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4 min readโขdecember 1, 2021
Saarah Hasan
Saarah Hasan
We draw Lewis Structures to get a visual representation of the molecules we examine, helping us determine ๐จ
Whether they are polar or nonpolar
The number and type of bonds they form, andย
The number of lone pairs of electrons in the molecule and where those lone pairs lie!
For example, here's the Lewis structure for water, aka H2O! ๐ง
Writing a Lewis structure for molecules of representative elements is based on the octet rule. Drawing them may get a little tricky. With some practice and understanding the directions to formatting Lewis Structures, weโll get the hang of it! ๐
Thereโs one main formula we need to keep in mind: S = N - A, where
S = the total number of electrons shared
S/2 = number of bonds
N = the sum of the number of electrons needed for all atoms involved to get a noble gas configuration. (N will be 2 for Hydrogen, 4 for Beryllium, 6 for Boron and Aluminum, and 8 for any other atom!)
A = the sum of the number of valence electrons available in each of the representative elements. Donโt forget to adjust this for ions!
U = the number of unshared electrons
U = A - Sย
U/2 = number of lone pairs
Choose a structure that is as symmetric as possible.
The least electronegative element is usually the central atom.
Use multiple bonds when needed as you place the number of shared electrons in the structure.
Place the unshared electrons as lone pairs to satisfy the octet rule for bonded atoms first! If you have any leftover, then place them on the central atom.
If there are more than one non-equivalent structure possible, use formal charges to choose between them.
Oxygens DO NOT bond to each other except in O2 and peroxides!
Formal charges are hypothetical charges on an atom in a molecule. We use them when we need to determine which Lewis structure is more likely. Here are our steps to do it:
Assign a formal charge for each atom.
Formal Charge = Group Number - (Number of Bonds + Number of Unshared Electrons)
Choose the better structure.
The formal charge on each atom is at or closer to 0.
The negative formal charges are on the more electronegative elements!
Resonance occurs when we have two or more Lewis structures that both accurately describe the bonding in a molecule. The true actual structure is a hybrid of the drawn structures.ย In organic chemistry, resonance comes up a lot more than in general chemistry. ๐ก
Radicals: Species that have an odd number of available electrons ๐ฒ
The weird structures of covalent compounds: Elements with an incomplete octet include Beryllium, Boron, and Aluminium ๐งฉ
Beryllium has 2 valence electrons, so it usually only forms 2 covalent bonds. We use 4 electrons as the number of needed electrons.
Boron and Aluminium have 3 valence electrons, so they usually form 3 covalent bonds. We use 6 electrons as the number of needed electrons.
Bigger valence shells: There are a lot of compounds where the central atom has to have more than 8 electrons to accommodate all bonded atoms. ๐
Letโs do some practice with Lewis Structures! ๐
Remember: S = N - A and U = A - Sย
N: (8) + (2 x 3) = 14
A: (1 x 3) + 5 = 8
S: 14 - 8 = 6
Number of bonds = S/2 = 6/2 = 3
U: 8 - 6 = 2
Number of lone pairs = U/2 = 2/2 = 1
Nitrogen is our central atom, and we draw 3 bonds connecting each hydrogen to it. Since our hydrogens are fulfilled with these bonds, we place one lone pair of electrons on the nitrogen. The one lone pair on our central atom lets us know that the molecule is polar. โ
N: (8) + (8 x 4) = 40
A: (4) + (7 x 4) = 32
S: 40 - 32 = 8
Number of bonds = S/2 = 8/2 = 4
U: 32 - 8 = 24
Number of lone pairs = U/2 = 24/2 = 12
Carbon is our central atom; we draw 4 bonds connecting each chloride to it. We add 3 lone pairs to each chloride, leaving us with no lone pairs to place on the carbon. โ
N: (2 x 2) + 8 = 12
A: (1 x 2) + (6) = 8
S: 12 - 8 = 4
Number of bonds = S/2 = 4/2 = 2
U: 8 - 4 = 4
Number of lone pairs = U/2 = 4/2 = 2
In this case, oxygen is our central atom; we draw one bond for each hydrogen atom to connect them. Since hydrogen is already fulfilled with these bonds, we add our two lone pairs to oxygen. โ
4 min readโขdecember 1, 2021
Saarah Hasan
Saarah Hasan
We draw Lewis Structures to get a visual representation of the molecules we examine, helping us determine ๐จ
Whether they are polar or nonpolar
The number and type of bonds they form, andย
The number of lone pairs of electrons in the molecule and where those lone pairs lie!
