1. Chapter 6- Chemical Bonding

2. Chemical Bonding
chemical bond- a mutual electrical attraction between the nuclei and the valence electrons of different atoms that binds the atoms together

3. ionic bond- a chemical bond that results from the electrical attraction between anions and cations

4. Formation of Ions from Metals
Ionic compounds result when metals react with nonmetals
Metals lose electrons to match the number of valence electrons of their nearest noble gas
Positive ions form when the number of electrons are less than the number of protons
Group 1 metals  ion 1+
Group 2 metals  ion 2+
Group 13 metals  ion 3+

5. Formation of Sodium Ion
Sodium atom Sodium ion
Na  – e  Na +
( = Ne)
11 p p+
11 e e-

6. Learning Check A. Number of valence electrons in aluminum
1) 1 e ) 2 e- 3) 3 e-
B. Change in electrons for octet
1) lose 3e ) gain 3 e ) gain 5 e-
C. Ionic charge of aluminum
1) ) ) 3+

7. Solution A. Number of valence electrons in aluminum 3) 3 e-
B. Change in electrons for octet
1) lose 3e-
C. Ionic charge of aluminum
3) 3+

8. Na + Cl  NaCl
sodium loses its one valence electron to form the cation Na+
This allows the atom to become like neon with eight electrons in it’s outer energy level.
chlorine gains that electron to form the anion Cl-
This allows the atom to become like argon with eight electrons in its outer energy level.
Na+ is attracted to Cl- (opposites attract) and an ionic bond forms

9. 1). Ionic bond – electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na+) and the Cl becomes (Cl-), charged particles or ions.

10. covalent bond- a chemical bond that results from the sharing of electron pairs between two atoms
nonpolar covalent bond- a covalent bond in which the bonding electrons are shared equally by the atoms forming the bond
-because their electronegativities are essentially equal
polar covalent bond- a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons
-because one atom has a greater electronegativity than the other

11. Nonpolar covalent bond
H · ·F
Fluorine’s greater electronegativity causes the shared electrons to move closer to it and creates areas of slight positive and negative charge forming a polar covalent bond
Nonpolar covalent bond
H· ·H
Equal electronegativities of the hydrogen atoms cause the pair of electrons to be shared equally and a nonpolar covalent bond to form.

12. 2. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons.
Oxygen Atom
Oxygen Atom
Oxygen Molecule (O2)

13. - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

14. Determining Bond Type
Find the absolute difference in the electronegativities of the bonding atoms.
The greater the difference, the greater the
% ionic character which makes it more like an ionic bond.
IF the absolute difference is
< 0.3 the bond is nonpolar covalent
> 0.3 and < 1.7 the bond is polar covalent
> 1.7 the bond is ionic

15. Chemical Bonding, continued
Section 1 Introduction to Chemical Bonding
Chapter 6
Chemical Bonding, continued
Sample Problem A
Use electronegativity values listed in Figure 20 from the previous chapter in your book, on page 161, and Figure 2 in your book, on page 176, to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl. In each pair, which atom will be more negative?

16. Chemical Bonding, continued
Section 1 Introduction to Chemical Bonding
Chapter 6
Chemical Bonding, continued
Sample Problem A Solution
The electronegativity of sulfur is 2.5. The electronegativities of hydrogen, cesium, and chlorine are 2.1, 0.7, and 3.0, respectively. In each pair, the atom with the larger electronegativity will be the more-negative atom.
Bonding between
Electroneg.
More-neg-
sulfur and
difference
Bond type
ative atom
hydrogen
2.5 – 2.1 = 0.4
polar-covalent
sulfur
cesium
2.5 – 0.7 = 1.8
ionic
chlorine
3.0 – 2.5 = 0.5

17. Chapter 6 Objectives Define molecule and molecular formula.
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Objectives
Define molecule and molecular formula.
Explain the relationships among potential energy, distance between approaching atoms, bond length, and bond energy.
State the octet rule.

18. Chapter 6 Objectives, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Objectives, continued
List the six basic steps used in writing Lewis structures.
Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both.
Explain why scientists use resonance structures to represent some molecules.

19. Chapter 6 Molecular Compounds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Molecular Compounds
A molecule is a neutral group of atoms that are held together by covalent bonds.
A chemical compound whose simplest units are molecules is called a molecular compound.

