1. Basic concepts in Chemical Bonding
Chapter 8
Basic concepts in Chemical Bonding

2. Lecture 26. Basic Concepts in Chemical Bonding.
The properties of substances are determined largely by the chemical bonds that hold them together. The two extreme types of chemical bond are the ionic bond, such as holds the Na+ and Cl- ions together in NaCl, or the covalent bonds that hold molecules together such as glucose.
red =
O atoms
H-atom
cyan =
C atom
Cl- ion
Na+ ion
Sodium chloride – ionic substance
no actual NaCl molecules
glucose molecule – molecule
persists even in solution

3. Lewis dot symbols: Lewis Symbols. Lewis suggested Lewis
dot diagrams as a way
of tracking the electrons
involved in bond formation.
These are only the valence
electrons, as the core
electrons do not participate
directly in chemical bonding.
Gilbert Lewis ( )

4. Lewis dot structures for elements:
The Lewis symbol for each element consists of the symbol for the element surrounded by dots for each valence electron: . . . .
H Li Be B . . .
C N O F

5. Lewis Dot Structures
Lewis dot structures present a simple approach to bonding that allows us to rationalize much molecular structure. The idea is that atoms share electrons in the valence shell to form the chemical bond, with one pair of electrons per bond. Note that each H-atom now has two electrons, which is the structure of He, the next inert gas.
Electron pair = single bond
Valence electrons
H-atom H-atom H2 molecule
(Each H-atom has one valence electron)

6. Lewis Dot Structures (contd.):
Two shared pairs of electrons
= double bond
O-atom O-atom O2 molecule
Periodic table
Oxygen has six valence electrons

7. The octet rule
Electrons are shared in forming bonds such that atoms have the same number of electrons in their valence shells as the nearest noble gas, including the electrons shared with the atom to which they are bonded.
O-atom O-atom O2 molecule
Each oxygen atom in the O2 molecule now has eight
valence electrons, including those it shares with the
other oxygen atom = number of electrons (8 = octet)
in the nearest inert gas = neon.

8. Some more examples of Lewis dot structures:
The N2 molecule:
triple bond
N-atom N-atom N2 molecule
Periodic table

9. 8.1 Chemical Bonds, Lewis Symbols, and the Octet rule.
Chemical bonding involves mainly the attempt to achieve the rare gas number of valence electrons, i.e. an octet. This can be achieved in several ways.
Ionic bond: Electrons are mainly the property of one of the two atoms forming the bond.
Covalent bond: Electrons are shared so that each atom has a noble gas electronic configuration.
Metallic bonds. Electrons are lost into the conduction band.

10. 8.2 Ionic Bonding.
This occurs between metallic elements from the left-hand side of the periodic table and non-metallic elements from the right hand side of the periodic table.
Note that Na gives up its lone valence electron to Cl, so that they both end up with an octet of electrons.

11. Energetics of Ionic Bond Formation:
We have seen that Na has a low ionization energy, and Cl has a high electron affinity. In fact the process:
Na(g) + Cl(g) = Na+(g) + Cl-(g)
is endothermic = 496 – 349 = kJ/mole.
What really stabilizes NaCl is the crystal lattice, where there are strong attractive forces between the positively charged Na+ ions and the negatively charged Cl- ions. This is the lattice energy, which is the energy required to break the NaCl lattice up into gaseous ions.
NaCl(s) = Na+(g) + Cl-(g) ΔHlattice = +788 kJ/mol

12. 8.3 Covalent bonding.
Here the two atoms share the electrons to achieve a covalent bond.
two pairs of electrons equally shared
between the two oxygen atoms

13. Multiple Bonds and bond order:
The sharing of a single pair of electrons consititutes a single bond. Sharing of two pairs of electrons constitutes a double bond, and sharing three pairs of electrons constitutes a triple bond.
H:H :O::O: :N:::N:
Single bond double bond triple bond
Bond order: a single bond has bond order = 1, a double bond has bond order = 2, and a triple bond has a bond order = 3.
.. ..

