1. Bonding: General Concepts
Chapter 8
Bonding: General Concepts

2. 8.1 Types of Chemical Bonds
8.2 Electronegativity
8.3 Bond Polarity and Dipole Moments
8.4 Ions: Electron Configurations and Sizes
8.5 Energy Effects in Binary Ionic Compounds
8.6 Partial Ionic Character of Covalent Bonds
8.7 The Covalent Chemical Bond: A Model
8.8 Covalent Bond Energies and Chemical Reactions
8.9 The Localized Electron Bonding Model
8.10 Lewis Structures
8.11 Exceptions to the Octet Rule
8.12 Resonance
8.13 Molecular Structure: The VSEPR Model

3. A Chemical Bond
Forces that hold groups of atoms together and make them function as a unit.
A bond will form if the energy of the aggregate is lower than that of the separated atoms.
Types of Chemical Bonds:
Ionic Bonding – electrons are transferred
Covalent Bonding – electrons are shared equally
Intermediate cases
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4. Ionic Compound
Cation is always smaller than atom from which it is formed.
Anion is always larger than atom from which it is formed.

5. Bond energy:
It is the energy required to break the bond.
Bond length:
It is the equilibrium distance where the energy of the system is minimum
How does a bonding force develop between two identical atoms ?

6. The Interaction of Two Hydrogen Atoms
When H atoms are brought close together, there are two unfavorable potential energy terms, proton-proton repulsion & electron-electron repulsion , and one favorable term, proton-electron attraction.
A bond will form (that is, the two H atoms will exist as a molecular unit)if the system can lower its total energy in the process .

7. Energy profile as a function of the distance between the nuclei of the hydrogen atoms.
The zero point of energy is :” the atoms at infinite separation .
At very short distances , when the atoms are very close together
Bond length is the distance at which the system has minimum energy .
The potential energy of each electron is lowered because of the increased attractive forces in this area

8. Covalent Bond No electron transfer
In the hydrogen molecule and in many other molecules in which electrons are shared by nuclei is called
Covalent bonding.
Covalent Bond
No electron transfer
Electrons are shared between two atoms, positioned between the two nuclei
electrons are shared equally between identical atoms
Example: H2, O2, etc.

9. Polar Covalent Bond
Unequal sharing of electrons between atoms in a molecule.
Results in a charge separation in the bond (partial positive and partial negative charge).
When a sample of hydrogen fluoride HF gas is placed in an electric field , the molecules tend to orient themselves, with the fluoride end closest to the positive pole and the hydrogen end closest to the negative pole.
H +δ- F -δ

10. Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms. The molecule is called “Dipolar”. H F
e- poor
e- rich F H

11. The Effect of an Electric Field on Hydrogen Fluoride Molecules
When no electric field is present, the molecules are randomly oriented.
(b) When the field is turned on, the molecules tend to line up with their
negative ends toward the positive pole and their positive ends toward the negative pole.

12. Electronegativity Electronegativity
to determine how polar a bond will be, we introduce
Electronegativity
Electronegativity
the ability of an atom in a molecule to attract shared electrons to itself
I.e., …how much an atom “wants” electrons within a bond
determined by Linus Pauling ( )

13. The general trend is electronegativity increase as we go right and up in the periodic table

14. The Relationship Between Electronegativity and Bond Type
Electronegativity difference increase  polarity increases

15. Example 8.1: :
Arrange (Order)the following bonds in order of increasing polarity:
H-H, O-H, Cl-H, S-H, and F-H
Solution:
The polarity of the bond increases as the difference in the electronegativity increases.
H-H < S-H < Cl-H < O-H < F-H
(2.1) (2.1) (2.5) (2.1) (3.0) (2.1) (3.5) (2.1) (4.0) (2.1)
Electronegativity difference
Polarity increases
Covalent bond polar covalent bond

16. Bond Polarity and Dipole Moments
A molecule with a center of negative charge and a center of positive charge is said to be dipolar or has a dipole moment.
For example hydrogen fluoride
HF behaves in electric field as if it had two centers of charge, H positive and F negative.

