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A Lewis structure is a graphic representation of the electron distribution around atoms. The reason for learning to draw Lewis structures is to predict the number and type of bonds that may be formed around an atom. A Lewis structure also helps to make a prediction about the geometry of a molecule. Chemistry students are often confused by the models, but drawing Lewis structures can be a straightforward process if the proper steps are followed. Be aware there are several different strategies for constructing Lewis structures. These instructions outline the Kelter strategy to draw Lewis structures for molecules.
Step 1: Find the Total Number of Valence Electrons
In this step, add up the total number of valence electrons from all the atoms in the molecule.
Step 2: Find the Number of Electrons Needed to Make the Atoms "Happy"
An atom is considered "happy" if the atom's outer electron shell is filled. Elements up to period four on the periodic table need eight electrons to fill their outer electron shell. This property is often known as the "octet rule".
Step 3: Determine the number of bonds in the molecule
Covalent bonds are formed when one electron from each atom forms an electron pair. Step 2 tells how many electrons are needed and Step 1 is how many electrons you have. Subtracting the number in Step 1 from the number in Step 2 gives you the number of electrons needed to complete the octets. Each bond formed requires two electrons, so the number of bonds is half the number of electrons needed, or:
(Step 2 - Step 1)/2
Step 4: Choose a Central Atom
The central atom of a molecule is usually the least electronegative atom or the atom with the highest valence. To find electronegativity, either rely on periodic table trends or else consult a table that lists electronegativity values. Electronegativity decreases moving down a group on the periodic table and tends to increase moving from left to right across a period. Hydrogen and halogen atoms tend to appear on the outside of the molecule and are rarely the central atom.
Step 5: Draw a Skeletal Structure
Connect the atoms to the central atom with a straight line representing a bond between the two atoms. The central atom can have up to four other atoms connected to it.
Step 6: Place Electrons Around Outside Atoms
Complete the octets around each of the outer atoms. If there are not enough electrons to complete the octets, the skeletal structure from step 5 is incorrect. Try a different arrangement. Initially, this may require some trial an error. As you gain experience, it will become easier to predict skeletal structures.
Step 7: Place Remaining Electrons Around the Central Atom
Complete the octet for the central atom with the remaining electrons. If there are any bonds left over from Step 3, create double bonds with lone pairs on outside atoms. A double bond is represented by two solid lines drawn between a pair of atoms. If there are more than eight electrons on the central atom and the atom is not one of the exceptions to the octet rule, the number of valence atoms in Step 1 may have been counted incorrectly.
This will complete the Lewis dot structure for the molecule.
Lewis Structures vs Real Molecules
While Lewis structures are useful, especially when you're learning about valence, oxidation states, and bonding, there are many exceptions to the rules in the real world. Atoms seek to fill or half-fill their valence electron shell. However, atoms can and do form molecules that are not ideally stable. In some cases, the central atom can form more than other atoms connected to it. Also, the number of valence electrons can exceed 8, especially for higher atomic numbers. Lewis structures are helpful for light elements but less useful for transition metals, including lanthanides and actinides. Students are cautioned to remember Lewis structures are a valuable tool for learning about and predicting the behavior of atoms in molecules, but they are imperfect representations of real electron activity.