Unless otherwise stated, all images in this file have been reproduced from: Blackman, Bottle, Schmid, Mocerino and Wille, Chemistry, 2012 (John Wiley & Sons) ISBN: 9 78047081 0866 1 Lewis structures Lecture 15 Ionic bonds Ionic solids Lecture 16 Covalent bonds Lewis structures Lecture 17 VSEPR Polyatomic molecules Earlier we described bonding in diatomic molecules in terms of the molecular orbitals. HF H2O σ* For polyatomic molecules, the MOs become much more complicated and often involve many atoms. n We need a simpler model to describe bonding. σ 3 Lewis model of bonding • Simple model to allow us to determine the connectivity in and structure of molecules • Phrased as a set of rules or algorithm for structure-writing • A chemical bond is thought of as a pair of electrons shared between two atoms; atoms share bonds to achieve a full valence shell (8 e−) • Developed in 1916 • before the nuclear structure of the atom was understood • before the wave nature of the electron was discovered • before molecular orbital theory was developed • Still widely used by chemists to represent molecules and bonds • Extremely useful but incomplete model → need to use it judiciously and understand its limitations as well as its strengths 4 Lewis structures of simple molecules A simple strategy for drawing molecular structures: Step 1: Count the total number of valence electrons in the molecule. Step 2: Draw single bonds between the atoms. (Assume 2 e− per bond.) Step 3: Count how many electrons are left and assign them as lone pairs to fill valence shells (usually 8 e−). Octet rule. Start with outer atoms. Step 4: If any octets are incomplete, share lone pairs to form multiple bonds. 5 Lewis structures of simple molecules Lewis structure of hypochlorous acid (HOCl): O has two lone pairs. Cl, typical of halogens, has 3. In polyatomic molecules, the centre atom is the one that needs to share the greatest number of electrons. Lewis structure of Oxygen molecule (O2): O2 has double bond. Unfilled valence shell Bond order = 2 Filled valence shells Each O atom has two lone pairs Worksheet questions 1 and 2 6 Resonance structures and delocalisation of electrons Sometimes the Lewis model allows several valid structures, e.g. O3: This structure satisfies the Lewis rules, but is usually disregarded because three-membered rings are not very stable (but they are known). These resonance structures are equivalent representations of bonds that are “in-between” single and double bonds. We use a double headed arrow, not equilibrium arrows ( ) for resonance. In ozone, this means that both O-O bonds are equivalent, and we have an average bond order of 3/2 or 1½. Remember that electrons are really delocalised between more than two nuclei in molecular orbitals. Resonance is simply a way to describe this in terms of Lewis structures. 7 Lewis structures of simple molecules A simple strategy for drawing molecular structures: Step 1: Count the total number of valence electrons in the molecule. Step 2: Draw single bonds between the atoms. (Assume 2 e− per bond.) Step 3: Count how many electrons are left and assign them as lone pairs to fill the valence shells (usually 8 e−). Octet rule. Start with outer atoms. Step 4: If any octets are incomplete, share lone pairs to form multiple bonds. Step 5: If there is more than one plausible structure, then we consider both (all) to be resonance structures. Worksheet question 3 8 Formal charges (i) Formal charge is the charge assigned to an atom in a molecule, assuming that electrons are shared equally between atoms. → It is the difference between the number of valence electrons of each atom, and the actual number of electrons associated with it. For example, an oxygen atom has 6 valence electrons. In the ozone molecule: 2 lone pair electrons 6 shared electrons FC = 6−2−½(6) = +1 6 lone pair electrons 2 shared electrons FC = 6−6−½(2) = −1 4 lone pair electrons 4 shared electrons FC = 6−4−½(4) = 0 Formal Valence electrons Lone pairs ½ (shared = − − Charge of free atom of electrons electrons) Worksheet question 4 9 Formal charges (ii) Formal charge can be used to decide the most likely Lewis structure from a number of possibilities. 