20. Visual Concepts
Chapter 6
Molecule

21. Chapter 6 Molecular Compounds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Molecular Compounds
The composition of a compound is given by its chemical formula.
A chemical formula indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.
A molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound.

22. Visual Concepts
Chapter 6
Molecule

23. Comparing Monatomic, Diatomic, and Polyatomic Molecules
Visual Concepts
Chapter 6
Comparing Monatomic, Diatomic, and Polyatomic Molecules

24. Formation of a Covalent Bond
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Formation of a Covalent Bond
The electron of one atom and proton of the other atom attract one another.
The two nuclei and two electrons repel each other.
These two forces cancel out to form a covalent bond at a length where the potential energy is at a minimum.

25. Formation of a Covalent Bond
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Formation of a Covalent Bond

26. Characteristics of the Covalent Bond
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Characteristics of the Covalent Bond
The distance between two bonded atoms at their minimum potential energy is the bond length.
In forming a covalent bond, the hydrogen atoms release energy. The same amount of energy must be added to separate the bonded atoms.
Bond energy is the energy required to break a chemical bond and form neutral isolated atoms.

27. Bond Length and Stability
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Bond Length and Stability

28. Bond Energies and Bond Lengths for Single Bonds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Bond Energies and Bond Lengths for Single Bonds

29. Characteristics of the Covalent Bond
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Characteristics of the Covalent Bond
When two atoms form a covalent bond, their shared electrons form overlapping orbitals.
This achieves a noble- gas configuration.
The bonding of two hydrogen atoms allows each atom to have the stable electron configuration of helium, 1s2.

30. The Octet Rule
octet rule- chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its outermost energy level
EXCEPTIONS:
hydrogen atoms are complete with two electrons (H2)
boron atoms are complete with 6 electrons (BF3)
some elements show expanded valence involving “d” orbitals (PF5 & SF6)

31. Electron-Dot Notation
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Electron-Dot Notation
To keep track of valence electrons, it is helpful to use electron-dot notation.
Electron-dot notation is an electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol. The inner-shell electrons are not shown.

32. Electron-Dot Notation, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Electron-Dot Notation, continued
Sample Problem B
Write the electron-dot notation for hydrogen.
Write the electron-dot notation for nitrogen.

33. N H Chapter 6 Electron-Dot Notation, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Electron-Dot Notation, continued
Sample Problem B Solution
A hydrogen atom has only one occupied energy level, the n = 1 level, which contains a single electron. H
The group notation for nitrogen’s family of elements is
ns2np3.
Nitrogen has five valence electrons. N

34. F F H H Chapter 6 Lewis Structures
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures
Electron-dot notation can also be used to represent molecules.
H H
The pair of dots between the two symbols represents the shared electron pair of the hydrogen-hydrogen covalent bond.
For a molecule of fluorine, F2, the electron-dot notations of two fluorine atoms are combined.
F F

35. pairs of electrons that are not shared in bonds.
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures
The pair of dots between the two symbols represents the shared pair of a covalent bond.
F F
In addition, each fluorine atom is surrounded by three
pairs of electrons that are not shared in bonds.
An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.

36. Lewis Structures :F-F: ¨ ¨
Lewis structures- are formulas in which atomic symbols represent nuclei and inner shell electrons, dot pairs adjacent to a single atom represent unshared electron pairs, and dashes between the atomic symbols represent covalent bonds between two atoms
¨ ¨
:F-F:
¨ ¨

37. H H F F Chapter 6 Lewis Structures
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures
The pair of dots representing a shared pair of electrons in a covalent bond is often replaced by a long dash.
H H
F F
example:
A structural formula indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule.
example: F–F H–Cl

38. Chapter 6 Lewis Structures
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures
The Lewis structures and the structural formulas for many molecules can be drawn if one knows the composition of the molecule and which atoms are bonded to each other.
A single covalent bond, or single bond, is a covalent bond in which one pair of electrons is shared between two atoms.

39. Lewis Structures, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures, continued
Sample Problem C
Draw the Lewis structure of iodomethane, CH3I.