17. Place a polar molecule in an electric field
the molecule will line up so that its “negative” end will line up with the positive pole and the “positive end” will line up with the negative pole
Dipole Moment
a polar molecule has a dipole moment
a polar molecule has a center of positive charge and a center of negative charge
Ex: H-F
The dipolar character of a molecule is represented by an arrow pointing to the negative charge center with the tail of the arrow indicating the positive center of charge.
Representation of dipole moment

18. Molecules with polar bonds and have net dipole moment:
Water molecule is polar molecule
The water molecule has a dipole moment .

19. The same type of behavior is observed for the NH3
The structure and charge distribution of the ammonia molecule
The dipole moment of the ammonia molecule oriented in an electric field
Polar molecule
It has dipole moment

20. Ex: A- Linear molecule: e.g: CO2 (O=C=O)
Some molecules have polar bonds, but are nonpolar
these molecules have no net dipole moment
due to the overall geometry of the molecule, the bond polarities cancel out, so the molecule has no net dipole moment
Ex:  A- Linear molecule: e.g: CO2 (O=C=O)

21. c. Tetrahedral molecules:
e.g: CCl4 , and CH4
Tetrahedral molecules with four identical bonds degrees apart

22. Example 8.2:
For each of the following molecules, show the direction of the bond polarities and indicate which ones have a dipole moment:
HCl
Cl2
CH4 ( tetrahedral molecule)
H2S (V-shaped molecule)
Answer:
HCl
The electronegativity of chlorine is greater than that of hydrogen (3.02.1). Thus the chlorine will be partially negative, and the hydrogen will be partially positive. The HCl molecule has a dipole moment.
Cl2
The two chlorine atoms share the electrons equally. No bond polarity occurs, and the Cl2 molecule has no dipole moment

23. CH4 ( tetrahedral molecule)
Carbon has a slightly higher electronegativity (2.5) than does hydrogen (2.1). this lead to small partial positive charges on the hydrogen atoms and small partial negative charge on the carbon. Bond polarities cancel. The molecule has no dipole moment.
H2S (V-shaped molecule)
The H2S molecule has a dipole moment.

24. Ions: Electron Configurations and sizes:
Atoms in stable compounds usually have a noble gas electron configuration.
In ionic compounds ;The nonmetals form anions, and the metals form cations to achieve a noble gas electron configuration ( Na+ - Cl-)
When two nonmetals react to form a covalent bond, they share electrons to complete the valence electron configuration of both atoms. That is, both nonmetals have noble gas electron configuration.

25. When metal react with nonmetal to form ionic compound, the valence electron configuration of the nonmetal achieves the electron configuration of the next noble gas atom and the valence orbital of the metal are emptied.
In this way both ions achieve noble gas electron configuration.

26. Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Atoms lose electrons so that cation has a noble-gas outer electron configuration.
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
H 1s1
H- 1s2 or [He]
Atoms gain electrons so that anion has a noble-gas outer electron configuration.
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]

27. Predicting formulas of ionic compounds
When we speak in this text of the stability of an ionic compound, we are referring to the solid state.
Group I forms ions
Group II forms ions
Group III forms ions
Group VI forms ions
Group VII forms -1 ions

28. Example of Ionic Compounds
MgO magnesium oxide is formed of Mg2+ and O2-
Mg :[Ne]4s Mg2+ :[Ne]
CaO formed from Ca2+ and O2-.
Ca :[Ar]4s Ca2+ :[Ar]
Al2O3 is formed of 2Al3+ and 3O2-.
Al:[Ne]3s23p1 Al3+ :[Ne]
O: 1s22s22p4
O2- 1s22s22p6 or [Ne]

29. Size of Ions
Positive ion (Cation)
formed by the loss of an electron from the parent atom
cation is smaller than the parent due to the loss of electron pair repulsion and the increased nuclear charge felt by each electron
Negative ion (Anion)
formed by the parent atom gaining an electron
anion is larger than the parent atom due to increased electron pair repulsion

