The basic definition of the Lewis structure
The Lewis structure is one of the most common Chemistry topics that learners pursuing a chemistry course need to understand. It basically means the diagrams that are drawn to show the relationship/bonding that occur between various atoms of a single molecule. The term is also used to refer to bonds drawn to show the bonds between lone pairs of electrons in a molecule. For any covalently bonded molecule, a lewis structure can be drawn to show the bonding in this molecule. Also, a lewis diagram can be drawn for coordination compounds. Other terms that may be used to refer to these bonds apart from lewis structures include; lewis dot diagrams, lewis dot structures, lewis electron dot structures, and electron dot structures. Using any of these terms is acceptable in chemistry to refer to the lewis structures defined above. These structures were named after a man referred to as Gilbert N Lewis.
Who was Gilbert N Lewis?
Gilbert N Lewis, was a physical chemist in the United States of America and is popularly known for his contributions in the field of chemistry and, specifically, for his work in offering deep insights into the bonds that are now named after him, the lewis structures. He is famously known for his discovery of the covalent bonds in molecules and his electron pairs concepts. He is known to have contributed the most to the valence bond theory in chemistry and is fondly remembered in the Chemistry world.
Gilbert lived between the years 1875 and 1946. At the highlight of his accomplished career in the field of chemistry, he was the dean at the University of California heading the college of chemistry. He was born in Massachusetts, studied in the United States, Germany and in the Philippines. He obtained a PHD in chemistry from the renowned Harvard University. He moved to the United States in 1912 to teach chemistry. He did immense work in the field of chemistry and mentored many chemists that would later receive Nobel prizes in Chemistry. Gilbert, however, was nominated for the award 41 times but never worn the award, something that ekes of controversy to-date. Controversy aside, however, he did immense work in Chemistry and will remain in the folklores of the Chemistry world.
A More scientific definition of the lewis structure
A lewis structure is a simplified representation of the valence shell electrons in a molecule. The structure is used to show how electrons are arranged around atoms in any molecule. In the diagrams, the electrons are usually shown as dots. On the other hand, the bonding electrons are shown as lines between two atoms. The aim of drawing up the lewis diagrams is to ensure you get the most suiting/best/recommended electron configuration. To do this, some of the issues that you need to be wary of are the Octet Rule and ensuring that the formal charges have been satisfied.
Side note: While keeping the aforementioned definition in mind, it is important to remember that drawing up of lewis structures is not an effort to explain the geometry of molecules involved, it is not an attempt to explain how the bonds in the molecules are formed, and it is not trying to explain how the electrons in the molecules are shared. The theory is limited and provides a simple theory on the electron structure.
The Lewis Theory
The lewis theory is named after Gilbert N Lewis. The theory focuses on valence electrons and informs that the outermost electrons are the most exposed to other atoms and are the most likely to form bonds with electrons from other atoms. The theory also informs that the outermost are the most energetic and, since they are furthest from the atom nucleus, the most likely to form bonds with other atoms.
Lewis Dot Diagrams
These diagrams are a useful way of representing valence electrons and they help show the electrons that are available for sharing in covalent bonds. In the diagram, the lavelce electrons are written as dots. These dots surround the element’s symbol. When drawing up the diagrams, a single dot is placed on one side of the symbol letter(s) until all the four sides have been filled with a single dot. After that, a second dot can be added for each additional electron for the sides till a maximum of two dots are used for the four sides. Pairing of electrons is a major concept in the lewis theory.
Unpaired electrons in lewis structures
The lewis theory largely revolves around pairing of electrons. Based on the definition above, you already have an idea of how the pairing looks like in a lewis structure. However, it is on the unpaired electrons that bonding really occurs. The unpaired electrons are responsible for the bonding and reactions between chemical elements. At these unpaired electrons, the diagram shows where electrons can be gained or lost during ionic compounds processes. The sharing of these electrons leads to molecular compounds.
An exception to the rule?
For noble gases, the lewis theory gets into interesting territory. With the noble gases, no bonding is possible. All the valence electrons in noble gases are in filled shells. This makes them unavailable for bonding purposes hence the unreactive nature of the noble gases.
Covalent bond in lewis structures
What is a covalent bond?
