• Document: Chapter 9: Chemical Bonding I: Lewis Theory
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Chemistry 1A: Chapter 9 Page |1 Chapter 9: Chemical Bonding I: Lewis Theory Homework: Read Chapter 9: Work out sample/practice exercises. Check for the MasteringChemistry.com assignment and complete before due date Chemical Bonding: How atoms are connected together and the three dimensional shapes of molecules are very important. Many chemicals need to have the right shape to fit into the correct receptor or react the expected way. Finding the correct “fit” will allow manmade drugs to do a certain job. Artificial sweeteners have a shape that fits our receptors on the tongue to fool our brain into believing we taste something sweet. The puzzle picture is showing the AIDS drug Indinavar as the missing piece depicting the protein HIV-protease. Pharmaceutical companies designed molecules that would disable HIV-protease by sticking to the molecule’s active site – protease inhibitors. To design such a molecule, researchers used bonding theories to simulate the shape of potential drug molecules and how they would interact with the protease molecule Why do chemical bonds form?  Chemical bonds form because they lower the potential energy between the particles that compose atoms  A chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms  You need to consider the following interactions: 1. nucleus–to–nucleus repulsions 2. electron–to–electron repulsions 3. nucleus–to–electron attractions Lewis dot structures: Chemistry 1A: Chapter 9 Page |2 A UC Berkeley chemistry professor in 1916, Gilbert Lewis, developed a method to teach his beginning chemistry students how to understand chemical bonding that represents valence electrons with dots for main group elements. Lewis structures allow us to predict many properties of molecules such as molecular stability, shape, size, polarity Types of Bonding: Types of Atoms Types of Bond Bond Characteristics Metal to Nonmetal Ionic Electrons Transfer Nonmetal to Nonmetal Covalent Electrons Shared Metals to Metals Metallic Sea of Electrons Ionic bonds occur between a cation (metal) and an anion (nonmetal). Ions are held together by electrostatic attraction, opposite charges ( +, − ) attracting each other. These attractions are quite strong and increase with increasing charges and decreasing ionic sizes. Ionic compounds must overcome large lattice energies to separate to melt or boil. Ionic compounds have high melting points due to a vast three dimensional network of attractions between ions. Covalent bonds occur when electrons are shared between nonmetal atoms. The length of a bond increases as the bond order decreases (triple < double < single). The amount of energy to break a chemical covalent bond in an isolated gas molecule is called the bond dissociation energy (BE or D). The strength of the bond energy increases with increasing bond order (single < double < triple bonds). Approximate Bond Energy values (±10%) can be found in reference sources. Note: Bond energy values assume starting and ending with gas substances. One can use these energies to find approximate enthalpies of reaction, (Hreaction = BEreactants – BEproducts). Covalent compounds have low melting points, the strong molecular bonds do not break apart when melting or boiling. Covalent compounds are made of discrete molecules held together by weak intermolecular attractions. Metallic bonds: The low ionization energy of metals allows them to lose electrons easily. The simplest theory of metallic bonding involves the metal atoms Chemistry 1A: Chapter 9 Page |3 releasing valence electrons to be shared by all atoms/ions in the metal. Metal cations are surrounded by a sea of electrons. The electrons are delocalized throughout the metal structure. Bonding results from attraction of the cations for the delocalized electrons. • Metallic solids conduct electricity well. As temperature increases, the electrical conductivity of metals decreases • Metallic solids conduct heat well • Metallic solids reflect light • Metallic solids are malleable and ductile • Metals generally have high melting points and boiling points, all but Hg are solids at room temperature Melting points of metal generally increase left-to-right across period Na (97.72 ºC) < Mg

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