1. Covalent bonding
Chapter 8

2. Molecular compounds
In nature, only the noble gas elements, such as helium and neon, exist as uncombined atoms.
They are monatomic; that is, they consist of single atoms.
Helium, which is less dense than air, is often used to inflate balloons.
But not all elements are monatomic.
O2 represents two oxygen atoms that are bonded together.

3. Sharing Electrons
The attractions that hold together the atoms in O2, H2O, CO2, and N2O cannot be explained by ionic bonding.
These bonds do not involve the transfer of electrons.
Recall that ionic bonds form when the combining atoms give up or accept electrons.
Another way that atoms can combine is by sharing electrons.
Atoms that are held together by sharing electrons are joined by a covalent bond.

4. Diatomic
An oxygen molecule is an example of a diatomic molecule—a molecule that contains two atoms.
Other elements found in nature in the form of diatomic molecules include hydrogen, nitrogen, and the halogens.
Br I N Cl H O F

5. Molecular compounds
Molecules can also be made of atoms of different elements.
A compound composed of molecules is called a molecular compound.
Water is an example of a molecular compound.
A molecular formula is the chemical formula of a molecular compound.
A molecular formula shows how many atoms of each element a substance contains.
The subscript written after an element’s symbol indicates the number of atoms of each element in the molecule.
If there is only one atom, the subscript 1 is omitted.

6. example Butane is also a molecular compound.
Butane is commonly used in lighters and household torches.
The molecular formula for butane is C4H10
According to this formula, one molecule of butane contains four carbon atoms and ten hydrogen atoms.

7. Molecular structure
A molecular formula reflects the actual number of atoms in each molecule.
A molecular formula does not tell you about a molecule’s structure.
In other words, it does not show either the arrangement of the various atoms in space or which atoms are covalently bonded to one another.
A variety of diagrams and molecular models can be used to show the arrangement of atoms in a molecule.
The arrangement of atoms within a molecule is called its molecular structure.

8. Practice
Acetylsalicylic acid, also known as aspirin, has a molecular formula of C9H8O4. What elements make up acetylsalicylic acid? How many atoms of each element are found in one molecule of acetylsalicylic acid?

9. Molecular compounds
Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds.
Many molecular compounds are gases or liquids at room temperature.
In contrast to ionic compounds, which are formed from a metal combined with a nonmetal, most molecular compounds are composed of atoms of two or more nonmetals.

10. Ionic vs covalent

11. Metal and a non-metal
In covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases.
Combinations of atoms of the nonmetals and metalloids in Groups 4A, 5A, 6A, and 7A of the periodic table are likely to form covalent bonds.
The combined atoms usually acquire a total of eight electrons, or an octet, by sharing electrons, so that the octet rule applies.

12. Single covalent bond
Two atoms held together by sharing one pair of electrons are joined by a single covalent bond.
Hydrogen gas consists of diatomic molecules whose atoms share only one pair of electrons, forming a single covalent bond.
An electron dot structure such as H:H represents the shared pair of electrons of the covalent bond by two dots.
The pair of shared electrons forming the covalent bond is also often represented as a dash, as in H—H for hydrogen.
A structural formula represents the covalent bonds as dashes and shows the arrangement of covalently bonded atoms.

13. Unshared electrons
A pair of valence electrons that is not shared between atoms is called an unshared pair, also known as a lone pair or a nonbinding pair.
In F2, each fluorine atom has three unshared pairs of electrons.

14. practice
Hydrochloric acid (HCl (aq)) is prepared by dissolving gaseous hydrogen chloride (HCl (g)) in water. Hydrogen chloride is a diatomic molecule with a single covalent bond. Draw the electron dot structure for HCl.

15. Double and Triple Covalent Bonds
A double covalent bond is a bond that involves two shared pairs of electrons.
The carbon dioxide (CO2) molecule contains two oxygens, each of which shares two electrons with carbon to form a total of two carbon–oxygen double bonds.
Similarly, a bond formed by sharing three pairs of electrons is a triple covalent bond.
A single nitrogen atom has five valence electrons; each nitrogen atom in the molecule must share three electrons to have the electron configuration of neon.

16. Electron dot structure
Diatomic Elements
Name
Chemical formula
Electron dot structure
Properties and uses
Fluorine F2
Greenish-yellow reactive toxic gas. Compounds of fluorine, a halogen, are added to drinking water and toothpaste to promote healthy teeth.
Bromine
Br2
Dense red-brown liquid with pungent odor. Compounds of bromine, a halogen, are used in the preparation of photographic emulsions.
Hydrogen H2
Colorless, odorless, tasteless gas. Hydrogen is the lightest known element.

