1. Chemistry Scholarship Atomic Structure and Bonding
Lewis structures and shapes of molecules

2. Lewis structure of atoms
The chemical symbol for the atom is surrounded by a number of dots corresponding to the number of valence electrons.

3. A recommended procedure might be:
Count the total number of valence electrons (N) needed to account for the atoms (based on the column of the atom in the periodic table) and charge (add one electrons for each negative charge; subtract one electron for each positive charge).
Draw the framework with single bonds. Some knowledge of the way the atoms are connected may be required.
Using lone pairs, complete octets around the non-central atoms.
Count the number of electrons depicted (two for each bond and two for each lone pair). If this number is less than N, then add electrons to the central atom until the total number of electrons depicted is N.
If the octet rule is not satisfied for the central atom and lone-pair electrons are nearby, use those electrons to make double or triple bonds to the central atom.
Check each atom to see if it has a formal charge. (Singly bonded oxygen will require a negative charge, for example.)

4. Lewis Structures for Ions of Elements
The chemical symbol for the element is surrounded by the number of valence electrons present in the ion. The whole structure is then placed within square brackets, with a superscript to indicate the charge on the ion.
Atoms will gain or lose electrons in order to achieve a stable, Noble Gas (Group 18), electronic configuration.
Negative ions (anions) are formed when an atom gains electrons.
Positive ions (cations) are formed when an atom loses electrons.

5. Question 1
Draw a Lewis structure for F2CO in which the central C atom obeys the octet rule, and answer the questions based on your drawing.
1. The number of lone pairs on the central C atom: 0
2. The central C atom forms 2 single bonds.

6. Exceptions to the Octet Rule
Considering the tremendous variety in properties of elements and compounds in the periodic system, it is asking a great deal to expect a rule as simple as Lewis’ octet theory to be able to predict all formulas or to account for all molecular structures involving covalent bonds.
Lewis’ theory concentrates on resemblances to noble- gas ns2np6 valence octets.
Therefore it is most successful in accounting for formulas of compounds of the representative elements, whose distinguishing electrons are also s and p electrons.
The octet rule is much less useful in dealing with compounds of the transition elements most of which involve some participation of d or f orbitals in bonding.
Even among the representative elements there are some exceptions to the Lewis theory.

7. These fall mainly into three categories:
Some stable molecules simply do not have enough electrons to achieve octets around all atoms. This usually occurs in compounds containing Be or B.
Elements in the third period and below can accommodate more than an octet of electrons. Although elements such as Si, P, S, Cl, Br, and I obey the octet rule in many cases, under other circumstances they form more bonds than the rule allows.
Free Radicals

8. 1. Electron Deficient Species
Good examples of this type of exception are provided by BeCl2 and BCl3.
Beryllium dichloride, BeCl2, is a covalent rather than an ionic substance.
Solid BeCl2 has a relatively complex structure at room temperature, but when it is heated to 750°C, a vapor which consists of separate BeCl2 molecules is obtained.
Since Cl atoms do not readily form multiple bonds, we expect the Be atom to be joined to each Cl atom by a single bond. The structure is

9. Instead of an octet the valence shell of Be contains only two electron pairs.
Similar arguments can be applied to boron trichloride, BCl3, which is a stable gas at room temperature.
We are forced to write its
structure as 
in which the valence shell of
boron has only three pairs of
electrons.

10. Molecules such as BeCl2 and BCl3 are referred to as electron deficient because some atoms do not have complete octets.
Electron-deficient molecules typically react with species containing lone pairs, acquiring octets by formation of coordinate covalent bonds.
Thus BeCl2 reacts with Cl– ions to form BeCl4–

11. BCl3 reacts with NH3 in the following way:

12. 2. Expanding the octet rule
An atom like phosphorus or sulfur which has more than an octet is said to have expanded its valence shell.
This can only occur when the valence shell has enough orbitals to accommodate the extra electrons.
For example, in the case of phosphorus, the valence shell has a principal quantum number n = 3.
An octet would be 3s23p6.
However, the 3d subshell is also available, and some of the 3d orbitals may also be involved in bonding.
This permits the extra pair of electrons to occupy the valence shell of phosphorus in PF5.

13. Expansion of the valence shell is impossible for an atom in the second period because there is no such thing as a 2d orbital.
The valence (n = 2) shell of nitrogen, for example, consists of the 2s and 2p subshells only.
Thus nitrogen can form NF3 (in which nitrogen has an octet) but not NF5.
Phosphorus, on the other hand, forms both PF3 and PF5, the latter involving expansion of the valence shell to include part of the 3d subshell.

14. Species with Expanded Octets
Examples of molecules with more than an octet of electrons are phosphorus pentafluoride (PF5) and sulfur hexafluoride (SF6).
Phosphorus pentafluoride is a gas at room temperature.
It consists of PF5 molecules in which each fluorine atom is bonded to the phosphorus atom.
Since each bond corresponds to a shared pair of electrons, the Lewis structure is
Instead of an octet the phosphorus atom has 10 electrons in its valence shell.
Sulfur hexafluoride (also a gas) consists of SF6 molecules. Its structure is

16. 3. Free Radicals
The majority of molecules or complex ions discussed in general chemistry courses are demonstrated to have pairs of electrons.
However, there are a few stable molecules which contain an odd number of electrons.
These molecules, called "free radicals", contain at least one unpaired electron, a clear violation of the octet rule.
Free radicals play many important roles a wide range of applied chemistry fields, including biology, medicine, and astro-chemistry.

