1. Basic Concepts of Chemical Bonding
Chapter 7
Basic Concepts of Chemical Bonding
Introduce people Me
Erica
David
Phil
David Lean
need to find out what Phil is doing (credentials)
Need to establish that we are A TEAM and that the Tutors are fantastic- so they don’t have problems when assessment arrives as a shock

2. Chemical Bonds Three basic types of bonds: Figure 7.1 Ionic Covalent
Electrostatic attraction between ions
Covalent
Sharing of electrons
Metallic
Metal atoms bonded to several other atoms
Figure 7.1

3. Ionic: Electrostatic attraction between ions
Mg+2O-2
K+2(Cr2O7)-2
The dichromate ion is a molecular ion involving 2 Chromiums stripped of their 6 outer shell electrons, surrounded by 7 oxygens. Collectively it holds 2 extra electrons or a charge of minus 2. This ion is relatively tightly bonded internally so it can behave as a single molecular anion. It is also a strong oxidizing agent and the Chromiums easily pick up some of their missing electrons for other things in solution or in the beaker like your fingers! Try figuring out which electrons are missing from the cations: K+, Mg+2 , Ni+2 and Cr+6. Write their electron configurations.
Ni+2O-2

4. Covalent: Shared pairs of electrons between adjacent atoms
C12H 22O 11 (s)
Br2 (l & g)
S8 (s)

5. Metallic: Shared conduction band electrons throughout crystal lattice
Mg0 (s)
Au20 (s)
Metals all look shiny and mirror like because the “sea” of conduction band electrons reflects weak photons of visible light. Most metals appear silvery like Mg or Ag. They are also conductive of heat and electrical current because electrons are free to travel across the lattice to equal energy states “next door”. Mg is a light metal with only 12 protons and 12 neutrons. Gold is very dense because the lattice consists of pairs of tightly bonded gold atoms, packed in a tight lattice of their own with a density of 19.3 g/cm3. As metals, they are electrically neutral like molecules of covalently bonded compounds.
Cu0 (s)

6. A neutral but active metal Na and a molecular gas Cl2 combine exothermically to give off light and heat and make the ionic compound NaCl. Na was oxidized and Cl was reduced.

7. Lewis Symbols
The Lewis symbol of an element consists of the chemical symbol for the element plus a dot for each valence electron.
Sulfur, for example, has the electron configuration [Ne]3s23p4. Therefore it has the following Lewis symbol: S
• • •
We only depict the (valence) electrons, of the unfilled outermost shell. The 12 inner core electrons (noble gas core) do not and never interact in chemical reactions, only the outer valence electrons are involved. The noble gas core electrons are too tightly held and too symmetric to break with the mere approach of protons in a neighboring atom.
which clearly depicts the six valence electrons.

8. Octet Rule
Atoms tend to gain, lose or share electrons until they are surrounded by eight valence electrons.
An octet of electrons consists of full s and p subshells in an atom.
Note: There are many exceptions to the octet rule,
but it remains useful for introducing many
important concepts of bonding.

9. Energetics of Ionic Bonding
As we have seen, it takes 495 kJ/mol to
remove the outer 3s1 electron from sodium.
We get 349 kJ/mol back by giving the 3p6
electron to chlorine.

10. Energetics of Ionic Bonding
These numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!
Figure 7.2

11. Energetics of Ionic Bonding
There must be a third piece to the puzzle.
What is as yet unaccounted for is the electrostatic attraction between the newly formed sodium cation and chloride anion.

12. Lattice Energy Q1Q2 d Eel = 
This third piece of the puzzle is the lattice energy:
The energy required to completely separate a mole of a solid ionic compound into its gaseous ions.
The energy associated with electrostatic interactions is governed by Coulomb’s law:
Lattice energy describes the attraction of neighboring ions in a lattice. In practice, the ion size changes depending on its environment. This is also called “Crystal Field Energy” or Madelung Constant for a particular compound and its structure.
Eel = 
Q1Q2 d

13. Lattice Energy Lattice energy increases with the charge on the ions.
It also increases with the decreasing size of ions.
Lattice energy increases with the proximity of an ion for its oppositely charged neighbors. If a cation could be held by the same anions but in 2 different coordinations 4 or 6, the 4-fold coordination would generally be stronger and have the higher lattice energy. The bigger the coordination number: 4, 6, 8, 12 etc. the weaker the lattice energy and easier it is to break ionic bonds, dissolve or melt a crystal.
Table 7.2

14. Energetics of Ionic Bonding
By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.
In practice it is easier to measure such things empirically (heat of melting, heat of dissolution etc.) than to figure out why they are that way!

