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Lewis Structures: An Introduction
Cristian M.on 27 February 2013
Transcript of Lewis Structures: An Introduction
Bonds are represented by a line between the two atoms that are bonded, rather than the two electrons in the bond. Valence electrons are determined from the atom's electron configuration; they are in the orbitals with the largest principal quantum numbers. In order to draw a Lewis structure, one must follow the procedure listed below. 1. Determine the number of valence electrons in the molecule. This, for atoms that aren't transition metals, is determined by the group number of the atoms: i.e., oxygen has 6 valence electrons because it is in group 6A. The sum of the atoms' valence electrons determines how many electrons need to be diagrammed in the molecule. 3. Bond each atom in the molecule initially with a single bond. 2. Determine which atom is to be the central atom and place the atoms. Molecules are usually symmetrical. Carbon is given highest priority in being the center atom. 5. Check to see that the atoms satisfy the octet rule*. This is where all the atoms have 8 valence electrons*. 4. Put the remaining electrons around the atoms in pairs. 6. If all the atoms satisfy the octet rule, proceed to step 9. If not, proceed to step 7. If the atom is charged, factor in the charge with the electrons. If the charge is positive, subtract that many electrons from the final count. For instance, if the molecule has a 2+ charge, subtract 2 electrons. If it is negatively charged, add that many electrons to the final count. For instance, if the molecule has a charge of 2-, add 2 electrons. 7. Some atoms have all 8 valence electrons and its neighbor(s) have less than that. Here double bonds and triple bonds must be formed so that the atoms share more electrons. These are shown by 2 or 3 lines on top of each other like an equal sign (=). 8. Repeat step 6. Take CO2, for instance.
4 + 2 * 6 = 16 Carbon is in group 4A Oxygen is in group 6A, and there are 2 oxygen atoms here Carbon is in the center, and the molecule is symmetrical. O C O After single bonding and adding in the lone pairs, this is what the structure looks like so far: BUT WAIT, the carbon atom only has 4 electrons. Why not try two double bonds? And there it is, the Lewis structure for CO2. Yep, this means that this rule doesn't always hold. 9. If the molecule is charged, surround the structure with brackets and indicate its charge at the top right. If not, leave it as it is. Atoms that Don't Satisfy the Octet Rule Atoms that Don't Satisfy the Octet Rule 1. Hydrogen and Helium These atoms only have a 1s orbital that can only hold 2 electrons. Thus, they obey the DUET RULE: they only share 2 electrons and so can only form one bond (unless it's helium, which rarely bonds). 2. Beryllium, Boron, and Aluminum These atoms tend to be electron deficient in their initial forming of compounds. Beryllium is satisfied with 4 valence electrons, Boron and Aluminum with 6. 3. Atoms in Periods 3 to 6 These atoms have d-orbitals, which enable them to expand their valence orbitals. Such atoms (S, Cl, etc.) may have 10 or even 12 valence electrons in a molecule. The halogens (group 7A elements) and Hydrogen will always form only single bonds. If you encounter such an atom, satisfy the octet rule for all the atoms and add the extras to said atom (making sure it is one that has d-orbitals); it should be in the middle. The process for drawing molecules involving such atoms is still the same as it would be without them. Aluminum Trichloride Methane Hexafluorophosphate Disclaimer: this presentation does not cover molecules with an odd number of valence electrons or the topics of resonance and formal charge. See http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/quantum.html for information on these topics; these are prior knowledge required for this topic. Draw the Lewis structure for the cyanide ion. Note that it has 10 electrons and it is negatively charged. Draw the Lewis structure for BrF5. Note that the halogens only make double bonds, bromine is in period 4, and the fluorine atoms cannot expand their valence shells.