1. Chapter 6 Table of Contents Section 1 Covalent Bonds
Covalent Compounds
Table of Contents
Section 1 Covalent Bonds
Section 2 Drawing and Naming Molecules
Section 3 Molecular Shapes

2. Chapter 6
Section 1 Covalent Bonds
Objectives
Explain the role and location of electrons in a covalent bond.
Describe the change in energy and stability that takes place as a covalent bond forms.
Distinguish between nonpolar and polar covalent bonds based on electronegativity differences.

3. Chapter 6 Objectives, continued
Section 1 Covalent Bonds
Objectives, continued
Compare the physical properties of substances that have different bond types, and relate bond types to electronegativity differences.

4. Chapter 6 Sharing Electrons
Section 1 Covalent Bonds
Sharing Electrons
When an ionic bond forms, electrons are rearranged and are transferred from one atom to another to form charged ions.
In another kind of change involving electrons, the neutral atoms share electrons.

5. Sharing Electrons, continued
Chapter 6
Section 1 Covalent Bonds
Sharing Electrons, continued
Forming Molecular Orbitals
A covalent bond is a bond formed when atoms share one or more pairs of electrons.
The shared electrons move within a space called a molecular orbital.
A molecular orbital is the region of high probability that is occupied by an individual electron as it travels with a wavelike motion in the three-dimensional space around one of two or more associated nuclei.

6. Chapter 6
Formation of a Covalent Bond

7. Chapter 6 Energy and Stability
Section 1 Covalent Bonds
Energy and Stability
Energy Is Released When Atoms Form a Covalent Bond

8. Energy and Stability, continued
Chapter 6
Section 1 Covalent Bonds
Energy and Stability, continued
Potential Energy Determines Bond Length
When two bonded hydrogen atoms are at their lowest potential energy, the distance between them is 75 pm.
The bond length is the distance between two bonded atoms at their minimum potential energy.
However, the two nuclei in a covalent bond vibrate back and forth. The bond length is thus the average distance between the two nuclei.

9. Chapter 6
Visual Concepts
Bond Length

10. Energy and Stability, continued
Chapter 6
Section 1 Covalent Bonds
Energy and Stability, continued
Bonded Atoms Vibrate, and Bonds Vary in Strength
The bond length is the average distance between two nuclei in a covalent bond.
At a bond length of 75 pm, the potential energy of H2 is –436 kJ/mol.
Thus 436 kJ of energy must be supplied to break the bonds in 1 mol of H2 molecules.
The energy required to break a bond between two atoms is the bond energy.
Bonds that have the higher bond energies (stronger bonds) have the shorter bond lengths.

11. Chapter 6
Visual Concepts
Bond Energy

12. Electronegativity and Covalent Bonding
Chapter 6
Section 1 Covalent Bonds
Electronegativity and Covalent Bonding
In covalent bonds between two different atoms, the atoms often have different attractions for shared electrons.
Electronegativity values are a useful tool to predict what kind of bond will form.

13. Electronegativity and Covalent Bonding, continued
Chapter 6
Section 1 Covalent Bonds
Electronegativity and Covalent Bonding, continued
Atoms Share Electrons Equally or Unequally
When the electronegativity values of two bonding atoms are similar, bonding electrons are shared equally.
A covalent bond in which the bonding electrons in the molecular orbital are shared equally is a nonpolar covalent bond.

14. Electronegativity and Covalent Bonding, continued
Chapter 6
Section 1 Covalent Bonds
Electronegativity and Covalent Bonding, continued
Atoms Share Electrons Equally or Unequally, continued
When the electronegativity values of two bonding atoms are different, bonding electrons are shared unequally.
A covalent bond in which the bonding electrons in the molecular orbital are shared unequally is a polar covalent bond.

15. Chapter 6
Predicting Bond Character from Electronegativity Differences

16. Electronegativity and Covalent Bonding, continued
Chapter 6
Section 1 Covalent Bonds
Electronegativity and Covalent Bonding, continued
Polar Molecules Have Positive and Negative Ends
In a polar covalent bond, the ends of the bond have opposite partial charges.
A molecule in which one end has a partial positive charge and the other end has a partial negative charge is called a dipole.
In a polar covalent bond, the shared pair of electrons is not transferred completely. Instead, it is more likely to be found near the more electronegative atom.

