Chemistry 2810 Lecture Notes Dr. M. Gerken Page33 3.4 Lewis dot diagrams and VSEPR structures Review Lewis structures and VSEPR from General Chemistry texts, and consult S-A-L: 2.1 - 2.2 and 3.1 One of the basic distinctions you must learn to make is between ionic and covalent compounds. You will do much better in this course, as well as in all other chemistry courses, if you know instinctively whether the material being discussed is one or the other. So how can you learn this? Short of sheer memory work for millions of compounds, it is very possible to learn this intuitive knowledge simply by developing the habit of asking yourself: Is this compound covalent (i.e. a molecule) or ionic (i.e. composed of two or more ions)? Even if the answer is not obvious, it can usually be deduced from the information given. Often it becomes very obvious if you stop and think about it. We start by considering simple binary compounds, for which this distinction is simple. A compound A B is generally considered ionic if the difference in electronegativity between A and B is 2 units. Thus for H-F, = (3.9 2.2) = 1.7, and HF is considered to be a (polar) covalent molecule. But Li F, 6c = (3.9 1.0) = 2.9, and thus LiF is ionic. Note however that the ionic character of LiF is predominantly observed in the bulk solid - gaseous LiF (at very high temperature) will contain some Li-F molecules. We now focus on the structure and symmetry of the common covalent molecules, including common covalent or molecular ions (also known as complex ions), for which there are chemical bonds within the ionic unit. An example of the latter is an ion such as the sulfate ion, SO42-, which has covalent S-O bonds. 3.4.1. Valence and Lewis diagrams In Chem. 1000 you learned how to write Lewis structures. The number of valence You were wondering... electrons is taken directly off the periodic table, and can be deduced from the group numbers Why can we ignore previous directly. (Using the new numbering sequence, for p-block elements, subtract 10.) The shells when counting the number of valence electrons includes all s electrons since the last noble gas configuration number of valence electrons? plus the electrons of the block in which the element finds itself. Completely filled orbitals (except s orbitals) sink to much lower energy, becoming unavailable for bonding to elements in the subsequent block. Although Lewis diagrams are not 100% reliable, they have the advantage of organizing thousands of varied chemical compounds into fast, easily understood diagrams which give a lot of useful information about the structure and reactivity of the compound. The essential postulate of this theory, first postulated in 1916 and still used today, is that bonds between atoms are due to shared electron pairs. Unshared electrons form lone pairs. Multiple bonds form between elements short of electrons. Double bonds have four shared electrons, triple bonds six. To write Lewis structures, follow the step-by-step guidelines given in the text (S-A-L) on p. 51-52. 1. Decide how many electrons are to be included in the diagram by adding together all the valence electrons provided by the atoms. Adjust for the ionic charge, if present (add electrons for anions, substract electrons for cations). 2. Write the chemical symbols with the right connectivity (this cannot be deduced from the Lewis theory). 3. Distribute the electrons in pairs so that there is one pair of electrons between each pair of bonded atoms, and then supply electron pairs (to form multiple bonds or lone pairs on terminal atoms) until each atom has an octet. If you are not sure about multiple bonds distribute electrons as lone pairs (first on terminal atoms until an octet is reached, then on central atoms remember elements from row three and lower can have more than eight electrons). 4. Include formal charges in your Lewis structure. The formal charge gives some indication of the electron distribution in the molecule, where this is not even. For each atom, count the sum of the number of lone pair electrons and one from each bond-pair. The difference between this count and the valence of the atom is its formal charge. 5. Find the Lewis structure with the smallest number of formal charges. To do this, change electron lone pairs to bonding electron pairs (forming the double or triple bond). 6. Resonance has to be invoked whenever there is more than one way to distribute the electrons according to the above rules without unreasonably increasing the number of formal charges. The true structure is said to be a blend or hybrid of the various resonance isomers. 7. Finally, there are many elements for which &quot;exceptions to the octet rule&quot; occur. The &quot;octet rule&quot; is only valid for elements of the second row and is, therefore, a limited, historical rule. The &quot;exceptions to the octet rule&quot; include (a) Be (4), B and Al (6 in some cases), as well as (b) all the &quot;heavy&quot; elements of period three and beyond, which may have 10 or 12 valence electrons about them. My rule of thumb in all such cases is to start from the outside and provide octets for the ligands first. If there are deficient or excess electrons at the central atom, verify that the atom is one of the ones mentioned here, and leave the diagram as produced. Let's do some examples: CO2, NO3-, SO32-, NSF3, XeF4, IF5, PF5, SF6. Chemistry 2810 Lecture Notes Dr. M. Gerken Page34 3.4.2. VSEPR model Just as Lewis structures give us a fast road to mapping the electrons of molecules, the Valence Shell Electron Pair Repulsion VSEPR model gives us a quick approach to determining molecular structure for many common main-group compounds. It is not much use for transition metal complexes, except those of the metals in their highest possible oxidation states. This concept, which is especially due to Prof. Ronald Gillespie of McMaster University (along with Prof. Nyholm of the U.K.), considers the electron pairs in molecules to be bound regions of negative charge, which naturally repel each other. The physical basis of this model is actually the Pauli exclusion principle. The basic arrangements which minimize electron pair repulsions are: basic shape hybridization of the central atom # of pairs 2 linear sp 3 <a href="/keyword/trigonal-planar/" >trigonal planar</a> sp2 4 tetrahedral sp3 5 trigonal bipyramidal dsp3 6 octahedral d2sp3 But since the central atom may have lone pairs, which do not contribute to the description of the shape of the molecule, there are several derivatives of the above. Within the derivatives, the choice of structure is such as to minimize 90 interactions in the order: LP/LP repulsions stronger than LP/BP repulsions, than BP/BP repulsions. The logic behind this is that LP are less constrained than BP, therefore are larger. This also accounts for deviations in bond angle in structures such as water and ammonia. Hybridization can also be used to re-configure the atomic orbitals of the atoms in the molecule according to the observed geometry. Note that when angles deviate from the ideal values, the extent of hybridization also changes. thus while CH4 has four ...