40. I H C Chapter 6 Lewis Structures, continued Sample Problem C Solution
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures, continued
Sample Problem C Solution
Determine the type and number of atoms in the molecule.
The formula shows one carbon atom, one iodine atom, and three hydrogen atoms.
Write the electron-dot notation for each type of atom in the molecule.
Carbon is from Group 14 and has four valence electrons. Iodine is from Group 17 and has seven valence electrons. Hydrogen has one valence electron. C I H

41. Lewis Structures, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures, continued
Sample Problem C Solution, continued
3. Determine the total number of valence electrons available in the atoms to be combined. C
1 × 4e– =
4e– I
1 × 7e–
7e– 3H
3 × 1e–
3e–
14e–

42. H H HC I or H C I H H H H C I Chapter 6 Lewis Structures, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures, continued
Sample Problem C Solution, continued
4. If carbon is present, it is the central atom. Otherwise, the least- electronegative atom is central. Hydrogen, is never central. H
H C I
5. Add unshared pairs of electrons to each nonmetal atom (except hydrogen) such that each is surrounded by eight electrons.
H H
HC I or H C I H H

43. C C or Chapter 6 Multiple Covalent Bonds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds
A double covalent bond, or simply a double bond, is a covalent bond in which two pairs of electrons are shared between two atoms.
Double bonds are often found in molecules containing carbon, nitrogen, and oxygen.
A double bond is shown either by two side-by-side pairs of dots or by two parallel dashes.
H H H H or
C C
C C
H H H H

44. C H or H or N N N N C H H C C Chapter 6 Multiple Covalent Bonds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds
A triple covalent bond, or simply a triple bond, is a covalent bond in which three pairs of electrons are shared between two atoms.
example 1—diatomic nitrogen:
N N
or N N
example 2—ethyne, C2H2:
H C
C H or H C
C H

45. Multiple Covalent Bonds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds
Double and triple bonds are referred to as multiple bonds, or multiple covalent bonds.
In general, double bonds have greater bond energies and are shorter than single bonds.
Triple bonds are even stronger and shorter than double bonds.
When writing Lewis structures for molecules that contain carbon, nitrogen, or oxygen, remember that multiple bonds between pairs of these atoms
are possible.

46. Multiple Covalent Bonds, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds, continued
Sample Problem D
Draw the Lewis structure for methanal, CH2O, which is also known as formaldehyde.

47. O H C Chapter 6 Multiple Covalent Bonds, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds, continued
Sample Problem D Solution
Determine the number of atoms of each element present in the molecule.
The formula shows one carbon atom, two hydrogen atoms, and one oxygen atom.
Write the electron-dot notation for each type of atom. Carbon is from Group 14 and has four valence electrons. Oxygen, which is in Group 16, has six valence electrons. Hydrogen has only one valence electron. C O H

48. Multiple Covalent Bonds, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds, continued
Sample Problem D Solution, continued
3. Determine the total number of valence electrons available in the atoms to be combined. C
1 × 4e– =
4e– O
1 × 6e–
6e– 2H
2 × 1e–
2e–
12e–

49. H H C O H H C O Chapter 6 Multiple Covalent Bonds, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds, continued
Sample Problem D Solution, continued
Arrange the atoms to form a skeleton structure for the molecule. Connect the atoms by electron-pair bonds. H
H C O
Add unshared pairs of electrons to each nonmetal atom (except hydrogen) such that each is surrounded
by eight electrons. H
H C O

50. Multiple Covalent Bonds, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds, continued
Sample Problem D Solution, continued
6a.Count the electrons in the Lewis structure to be sure that the number of valence electrons used equals the number available.
The structure has 14 electrons. The structure has two valence electrons too many.
6b.Subtract one or more lone pairs until the total number of valence electrons is correct.
Move one or more lone electron pairs to existing bonds until the outer shells of all atoms are compl filled.
etely

51. H H H C O H C O or Chapter 6 Multiple Covalent Bonds, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds, continued
Sample Problem D Solution, continued
Subtract the lone pair of electrons from the carbon atom. Move one lone pair of electrons from the oxygen to the bond between carbon and oxygen to form a double bond. H
H C O H
H C O or

52. Drawing Lewis Structures with Many Atoms
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Drawing Lewis Structures with Many Atoms

53. Drawing Lewis Structures with Many Atoms
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Drawing Lewis Structures with Many Atoms