17. Electronegativity and Covalent Bonding, continued
Chapter 6
Section 1 Covalent Bonds
Electronegativity and Covalent Bonding, continued
Polar Molecules Have Positive and Negative Ends, continued
The symbol  is used to mean partial.
+ is used to show a partial positive charge
– is used to show a partial negative charge charge
example: H+F–
Because the F atom has a partial negative charge, the electron pair is more likely to be found nearer to the fluorine atom

18. Comparing Polar and Nonpolar Covalent Bonds
Chapter 6
Visual Concepts
Comparing Polar and Nonpolar Covalent Bonds

19. Polarity Is Related to Bond Strength
Chapter 6
Section 1 Covalent Bonds
Polarity Is Related to Bond Strength
In general, the greater the electronegativity difference, the greater the polarity and the stronger the bond.

20. Electronegativity and Bond Types
Chapter 6
Section 1 Covalent Bonds
Electronegativity and Bond Types
Differences in electronegativity values provide one model that can tell you which type of bond two atoms will form.
Another general rule states:
A covalent bond forms between two nonmetals.
An ionic bond forms between a nonmetal and a metal.

21. Properties of Substances Depend on Bond Type
Chapter 6
Section 1 Covalent Bonds
Properties of Substances Depend on Bond Type
The type of bond that forms (metallic, ionic, or covalent) determines the properties of the substance.
The difference in the strength of attraction between the basic units of ionic and covalent substances causes these types of substances to have different properties.

22. Chapter 6
Properties of Substances with Metallic, Ionic, and Covalent Bonds

23. Section 2 Drawing and Naming Molecules
Chapter 6
Objectives
Draw Lewis structures to show the arrangement of valence electrons among atoms in molecules and polyatomic ions.
Explain the differences between single, double, and triple covalent bonds.
Draw resonance structures for simple molecules and polyatomic ions, and recognize when they are required.

24. Chapter 6 Objectives, continued
Section 2 Drawing and Naming Molecules
Chapter 6
Objectives, continued
Name binary inorganic covalent compounds by using prefixes, roots, and suffixes.

25. Lewis Electron-Dot Structures
Section 2 Drawing and Naming Molecules
Chapter 6
Lewis Electron-Dot Structures
Valence electrons are the electrons in the outermost energy level of an atom.
A Lewis structure is a structural formula in which valence electrons are represented by dots.
In Lewis structures, dot pairs or dashes between two atomic symbols represent pairs in covalent bonds.

26. Lewis Electron-Dot Structures, continued
Section 2 Drawing and Naming Molecules
Chapter 6
Lewis Electron-Dot Structures, continued
Lewis Structures Show Valence Electrons
As you go from element to element across a period, you add a dot to each side of the element’s symbol.

27. Lewis Electron-Dot Structures, continued
Section 2 Drawing and Naming Molecules
Chapter 6
Lewis Electron-Dot Structures, continued
Lewis Structures Show Valence Electrons, continued
You do not begin to pair dots until all four sides of the element’s symbol have a dot.

28. Lewis Electron-Dot Structures, continued
Section 2 Drawing and Naming Molecules
Chapter 6
Lewis Electron-Dot Structures, continued
Lewis Structures Show Valence Electrons, continued
An element with an octet of valence electrons has a stable configuration.
The tendency of bonded atoms to have octets of valence electrons is called the octet rule.

29. Chapter 6
Visual Concepts
The Octet Rule

30. Lewis Electron-Dot Structures, continued
Section 2 Drawing and Naming Molecules
Chapter 6
Lewis Electron-Dot Structures, continued
Lewis Structures Show Valence Electrons, continued
When two chlorine atoms form a covalent bond, each atom contributes one electron to a shared pair.
An unshared pair, or a lone pair, is a nonbonding pair of electrons in the valence shell of an atom.

31. Lewis Electron-Dot Structures, continued
Section 2 Drawing and Naming Molecules
Chapter 6
Lewis Electron-Dot Structures, continued
Lewis Structures Show Valence Electrons, continued
A single bond is a covalent bond in which two atoms share one pair of electrons
The electrons can pair in any order. However, any unpaired electrons are usually filled in to show how they will form a covalent bond.

32. Chapter 6 Multiple Bonds
Section 2 Drawing and Naming Molecules
Chapter 6
Multiple Bonds
For O2 to make an octet, each atom needs two more electrons. The two atoms share four electrons.
A double bond is a covalent bond in which two atoms share two pairs of electrons.

33. Multiple Bonds, continued
Section 2 Drawing and Naming Molecules
Chapter 6
Multiple Bonds, continued
For N2 to make an octet, each atom needs three more electrons. The two atoms share six electrons.
A triple bond is a covalent bond in which two atoms share three pairs of electrons.

34. Chapter 6 Resonance Structures
Section 2 Drawing and Naming Molecules
Chapter 6
Resonance Structures
Some molecules, such as ozone, O3, cannot be represented by a single Lewis structure.
When a molecule has two or more possible Lewis structures, the two structures are called resonance structures.