1 S: FC = 6−6−½(2) = −1 C: FC = 4−0−½(8) = 0 N: FC = 5−2−½(6) = 0 2 S: FC = 6−4−½(4) = 0 C: FC = 4−0−½(8) = 0 N: FC = 5−4−½(4) = −1 3 S: FC = 6−2−½(6) = +1 C: FC = 4−0−½(8) = 0 N: FC = 5−6−½(2) = −2 Step 6: Minimise the formal charges on all atoms. • Structure 3 has larger formal charges → unlikely • Structures 1 and 2 have the same formal charges (just in different places) → compare electronegativities • N is more electronegative than S, so structure 2 is more likely than structure 1 • But, both 1 and 2 are seen in nature 10 Lewis structures of simple molecules A simple strategy for drawing molecular structures: Step 1: Step 2: Step 3: Step 4: Step 5: Step 6: Count the total number of valence electrons in the molecule. Draw single bonds between the atoms. (Assume 2 e− per bond.) Count how many electrons are left and assign them as lone pairs to fill the valence shells (usually 8 e−). Octet rule. Start with outer atoms. If any octets are incomplete, share lone pairs to form multiple bonds. If there is more than one plausible structure, then we consider both (all) to be resonance structures. Worksheet question 5 Minimise the formal charges on all atoms. 11 Breaking the octet rule? (i) Atoms can have a maximum of 4 electron pairs (8 e−) surrounding them (H can only have 1 electron pair, 2 e−). But, B and Be may have fewer than 8 valence electrons, e.g. BF3. Elements below the second period may bind more than 4 other atoms. If you follow the rules for these molecules, it is looks like you have to break the octet: Group 15 (P, As, Sb) up to 5 atoms Group 16 (S, Se, Te) up to 6 atoms Group 17 (Cl, Br, I) up to 7 atoms 10 valence e− for P? 12 valence e− for S? 12 valence e− for Br? In reality, these molecules do obey the octet rule → the links here comprise < 2 e− (e.g. SF6 has only 4 bonding MOs to cover the 6 links → each BO = 2/3.) 12 Hypercoordinate molecules 1. Count the total number of valence electrons 2. Form bonds as usual, but make sure you don’t exceed the octet for any atom 3. Assign remaining electrons as lone pairs 4. If all octets are full, un-share bonding pairs to make space for left over electrons 5. This gives structures which appear to have electrostatic interactions between charged ions (partial ionic bonding) 6. But… all possible resonance structures must be considered 7. Overall, two bonds are shared over three linkages → on average, bond order is 2/3 7 + 7 3 = 28 Consistent with MO theory – only two true bonding MOs Worksheet Question 6 13 Breaking the octet rule? (ii) For SO2, you might come across either of the following structures: Analogous structure to O3 CO2-like structure O: 6 valence electrons 6 lone pair electrons 2 shared electrons formal charge = -1 O: 6 valence electrons 4 lone pair electrons O: 6 valence electrons S: 6 valence electrons 4 shared electrons 2 lone pair electrons 4 lone pair electrons formal charge = 0 8 shared electrons 4 shared electrons formal charge = 0 formal charge = 0 S: 6 valence electrons 2 lone pair electrons • Assumes breaking the octet is ok for S 6 shared electrons formal charge = +1 • Uses this to eliminate formal charges • This is a much better approximation to the MO description of this molecule • This is what you will see in most text books and in a lot of scientific writing Worksheet question 7 14 Summary of the basic principles of the Lewis model Lewis structures allow us to draw out plausible or reasonable structures for connectivity in molecules and ions and to get some insight into bond orders and electron distributions. Octet rule: atoms share electrons in order to have filled valence shells, i.e. adopt the electron configuration of the following noble gas. • A filled valence shell for H contains 2 electrons. • A filled valence shell for all other main block elements contains 8 electrons. • Elements in the 3rd, 4th etc. rows may have extra linkages, but they still obey the octet rule. • Elements at the beginning of a row may not achieve a full valence shell. Note: Bonding electrons are counted towards the valence shells of both bonding partners. 15 Summary of rules Step 1: Count the total number of valence electrons in the molecule. Step 2: Draw single bonds between atoms. Assume 2 e− per bond – do not exceed octets! Step 3: Count how many electrons are left and assign them as lone pairs to fill valence shells (8 e− except for H, Be, B). Start with outer atoms. Step 4a: If any octets are incomplete, share lone pairs → multiple bonds. Step 4b: If electrons are left over, un-share bonding pairs → lone pairs. Step 5: If there is more than one plausible structure, consider both (all) resonance structures. Step 6: Minimise the formal charges on all atoms. (But don’t exceed octets!) 16 Learning outcomes After this lecture you should be able to: • Draw out plausible Lewis structures for simple polyatomic molecules. • Assign bond orders based on sharing of electrons, resonance structures and formal charges. • Explain the relationship between resonance and electron delocalisation in molecular orbitals. Further Reading: Blackman 5.3 (but ignore everything they say about expansion of the octet) 17