For example, here's the Lewis structure for water, aka H2O! ๐ง
Writing a Lewis structure for molecules of representative elements is based on the octet rule. Drawing them may get a little tricky. With some practice and understanding the directions to formatting Lewis Structures, weโll get the hang of it! ๐
Thereโs one main formula we need to keep in mind: S = N - A, where
S = the total number of electrons shared
S/2 = number of bonds
N = the sum of the number of electrons needed for all atoms involved to get a noble gas configuration. (N will be 2 for Hydrogen, 4 for Beryllium, 6 for Boron and Aluminum, and 8 for any other atom!)
A = the sum of the number of valence electrons available in each of the representative elements. Donโt forget to adjust this for ions!
U = the number of unshared electrons
U = A - Sย
U/2 = number of lone pairs
Choose a structure that is as symmetric as possible.
The least electronegative element is usually the central atom.
Use multiple bonds when needed as you place the number of shared electrons in the structure.
Place the unshared electrons as lone pairs to satisfy the octet rule for bonded atoms first! If you have any leftover, then place them on the central atom.
If there are more than one non-equivalent structure possible, use formal charges to choose between them.
Oxygens DO NOT bond to each other except in O2 and peroxides!
Formal charges are hypothetical charges on an atom in a molecule. We use them when we need to determine which Lewis structure is more likely. Here are our steps to do it:
Assign a formal charge for each atom.
Formal Charge = Group Number - (Number of Bonds + Number of Unshared Electrons)
Choose the better structure.
The formal charge on each atom is at or closer to 0.
The negative formal charges are on the more electronegative elements!
Resonance occurs when we have two or more Lewis structures that both accurately describe the bonding in a molecule. The true actual structure is a hybrid of the drawn structures.ย In organic chemistry, resonance comes up a lot more than in general chemistry. ๐ก
Radicals: Species that have an odd number of available electrons ๐ฒ
The weird structures of covalent compounds: Elements with an incomplete octet include Beryllium, Boron, and Aluminium ๐งฉ
Beryllium has 2 valence electrons, so it usually only forms 2 covalent bonds. We use 4 electrons as the number of needed electrons.
Boron and Aluminium have 3 valence electrons, so they usually form 3 covalent bonds. We use 6 electrons as the number of needed electrons.
Bigger valence shells: There are a lot of compounds where the central atom has to have more than 8 electrons to accommodate all bonded atoms. ๐
Letโs do some practice with Lewis Structures! ๐
Remember: S = N - A and U = A - Sย
N: (8) + (2 x 3) = 14
A: (1 x 3) + 5 = 8
S: 14 - 8 = 6
Number of bonds = S/2 = 6/2 = 3
U: 8 - 6 = 2
Number of lone pairs = U/2 = 2/2 = 1
Nitrogen is our central atom, and we draw 3 bonds connecting each hydrogen to it. Since our hydrogens are fulfilled with these bonds, we place one lone pair of electrons on the nitrogen. The one lone pair on our central atom lets us know that the molecule is polar. โ
N: (8) + (8 x 4) = 40
A: (4) + (7 x 4) = 32
S: 40 - 32 = 8
Number of bonds = S/2 = 8/2 = 4
U: 32 - 8 = 24
Number of lone pairs = U/2 = 24/2 = 12
Carbon is our central atom; we draw 4 bonds connecting each chloride to it. We add 3 lone pairs to each chloride, leaving us with no lone pairs to place on the carbon. โ
N: (2 x 2) + 8 = 12
A: (1 x 2) + (6) = 8
S: 12 - 8 = 4
Number of bonds = S/2 = 4/2 = 2
U: 8 - 4 = 4
Number of lone pairs = U/2 = 4/2 = 2
In this case, oxygen is our central atom; we draw one bond for each hydrogen atom to connect them. Since hydrogen is already fulfilled with these bonds, we add our two lone pairs to oxygen. โ
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