Lecture 41 Molecular Structures Ozgur Unal 1

 Molecular formula for compounds do not show how atoms are bonded together in a molecule.  In order to show the structure of a molecule we can use different representations:  Space-filling molecular model  Ball-and-stick molecular model  Structural formula  Lewis structure  Check out Figure 8.13  Structural formula and Lewis structure are very useful to show how atoms are bonded. 2

Steps to draw the Lewis structure of molecules: 1- Predict the location of certain atoms. 2- Determine the total number of electrons available for bonding. 3- Determine the number of electron pairs available for bonding. 4- Place the bonding pairs. 5- Determine the number of remaining electron pairs. 6- Determine whether each atom satisfies the octet rule. 3

 Example: Ammonia is a raw material used in the manufacture of many materials, including fertilizers, cleaning products, and explosives. Draw the Lewis structure for ammonia. 4  Example: A nitrogen trifluoride molecule contains numerous lone pairs. Draw its Lewis structure.

 Example: Carbon dioxide is a product of all cellular respiration. Draw the Lewis structure of carbon dioxide. 5  Example: Draw the Lewis structure of ethylene, C2H4.  Example: A molecule of carbon disulfide contains both lone pairs and multi-covalent bonds. Draw its Lewis structure.

 Although the unit acts as an ion, the atoms with polyatomic ion are covalently bonded.  Calculate the total number of valence electrons  Subtract the charge of the ion in order to find the total number of electrons available for bonding. 6  Example: PO valence electrons from P 4*6 = 24 valence electrons from O4 3 electrons from the charge  Total 32 electrons available for bonding.

 Example: Draw the Lewis structure for NH  Example: Draw the Lewis structure for Cl O4 -.  Example: Draw the Lewis structure for HC O3 -.

Lecture 42 Resonance Structures Ozgur Unal 8

 Using the same sequence of atoms, it is possible to have more than one correct Lewis structure when a molecule or polyatomic ion has both double bond and a single bond.  Example: NO3 -  Check out Figure 8.14  Resosnance is a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion. 9  Example: NO2 -, SO3 -2, CO3 -2, O3  Each molecule or ion that undergoes resonance behaves as if it has only one structure.  Experimentally measured bond lengths show that the bonds are identical to each other.  They are shorter than single bonds but longer than double bonds.

 Example: Draw the Lewis resonance structure for NO  Example: Draw the Lewis resonance structure for S O2.  Example: Draw the Lewis resonance structure for O3.

 Some molecules do not obey the octet rule.  There are several reasons for these exceptions:  Odd number of valence electrons  Suboctet and coordinate covalent bonds  Expanded octets 11  Odd number of valence electrons:  Example: NO2, ClO2 and NO

12  Suboctet and coordinate covalent bonds:  Suboctet  Stable configurations with fewer than 8 electrons present around an atom.  Example: BF3  A coordinate covalent bond forms when one atom donates both of the electrons to be shared with an atom or ion that needs two electrons.  Example: BF3 + NH3  Atoms or ions with lone pairs often form coordinate covalent bonds with atoms or ions that need two more electrons.

13  Expanded Octets:  These are molecules with central atoms that contain more than 8 valence electrons.  This electron arrangement is referred to as an expanded octet.  An expanded octet can be explained by considering the d orbital that occurs in the energy levels of elements in period three or higher.  Example: PCl5, SF6  Example: Draw the Lewis structure for Xenon tetrafluoride.  Example: Draw the Lewis structure for ClF3.