1. Chapter 8 Basic Concepts of Chemical Bonding
CHEMISTRY The Central Science 9th Edition
Chapter 8 Basic Concepts of Chemical Bonding
Chapter 8

2. 8.1: Chemical Bonds, Lewis Symbols, and the Octet Rule
Chemical bond: attractive force holding two or more atoms together
Covalent bond: electrons are shared
Usually found between nonmetals
Ionic bond results from the transfer of electrons from a metal to a nonmetal
Metallic bond: attractive force holding pure metals together
Chapter 8

3. Valence electrons are represented as dots around the symbol for the element
Electrons available for bonding are indicated by unpaired dots
Text, P. 300
Chapter 8

4. All noble gases except He have an ns2np6 configuration
The Octet Rule
All noble gases except He have an ns2np6 configuration
Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs)
Chapter 8

5. Na(s) + ½Cl2(g)  NaCl(s) DHºf = -410.9 kJ
8.2: Ionic Bonding
Consider the reaction between sodium and chlorine:
Na(s) + ½Cl2(g)  NaCl(s) DHºf = kJ
Chapter 8

6. NaCl forms a very regular structure
Regular arrangement of Na+ and Cl- in 3D
Ions are packed as closely as possible
Chapter 8

7. Energetics of Ionic Bond Formation
Lattice energy: the energy required to completely separate an ionic solid into its gaseous ions
NaCl(s)  Na+(g) + Cl-(g) is endothermic (H = +788 kJ/mol)
Lattice energy depends on the charges on the ions and the sizes of the ions
Chapter 8

8. Lattice energy (and thus stability) increases as
The charges on the ions increase
The distance between the ions decreases
High lattice energies make ionic compounds hard and brittle with high melting points
Chapter 8

9. Text, P. 303

10. 8.3: Covalent Bonding
When two similar atoms bond, neither of them wants to lose or gain an electron to form an octet
Similar electron affinities
They share pairs of electrons to each obtain an octet
Example: H + H  H2
Chapter 8

11. Lewis Structures
Covalent bonds can be represented by the Lewis symbols of the elements:
In Lewis structures, each pair of electrons in a bond is represented by a single line:
Chapter 8

12. Bond distances decrease from single  triple bonds
Multiple Bonds
It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds):
One shared pair of electrons = single bond
Two shared pairs = double bond
Three shared pairs = triple bond
Bond distances decrease from single  triple bonds
Chapter 8

13. 8.4: Bond Polarity and Electronegativity
Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons
Unequal sharing of electrons results in polar bonds
Chapter 8

14. Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F)
Electronegativity
Electronegativity: The ability of one atom in a molecule to attract electrons to itself
Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F)
Electronegativity increases
across a period
decreases down a group
Chapter 8

15. Electronegativity
Text, P. 310

16. Electronegativity and Bond Polarity
Difference in electronegativity is a gauge of bond polarity: (NO NOTES)
differences around 0: non-polar covalent bonds (equal or almost equal sharing of electrons)
differences around 2: polar covalent bonds (unequal sharing of electrons)
differences around 3: ionic bonds (transfer of electrons)
Chapter 8

17. There is no sharp distinction between bonding types
The positive end (or pole) in a polar bond is represented + and the negative pole -
Chapter 8

18. There is more electron density on F than on H
Dipole Moments
Consider HF:
Polar bond
There is more electron density on F than on H
Since there are two different “ends” of the molecule, we call HF a dipole
Dipole moment, m, is the magnitude of the dipole
Chapter 8

19. Bond Types and Nomenclature Ionic Molecular
MgH2
Magnesium hydride
H2S
Hydrogen sulfide
FeF2
Iron(II) fluoride
OF2
Oxygen difluoride
Mn2O3
Manganese(III) oxide
Cl2O3
Dichlorine trioxide
The least EN element is first
The more EN element ends in –ide
Use prefixes for molecular compounds
Chapter 8

