1. Basic Concepts in Chemical Bonding

2. Introduction
Attractive forces that hold atoms together in compounds are called chemical bonds.
The electrons involved in bonding are usually those in the outermost (valence) shell.

3. Introduction Chemical bonds are classified into two types:
Ionic bonding results from electrostatic attractions among ions, which are formed by the transfer of one or more electrons from one atom to another.
Covalent bonding results from sharing one or more electron pairs between two atoms.

4. Lewis Dot Representations of Atoms
The Lewis dot representation (or Lewis dot formulas)
convenient bookkeeping method for valence electrons
electrons that are transferred or involved in chemical bonding

5. Lewis Dot Representations of Atoms

6. Lewis Dot Representations of Atoms
elements in same group have same Lewis dot structures

7. Lewis Dot Representations of Atoms
elements in same group have same Lewis dot structures

8. Ionic Bonding
We can also use Lewis formulas to represent the neutral atoms and the ions they form.

9. Ionic Bonding Li+ ions contain two electrons
same number as helium
F- ions contain ten electrons
same number as neon
Li+ ions are isoelectronic with helium
F- ions are isoelectronic with neon
Isoelectronic species contain the same number of electrons.

10. Ionic Bonding
potassium reacts with bromine
equation for the reaction

11. Ionic Bonding potassium reacts with bromine
2 K(s) + Br2 (l) ® 2 K+ Br-(s)
electronically this is occurring
4s p
K [Ar] ­
Br [Ar] ­¯ ­¯ ­
becomes

12. Ionic Bonding
4s p
K [Ar]
Br ­¯ ­¯ ­¯ ­¯ [Kr]
Lewis dot formulas

13. Ionic Bonding cations become isoelectronic with preceding noble gas
anions become isoelectronic with following noble gas

14. Covalent Bonding covalent bonds form when atoms share electrons
share 2 electrons - single covalent bond
share 4 electrons - double covalent bond
share 6 electrons - triple covalent bond
attraction is electrostatic in nature
lower potential energy when bound

15. Covalent Bonding

16. Covalent Bonding

17. Covalent Bonding Lewis dot representation H molecule formation
HCl molecule formation

18. Covalent Bonding extremes in bonding pure covalent bonds
electrons equally shared by the atoms
pure ionic bonds
electrons are completely lost or gained by one of the atoms
most compounds fall somewhere between these two extremes

19. Lewis Dot Formulas for Molecules and Polyatomic Ions
homonuclear diatomic molecules
hydrogen, H2
fluorine, F2
nitrogen, N2

20. Lewis Dot Formulas for Molecules and Polyatomic Ions
heteronuclear diatomic molecules
hydrogen halides
hydrogen fluoride, HF
hydrogen chloride, HCl
hydrogen bromide, HBr

21. Lewis Dot Formulas for Molecules and Polyatomic Ions
water, H2O

22. Lewis Dot Formulas for Molecules and Polyatomic Ions
ammonia molecule , NH3

23. Lewis Dot Formulas for Molecules and Polyatomic Ions
ammonium ion , NH4+

24. The Octet Rule
Elements try to a achieve noble gas configuration in most of their compounds.
Lewis dot formulas are based on the octet rule.

25. N - A = S rule N = # of electrons needed to be noble gas
usually 8
2 for H
A = # of electrons available in outer shells
equal to group #
8 for noble gases

26. N - A = S rule
for ions
add one e- for each negative charge
subtract one e- for each positive charge
central atom in a molecule or polyatomic ion is determined by:
atom that requires largest number of electrons
for two atoms in same group - less electronegative element is central

27. Drawing Lewis Dot Formulas
Write Lewis dot and dash formulas for hydrogen cyanide, HCN.
N = = 18
A = = 10
S =

28. Drawing Lewis Dot Formulas
Write Lewis dot and dash formulas for the sulfite ion, SO32-.
N = 8 + 3(8) = 32
A = 6 + 3(6) + 2 = 26
S = 6

29. Drawing Lewis Dot Formulas
Write Lewis dot and dash formulas for sulfur trioxide, SO3.
N = 8 + 3(8) = 32
A = 6 + 3(6) = 24
S =8

30. Resonance three possible structures for SO3
two or more dot structures necessary to show bonding
invoke resonance
Double-headed arrows are used to indicate resonance formulas.