17. Polyatomic ion
A polyatomic ion, such as NH4+, is a tightly bound group of atoms that has a positive or negative charge and behaves as a unit.
Most polyatomic cations and anions contain covalent bonds.
Therefore, compounds containing polyatomic ions include both ionic and covalent bonding.

18. Common Polyatomic ions

19. practice
The H3O+ ion forms when a hydrogen ion is attracted to an unshared electron pair in a water molecule. Draw the electron dot structure of the hydronium ion.

20. Predicting Bond Type

21. LDS Mechanics
Atoms are represented by atomic symbols surrounded by valence electrons.
Lone Pair (6 x)
Electron pairs between atoms indicate bond formation.
Bonding Pair

22. LDS Mechanics (cont.) Three steps for “basic” Lewis structures:
Sum the valence electrons for all atoms to determine
total number of electrons.
Use pairs of electrons to form a bond between each
pair of atoms (bonding pairs).
Arrange remaining electrons around atoms (lone pairs)
to satisfy the “octet rule” (“duet” rule for hydrogen).

23. LDS Mechanics (cont.)
An example: Cl2O
20 e-
16 e- left

24. LDS Mechanics (cont.)
An example: CH4
8 e-
0 e- left
Done!

25. LDS Mechanics (cont.) An example: CO2 16 e- 12 e- left 0 e- left
Octet Violation
0 e- left
CO double bond

26. LDS Mechanics (cont.)
An example: NO+ + +
10 e-
8 e- left +

27. LDS Mechanics (cont.)
NO3-
24 e-

28. Resonance Structures
We have assumed up to this point that there is one correct Lewis structure.

29. Resonance Structures (cont.)
The classic example: O3.
Both structures are correct!

30. Resonance Structures (cont.)
In this example, O3 has two resonance structures:
Conceptually, we think of the bonding being an average of these two structures.
Electrons are delocalized between the oxygens.

31. Resonance Structures (cont.)
NO3- is a classic example of resonance:

32. Shapes and Polarity of Molecules

33. VSEPR In the valence-shell electron-pair repulsion theory
(VSEPR), the electron groups around a central atom
are arranged as far apart from each other as possible
have the least amount of repulsion of the negatively charged electrons
have a geometry around the central atom that determines molecular shape

34. Shapes of Molecules The three-dimensional shape of a molecule
is the result of bonded groups and lone pairs of electrons around the central atom
is predicted using the VSEPR theory (valence-shell-electron- pair repulsion)

35. Guide to Predicting Molecular Shape (VSEPR Theory)

36. Two Electron Groups In a molecule of BeCl2,
there are two electron groups bonded to the central atom, Be (Be is an exception to the octet rule)
: Cl : Be : Cl :
to minimize repulsion, the arrangement of two electron groups is 180°, or opposite each other
the shape of the molecule is linear

37. Two Electron Groups with Double Bonds
In a molecule of CO2,
there are two electron groups bonded to C (electrons in each double bond are counted as one group)
repulsion is minimized with the double bonds opposite each other at 180°
the shape of the molecule is linear

38. Three Electron Groups In a molecule of BF3,
three electron groups are bonded to the central atom B (B is an exception to the octet rule) ..
: F:
: F : B : F :
repulsion is minimized with 3 electron groups at angles of 120°
the shape is trigonal planar

39. Two Electron Groups and One Lone Pair
In a molecule of SO2,
S has 3 electron groups; 2 electron groups bonded to O atoms and one lone pair
:O :: S : O : ..
repulsion is minimized with the electron groups at angles of 120°, a trigonal planar arrangement
the shape is Bent (120°), with two O atoms bonded to S
● ●

40. Four Electron Groups In a molecule of CH4,
there are four electron groups around C
repulsion is minimized by placing four electron groups at angles of 109°, which is a tetrahedral arrangement
the four bonded atoms form a tetrahedral shape

41. Three Bonding Atoms and One Lone Pair
In a molecule of NH3,
three electron groups bond to H atoms, and the fourth one is a lone (nonbonding) pair
repulsion is minimized with 4 electron groups in a tetrahedral arrangement
the three bonded atoms form a pyramidal (~109°) shape

42. Two Bonding Atoms and Two Lone Pairs
In a molecule of H2O,
two electron groups are bonded to H atoms and two are lone pairs (4 electron groups)
four electron groups minimize repulsion in a tetrahedral arrangement
the shape with two bonded atoms is bent (~105°)

43. Shapes with Two and Three Electron Groups

44. Shapes with Four Electron Groups

45. Polar Bonds Polar bonds have a positive and negative end.
The charge difference is due to the difference in electronegativity of the two bonded atoms.
Electronegativity is a measure of how much an atom wants to gain electrons. F is the most electronegative atom on the Periodic Table.