This is a question you are most likely to run into while studying lewis structures and the lewis theory in general. The theory generally explains bonding that occurs between elements when valence electrons are shared. A covalent bond is a bond that happens when two non-metal elements combine and since each element is strong and cannot lose its singular valence electron, the elements end up sharing the valence electron in the molecular compound leading to a bond between the elements referred to as the covalent bond. These shared electrons will act as if they belong to either atom in the bond. Usually, a line is used to depict these electrons in a diagram. In the lewis structures, a line is used to show two electrons.
The Octet Rule in Lewis structures
As mentioned above, the octet rule is one of the things that are followed when drawing up lewis structures.SO what exactly is the octet rule?
In the octet rule, all atoms have eight electrons surrounding the atom. With the exception of hydrogen that has two electrons, the rest will show eight electrons in the lewis diagrams (with some exceptions that will be mentioned below).
Exceptions to the Octet rule
In most cases, the octet rule will apply when you are drawing up lewis structures. However, there are three exceptions when the octet rule will not apply. These exceptions include:
- BCI3 molecules: In these molecules, you will find that one or more atoms will have less than eight electrons; hence the octet rule will not apply in their cases.
- NO molecules: These molecules possess an odd number of electrons.
- Molecules containing more than eight electrons. A good example of such a molecule is the SF6.
How to Draw Lewis structures
After learning basics on the lewis structures, defining them and understanding what makes them tick, this section will teach you how to draw them. First though, you need to understand the rules to follow when drawing these structures. These include:
- Count the number of electrons (valence electrons) in the polyatomic ion or molecule.
- Use relativity of atoms when deciding where to place the atoms in the diagram. For instance, the least electronegative element is placed in the centre, or the atom with the higher period number is placed in the centre in case of atoms belonging to the same group. The outer atoms are referred to as terminal atoms.
- Draw a bond from one atom to the other (Specifically, from each outer atoms to the central atom). Be careful to use only two electrons in each bond.
- Distribute remaining electrons in pairs till each atom has eight electrons (with the exception of H that contains 2 electrons). When deciding on placement of lone pairs of electrons, place them first on the terminal atoms then lastly on the central atoms. Be careful to ensure that you have the same number of electrons as you had in the 1st
- This last step is to ensure that all atoms have the octet number. If an atoms fails to have an octet, you will have to move lone pairs of electrons from terminal atoms in between the terminal atom and the central atom. This results in double or triple bond placement. The formal charge Guideline is used to determine this placement.
Formal charge guideline in lewis structures
The formal charge guideline is a formula that is used beside the octet rule to determine the placement of bonds in lewis structures. The following formal charge formula is used in this case:
Formal charge= Valence – (1/2 bonding e-) – (lone pair e-)
- Formal charge: The charge an atom would have if electrons that bond were shared equally
- The sum of the formal charges must equal that on the species.
- Smaller formal charges are usually more stable compared to those of larger charges.
- Like charges on atoms that are adjacent on the structure should be avoided.
- Preferably, a more negative formal charge needs to be on a more electronegative atom.
Examples of popular lewis structures
In this section, we will learn how lewis structures for popular molecules are drawn. Some of these molecules include:
- Methane, CH4
- Ammonia, NH3
- Water, H20
- Hydronium Ion, H3O+
- Hydrogen Cyanide, HCN
- Carbon Dioxide, CO2
- Carbon Tetrachloride, CCI4
- Phosgene, COCI2
Methane, CH4 Lewis Structure
There are 8 valence electrons. Represented as (4 + 4×1)
Start by placing the C in the center, and then follow with connecting the four H’s to it:
All valence electrons used here
The octet rule satisfied
All of the atoms have zero formal charges.
Ammonia, NH3 Lewis Structure
8 valence electrons in this case
Represented by (5 + 3×1)
Place the N in the center, and connect the three H’s to it:
Six of the 8 valence electrons used.
You cannot use last two electrons here as it would violate the octet rule. They go to the central atom, N
All valence electrons used. Octet rule satisfied. All atoms have zero formal charge.
Water, H20 Lewis Structure
We have 8 valence electrons here
These are represented as (2×1 + 6).
You start by placing the O in the center and then connecting the two H’s to it:
Four electrons used here. The other four go to the O atom. They go to the O in two pairs.
All valence electrons used. Octet rule satisfied. All atoms have zero formal charge.
Hydronium Ion, H3O+ Lewis Structure
We have 8 valence electrons here.
