Lewis Structures (The Localized Electron Model)

Transcription

1 Lewis Structures (The Localized Electron Model) G. N. Lewis Using electron-dot symbols, G. N. Lewis developed the Localized Electron Model of chemical bonding (1916) in which valence electrons exist as lone pairs or as individual electrons seeking to form a pairing in order to achieve an octet. Later, Linus Pauling would expand the Localized Electron Model to include and Orbital Hybridization, Collectively known as Valence Bond Theory (1930). In 1957 VSEPR Theory was added to predict molecular geometry, also describing any resulting molecular polarity in molecules. Linus and Ava Helen Pauling in Munich, with Walter Heitler (left) and Fritz London (right) Localized Electron Model In Lewis s Localized Electron Model, molecules are described as being composed of atoms that are bound together by sharing pairs of electrons. He was able to show that the arrangement of atoms in molecules could be predicted based on the arrangements of valence electrons of all atoms involved in the molecule. Walter Heitler and Fritz London (1927) were the first to solidify Lewis s idea by linking atomic orbital overlap to Schrödinger s wave equation (1925) to show how two hydrogen atom wave functions join together to form a covalent bond. According to Lewis Theory, there are two types of valence electrons: Non-bonding (or unshared) pairs Bonding single (or unpaired) electrons Boron has three unpaired electrons therefore it can form three covalent bonds Bromine has three unshared pairs and one unpaired electron, therefore it can only form one covalent bond. What about nitrogen? We have seen how we can build models of molecules by combining atoms according to electron dot structures... : Br : + 3 = N Br Today, we are going to learn a process by which we will be able to draw a model of any molecule. : :.. :.. Br.. : 1

2 Lewis Structures Lewis structures are representations of molecules showing all electrons, bonding and nonbonding. PCl (7) = Find the sum of valence electrons of all atoms in the molecule from the group number or electron dot structure. Keep track of the electrons: 26 6 = Build a reasonable skeletal structure for the molecule using only single bonds. The central atom should be the least electronegative element that isn t hydrogen. Keep track of the electrons: 26 6 = 20 Things to consider when building primary skeleton: Oxygen never bonds to itself, except in O 2 and O 3 Carbon atoms are usually bonded to each other In molecules containing both H and O, hydrogen is usually bonded to oxygen Keep track of the electrons: 26 6 = = 2 3. Subtract the total number of electrons used in the primary bonds from the available valence electrons. 4. Fill the octets of the outer atoms by adding unshared pairs Keep track of the electrons: 26 6 = = 2 2 = 0 5. Fill the octet of the central atom. 6. Check to see that all atoms have and octet and that the correct number of valence electrons were used 2

3 7. If you run out of electrons before the central atom has an octet Example: Try building a Lewis structure for HCN 5. form multiple bonds until it does. 1. Let s try drawing the Lewis Structures for the following molecules: A. Carbon tetrachloride B. Ammonia C. Oxygen D. Carbon dioxide E. Dihydrogen carbon monoxide F. Ethanal (C 2 H 4 O) Polyatomic ions are formed from a class of molecules called Acids, or in some rare cases, from Bases. Polyatomic ions are formed as acids or bases loose or gain hydrogen atoms. Hydrogen nitrate For example: NO 3 - nitrate ion Hydrogen nitrate looses a hydrogen proton when placed in water, resulting in the formation of the nitrate ion (notice the 1- charge) 2. Let s draw the Lewis structure for dihydrogen sulfate and for the sulfate anion formed when dihydrogen sulfate is placed in water. Lewis structures for polyatomic ions must account for the loss or gain of valence electrons Cations decrease valence electrons by amount of charge Anions increase valence electrons by amount of charge Lewis structures for polyatomic ions are written in brackets [ ] with the charge denoted as a superscript. 3

4 3.Try drawing the Lewis structure for hydrogen nitrate and the nitrate ion. 4. Draw the Lewis Structure for ozone, O 3. You may notice more than one Lewis structure can be drawn for these species. Notice that two L.S. can be drawn correctly for ozone, O 3 RESONANCE theory, developed by Lewis (1928), is a key component of valence bond theory and arises when no single conventional model using only even number of electrons shared exclusively by two atoms can actually represent the observed molecule. involves modeling the structure of a molecule as an intermediate, or average, between several simpler but incorrect structures. One Lewis structure cannot accurately depict a molecule such as ozone. We use multiple structures, resonance structures, to describe the molecule. is denoted by a double headed arrow separating the different Lewis Structures: But this is at odds with the true, observed structure of ozone, in which both O O bonds are the same length. In truth, the electrons that form the second C O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. They are not localized, but rather are delocalized. 4

5 Just as green is a synthesis of blue and yellow Observe HCO 2- : ozone is a synthesis of these two resonance structures. In truth, the electrons that form the second C O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. They are not localized, but rather are delocalized. 5. Draw all three resonance structures for the nitrate ion. 5