1. Chemical Bonding I: Basic Concepts
Chemistry
Second Edition
Julia Burdge
Lecture PowerPoints
Jason A. Kautz
University of Nebraska-Lincoln 8
Chemical Bonding I: Basic Concepts 1
Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. Chemical Bonding I: Basic Concepts8
8.1 Lewis Dot Symbols
8.2 Ionic Bonding
Lattice Energy
The Born-Haber Cycle
8.3 Covalent Bonding
Lewis Structures
Multiple Bonds
Comparison of Ionic and Covalent Compounds
8.4 Electronegativity and Polarity
Electronegativity
Dipole Moment and Partial Charges
8.5 Drawing Lewis Structures
8.6 Lewis Structures and Formal Charge
8.7 Resonance
8.8 Exceptions to the Octet Rule
Incomplete Octets
Odd Numbers of Electrons
Expanded Octets
8.9 Bond Enthalpy

3. Na: 1s22s22p63s1 Na+: 1s22s22p6 Cl: 1s22s22p63s23p5
Lewis Dot Symbols
8.1
Atoms combine in order to achieve a more stable electron configuration.
Maximum stability results when a chemical species is isoelectronic with a noble gas.
Na: 1s22s22p63s1
Na+: 1s22s22p6
10 electrons total, isoelectronic with Ne
Cl: 1s22s22p63s23p5
Cl‒: 1s22s22p63s23p6
18 electrons total, isoelectronic with Ar

4. B • Boron 1s22s22p1 B • B • B • Lewis Dot Symbols
When atoms form compounds, it is their valence electrons that actually interact.
A Lewis dot symbol consists of the element’s symbol surround by dots.
Each dot represents a valence electron. B •
Boron 1s22s22p1
3 valence electrons
Lewis dot symbol for boron B • B • B •
other reasonable Lewis dot symbols for boron

5. Lewis Dot Symbols
Lewis dot symbols of the main group elements.

6. 5 valence electrons; first pair formed in the Lewis dot symbol
Lewis Dot Symbols
Dots are not paired until absolutely necessary. B •
1s22s22p1 C •
1s22s22p2 N •• •
1s22s22p3
5 valence electrons; first pair formed in the Lewis dot symbol Na •
For main group metals such as Na, the number of dots is the number of electrons that are lost. O •• •
For nonmetals in the second period, the number of unpaired dots is the number of bonds the atom can form.

7. Na • Na 1s22s22p63s1 Na+ Na+ 1s22s22p6 O •• • O •• • O 1s22s22p4
Lewis Dot Symbols
Ions may also be represented by Lewis dot symbols.
Remember the charge Na •
Na 1s22s22p63s1
Na+
Na+ 1s22s22p6
Valence electron lost in the formation of the Na+ ion.
Core electrons not represented in the Lewis dot symbol O •• • 2‒ O •• •
O 1s22s22p4
O2‒ 1s22s22p6

8. K • K+ + e− + •• • I e− K • K+ + •• I 8.2 Ionic Bonding
Ionic bonding refers to the electrostatic attraction that holds oppositely charge ions together in an ionic compound. K • K+ + e− + •• • − I e− K • K+ + •• − I
The attraction between the cation and anion draws them together to form KI

9. NaCl(s)  Na+(g) + Cl−(g) Hlattice = +788 kJ/mol
Ionic Bonding
Lattice energy is the amount of energy required to convert a mole of ionic solid to its constituent ions in the gas phase. − +
NaCl(s)  Na+(g) + Cl−(g) Hlattice = +788 kJ/mol

10. Ionic Bonding
The magnitude of lattice energy is a measure of an ionic compound’s stability.
Lattice energy depends on the magnitudes of the charge and on the distance between them.
  d Q1 Q2
Q = amount of charge
d = distance of separation

11. Ionic Bonding
The magnitude of lattice energy is a measure of an ionic compound’s stability.
Lattice energy depends on the magnitudes of the charge and on the distance between them.

12. Ionic Bonding

13. Ionic Bonding
A Born-Haber cycle is a cycle that relates the lattice energy of an ionic compound to quantities that can be measured.