They are represented as (3×1 + 6 – 1).
Start by placing the O in the center, and then proceed to connecting the three H’s to it:
6 valence electrons used here. The other two go to oxygen.
All valence electrons used. Octet rule satisfied.
The formal charge on this case for the oxygen atom is 1+ (8 – ½·6 – 2). This is represented as:
Hydrogen Cyanide, HCN Lewis Structure
We have 10 valence electrons here.
These are represented as (1 + 4 + 5).
Start by placing the C in the center, and then proceed by connecting the H and N to it:
Four valence electrons used here. The rest (six) are used on the N as shown below:
As seen above, Octet rule not satisfied on C. Also, a 2+ formal charge is seen on the C. the formal charge is represented as(4 – ½·4 – 0). On the N, the formal charge is 2- represented as (5 – ½·2 – 6). These are shown in the diagram below:
To correct the situation and satisfy octet rule, move two pairs of electrons from the N to make a triple bond in between the C and N as shown below:
The octet rule satisfied. Zero formal charges now.
Carbon Dioxide, CO2 Lewis Structure
We have 16 valence electrons in this case.
They are mathematically represented as (4 + 2×6).
Start by placing the C in the center, and then proceed by connecting the two O’s to it. Lastly, finish by placing the remaining valence electrons on the O’s as shown below:
16 valence electrons used in this case. Octet rule clearly not satisfied. Additionally, lots of formal charges in the structure as shown below:
However, we can satisfy the octet rule and the formal charges guideline by; moving a pair of electrons from each oxygen atom in between the carbon and oxygen atoms as shown:
Now we have: Satisfied the Octet rule; and we have zero formal charges.
Carbon Tetrachloride, CCI4 Lewis Structure
We have 32 valence electrons in this case.
They are mathematically represented as (4 + 4×7).
Start by placing the C in the center, and then proceed by connecting the four Cl’s to it as shown below:
Here we use 8 electrons. We remain with 24. These are paired and placed on the CI’s as shown:
By doing this we satisfy the Octet rule and use up all the valence electrons. In addition, all atoms have zero formal charges.
Phosgene, COCI2 Lewis Structure
We have 24 valence electrons in this case:
They are represented mathematically as (4 + 6 + 2×7).
We start by placing the C in the center, and then proceed by connecting the O and the two Cl’s to it.
The rest of the valence electrons are placed on the O and CI atoms as shown below:
As we can see, the Octet rule is not satisfied on the C. We need to move a pair of electrons from the O or we can also move one of the CI’s which will then result in a double bond. This will make a carbon-chlorine with a double bond that satisfied the octet rule but with formal charges. In addition, there would a positive charge on the CI atom. However, making a carbon-oxygen bond would result in zero formal charges and satisfaction of the octet rule. All these can be seen in the diagrams with the last structure representing both satisfaction of the octet rule and formal charges of zero.
Nitrogen Monoxide, NO lewis structures
Nitrogen monoxide, NO, is one of those that form part of the exception to the octet rule. To make a NO lewis structure, follow the following procedure:
In the case of NO, we have 11 valence electrons.
These are mathematically represented as (5 + 6).
As we can see from the structure, while the formal charges are zero, the octet rule is not satisfied on the N element. There is no way to satisfy the Octet rule with the Nitrogen Oxide, NO. This is a free radical and known to be extremely reactive in nature.
NOF lewis structure
This is a compound with three elements: Nitrogen, Oxygen and Fluorine. Nitrogen has 5 valence electrons, Oxygen has 6 valence electrons, and Fluorine has 7. The total number of valence electrons for this compound is 18. Nitrogen is the least electronegative compared to the other two and the Nitrogen atoms will hence to the centre.
Put Oxygen on one side and Fluorine on the other side.
In the first step, put two valence electrons between the elements. That will use up 4 valence electrons.
In the second step, put pairs of valence electrons around the outside and then the last two to the centre to use up all 18 electrons.
On checking, we have an octet for Fluorine and Oxygen have 8 each but nitrogen has 6. Move two electrons from oxygen to face the centre so that oxygen now shares two of its 8 with nitrogen.
On checking again, now all three elements have octets. The formal charges are also Zero and so both the octet rule and the Zero formal charge guidelines have been met.
Lewis structures Homework Help
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