Notes
2.1 Molecular Structures
Sections
Molecular Structures
Here, we will practice drawing molecules in different styles.
Condensed structures present the ordering of atoms from molecular formulas.
Lewis structures explicitly show the bonds connecting the atoms.
Partially condensed structures are a mixture of condensed structures and Lewis structures.
What matters here is that we want our molecular structures to contain as many neutral atoms as possible.
We already have guidelines for spotting neutral atoms from:
Their valency (their expected number of bonds)
Their number of nonbonding electrons (recall our studies of formal charges).
Drill 1
For the first drill:
Draw a Lewis structure beginning from the condensed formula (CH3 CH2)2 CH CH2 CH2 OH.
In order to later get to the Lewis structure, first we'll go through a series of steps to correctly bond all of the atoms in a partially condensed structure.
Like reading a book, we first connect atoms with the leftmost groups: CH3 CH2.
This particular set of atoms is called an ethyl group.
The ordering of the atoms tells us that the first carbon has three hydrogens bonded to it, and it is connected to the second carbon with two bonded hydrogens.
The two ethyl groups in this molecule are attached to the same carbon.
This carbon has two more bonds: one to a hydrogen, and one to a CH2 group.
The carbon of this CH2 bridges to a second CH2.
From there, we connect to the oxygen of a hydroxyl group.
To begin drawing the Lewis structure, re-draw the atoms in the same orientation.
Connect them together with single bonds.
All the carbons are tetravalent with four single bonds, which makes them neutral.
All hydrogens are monovalent, as is predicted for a neutral hydrogen.
Oxygen is divalent, as we would predict. Unlike with carbon and hydrogen, we have to assign nonbonding electrons to ensure that the oxygen atom is neutral.
How many electrons should surround the oxygen?
Take the number of bonds on each oxygen, 2, and subtract it from the number of valence electrons, 6.
The resulting number, 4, is the number of nonbonding electrons.
In this way, we can account for the electrons in Lewis structures by adding double bonds, triple bonds, or lone pairs to adjust the electrons surrounding individual atoms.
Drill 2
For the second drill, draw a Lewis structure from the condensed formula CH3 CH2 CH2 COO CH3.
The condensed structure in this molecule begins with a carbon chain of CH3 CH2 CH2.
Add a bond to the next carbon.
This carbon is connected to two oxygens: one above, and one to the right.
The oxygen on the right is bonded to a methyl group, forming an alkoxy group.
At this point, all carbons have four bonds and all hydrogens have one bond.
Based on our valency predictions for oxygen, the oxygen that only has a single bond to carbon will not be neutral.
Therefore, add a double bond to this oxygen to form a carbonyl group.
Now that we have the arrangement of the atoms, we are ready to form the Lewis structure.
Explicitly draw all atoms and their connecting bonds.
We have already checked the valencies of the carbon, hydrogen, and oxygen atoms; however, we don't know if some have lone pairs.
Once again, we see that each oxygen should be surrounded by four unshared electrons to achieve a total electron count equal to its six valence electrons.
The most difficult part of this structure was deciding the placement and number of bonds to each oxygen. Just like the hydroxyl and methyl groups, the –COO- group is a distinctive set of atoms with a unique function.
In this case, the –COO CH3 is a carboxylic ester group.
Drill 3
For the last drill, begin with the condensed formula: H2N CH2 CH2 CH CH2
Here, the condensed structure begins with an amine group. The nitrogen is trivalent, with two bonds to hydrogens and one bond to carbon.
Next are two bridging methylene groups.
The next connection is to a carbon that is attached to a hydrogen and a terminal carbon.
Showing the bond between these last two carbons, we see that each is only bonded to three other atoms.
To achieve tetravalency, make the connection a double bond instead of a single bond. When two carbons are connected by a double bond, it is known as an alkene group.
All carbons now have four bonds and are neutral.
We can also see that the hydrogens, which are all monovalent, are neutral.
At this point, let's re-draw the atoms to form the basis for the Lewis structure.
Show all the single and double bonds.
There is only one atom that requires an electron count: the nitrogen.
A neutral nitrogen has five electrons.
Subtract the three bonds from the five valence electrons of nitrogen to determine the number of nonbonding electrons: 2.