14. Ionic Bonding Calculate the lattice energy of rubidium iodide (RbI).
Solution:
Step 1: Combine the pertinent thermodynamic data.
Rb+(g) + e− + Cl(g)
EA1(Cl)
IE1(Na)
Rb+(g) + Cl−(g)
Rb(s) + I(g)
Energy
ΔHf°[I(g)]
Rb(g) + ½I2(g)
ΔHf°[Rb(g)]
Rb(s) + ½I2(g)
Lattice energy
of RbI
ΔHf°[RbI(s)]
RbI(s)

15. Ionic Bonding Calculate the lattice energy of rubidium iodide (RbI).
Solution:
Step 2: Use Hess’s law to calculate the lattice energy.
Lattice energy = ΔHf°[Rb(g)] + ΔHf°[I(g)] + IE1(Rb) + |ΔHf°[RbI(s)]| − EA(I)
ΔHf°[Rb(g)]
ΔHf°[I(g)]
ΔHf°[RbI(s)]
RbI(s)
Rb(s) + ½I2(g)
Rb(g) + ½I2(g)
Rb(s) + I(g)
Rb+(g) + e− + Cl(g)
Rb+(g) + Cl−(g)
IE1(Na)
Energy
Lattice energy
of RbI
EA1(Cl)

16. Lattice energy = 629 kJ/mol
Ionic Bonding
Calculate the lattice energy of rubidium iodide (RbI).
Solution:
Step 3: Using the data found in the textbook, calculate the lattice energy.
Lattice energy = ΔHf°[Rb(g)] + ΔHf°[I(g)] + IE1(Rb) + |ΔHf°[RbI(s)]| − EA(I)
Lattice energy = 85.8 kJ/mol kJ/mol kJ/mol kJ/mol − 295 kJ/mol
Lattice energy = 629 kJ/mol
Step
Energy (kJ/mol)
ΔHf°[Rb(g)]
85.8
ΔHf°[I(g)]
106.8
IE1(Rb)
403
ΔHf°[RbI(s)]
−328
EA(I)
295

17. H• •H H H • + 8.3 Covalent Bonding
The Lewis theory of bonding states that a chemical bond involves atoms sharing electrons.
When electrons are shared between atoms, a covalent bond results.
Each electron in the bond is attracted to both nuclei. H• •H +
H H •
Each H counts both electrons to “feel” as though it has a He configuration

18. • F• •F F F + Covalent Bonding F 1s22s22p5
According to the octet rule atoms will lose, gain, or share electrons in order to achieve a noble gas electron configuration.
Only valence electrons contribute to bonding. • F• •F +
F F
Each F counts both shared electrons to “feel” as though it has a Ne configuration.
F 1s22s22p5
valence electrons

19. Covalent Bonding
Pairs of valence electrons not involved in bonding are called lone pairs.
F F •

20. Shared electrons shown as dashes (bonds)
Covalent Bonding
A Lewis structure is a representation of covalent bonding.
Shared electron pairs are shown either as dashes or as pairs of dots.
Lone pairs are shown as pairs of dots on atoms.
H with 2 e−
H with 2 e− H O •• H O ••
Shared electrons shown as dashes (bonds)
O with 8 e−

21. •• O C H O •• O C •• = H O •• Covalent Bonding
In a single bond, atoms are held together by one electron pair.
In a double bond, atoms share two pairs of electrons.
C with 8 e−
O with 8 e−
O with 8 e− •• O C H O ••
2 shared pairs of electrons result in double bonds
one shared pair of electrons results in a single bond O C •• = H O ••

22. 3 shared pairs of electrons result in triple bonds
Covalent Bonding
A triple bond occurs when atoms are held together by three electron pairs.
each N has 8 e− •• N N •• ≡
3 shared pairs of electrons result in triple bonds

23. Covalent Bonding
Bond length is defined as the distance between the nuclei of two covalently bonded atoms.

24. N≡N 1.10 Å N=N 1.24 Å N−N 1.47 Å < <
Covalent Bonding
Multiple bonds are shorter than single bonds.
For a given pair of atoms:
~triple bonds are shorter than double bonds
~double bonds are shorter than single bonds
Multiple bonds are stronger than single bonds.
N≡N
1.10 Å
N=N
1.24 Å
N−N
1.47 Å
<
<

25. Covalent Bonding
There are two types of attractive forces in covalent molecules:
Intramolecular bonding force that holds the atoms together
Intermolecular forces between different molecules.
Intermolecular forces are weak
compared to intramolecular
forces.
Covalent compounds tend to
be gases, liquids, or low-melting
solids.