15. Energetics of Ionic Bonding
These phenomena also help explain the “octet rule”.
Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.

16. Electron Configuration of Ions of the Representative Elements
Na 1s22s22p63s1 = [Ne]3s1
Na+ 1s22s22p63s1 = [Ne]
Cl 1s22s22p63s23p5 = [Ne] 3s23p5
Cl_ 1s22s22p63s23p6 = [Ne] 3s23p6 = [Ar]

17. Electron Configuration of Ions of the Representative Elements
In polyatomic ions, two or more atoms are bound together by predominantly covalent bonds and the whole acts as a charged species when the ion forms an ionic compound with an ion of opposite charge.
E.g. NH4+ and NH4Cl
CO32- and Na2CO3

18. Electron Configuration of Ions of the Representative Elements
In forming ions, transition metals lose the valence-shell s electrons first, then as many d electrons as are required to reach the charge of the ion.
For example, Fe has the electron configuration [Ar]3d64s2. To form the Fe2+ ion, the two 4s electrons are lost giving the [Ar]3d6 electron configuration.
To form the Fe3+ ion, the two 4s electrons and one d electron are lost giving the [Ar]3d5 electron configuration.

19. Covalent Bonding In these bonds, atoms share electrons.
There are several electrostatic interactions in these bonds:
Attractions between negative electrons and positive nuclei
Repulsions between -electrons
Repulsions between +nuclei
The approach of 2 atoms creates an energy minimum between them where electrons can be shared or held closer than in single atoms. Bond lengths and strengths are greater than electron energy levels in single atoms. The attraction of a neighboring positive nucleus promotes the electon to an excited state enabling it to participate in a covalent bond.
Figure 7.4

20. Lewis Structures
Lewis structures are representations of molecules showing all electrons - bonding and nonbonding.
--- denotes 1 pair of electrons.

21. Lewis Structures: Multiple Bonds
When two electron pairs are shared, two lines are drawn, i.e. double bond.
O :: C :: O or O = C = O
When three electron pairs are shared, three lines are drawn, i.e. triple bond.
Double and triple bonds are shorter and stronger and less capable of rotation than single bonds.
:N ::: N: or :N ≡ N:

23. Bond Polarity and Electronegativity
Nonpolar covalent bond
the electrons are shared equally between the two atoms, e.g. in Cl2 and N2
Polar covalent bond
one of the atoms exerts a greater attraction for the bonding electrons than the other, e.g. in H2O

24. Bond Polarity and Electronegativity
The ability of an atom in a molecule to attract electrons to itself.
On the periodic chart, electronegativity increases as you go:
from left to right across a row
from the bottom to the top of a column
Figure 7.5

25. Dipole Moments
When two atoms share electrons unequally, a bond dipole results.
The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:
 = Qr
It is measured in debyes (D).
Figure 7.7

26. Polar Covalent Bonds
Although atoms often form compounds by sharing electrons, the electrons are not always shared equally.
Figure 7.6
Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does.
Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

27. Polar Covalent Bonds
The greater the difference in electronegativity, the more polar is the bond.
Table 7.3
Figure 7.8

28. Explaining Polar Molecules: HCl
O.N.:
Cl = -1
H = +1
Protons (group number) – electons = oxidation number after assigning bonding electrons to the more electronegative element.

29. Drawing Lewis Structures
Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.
If it is an anion, add one electron for each negative charge.
If it is a cation, subtract one electron for each positive charge.
PCl3
(3 x 7) = 26

30. Drawing Lewis Structures
Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond, a dash (representing two electrons).

31. Drawing Lewis Structures
Complete the octets around all the atoms bonded to the central atom.

32. Drawing Lewis Structures
Place any leftover electrons on the central atom.
Remember we have 26 electrons

33. Drawing Lewis Structures
If there are not enough electrons to give the central atom an octet, try multiple bonds.

34. Drawing Lewis Structures - Formal Charge
Assign formal charges:
For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms.
Subtract that from the number of valence electrons for that atom: The difference is its formal charge.

