#### Related to [[Atomic bonding]] --- ## Definition >[!example] >#### Lewis dot structures are representations of molecules that are created by connecting atoms with **bonds** *(lines)* and **dots** *(electrons or lone pairs)* > >![[lewis structures.png]] > A related definition that is good to remember is **steric number**. The **steric number of an atom is the number of "things" attached to a central atom.** --- ## Steps - First identify the total number of [[Valence Electrons]] in the **molecule**. - Draw bonds between atoms to satisfy the [[Octet Rule]]. - Fill in the remaining [[Electrons]] by adding lone pairs first to the outside atoms starting from left to right and then working your way in/adding bonds as needed to satisfy the octet rule. > *Remember once bonded hydrogen can accept* ***NO MORE ELECTRONS.*** - Draw resonant structures >[!tip] >- Resonant structures are alternative versions of a given structures that **also satisfy the rules.** >- Many times it may just be a double bond switched from one atom to another. > ></br> > > ![[Lewis dot structures resonant structure example.png|800]] - Enclose diagram in brackets and indicate a charge *if there is one*. #### What to do with charges/ions? - Add or remove electrons from the structure and make sure that all the rules still work. In fact some structures only work with an additional electron or with one electron removed. --- ## Octet rule and exceptions - This rule states how many bonds a given atom **wants**, once it has those bonds we say it *satisfies the octet rule*. - One bond is represented by one line *(double bonds are represented by two and so on)*. Bonds are also represented by dots between two atoms where each bond is made up of two dots *or electrons*. - Remember there are many exception and weird applications of this rule. >[! Octet rule for common atoms] ><br> > >![[Lewis dot structures octet rule diagram.png | 800]] ## Exceptions to the octet rule - Like stated above there are many cases in which the octet rule will seemingly be broken for no apparent reason. However there are generally three reasons why a compound wont want to bond how you expect it to. - Starting on the 3rd row elements have a **high energy d orbital**. Electrons on that orbital can participate in bonding which is why elements like Xe can form *so many bonds*. - **Boron wants 3 bonds, phosphorous wants 5, sulfur wants 6 bonds, chlorine, fluorine, bromine and iodine all want one bond. Xe usually forms 4-6 bonds but this can vary. >[!cite] >#### For covalently bonded molecules ><br> > >![[Lewis dot structures 10.png |825]] >![[Lewis dot structures 11.png|825]] >![[Lewis dot structures 12.png|825]] --- ## Lewis dot structures for other applications - To show the formation of ions and such use the following equations. - These are applicable whether dots are used or not. >[! Ion formation with LDS] ><br> > >![[Lewis dot structures ex4.png | 700]] --- ## Formal charge - Individual atoms within a molecule may have a certain charge *either negative or positive*, generally this is determined by the amount of bonds and lone pairs they have. - If an atom has more valence electrons when bonded then it usually does *given by the periodic table* it becomes a negatively charged **anion**. If it has less valence electrons than usual it becomes a positively charged **cation**. - Each bond counts as **only one electron!**. - The magnitude of the charge increases **in accordance with the number of bonds** or lack thereof. >[!warning] >- In general the higher the formal charge in either direction the more unstable the molecule! >[!quote] ></br> > >![[Formal charge equation.png]] >![[Formal charge example 2.png|500]] --- #mainpage