30. Ions containing the same number of electrons.
Isoelectronic ions
Ions containing the same number of electrons.
The number of electrons and the number or protons are affect on the size of ions. But in case of isoelectronic ions, ions have the same number of electrons.
ions with the same number of electrons
Na+ , Mg+2, Al+3, F-, O-2, N-3
all contain 10 electrons
all have the same electron configuration as Ne
in terms of size, N-3>O-2>F->Na+>Mg+2>Al+3
the ion with the greater number of protons in an isoelectronic series will be the smallest due to the greater nuclear charge pulling the electrons in closer

31. 8O-2
10 electrons s2 2s2 2p6
9F-
11Na+
12Mg+2
13Al+3
The attraction force between the 10 electrons and the positive charge on the nucleus increase with increase the nuclear charge ‘Z’ increases. Therefore the size of ions decreases as the nuclear charge ‘Z’ increases. Also we can see of isoelectronic ions decreases with increasing atomic number
O-2> F- > Na+ > Mg+2 > Al+3 size of isoelectronic ions

32. Example :Arrange the ions Se-2, Br-, Rb+, and Sr+2 in order of decreasing size
Answer: This is an isoelectric series of ions with the krypton electron configuration. Since these ions all have the same number of electrons,
Z (ion) (no. of electrons)
Their sizes will depend on the nuclear charge .
Z (atom) 34 (Se2-) (Br-) (Rb+) (Sr2+) (no. of protons) Since the nuclear charge is greatest for (Sr2+) , it is the smallest of these ions.
The (Se2-) ion is largest. (Z increases and size decreases)
Ion Se2- > Br- > Rb+ > Sr2+

33. Example 8.4
Choose the largest ion in each of the following groups:
Li+, Na+, K+, Rb+, Cs+ (Group 1A)
Ba2+, Cs+, I-, Te2- (isoelectronic series)
Answer:
Since size increases down a group Cs+ is largest.
This is an isoelectronic series of ions with xenon electron configuration (size decreases with increasing Z):
Ion Te2- > I- > Cs+ > Ba2+
Z (atom) (no. of protons)
Z (ion) (no. of electrons)

34. The Localized Electron Bonding Model
The Covalent Chemical Bond: A model:
The Localized Electron Bonding Model
The model assumes that A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.
electron pairs are localized on a particular atom or in the space between two atoms.

35. Localized Electron Bonding Model
Lone pairs
electrons localized on a particular atom
Bonding pairs (or shared pairs)
electrons localized in the space between two atoms
On to Lewis structures!

36. Lewis Structure
The Lewis structure of a molecule Shows how valence electrons are arranged among atoms in a molecule.
Octet rule, atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas.

37. Hydrogen forms stable molecules where it shares 2 electrons
Hydrogen forms stable molecules where it shares 2 electrons. When 2 hydrogen atoms ,each with 1 electron , combine to form the H2 molecule .
H H
H : H
By sharing electrons , each H in H2 , in effect, has 2 electrons ,that is , each H has a filled valence shell .

38. In covalent bond formation, atoms go as far as possible toward completing their octets (duplets) by sharing electron pairs.
Each fluorine atom also has three pairs of electrons not involved in bonding. These are the lone pairs .

39. Steps for Writing Lewis Structures
Sum the valence electrons from all the atoms.
Use a pair of electrons to form a bond between each pair of bound atoms.
Atoms usually have noble gas configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen).
The central atom is usually written first in the formula [the central atom is usually the least electronegative atom, exceptions in H2O, and NH3 where O and N are the central atoms]
Hydrogen is never the central atom

40. Sum of valence electrons = 1 + 1 + 6 = 8
Examples:
Lewis structure of water
1- H2O (water): 1H and 8O
Sum of valence electrons = = 8
single covalent bonds
8e- H
2e-
2e- + O + H O H or
2- CO2 (carbon dioxide): 6C and 8O
Sum of valence electrons = = 16
Double bond – two atoms share two pairs of electrons
double bonds or
8e-
8e-
8e- O C O C
double bonds

41. 3- CN- (Cyanide ion): 6C and 7N
Sum of valence electrons = = 10
Triple bond – two atoms share three pairs of electrons
8e- c N
8e- or C N
triple bond
triple bond
4. NO+ (Nitrogen oxide ion) 7N and 8O
Sum of valence electrons = – 1 = 10
[NO]+

42. 5. N2 (Nitrogen) N and 7N
Sum of valence electrons = = 10
NN
CO2 O C ..