35. Multiple Bonds, continued
Section 2 Drawing and Naming Molecules
Chapter 6
Multiple Bonds, continued
Naming Covalent Compounds
The first element named is usually the first one written in the formula. It is usually the less-electronegative element.
The second element named has the ending -ide.
Unlike the names for ionic compounds, the names for covalent compounds must often distinguish between two different molecules made of the same elements.

36. Naming Covalent Compounds, continued
Section 2 Drawing and Naming Molecules
Chapter 6
Naming Covalent Compounds, continued
This system of prefixes is used to show the number of atoms of each element in the molecule.

37. Naming Covalent Compounds, continued
Section 2 Drawing and Naming Molecules
Chapter 6
Naming Covalent Compounds, continued
Prefixes can be used to show the numbers of each type of atom in diphosphorus pentasulfide.

38. Naming Compounds Using Numerical Prefixes
Chapter 6
Visual Concepts
Naming Compounds Using Numerical Prefixes

39. Determining Molecular Shapes
Chapter 6
Section 3 Molecular Shapes
Determining Molecular Shapes
The three-dimensional shape of a molecule is important in determining the molecule’s physical and chemical properties.
A Lewis Structure Can Help Predict Molecular Shape
You can predict the shape of a molecule by examining the Lewis structure of the molecule.

40. Determining Molecular Shapes, continued
Chapter 6
Section 3 Molecular Shapes
Determining Molecular Shapes, continued
A Lewis Structure Can Help Predict Molecular Shape, continued
The valence shell electron pair repulsion (VSEPR) theory is a theory that predicts some molecular shapes based on the idea that pairs of valence electrons surrounding an atom repel each other.

41. Determining Molecular Shapes, continued
Chapter 6
Section 3 Molecular Shapes
Determining Molecular Shapes, continued
Electron Pairs Can Determine Molecular Shape
According to the VSEPR theory, the shape of a molecule is determined by the valence electrons surrounding the central atom.
Electron pairs are negative, so they repel each other.
Therefore, the shared pairs that form different bonds repel each other and remain as far apart as possible.

42. Determining Molecular Shapes, continued
Chapter 6
Section 3 Molecular Shapes
Determining Molecular Shapes, continued
Electron Pairs Can Determine Molecular Shape, continued
For CO2, the two double bonds around the central carbon atom repel each other and remain far apart.
For BF3, the three single bonds around the central fluorine atom will be at a maximum distance apart.

43. Determining Molecular Shapes, continued
Chapter 6
Section 3 Molecular Shapes
Determining Molecular Shapes, continued
Electron Pairs Can Determine Molecular Shape, continued
The four shared pairs of electrons in CH4 are farthest apart when each pair is positioned at the corners of a tetrahedron.

44. VSEPR and Lone Electron Pairs
Chapter 6
Chapter 6
Visual Concepts
VSEPR and Lone Electron Pairs

45. VSEPR and Basic Molecular Shapes
Chapter 6
Visual Concepts
VSEPR and Basic Molecular Shapes

46. Predicting Molecular Shapes
Chapter 6
Section 3 Molecular Shapes
Predicting Molecular Shapes
Sample Problem D
Determine the shape of H2O.

47. Predicting Molecular Shapes
Chapter 6
Section 3 Molecular Shapes
Predicting Molecular Shapes
Sample Problem D Solution
Draw the Lewis structure for H2O.
Count the number of shared and unshared pairs of electrons around the central atom.
H2O has two shared pairs and two unshared pairs.

48. Predicting Molecular Shapes
Chapter 6
Section 3 Molecular Shapes
Predicting Molecular Shapes
Sample Problem D Solution, continued
Find the shape that allows the shared and unshared pairs of electrons to be as far apart as possible.
The water molecule will have a bent shape.

49. Molecular Shape Affects a Substance’s Properties
Chapter 6
Section 3 Molecular Shapes
Molecular Shape Affects a Substance’s Properties
Shape Affects Polarity
One property that shape determines is the polarity of a molecule.
The polarity of a molecule that has more than two atoms depends on the polarity of each bond and the way the bonds are arranged in space.

50. Molecular Shape Affects a Substance’s Properties, continued
Chapter 6
Section 3 Molecular Shapes
Molecular Shape Affects a Substance’s Properties, continued
Shape Affects Polarity, continued
If two dipoles are arranged in opposite directions, they will cancel each other.
If two dipoles are arranged at an angle, they will not cancel each other.
Compare the molecules of nonpolar carbon dioxide, CO2, which has a linear shape, and polar water, H2O, which has a bent shape.

51. Chapter 6
Molecular Shape Affects Polarity