20. 8.5: Drawing Lewis Structures
Calculate valence electrons needed to achieve noble gas configuration for each element in molecule (N)
Calculate the actual # of valence electrons for each element in molecule (A) (add e-/subtract e- for polyatomic ions)
Calculate the number of shared electrons in molecule (S): S = N – A
Develop a skeletal structure (least EN element central; symmetry)
Place S electrons in the structure as shared pairs
Place remaining A electrons on atoms to make octets (H and He)
Chapter 8

21. To determine which structure is most reasonable, use formal charge
It is possible to draw more than one Lewis structure with the octet rule obeyed for all the atoms
To determine which structure is most reasonable, use formal charge
Formal charge is the charge on an atom that it would have if all the atoms had the same electronegativity
Chapter 8

22. To calculate formal charge:
FC = valence electrons - number of bonds - nonbonding electrons
(For a polyatomic ion, sum equals the charge of the ion)
Chapter 8

23. There are 4 valence electrons
Consider:
For C:
There are 4 valence electrons
In the Lewis structure there are 2 nonbonding electrons and 3 bonds
Formal charge: 4 – = -1
Chapter 8

24. There are 5 valence electrons
Consider:
For N:
There are 5 valence electrons
In the Lewis structure there are 2 nonbonding electrons and 3 bonds
Formal charge = 5 – = 0
The sum of the formal charges on the ion equals the charge of the ion ( = -1)
We write:
Chapter 8

25. The most stable structure has:
the lowest formal charge on each atom (close to zero)
the most negative formal charge on the most electronegative atom
no adjacent atoms with FC of the same sign
Chapter 8

26. Lewis structures: N2, SiF4, CO2, CS2, C3H8, (SO4)-2, (NH4)+, H2SO4, SOCl2, NH2OH
Sample Problems # 47, 51
Chapter 8

27. 8.6: Resonance Structures
Some molecules are not well described by a single Lewis Structure
Resonance structures are attempts to represent a real structure that is a mix between several extreme possibilities
Chapter 8

28. Common examples: O3, SO3, NO3-, NO2, and benzene
Chapter 8

29. Resonance in Benzene
In the resonance structures for benzene there are single bonds between each pair of C atoms and the 6 additional electrons are delocalized over the entire ring:
Benzene belongs to a category of organic molecules called aromatic compounds (due to their odor)
Chapter 8

30. 8.7: Exceptions to the Octet Rule
There are three classes of exceptions to the octet rule:
1. Odd Number of Electrons
Few examples. Generally molecules such as ClO2, NO, and NO2 have an odd number of electrons
dimerization occurs
Chapter 8

31. 2. Less than an Octet Relatively rare
Molecules with less than an octet are typical for compounds of Groups 2 and 13
Be: N=4
B: N=6
Chapter 8

32. This is the largest class of exceptions
3. More than an Octet
This is the largest class of exceptions
Atoms from the 3rd period onward can accommodate more than an octet
the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density
Lewis Structures: If everything has an octet and “A” electrons remain, put them on the central atom
Chapter 8

33. Lewis structures for: NO, BeCl2, BCl3, SF6, KrF2, PF5, (I3)-
Sample Problems # 63, 65
Chapter 8

34. 8.8: Strengths of Covalent Bonds
The energy required to break a covalent bond is called the bond dissociation enthalpy, D
Bond enthalpies can either be positive or negative
Chapter 8

35. Text, P. 329

36. Bond Enthalpies and the Enthalpies of Reactions
Bond enthalpies can be used to calculate the enthalpy for a chemical reaction
In any chemical reaction bonds need to be broken and then new bonds get formed
Chapter 8

37. Mathematically, if Hrxn is the enthalpy for a reaction, then
Sample Problem # 69
Chapter 8

38. Bond Enthalpy and Bond Length
Multiple bonds are shorter than single bonds
Multiple bonds are stronger than single bonds
As the number of bonds between atoms increases, the atoms are held closer and more tightly together (bond enthalpy increases)
Chapter 8

39. End of Chapter 8: Basic Concepts of Chemical Bonding