35. Drawing Lewis Structures
The best Lewis structure is the one:
with the fewest charges
that puts a negative charge on the most electronegative atom.

36. Resonance Structures
This is the Lewis structure we would draw for ozone, O3.

37. Resonance Structures
This is at odds with the true, observed structure of ozone, in which:
both O—O bonds are the same length, and
both outer oxygens have a charge of ½.
Figure 7.10

38. Resonance Structures
One Lewis structure cannot accurately depict a molecule such as ozone.
We use multiple structures called resonance structures, to describe the molecule.

39. Resonance Structures
Just as green is a synthesis of blue and yellow, ozone is a synthesis of these two resonance structures.
Figure 7.11

40. Resonance Structures
In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon.
They are not localised, but rather are delocalised.

41. Exceptions to the Octet Rule
There are three types of ions or molecules that do not follow the octet rule:
Molecules and polyatomic ions containing an odd number of electrons.
Molecules and polyatomic ions in which an atom has fewer than an octet of valence electrons.
Molecules and polyatomic ions in which an atom has more than an octet of valence electrons.

42. Exceptions to the Octet Rule Odd Number of Electrons
Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons, e.g. NO contains = 11 valence electrons.
N = O .. .
and

43. Exceptions to the Octet Rule Less than an Octet of Valence Electrons
Consider BF3
Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine.
This would not be an accurate picture of the distribution of electrons in BF3.

44. Exceptions to the Octet Rule Less than an Octet of Valence Electrons
Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

45. Exceptions to the Octet Rule Less than an Octet of Valence Electrons
If filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.

46. Exceptions to the Octet Rule More than an Octet of Valence Electrons
The only way PCl5 can exist is if phosphorus has 10 electrons around it.
It is allowed to expand the octet of atoms on the 3rd row or below:
Presumably d orbitals in these atoms participate in bonding.

47. Exceptions to the Octet Rule More than an Octet of Valence Electrons
Even though we can draw a Lewis structure for the phosphate ion that has only eight electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

48. Exceptions to the Octet Rule More than an Octet of Valence Electrons
This eliminates the charge on the phosphorus and the charge on one of the oxygens, i.e. when the central atom is on the 3rd row or below and expanding its octet eliminates some formal charges, do so.

49. Strengths of Covalent Bonds
Most simply, the strength of a bond is measured by determining how much energy is required to break the bond.
This is the bond enthalpy.
The bond enthalpy for a Cl—Cl bond,
D(Cl—Cl), is measured to be 242 kJ/mol.

50. Average Bond Enthalpies
This table lists the average bond enthalpies for many different types of bonds.
Average bond enthalpies are positive, because bond breaking is an endothermic process.
Table 7.4
NOTE: These are average bond enthalpies, not absolute bond enthalpies; the C—H bonds in methane, CH4, will be a bit different than the C—H bond in chloroform, CHCl3.

51. Bond Enthalpies and Enthalpies of Reaction
Yet another way to estimate H for a reaction is to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed, i.e.
Hrxn = (bond enthalpies of bonds broken)
- (bond enthalpies of bonds formed)

52. Bond Enthalpies and Enthalpies of Reaction
CH4(g) Cl2(g)  CH3Cl(g) HCl(g)
In this example, one
C—H bond and one
Cl—Cl bond are broken
and one C—Cl and one
H—Cl bond are formed.
Figure 7.12

53. Bond Enthalpies and Enthalpies of Reaction
(bond enthalpies of bonds broken) - (bond enthalpies of bonds formed)
Hrxn = [D(C—H) + D(Cl—Cl)]  [D(C—Cl) + D(H—Cl)]
= [(413 kJ) + (242 kJ)]  [(328 kJ) + (431 kJ)]
= (655 kJ)  (759 kJ)
= 104 kJ

54. Bond Enthalpy and Bond Length
We can also measure an average bond length for different bond types.
As the number of bonds between two atoms increases, the bond length decreases.
Table 7.5

55. (1,2,3-Trinitroxypropane)
Nitroglycerin
(1,2,3-Trinitroxypropane) +
Diatomaceous Earth
Sodium Carbonate=
Dynamite