14. Examples of Lewis dot diagrams:
Methane, CH4:
One shared
pair of electrons
= single bond
Carbon has four
valence electrons (red)
Hydrogens
achieve two
electrons like He
Carbon achieves
octet of electrons
single line
= single bond

15. Examples of Lewis dot diagrams:
Carbon dioxide: (CO2)
Carbon has four
valence electrons (red)
oxygens have six
valence electrons (black)
O=C=O
double line =
double bond
two shared
pairs of electrons
= double bond
Carbon and both oxygens
achieve an octet of electrons

16. Examples of Lewis dot diagrams:
Sulfur dioxide: (SO2)
single
bond?
double
bond?
O=S-O
(or O-S=O ?)
SO2 is an example where a
molecule can be written in
two ways and actual structure
is the average of the two. This
is called RESONANCE (see later)
actual structure is average
of the two (bond order = 1½) :

17. Note that as the bond order between two atoms goes up, the bond length gets shorter. Stronger bonds tend to be shorter:
increasing bond order:
single double triple
N-N N=N N≡N
1.47 Å Å Å
bonds get shorter
bonds get stronger

18. Table. Average bond lengths (Å) for some single and multiple bonds:
Bond Length Bond Length Bond Length
C-C C-O N-O 1.36
C=C C=O N=O 1.22
C≡C C≡O O-O 1.48
C-N N-N O=O 1.21
C=N N=N O-H 0.96
C≡N N≡N C-F 1.38
C-H N-H C-Cl 1.78
H-H C-Br C-I 2.14

19. remember that the atomic radii decrease along a period in the P.T.
Among the following examples, which
bond is shortest?
H-H S
C-H
Cl-Cl
C C
Bond length depends on (a) radii of the bonded atoms
remember that the atomic radii decrease along a period in the P.T.
(b) the number of bonds between atoms

20. 8.4 Bond Polarity and Electronegativity.
The concept of electronegativity was developed by Linus Pauling. Electronegativity is the ability of an element to attract electrons to itself in a molecule. Electronegativity increases across the periodic table and is at a maximum in the top right hand corner at fluorine, and is at a minimum at the bottom left hand corner at Cesium.
Linus Carl Pauling
( )

21. Electronegativities of the Elements

22. Electronegativities of some main group elements:H
2.1
Li Be B C N O F
Na Mg Al Si P S Cl
K Ca Ga Ge As Se Br
Rb Sr In Sn Sb Te I

23. Electronegativity (EN) and bond polarity:
The greater the difference in EN the more ionic the bond. For EN differences less than 0.5 we can say the bond is covalent, and for differences greater than 2.0 we can say the bond is ionic, for it is polar covalent:
Molecule: F HF LiF
______________________________________________________
EN diff: = = = 3.0
Type: non-polar polar ionic
covalent covalent

24. Electronegativity and bonding:
Some typical ranges for EN differences are:
EN difference bonding type Example EN difference
range
____________________________________________________________________________
> ionic LiF = 3.0
ionic NaCl = 2.1
polar covalent HF = 1.9
< covalent F-F = 0.0
covalent C-H = 0.4
covalent Li-Li = 0.0
covalent Au-C = 0.1
______________________________________________________________________________

25. Bond polarity: F2 HF LiF covalent covalent
With greater EN difference, the electron density is pulled
onto the more electronegative of the two atoms forming
the molecule:
non-polar polar ionic
covalent covalent
electron
density
equally
shared
electron
density
largely
on F
F HF LiF

26. S Among the following examples, which bond is most polar?
all equally non-polar S
C-F (C-H is non-polar)
P-Cl (furthest apart in
Periodic Table  largest DEN)
all equally non-polar

27. 8.6 Resonance structures: Ozone (O3)
bond order = 1½
double arrow
= resonance
O-O bonds = 2.78 Å
The ozone molecule can be written with two
equivalent Lewis dot structures. In such a situation the actual structure is the average of these two structures, with the two O-O bond lengths equal. O O O
The ozone molecule

28. Resonance structures – the nitrite anion: (NO2-)
In drawing up a Lewis dot diagram, if we are dealing with
an anion, we must put in an extra electron for each
negative charge on the anion:
negative charge
on anion
One extra electron
in Lewis dot
diagram because
of single negative
charge on anion
Bond order
= 1½
Two resonance structures average structure

