1. Chapter 8 Concepts of Chemical Bonding
Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten
Chapter 8 Concepts of Chemical Bonding
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice Hall, Inc.

2. 1. WHY do Atoms form Chemical Bonds?
Atoms have lower potential energy when they are bonded to other atoms than when they are independent particles and this is a GOOD thing!
Matter wants to be in its LOWEST POSSIBLE ENERGY state because they are more stable and less reactive

3. 1. HOW do Atoms form Chemical Bonds?
Only the valence (outermost) electrons of atoms are involved in bonding (never inner core electrons)
These electrons are the furthest from the nucleus and least attracted to the atom making them easiest remove/attract away from the atomic nucleus
In a chemical bond, when atoms get close to one another, there are both attractive and repulsive forces at work:

4. Attractive Forces– Repulsive Forces–
the nucleus of one atom is attracted to the electrons of the other atom (causes a decrease in total potential energy = good)
Repulsive Forces–
the two sets of nuclei and/or two sets of electron clouds begin to repel one another (causes an increase in total potential energy = bad)

5. Bond Length: the ideal distance between bonded atoms in which attractive forces are maximized and repulsive forces are minimized (this distance depends on the identities of the atoms bonding)
Bond Strength: the more attracted two atoms are to one another, the greater the bond strength (this can be quantifiably measured by calculating the bond enthalpy (aka: bond energy)
Bond Enthalpy: the energy required to break a chemical bond (unit: kJ/mol)

6. Bond Type, Bond Length, & Bond Enthalpy
Chemical Bond
Bond Length (pm)
Bond Enthalpy (kJ/mol)
Cl-Cl
199
243
Br-Br
229
193
H-Cl
127
432
H-Br
141
366
Relationship:
As bond length
____________, bond strength and bond enthalpy
____________.
decreases
increases

7. 2-4. Types of Chemical Bonds
Three types of chemical bonds:
Ionic
Covalent
Metallic

8. 18. Table of Electronegativies:
Keep in mind that electronegativities are approximate measures of the relative tendencies of these elements to attract electrons to themselves in a chemical bond. The greater the atom’s electronegativity, the greater its ability to attract electrons to itself.

9. 2. Ionic Bonding
nonmetal atom gains electrons from a metal atom due to a large difference in electronegativity (∆EN > 1.7).
This complete transfer of electrons results in a cation and anion which are then electrostatically attracted to one another (opposite charges) resulting in a very strong bond between the ions.
Example:
Na–Cl
∆EN = │1.0 – 3.0│= 2.0
Metal transfer Nonmetal
(low EN) (high EN) e–

10. All ionic compounds have a particular arrangement known as a crystal lattice structure.
This structure organizes the cations and anions so that they surround one another giving the compound its lowest potential energy (most stable arrangement)

11. Lattice Energy–
the energy released when 1 mole of an ionic crystalline compound is formed from gaseous ions (unit: kJ/mol)
While bond enthalpy is energy REQUIRED to BREAK a chemical bond, lattice energy is energy RELEASED when 1 mole of ionic bonds are FORMED.
The strong attraction of (+)/(–) ions is described by the lattice energy in the following figure. All lattice energies are negative values (aka: energy released)

12. 14. Lattice Energies of Some Common Ionic Compounds
Which ionic compound in the table has the strongest ionic bonding?
Why do you think that is?
The higher (more negative) the lattice energy, the stronger the ionic bond and the more stable the compound!
Compound
Lattice Energy
(kJ/mol)
NaCl
-787.5
NaBr
-751.4
CaF2
LiCl
-861.3
LiF
-1032
MgO
-3760
KCl
-715

13. 3. Covalent Bonding
two nonmetal atoms sharing electrons in order to achieve their octets since nonmetals have high electronegativities

14. Sharing the Electrons….
Electronegativity (ability of an atom to attract electrons toward itself) plays a crucial role in determining the type of chemical bond formed b/c not all of the atoms in a bond will have the same electronegativity.
The difference in two atom’s EN values will determine the type (and polarity) of the bond formed.
Bonding between atoms of different elements is rarely 100% ionic or 100% covalent.

15. Polar Covalent Bonding
Unequal sharing of electrons between two nonmetal atoms due to an intermediate difference in electronegativity (∆EN = ).
This means that the electrons in a polar bond are pulled closer to (and spend more time near) the more EN atom.
Example:
P–O
∆EN = │2.1 – 3.5│= 1.4

16. Nonpolar Covalent Bonding
Equal sharing of electrons between two nonmetal atoms with identical or very similar electronegativities (∆EN < 0.39).
This means that no one atom attracts electrons better than the other
Example:
F –F
∆EN = │4.0 – 4.0│= 0
S– Br
∆EN = │2.5 – 2.8│= 0.3

17. 3. Metallic Bonding
metal atoms have typically low electronegativities and tend to lose electrons forming cations. When these metals bond together, there is an electrostatic attraction between these lost/delocalized electrons which forms a sea of randomly moving electrons
Example:
Cu–Cu
Cu– Zn “brass”

18. Difference in Electronegativity
ΔEN & Bond Type
Difference in Electronegativity
Bond Type
< 0.39
Nonpolar Covalent
Polar Covalent
> 1.7
Ionic
Increase in ionic character of a bond

19. 5. Strength of Chemical Bonds
Ionic bonds are the strongest (due to the presence of ions with full charges)
Polar covalent bonds are stronger than nonpolar covalent bonds (due to the partial charges in the bond that result in polar bonds from larger ∆EN)

