Chemistry 445 Inorganic Chemistry Lecture 1. Lewis Dot Structures VSEPR G. N. Lewis was probably the best chemist who never won the Nobel Prize Gilbert Newton Lewis (1875-1946)

Lewis Dot Structures (revision) 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 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)

Lewis Dot Structures (contd.): Two shared pairs of electrons = double bond O-atom O-atom O2 molecule Periodic table 1 2 3 4 5 6 7 8 Oxygen has six valence electrons

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.

8.1 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.

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.

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

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. Fractional bond orders such as 1½ or 1⅓ are also possible, as discussed below. .. ..

Some more examples of Lewis dot structures: The N2 molecule: triple bond N-atom N-atom N2 molecule Periodic table 1 2 3 4 5 6 7 8

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

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

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) :

Slightly different Lewis dot representations: One can also represent molecules/ions with a combination of dots and lines for bonds, remembering that each line represents a shared pair of electrons, e.g. the phosphate anion:

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

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

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

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

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

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

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

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

Some compounds greatly exceed an octet of electrons: IF7 XeF6 (both I and Xe have 14 valence e’s) (Think about [XeF8]2-)