29. The nitrate anion: average bond order (B.O.)= 2 + 1 + 1 = 1⅓ 3
to work out bond order,
pick the same bond in
each structure and
average the bond order
for that bond
Number of canonical structures

30. Resonance in benzene. There are two canonical structures
for benzene, which
means that the C to C
bonds have a bond
order of (2+1)/2 = 1.5.
The benzene ring has
a very high stability
due to this resonance,
which is called
aromaticity. or
Short-hand versions for the benzene ring

31. 8.7. Exceptions to the octet rule.
BF3. This can be written as F2B=F with three resonance structures. To complete its octet, BF3 readily reacts with e.g. H2O to form BF3.H2O. The actual structure of BF3 appears not to involve a double bond and does not obey the octet rule:
Best repre-
sentation of
BF3 with B
having only
6 electrons
in its valence
shell
Possible resonance
structure for BF3,
but is not important
as this would
involve the very
electronegative
F donating e’s to B

32. Exceptions to the octet rule: free radicals
There are some molecules that do not obey the octet rule because they have an odd number of electrons. Such molecules are very reactive, because they do not achieve an inert gas structure, and are known as free radicals. Examples of free radicals are chlorine dioxide, nitric oxide, nitrogen dioxide, and the superoxide radical:
odd electrons
nitric oxide chlorine dioxide

33. Exceptions to the Octet rule: Heavier atoms (P, As, S, Se, Cl, Br, I) may attain more than an octet of electrons:
Example: PF5.
In PF5, the P atom has ten electrons in its valence shell, which occurs commonly for heavier non-metal atoms:
leave off F
electrons not
shared with P F F P F F
P has
10 valence
electrons F
PF5

34. Many phosphorus compounds do obey the octet rule:
PF3 and [PO4]3- :
three blue electrons are
from charge on anion

35. Some compounds greatly exceed an octet of electrons:
IF XeF6
(both I and Xe have 14 valence e’s)
(Think about [XeF8]2-)

36. 8.8 Strengths of Covalent Bonds.
The strength of a covalent bond is measured by the energy required to break that bond into atoms:
Cl-Cl(g) → 2 Cl(g) ΔH = 242 kJ/mol
The strength of the C-H bond comes from:
CH4(g) → C(g) H(g) ΔH = 1660 kJ/mol
Divide by 4 = 415 kJ/mol ( a ‘best’ value is 413 kJ/mol)

37. Table of bond enthalpies (kJ/mol):
Single bonds:
C-H 413 C-F 485 N-H 391 F-F C-C 348 C-Cl 328 N-N 163 Cl-Cl 242
C-N 293 C-Br 276 O-H 463 Br-Br 193
C-O 358 C-I 240 O-O 146 I-I
H-H H-F 567 HCl 431 H-I 299
Multiple bonds:
C=C 614 C≡N 891 N=N 418 O=O 495
C≡C 839 C=O 799 N≡N 941 S=O 523
C=N 615 C≡O N=O 607 S=S 418

38. The shortest bonds tend to be the hardest to break in
Which of the following bonds do you think would be most difficult (require the largest energy) to break? S
H-H
C-F
Cl-Cl
C C
The shortest bonds tend to be the hardest to break in
a series of similar bonds. Bonds involving F also tend
to be very strong, because greater electronegativity
difference leads to stronger bonds.

39. Bond enthalpies and the enthalpy of reactions.
The enthalpy of a reaction can be calculated by application of Hess’s Law (ΔH for overall reaction is sum of H for the individual steps).
ΔHrxn = Σ(bond enthalpies of bonds broken) – Σ(bond enthalpies of bonds formed)

40. Example:
Consider the reaction between CH4 and Cl2 to give methyl chloride:
CH3-H Cl-Cl → CH3-Cl H-Cl
Bonds broken: C-H = 413 kJ/mol
Cl-Cl = 242 kJ/mol
Bonds formed: C-Cl = 328 kJ/mol
H-Cl = 431 kJ/mol
ΔHrxn = = -104 kJ/mol

41. Estimate DHo for the following reaction:
CH4 (g) O2 (g) → CO2 (g) H2O (g)
DHorxn = Σ n x Dbroken – Σ m x Dformed
DHorxn = [4 x x 498] – [2 x x 463] = kJ