26. Covalent Bonding

27. Electronegativity and Polarity
8.4
There are two extremes in the spectrum of bonding:
covalent bonds occur between atoms that share electrons
ionic bonds occur between a metal and a nonmetal and involve ions
Bonds that fall between these extremes are polar.
In polar covalent bonds, electrons are shared but not shared equally.
M:X
Pure covalent bond
Neutral atoms held together by equally shared electrons
Mδ+Xδ−
Polar covalent bond
Partially charged atoms held together by unequally shared electrons
M+X−
Ionic bond
Oppositely charged ions held together by electrostatic attraction

28. Electronegativity and Polarity
Electron density maps show the distributions of charge.
Electrons are shared equally
nonpolar covalent
Electrons are not shared equally and are more likely to be associated with F;
polar covalent
Electrons are not shared but rather transferred from Na to F;
ionic

29. Electronegativity and Polarity
Electronegativity is the ability of an atom in a compound to draw electrons to itself.

30. Electronegativity and Polarity
Electronegativity is the ability of an atom in a compound to draw electrons to itself.

31. Electronegativity and Polarity
There is no sharp distinction between the different types of bonds:
nonpolar covalent and polar covalent
polar covalent and ionic
The following rules help distinguish among them:
A bond between atoms whose electronegativites differ by less than 0.5 is general considered purely covalent or nonpolar.
A bond between atoms who’s electronegativies differ by the range of 0.5 to 2.0 is generally considered polar covalent.
A bond between atoms whose electronegativities differ by 2.0 or more is generally considered ionic.

32. Electronegativity and Polarity
An arrow is used to indicate the direction of electron shift in polar covalent molecules.
Regions where electrons spend little time
Regions where electrons spend a lot of time H− F •• H− F •• δ+ δ−

33. Electronegativity and Polarity
A quantitative measure of the polarity of a bond is its dipole moment (μ).
Q is the charge
r is the distance between the charges
Expressed in debye units (D)
μ = Q x r
1 D = x 10−30 C•m

34. Electronegativity and Polarity
A quantitative measure of the polarity of a bond is its dipole moment (μ).
μ = Q x r

35. Drawing Lewis Structures
8.5
Follow these steps when drawing Lewis structure for molecules and polyatomic ions.
1) Draw the skeletal structure of the compound. The least
electronegative atom is usually the central atom.
2) Count the total number of valence electrons present; add electrons for
negative charges and subtract electrons for positive charges.
3) For each bond in the skeletal structure, subtract two electrons from
the total valence electrons.
4) Use the remaining electrons to complete octets of the terminal atoms.
Complete the octets of the most electronegative atom first.
5) Place any remaining electrons on the central atom.
6) Complete the octet of the central atom by forming multiple bonds.

36. Drawing Lewis Structures

37. Drawing Lewis Structures
Draw the Lewis Structure for ClO3−.
Solution:
Step 1 Draw the skeletal structure of the compound. The least
electronegative atom is usually the central atom.
chlorine’s electronegativity = 3.0
oxygen’s electronegativity = 3.5

38. Drawing Lewis Structures
Draw the Lewis Structure for ClO3−.
Solution:
Step 2 Count the total number of valence electrons present; add electrons for negative charges and subtract electrons for positive charges.
Cl: 1 x 7 = 7 e−
O: 3 x 6 = 18 e−
Charge: 1 e−
Total: 26 valence e−

39. Drawing Lewis Structures
Draw the Lewis Structure for ClO3−.
Solution:
Step 3 For each bond in the skeletal structure, subtract two electrons
from the total valence electrons.
3 bonds x 2 e− = 6 e−
26 – 6 = 20 e− left over

40. Drawing Lewis Structures
Draw the Lewis Structure for ClO3−.
Solution:
Step 4 Use the remaining electrons to complete octets of the terminal
atoms. Complete the octets of the most electronegative atom
first.
6 electrons used in single bonds
20 electrons must be placed in the structure starting with terminal atoms
The structure now has a total of 24 valence electrons of the available 26.

41. Drawing Lewis Structures
Draw the Lewis Structure for ClO3−.
Solution:
Step 5 Place any remaining electrons on the central atom.
The structure now has a total of 24 valence electrons of the available 26.
2 e− remain to be added to the structure.

42. Drawing Lewis Structures
Draw the Lewis Structure for ClO3−.
Solution:
Step 5 Complete the octet of the central atom by forming multiple bonds.
All atoms have an octet in this structure.