20. 20. Partial Charge in Polar Covalent Bonds
The more EN element in a polar covalent bond will attract the e– closer to itself, resulting in a partial negative charge on that atom
The opposite end of the bond (less EN atom) is thus deprived of e– near it, so it forms a partial positive charge
Note: only ionic bonds have full charges (transfer of e–) and nonpolar bonds have NO charges (equal sharing)
Two ways to indicate the polarity of a bond:

22. 22. Atoms tend to form bonds that follow the octet rule
22. Atoms tend to form bonds that follow the octet rule. What is the octet rule?
Octet Rule-
chemical compounds tend to form so that each atom (either by gaining, losing, or sharing electrons) has an octet (8 valence electrons) in its highest energy level
*Everyone wants to be like the Noble Gases!
*This rule is responsible for ALL chemical bonding between atoms!!!
In order to fulfill the octet rule through covalent bonding atoms share electrons

23. 23. Lewis Structures: Representations of Covalent Compounds
A way to show which atoms are bonded to which in a covalent compound
Shows how many electrons are covalently shared between each of the atoms

24. 24. Using Lewis Structures to Represent Covalent Compounds
Show how two fluorine atoms form a single covalent bond to fulfill the octet rule:
Covalent = nonmetals sharing electrons
Octet rule = each atom wants 8 valence electrons
By sharing, each atom can count the e– as a part of their complete octet!!!

25. 24. Using Lewis Structures to Represent Covalent Compounds
The pair of electrons in between the two fluorine atoms represents the shared pair of electrons involved in the covalent bond (aka: bonding electrons). This pair is often replaced by a long dash which represents the covalent bond
The unshared or lone pair electrons are those not involved in the bonding. They belong exclusively to one atom. (aka: nonbonding electrons)

26. 25. Multiple Bonds in Covalent Molecules
Single bond:
sharing of ONE pair of e– (2 total electrons)
Double bond:
sharing of TWO pairs of e– (4 total electrons)
Triple bond:
sharing of THREE pairs of e– (6 total electrons)

27. 26: Bond lengths & energies of single, double, & triple bonds
Bond Length (pm)
Bond Energy (kJ/mol)
C-C
(single)
154
346
C=C
(double)
134
612
(triple)
120
835

28. 27. What happens to length, strength, & bond energy when a multiple bond is formed?
Bond length decreases
Bond strength increases
Bond energy/enthalpy increases
Single Bonds  longest & weakest
Triple Bonds  shortest & strongest

29. Rules for drawing Lewis Structures
Packet pg. 4

30. Writing Lewis Structures
Determine the number of atoms of each element present in the molecule
1 P atom
3 Cl atoms
PCl3

31. Writing Lewis Structures
2. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.
For anions  add one electron for each negative charge.
For cations  subtract one electron for each positive charge.
PCl3
(7) = 26

32. Writing Lewis Structures
3. The central atom is the least electronegative element but NEVER hydrogen.
Connect the outer atoms to the central with single bonds (a dash)
Keep track of the electrons:
26  6 = 20

33. Writing Lewis Structures
4. Complete the octets of the outer atoms bonded to the central
Keep track of the electrons:
26  6 = 20  18 = 2

34. Writing Lewis Structures
5. Place any leftover electrons on the central atom, even if doing so results in more than an octet around the central atom
Expanded Octet – central atom has more than 8 electrons; this is allowed for any atom in the 3rd row & after
Writing Lewis Structures
Keep track of the electrons:
26  6 = 20  18 = 2  2 = 0

35. Writing Lewis Structures
If there are not enough electrons to give the central atom a complete octet, multiple bonds are needed
Remove a nonbonding pair of electrons from an outer atom
Draw a double bond connecting the central atom and that outer atom
If need be, a triple bond may be formed by removing another pair of nonbonding electrons from the same outer atom
Packet pg. 5

36. 28. What is resonance & when does it occur?
Resonance- refers to the bonding in molecules or ions that CANNOT be correctly represented by a SINGLE Lewis Structure
Example: Nitrite ion, NO2–
Packet pg. 5

37. Metallic Bonding

38. 29. What is a metallic bond?
Metallic Bond – chemical bonding that results from the attraction between metal atoms and the surrounding “sea of electrons”
“Sea of Electrons” – valence electrons that can move freely throughout a network of metal atoms (they are not attached to any one atom

39. Metallic Bonding & Sea of Electrons

40. 30. What are some unique properties of metals that are related to metallic bonds?
The freedom of electrons to move in a network of metal atoms accounts for:
1) high electrical & thermal conductivity
2) strong absorbers & reflectors of light
= lustrous
3) malleable & ductile

41. 31. Enthalpy of Vaporization-
The amount of energy, as heat, required to vaporize a metal is a measure of the strength of the metallic bonds that hold it together
Generally, the more valence electrons the metal has, the stronger the metallic bonding (refer to figure 4.3 on next slide).
*The higher the enthalpy of vaporization, the stronger the metallic bond (more energy is required to vaporize the metal)

42. 32. Figure 4.3 – Enthalpies of Vaporization of Metals (kJ/mol)
Group 1
Group 2
Group 13
Li = 146
Be = 297
Na = 97
Mg = 128
Al = 294
K = 77
Ca = 155
Sc = 333
Rb = 76
Sr = 137
Y = 365
Cs = 64
Ba = 140
La = 402