Hence, draw one lone pair on the nitrogen.
Full-Length Text
Here, we will practice drawing molecules in different styles.
Write that condensed structures present the ordering of atoms from molecular formulas.
Write that Lewis structures explicitly show the bonds connecting the atoms.
Write that partially condensed structures are a mixture of condensed structures and Lewis structures.
What matters here is that we want our molecular structures to contain as many neutral atoms as possible.
We already have guidelines for spotting neutral atoms from:
Their valency (their expected number of bonds)
Their number of nonbonding electrons (recall our studies of formal charges).
For the first drill:
Draw a Lewis structure beginning from the condensed formula (CH3 CH2)2 CH CH2 CH2 OH.
In order to later get to the Lewis structure, first we'll go through a series of steps to correctly bond all of the atoms in a partially condensed structure.
Like reading a book, we first connect atoms with the leftmost groups: CH3 CH2. This particular set of atoms is called an ethyl group.
The ordering of the atoms tells us that the first carbon has three hydrogens bonded to it, and it is connected to the second carbon with two bonded hydrogens.
The two ethyl groups in this molecule are attached to the same carbon.
This carbon has two more bonds: one to a hydrogen, and one to a CH2 group.
The carbon of this CH2 bridges to a second CH2.
From there, we connect to the oxygen of a hydroxyl group.
To begin drawing the Lewis structure, re-draw the atoms in the same orientation.
Connect them together with single bonds.
All the carbons are tetravalent with four single bonds, which makes them neutral.
All hydrogens are monovalent, as is predicted for a neutral hydrogen.
Oxygen is divalent, as we would predict. Unlike with carbon and hydrogen, we have to assign nonbonding electrons to ensure that the oxygen atom is neutral.
How many electrons should surround the oxygen?
Take the number of bonds on each oxygen, 2, and subtract it from the number of valence electrons, 6.
The resulting number, 4, is the number of nonbonding electrons.
In this way, we can account for the electrons in Lewis structures by adding double bonds, triple bonds, or lone pairs to adjust the electrons surrounding individual atoms.
For the second drill, draw a Lewis structure from the condensed formula CH3 CH2 CH2 COO CH3.
The condensed structure in this molecule begins with a carbon chain of CH3 CH2 CH2.
Add a bond to the next carbon.
This carbon is connected to two oxygens: one above, and one to the right.
The oxygen on the right is bonded to a methyl group, forming an alkoxy group.
At this point, all carbons have four bonds and all hydrogens have one bond.
Based on our valency predictions for oxygen, the oxygen that only has a single bond to carbon will not be neutral.
Therefore, add a double bond to this oxygen to form a carbonyl group.
Now that we have the arrangement of the atoms, we are ready to form the Lewis structure.
Explicitly draw all atoms and their connecting bonds.
We have already checked the valencies of the carbon, hydrogen, and oxygen atoms; however, we don't know if some have lone pairs.
Once again, we see that each oxygen should be surrounded by four unshared electrons to achieve a total electron count equal to its six valence electrons.
The most difficult part of this structure was deciding the placement and number of bonds to each oxygen. Just like the hydroxyl and methyl groups, the –COO- group is a distinctive set of atoms with a unique function.
In this case, the –COO CH3 is a carboxylic ester group.
For the last drill, begin with the condensed formula: H2N CH2 CH2 CH CH2
Here, the condensed structure begins with an amine group. The nitrogen is trivalent, with two bonds to hydrogens and one bond to carbon.
Next are two bridging methylene groups.
The next connection is to a carbon that is attached to a hydrogen and a terminal carbon.
Showing the bond between these last two carbons, we see that each is only bonded to three other atoms.
To achieve tetravalency, make the connection a double bond instead of a single bond. When two carbons are connected by a double bond, it is known as an alkene group.
All carbons now have four bonds and are neutral.
We can also see that the hydrogens, which are all monovalent, are neutral.
At this point, let's re-draw the atoms to form the basis for the Lewis structure.
Show all the single and double bonds.
There is only one atom that requires an electron count: the nitrogen.
A neutral nitrogen has five electrons.
Subtract the three bonds from the five valence electrons of nitrogen to determine the number of nonbonding electrons: 2.
Hence, draw one lone pair on the nitrogen.