43. Exceptions to the Octet Rule
The Incomplete Octet (Be &B)
Some atoms are satisfied with less than an octet
Be (1s22s2) is stable with only four valence electrons
- Boron (1s2 2s2 2p1) also tends to form compounds
with less than eight electrons
Odd-Electron Molecules:
- Some molecules have an odd number of electrons
can't satisfy octet rule; usually N has the odd number

44. The Expanded Octet:
– Atoms in and beyond the 3rd period can
have more than eight electrons when in a
compound
– Thus, when drawing Lewis electron-dot
formulas, extra electrons go on the central
atom.
Example: SF6 [6 + 6 (7) = 48 valence electrons ]
I-3: 3(7) + 1 = 22 valence electrons
PCl5 : 5 + 5(7) = 40 electrons

45. Consider the Lewis structure for the nitrate ion NO-3:
Resonance: Blending of Structures
Consider the Lewis structure for the nitrate ion NO-3:
Based on Lewis structure it should be two types of N…O bonds
. Experiments show that NO-3 has only one type of N…O bond with length and strength between those expected for a single and double bonds.
. This can be explained on the basis of the resonance structures of NO-3:

46. All valid. We cannot find two bond lengths (hypothetical N-O vs N=O
Physical evidence shows that NO3- has three equivalent bonds
The correct description of NO3- is not one of the three Lewis structures, but an “average” of the three Lewis structures

47. resonance structure: one of two or more Lewis structures representing a single molecule that cannot be described fully with only one Lewis structure
Example: Given the Lewis structure for ozone, O3

48. Formal Charge
An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.
1. For neutral molecules, sum of formal charges must equal zero.
2. For ions, the sum of the formal charges must equal charge.

49. Consider the Lewis structure for POCl3
Consider the Lewis structure for POCl3. Assign the formal charge for each atom in the molecule. 
P: 5 – (4+0) = +1
O: 6 – (6+1) = –1
Cl: 7 – (6+1) = 0

50. VSEPR Model (Valence-Shell Electron-Pair Repulsion model)
This model is useful in predicting the geometries of molecules formed from nonmetals.
The structures of molecules play a very important role in determining their chemical properties .
The main postulate of this model is that “ The structure around a given atom is determined principally by minimizing electron-pair repulsions.”
The bonding and nonbonding pairs around a given atom will be positioned as far apart as possible.

51. Number of electron pairs
Example
Structure
Number of electron pairs
Linear 2
Trigonal planar 3
Tetrahedral 4
Trigonal bipyrimidal 5
Octahedral 6

52. Predicting a VSEPR Structure
1. Draw Lewis structure for the molecule.
2. Put the electron pairs as far apart as possible.
3. Determine positions of atoms fro the way electron
pairs are shared.
4. Determine the name of molecular structure from
positions of the atoms.

53. The Bond Angles In the CH4, NH3, and H2O Molecules

54. < bonding-pair vs. bonding pair repulsion lone-pair vs. lone pair
lone-pair vs. bonding
<

56. Example:
Predict the molecular structure and bond angles for the following molecule or ion:
HCN
PH3
CHCl3
NH4+
H2CO2
CO2

57. Answer:
a) HCN
Draw the Lewis structure
H-CN:
2- count the electron pairs
2 electron pairs as single bonds
2 effective electrons as triple bond
3-determine the position of atoms
4- determine the name of molecular structure
Linear , 180o

58. b) PH3 trigonal pyramid  109.5o
C)CHCl3 tetrahedral, 109.5o
d) NH tetrahedral, 109.5o
e) H2CO trigonal planar, 120 o
f)CO2 linear, 180 o