31. Resonance Structures
Some molecules are not described by Lewis Structures very well.
Typically, structures with multiple bonds can have similar structures with the multiple bonds between different pairs of atoms.

32. Resonance flaw in our representations of molecules
no single or double bonds in SO3
all bonds are the same
best picture
incorrect shape

33. Limitations of the Octet Rule for Lewis Formulas
species that violate the octet rule
N - A = S rule does not apply
covalent compounds of Be, B, N
covalent compounds with expanded octets
species containing an odd number of electrons

34. Limitations of the Octet Rule for Lewis Formulas
In cases where the octet rule does not apply, the substituents attached to the central atom nearly always attain noble gas configurations.
The central atom does not.

35. Limitations of the Octet Rule for Lewis Formulas
Write dot and dash formulas for BBr3.

36. Limitations of the Octet Rule for Lewis Formulas
Write dot and dash formulas for AsF5.

37. Formal Charges
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, we use formal charge.

38. Formal Charges
The charge on an atom that it would have if all the atoms had the same electronegativity.
To calculate formal charge, electrons are assigned as follows:
All nonbonding electrons are assigned to the atom on which they are found.
Half the bonding electrons are assigned to each atom in a bond.

39. Formal Charges
valence electrons – (½ number of bonds + lone pair electrons)
Consider:

40. Formal Charges For C: There are 4 valence electrons.
In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure.
Formal charge: = -1

41. Formal Charges Consider: For N: We write:
There are 5 valence electrons.
In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure.
Formal charge = = 0
We write:

42. Formal Charges The most stable structure has:
the smallest formal charge on each atom,
the most negative formal charge on the most electronegative atoms.

43. Polar and Nonpolar Covalent Bonds
unequally shared electrons
assymmetrical charge distribution
different electronegativities

44. Electronegativity

45. Polar and Nonpolar Covalent Bonds

46. Polar and Nonpolar Covalent Bonds
Electron density map of HF
blue areas - low electron density
red areas - high electron density
polar molecules have separation of centers of negative and positive charge

47. Polar and Nonpolar Covalent Bonds

48. Polar and Nonpolar Covalent Bonds
Electron density map of HI
blue areas - low electron density
red areas - high electron density
notice that the charge separation is not as big as for HF
HI is only slightly polar

49. Polar and Nonpolar Covalent Bonds

50. Continuous Range of Bonding Types
all bonds have some ionic and some covalent character
HI is about 17% ionic
greater the electronegativity differences the more polar the bond

51. Continuous Range of Bonding Types
molecules whose centers of positive and negative charge do not coincide are polar
have a dipole moment
dipole moment has the symbol m
m is the product of the distance,d, separating charges of equal magnitude and opposite sign, and the magnitude of the charge, q

52. Dipole Moments
molecules that have a small separation of charge have small m
molecules that have a large separation of charge have a large m
example
HF and HI

53. Dipole Moments in molecules
some nonpolar molecules have polar bonds
2 conditions must hold for a molecule to be polar

54. Dipole Moments
There must be at least one polar bond present or one lone pair of electrons.
The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments do not cancel one another.

55. Polar and Nonpolar Covalent Bonds
electrons are shared equally
symmetrical charge distribution
must be the same element to share exactly equally H2 N2

56. Bond Energies
Bond energy - amount of energy required to break the bond and separate the atoms in the gas phase 61 61 61 61

57. Bond Energies Table of average bond energies
Molecule Bond Energy (kJ/mol) F Cl Br
O2 (double bond) 498
N2 (triple bond) 946 62 62 62 62

58. Bond Energies Bond energies can be calculated from other DHo298 values63 63 63 63

59. Bond Energies
Example: Calculate the bond energy for hydrogen fluoride, HF. 64 64 64 64

60. Bond Energies
Example: Calculate the bond energy for hydrogen fluoride, HF. 67 67 67 67

61. Bond Energies
Example: Calculate the average N-H bond energy in ammonia, NH3. 68 68 68 68

62. Bond Energies 69 69 69 69

63. Bond Energies
In gas phase reactions DHo values may be related to bond energies of all species in the reaction. 70 70 70 70

64. Bond Energies
Example: Use the bond energies to estimate the heat of reaction at 25oC for the reaction below. 71 71 71 71

65. Bond Energies
Example: Use the bond energies to estimate the heat of reaction at 25oC for the reaction below. 72 72 72 72