Dr. Saidane Chem. 152 Molecular Shapes Skills you should have mastered Conceptual 1. 2. 3. 4. Distinguish between single, double and triple bonds. Define a polyatomic ion. Define resonance. Explain the significance of electronegativity and what the ionic or covalent character of a bond means. 5. Explain how the polarity of a bond depends on the electronegativities of the two atoms in the bond. 6. Relate the properties of a molecule to its shape. 7. Explain the VSEPR theory. 8. Explain how the shapes of molecules affect their polarity. 9. Distinguish between polar bond and polar molecule. Problem-solving 1. 2. 3. 4. 5. Write Lewis structures for an ionic, metallic, and covalent compound. Determine the formula of an ionic compound using the Lewis structure and the crisscross method. Write Lewis structures and resonance forms of polyatomic ions and molecular compounds. Use a table of electronegativities to predict the polarity of bonds. Using the VSEPR theory and the Lewis structure, predict the shape, the bond angles, and the polarity of molecules with single bonds, multiple bonds and lone pairs. Lewis Structures for Covalent Compounds Nonmetals share electrons with one another until they have completed their octet (or duplet for hydrogen). A Lewis structure of a covalent compound shows the arrangements of electrons as shared pairs (or lines) and lone pairs (pairs of dots). To write the Lewis structure of a covalent compound we need to follow these steps: 1. Calculate the total number of valence electrons in the compound. Ex: CH4 has 8 valence electrons (4 electrons for C and one electron for each H). 2. Arrange the atoms in the compound. Choose the central atom and place the remaining atoms symmetrically around the central atom. Ex: SO2 is OSO not SOO. 3. Place one electron pair (or a line) between each pair of bonded atoms. Each bonded pair of atoms must have at least a single bond between them. 4. Complete the octet of each atom (duplet for H) by placing the remaining valence electron as electron pairs around the atoms. If there are not enough electron pairs form multiple bonds. Exception to the Octet Rule Some elements do not have an octet in their Lewis structures. Hydrogen and boron are the elements than can have a valence shell with less than eight electrons. H has two electrons (duplet). Boron tends to form bonds in which six electrons surround it, such as in BF3. Resonance Structures In some Lewis structures of covalent compounds with multiple bonds, the multiple bonds can be written in several equivalent locations. For each Lewis structure of the compound with a different location of the multiple bond we have a resonance structure and each one of them is called resonance hybrid. SHAPES OF MOLECULES AND IONS To a chemist, one of the most interesting aspects of a molecule is its shape or molecular geometry, which is the arrangement in space of the atoms bonded to each other. Lewis structures do not directly show the three-dimensional arrangements of atoms in a molecule, which play a key role in determining chemical reactivity. For example, C3H8 have different Lewis structures called isomers. Each isomer has different physical properties (smell, boiling point, etc.) and chemical properties (react differently when placed in the presence of the same compound). To account for these molecular shapes, we need just one addition to Lewis’s model: the VSEPR model. The VSEPR Model The VSEPR model is the Valence-Shell Electron-Pair Repulsion Model. This model is based on the fact that regions of high electron concentration repel one another. Bonding electrons and lone pairs take up positions as far from one another as possible, to minimize their repulsion and therefore maximize their separations. In this model, we focus on the central atom of a molecule and we imagine that all the electrons involved in bonds to this atom and the lone pairs belonging to the central atom lie on the surface of an invisible sphere that surrounds the central atom. These bonding electrons and lone pairs are regions of high electron concentration, and they repel one another. To minimize these repulsions, these regions move as far as possible on the surface of the sphere. We identify the shape of the molecule by noting where the atoms lie and look up the name of the shape. Multiple Bonds in the VSEPR Model When using the VSEPR model to predict molecular shapes, we do not need to distinguish between single and multiple bonds. A multiple bond is treated as another region of high electron concentration just as a single bond. Molecules With Lone Pairs on The Central Atom If a molecule has lone pairs on the central atom, they occupy a region around the central atom and therefore contribute to the shape of the molecule. However only the positions of atoms are considered when describing the shape of a molecule, and therefore lone pairs are ignored when we name the shape of the molecule. The Distorting Effect Of Lone Pairs According to the VSEPR model, lone pairs have a more strongly repelling effect than bonding pairs. In other words it is best for lone pairs to be as far as possible from each other as possible. It is also best for atoms to be far from lone pairs, even though that might bring them close to other atoms. The strengths of repulsions are in the order: lone pair-lone pair > lone pair-bonding air > bonding pair-bonding-pair. How To Use the VSEPR Model The general procedure for predicting the shape of a molecule is as follows: 1. Write the Lewis structure of the molecule. Treat a multiple bond equivalent of a single electron pair. Identify the central atom and assign the letter A to the central atom. 2. Identify the number of bonds (bonding domain) and lone pairs (nonbonding domain) around the central atom. Each lone pair will be assigned the letter E, and each bond will be assigned the letter X. 3. Classify the molecule as an AXnEm, n is the number of bonds and m the number of lone pairs around the central atom. Determine the shape according to the class. 4. Allow the molecule to distort so that lone pairs are as far as possible from one another and from bond pairs. The shape of the molecule is described in the following table: n m Shape 2 0 Linear 3 0 Trigonal Planar 2 1 Bent 4 0 Tetrahedral 3 1 Trigonal Pyramidal 2 2 Bent CHARGE DISTRIBUTION IN MOLECULES We need to know where electrons are most likely to be found in molecules if we want to understand a molecule’s chemical properties. Polar Bonds A polar covalent bond is a bond between two atoms that share unequally the bonding electrons arising from their difference in electronegativity. In this case an atom will have a partial positive charge and the other atom in the bond will have a partial negative charge. Partial charges give rise to electric dipole moment. The size of an electric dipole is a measure of the magnitude of the partial charges. Polar Molecules A polar molecule is a molecule with a non-zero dipole moment. A nonpolar molecule is a molecule that has zero electric dipole moment. The shape of a molecule governs whether or not it is polar. Highly symmetric arrangements of polar bonds result in molecules that are nonpolar, because in these molecules the bond dipoles cancel.