the atoms.
The ability of an atom to attract electrons in a covalent bond is the atom’s electronegativity.
Molecular compounds have lower melting points and boiling points than ionic compounds.