17. Three well-known examples of such molecules are nitrogen (II) oxide, nitrogen(IV) oxide, and chlorine dioxide.
The most plausible Lewis structures for these molecules are

18. Free radicals are usually more reactive than the average molecule in which all electrons are paired.
In particular they tend to combine with other molecules so that their unpaired electron finds a partner of opposite spin.
Since most molecules have all electrons paired, such reactions usually produce a new free radical.
This is one reason why automobile emissions which cause even small concentrations of NO and NO2 to be present in the air can be a serious pollution problem.

19. When one of these free radicals reacts with other automobile emissions, the problem does not go away.
Instead a different free radical is produced which is just as reactive as the one which was consumed.
To make matters worse, when sunlight interacts with NO2, it produces two free radicals for each one destroyed:

20. Expanded octets sometimes form resonance structures
Resonance structures are all the possible Lewis structures for a molecule
Formal charge is a technique to identify which resonance structure is the most stable.
The most stable Lewis structure will be where the formal charges are evenly distributed throughout the molecule.

21. Definition of formal charge
The formal charge of an atom in a molecule or ion is the charge it would have if all bonding electrons were shared equally.
To decide which structure is the most likely, assign ‘formal charge (FC)’ to each atom

22. When determining Lewis structure for a molecule, the structure is chosen in which the formal charge on each of the atoms is minimised.
The structure with the smallest separation of formal charges will be the most stable structure.
The sum of all the formal charges should equal the total charge of the molecule.

23. In calculating formal charge, we assume that all bonding electrons are shared equally between the atoms in the bond and therefore that each atom gets half of them

24. Example: Sulfur trioxide
The central sulfur atom has 8, 10 or 12 electrons and SO3 has a total of 24 valence electrons

27. In structure 1 The S atom: FC of S = 6 – 0 – 8/2 = +2
The single-bonded O atoms
FC of O = 6 – 6 – 2/2
= -1
The double-bonded O atom
FC of O = 6 – 4 – 4/2
= 0

28. In structure 2 The S atom: FC of S = 6 – 0 – 10/2 = +1
The single-bonded O atom
FC of O = 6 – 6 – 2/2
= -1
The double-bonded O atoms
FC of O = 6 – 4 – 4/2
= 0

29. In structure 3: The S atom: FC of S = 6 – 0 – 12/2 = 0
The double-bonded O atoms
FC of O = 6 – 4 – 4/2
Most stable resonance structure

30. Now study the Lewis diagram and resonance structures of the SO42- ion:

31. Consider CNS-
1. Find the Lewis Structure of the molecule. Remember the Lewis Structure rules.)
2. Resonance: All elements want an octet, and we can do that in multiple
ways by moving the terminal atom's electrons around (bonds too).

32. .
                                                                        
Assign Formal Charges
Formal Charge = (number of valence electrons in free orbital) - (number of lone-pair electrons) - (   number bond pair electrons)
Remember to determine the number of valence electron each atom has before assigning Formal Charges
C = 4 valence e-, N = 5 valence e-, S = 6 valence e-, also add an extra electron for the (-1) charge. The total of valence electrons is 16.

33. .)
                                                                                                                                                                                                                                                                                                                         

34. The most electronegative atom usually has the negative formal charge, while the least electronegative atom usually has the positive formal charges.

35. Practise!!
Further examples:

36. Shapes of Molecules
In any molecule, electron pairs mutually repel each other.
To reduce this repulsion, the molecule adopts a shape which allows electron pairs (bonding and non-bonding) to be as far as possible from each other (VSEPR theory).
The shape of the molecule depends on both the number of atoms linked to the central atom and the number of regions of electron density (bonded and non-bonded electron pairs) around the central atom.

37. VSEPR table
The bond angles in the table below are ideal angles from the simple VSEPR theory, followed by the actual angle for the example given in the following column where this differs.
For many cases, such as trigonal pyramidal and bent, the actual angle for the example differs from the ideal angle, but all examples differ by different amounts.
For example, the angle in H2S (92°) differs from the tetrahedral angle by much more than the angle for H2O (104.5°) does.

38. Bonding electron pairs
Lone pairs
Electron domains
Shape
Ideal bond angle (example's bond angle)
Example
Image 2
linear
180°
CO2 3
trigonal planar
120°
BF3 1
bent
120° (119°)
SO2 4
tetrahedral
109.5°
CH4
trigonal pyramidal
107°
NH3
angular
104.5°
H2O 5
trigonal bipyramidal
90°, 120°, 180°
PCl5
seesaw
180°, 120°, 90° (173.1°, 101.6°)
SF4
T-shaped
90°, 180° (87.5°, < 180°)
ClF3
XeF2 6
octahedral
90°, 180°
SF6
square pyramidal
90° (84.8°), 180°
BrF5
square planar
XeF4

39. Polarity: Polar and non-polar molecules
have at least one polar covalent bond and the bonds are arranged asymmetrically around the central atom.
the charge is not symmetrically distributed, the dipole moment of the polar bonds do not cancel out (thus have a net dipole on molecule)
Centre of negative charge do not coincide with centre of positive charge

40. Non-polar molecules:
have no polar bonds (e.g. O2); OR
have polar bonds which are arranged symmetrically so that the dipole moments of the polar bonds cancel out (e.g. CF4 has 4 equivalent polar C-F bonds in a symmetric tetrahedral shape)
Centre of negative charge and centre of positive charge coincide and are therefore cancelled

41. Properties “Like dissolve like”
polar molecular substances dissolve in polar solvents (e.g. ethanol dissolves in water)
non- polar substances dissolve in non-polar solvents (e.g. wax, a hydrocarbon, and solid iodine both dissolve in cyclohexane, but not in water).
MUCH MORE TO THIS – see you L3 notes