43. Lewis Structures and Formal Charge
8.6
Formal charge can be used to determine the most plausible Lewis Structure when more than one possibility exists for a compound.
To determine associated electrons:
1) All the atom’s nonbonding electrons are associated with the atom.
2) Half the atom’s bonding electrons are associated with the atom.
Formal charge = valence electrons – associated electrons

44. Lewis Structures and Formal Charge
Determine the formal charges on each oxygen atom in the ozone molecule (O3). O •• = −
4 unshared + 4 shared = 6 e− 2
2 unshared + 6 shared = 5 e− 2
6 unshared + 2 shared = 7 e− 2
Valence e− 6 6 6
e− associated with atom 6 5 7
Difference (formal charge) +1 −1

45. Lewis Structures and Formal Charge
When there is more than one possible structure, the best arrangement is determined by the following guidelines:
1) A Lewis structure in which all formal charges are zero is preferred.
2) Small formal charges are preferred to large formal charges.
3) Formal charges should be consistent with electronegativities. O C •• = O C •• ≡ −
Formal charge +1 −1
better structure (based on formal charge)

46. Lewis Structures and Formal Charge
Based on formal charge, identify the best and the worst structures for the isocyanate ion below:
Solution:
Step 1 Assign formal charges on each atom using the formula
Formal charge = valence electrons – associated electrons
Ve− 4 5 6 4 5 6 4 5 6
Ae− 6 4 6 5 4 7 7 4 5 FC −2 +1 −1 +1 −1 −3 +1 +1

47. Lewis Structures and Formal Charge
Based on formal charge, identify the best and the worst structures for the isocyanate ion below:
Solution:
Step 2 Determine the best and worst structure FC −2 +1 −1 +1 −1 −3 +1 +1
best structure:
small formal charges
Formal charges more consistent with electronegativities
worst structure:
Large formal charges
Formal charges inconsistent with electronegativities

48. Resonance
8.7
A resonance structure is one of two or more Lewis structures for a single molecule that cannot by represented accurately by only one Lewis structure.
Resonance structures are a human invention.
Resonance structures differ only in the positions of their electrons O •• = − O •• − =

49. Exceptions to the Octet Rule
8.8
Exceptions to the octet rule fall into three categories:
1) The central atom has fewer than eight electrons due to a shortage of
electrons.
2) The central atom has fewer than eight electrons due to an odd number of electrons.
only 4 total valence electrons in the system H Be O N •• = − •
17 valence electrons in the system

50. Exceptions to the Octet Rule
Exceptions to the octet rule fall into three categories:
3) The central atom has more than eight electrons.
Sulfur has 6 bonds corresponding to 12 electrons

51. Bond Enthalpy
8.9
The bond enthalpy is the enthalpy change associated with breaking a bond in 1 mole of gaseous molecule.
The enthalpy for a gas phase reaction is given by:
H2(g) → H(g) + H(g) ΔH° = kJ/mol
ΔH° = ΣBE(reactants) – ΣBE(products)
ΔH° = total energy input – total energy released
bonds broken
bonds formed

52. Bond Enthalpy
Bond enthalpy change in an exothermic reaction:

53. Bond Enthalpy
Bond enthalpy change in an endothermic reaction:

54. CH4(g) + Br2(g) → CH3Br(g) + HBr(g)
Bond Enthalpy
Calculate the enthalpy of reaction for the following:
CH4(g) + Br2(g) → CH3Br(g) + HBr(g)
Step 1: Draw the Lewis structures to determine what bonds should be broken and what bonds are to be formed.
ΔH° = ΣBE(reactants) – ΣBE(products)

55. Bond enthalpy (kJ/mol)
Step 2: Bonds to break: 4 C−H and 1 Br−Br
Bonds to form: 3 C−H, 1 C−Br, and 1 H−Br
Bond
Bond enthalpy (kJ/mol)
C−H
414
Br−Br
192.5
C−Br
276
H−Br
366.1
ΔH° = ΣBE(reactants) – ΣBE(products)
ΔH° = [4BE(C−H) + 1BE(Br−Br)] – [3BE(C−H) + 1BE(C−Br) + 1BE(H−Br)]
ΔH° = [4(414 kJ/mol) + 1(192.5 kJ/mol)] – [3(414 kJ/mol) + 1(276 kJ/mol) + 1(366.1 kJ/mol)]
ΔH° = -36 kJ/mol 55

56. Chapter Summary: Key Points8
Lewis Dot Symbols
Ionic Bonding
Lattice Energy
Born-Haber Cycle
Covalent Bonding
Octet Rule
Lewis Structures
Bond Order
Bond Polarity
Electronegativity
Dipole Moment
Drawing Lewis Structures
Formal Charge
Resonance Structures
Incomplete Octets
Odd Numbers of Electrons
Expanded Octets
Bond Enthalpy