46. Polar Bonds
The direction of greatest electronegativity is shown with an arrow called a dipole.
The dipole point toward the atom with the greatest electronegativity (closest to F)

47. Polar Molecules
Polar Molecules are molecules which have an uneven distribution of charge. One side of the molecule is negative while one side of the molecule is slightly positive.
Non-polar molecules are molecules in which there is no net separation of charge. The electrons are evenly distributed. There is no net separation of charge.

48. Polarity of Molecules Cl Be Cl Net Dipole H C N
Non-polar molecules – the dipoles are all equal and opposite. They cancel out and the molecule is non-polar.
Cl Be Cl
Polar molecules – the dipoles are not equal and opposite so the molecule has a net dipole and is polar.
Net Dipole
H C N
No Net Dipole

49. Polarity of Molecules
There are two steps in determining the polarity of a molecule:
Step 1: Use electronegativities to determine the direction of the dipoles for each bond making up the molecule. (F is the most electronegative)
Step 2: Determine if there is a net dipole by looking at the shape of the molecule.

50. Polarity of Molecules Step 2: Determine the shape of the molecule
To determine the molecular shape of a molecule we must first determine its Lewis dot diagram.
According to VSEPR theory, since the carbon atom is surrounded by two electron clouds (remember multiple bonds only count as one cloud), the shape of this molecule must be linear.

51. Polarity of Molecules
Step 2: Determine the shape of the molecule (continued)
Once we know the shape of the molecule, we must analysize how the electrons are distributed to determine if there is an even distribution (non-polar) or uneven distribution (polar).
In this case we know that the oxygen is more electronegative (closer to F) than the carbon and therefore should pull the electrons out away from the carbon.

52. But, the dipoles are equal and opposite so the molecule is non-polar.
If we look at the charge distribution in each bond, we get the following:
No Net Dipole
Since the oxygen is more electronegative than the carbon, the electrons will be pulled toward the oxygen atoms and away from the carbon atoms.
But, the dipoles are equal and opposite so the molecule is non-polar.

53. Example #2 – SO2
To determine the molecular shape of a molecule we must first determine its Lewis dot diagram.
According to VSEPR theory, since the sulfur atom is surrounded by three electron clouds (remember multiple bonds only count as one cloud), the shape of this molecule must be bent.

54. If we look at the charge distribution in each bond, we get the following:
Net dipole
Since the polarity of the bonds and the shape of the molecule result in an uneven distribution of charge – SO2 is a polar molecule.

55. Now that you have seen how to apply the two steps to determine the polarity of molecules, see if you can predict the polarity of the following:
1. H2O PH3
2. CCl4
3. Ammonia (NH3)
4. SO3
5. CH3Cl

56. Intermolecular Forces
The forces with which molecules attract each other.
Intermolecular forces are weaker than ionic or covalent bonds (intramolecular forces).
Intermolecular forces are responsible for the physical state of a compound (solid, liquid or gas).

57. Intermolecular Forces
Van der Waals Forces
Dispersion Forces
Dipole Interactions
Hydrogen Bonds

58. Van der Waals Forces –Dispersion Forces
They are the weakest attractions between molecules.
All covalent compounds with electrons have dispersion forces.

59. Van der Waals Forces-Dispersion Forces
Caused by the motion of electrons.
Increase as the number of electrons increases.
Weakest of all intermolecular forces.

60. Van der Waals-Dipole Interactions
Electrostatic interaction between the oppositely charged regions of polar molecules (dipoles).
Positive atom on one polar molecule attracts to the negative atom of another polar molecule.

61. Dipole
A polar molecule that has two poles.

62. Which Molecules have Dipole Interactions?
Polar
Non-polar

63. Hydrogen Bonding
Hydrogen bonding is the attraction between a hydrogen atom of a molecule to an unshared pair of electrons in another molecule.
Hydrogen bonding occurs in molecules where hydrogen is covalently bonded to a very electronegative element.
Hydrogen bonding occurs in molecules containing N, O, F.

64. Hydrogen Bonding, Continued
Hydrogen bonds are the strongest of all intermolecular forces.
Hydrogen bonds are possible because in hydrogen atoms there is no shielding of the nucleus.

65. Hydrogen Bonding