Question
Draw the Lewis structure for CH3NCO and discuss its resonance forms and reactivity.
Solution
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To draw the Lewis structure for CH_3NCO, we need to follow a series of steps to ensure that we account for all the valence electrons and satisfy the octet rule for each atom where applicable. CH_3NCO is also known as methyl isocyanate, and it has a structure where a methyl group (CH_3) is attached to a nitrogen atom, which is double-bonded to a carbon atom, which in turn is double-bonded to an oxygen atom.
To draw the Lewis structure for CH_3NCO, we need to follow a series of steps to ensure that we account for all the valence electrons and satisfy the octet rule for each atom where applicable. CH_3NCO is also known as methyl isocyanate, and it has a structure where a methyl group (CH_3) is attached to a nitrogen atom, which is double-bonded to a carbon atom, which in turn is double-bonded to an oxygen atom.
Here are the steps to draw the Lewis structure for CH_3NCO:
1. Count the Total Number of Valence Electrons:
- Carbon (C) has 4 valence electrons.
- Hydrogen (H) has 1 valence electron each, but there are three hydrogens, so 3 valence electrons in total.
- Nitrogen (N) has 5 valence electrons.
- Oxygen (O) has 6 valence electrons.
- Adding these together, the total number of valence electrons is
2. Determine the Central Atom:
- Carbon is less electronegative than nitrogen and oxygen, so it will be the central atom to which the other atoms are attached.
3. Sketch the Skeleton Structure:
- Place the carbon atom in the center.
- Attach the three hydrogen atoms to the carbon atom.
- Attach the nitrogen atom to the carbon atom.
- Attach the oxygen atom to the carbon atom on the opposite side of the nitrogen.
4. Distribute the Valence Electrons:
- Start by placing a pair of electrons between each bonded atom to represent a chemical bond.
- Each hydrogen atom needs only 2 electrons to be stable, so after forming single bonds with carbon, no more electrons are placed around hydrogen.
- Carbon now has 3 single bonds, using 6 of its 4 valence electrons, leaving it with 2 electrons, which will form a double bond with either nitrogen or oxygen.
- Nitrogen has 5 valence electrons and needs 8 for a full octet, so it will form a double bond with carbon.
- Oxygen has 6 valence electrons and also needs 8 for a full octet, so it will form a double bond with carbon.
5. Complete the Octets:
- Place the remaining electrons around the nitrogen and oxygen to complete their octets.
6. Check the Formal Charges:
- Calculate the formal charge for each atom to ensure the most stable structure.
- The formal charge is calculated using the formula:
7. Draw Resonance Forms if Applicable:
- Since both nitrogen and oxygen can form double bonds with carbon, there are resonance structures where the double bond can be between carbon and nitrogen or carbon and oxygen.
The Lewis structure for CH_3NCO with the lowest formal charges and the most common resonance form is:
```
H
|
H-C-N=C=O
|
H
```
In this structure, the carbon atom is double-bonded to the oxygen atom and the nitrogen atom, forming a linear arrangement around the carbon atom. The nitrogen atom has a lone pair of electrons, and the oxygen atom has two lone pairs of electrons.
Resonance forms for CH_3NCO involve shifting the pi (π) electrons between the nitrogen and carbon atoms and between the carbon and oxygen atoms. However, the resonance form shown above is the most stable due to the electronegativity of oxygen, which makes it more favorable for oxygen to have a double bond with carbon than nitrogen.
Reactivity of CH_3NCO is influenced by the presence of the isocyanate group (NCO). The carbon in the isocyanate group is electrophilic, meaning it is susceptible to attack by nucleophiles. This reactivity can lead to various chemical reactions, such as nucleophilic addition or substitution, which are important in the synthesis of urethanes and other organic compounds.
Here are the steps to draw the Lewis structure for CH_3NCO:
1. Count the Total Number of Valence Electrons:
- Carbon (C) has 4 valence electrons.
- Hydrogen (H) has 1 valence electron each, but there are three hydrogens, so 3 valence electrons in total.
- Nitrogen (N) has 5 valence electrons.
- Oxygen (O) has 6 valence electrons.
- Adding these together, the total number of valence electrons is
2. Determine the Central Atom:
- Carbon is less electronegative than nitrogen and oxygen, so it will be the central atom to which the other atoms are attached.
3. Sketch the Skeleton Structure:
- Place the carbon atom in the center.
- Attach the three hydrogen atoms to the carbon atom.
- Attach the nitrogen atom to the carbon atom.
- Attach the oxygen atom to the carbon atom on the opposite side of the nitrogen.
4. Distribute the Valence Electrons:
- Start by placing a pair of electrons between each bonded atom to represent a chemical bond.
- Each hydrogen atom needs only 2 electrons to be stable, so after forming single bonds with carbon, no more electrons are placed around hydrogen.
- Carbon now has 3 single bonds, using 6 of its 4 valence electrons, leaving it with 2 electrons, which will form a double bond with either nitrogen or oxygen.
- Nitrogen has 5 valence electrons and needs 8 for a full octet, so it will form a double bond with carbon.
- Oxygen has 6 valence electrons and also needs 8 for a full octet, so it will form a double bond with carbon.
5. Complete the Octets:
- Place the remaining electrons around the nitrogen and oxygen to complete their octets.
6. Check the Formal Charges:
- Calculate the formal charge for each atom to ensure the most stable structure.
- The formal charge is calculated using the formula:
7. Draw Resonance Forms if Applicable:
- Since both nitrogen and oxygen can form double bonds with carbon, there are resonance structures where the double bond can be between carbon and nitrogen or carbon and oxygen.
The Lewis structure for CH_3NCO with the lowest formal charges and the most common resonance form is:
```
H
|
H-C-N=C=O
|
H
```
In this structure, the carbon atom is double-bonded to the oxygen atom and the nitrogen atom, forming a linear arrangement around the carbon atom. The nitrogen atom has a lone pair of electrons, and the oxygen atom has two lone pairs of electrons.
Resonance forms for CH_3NCO involve shifting the pi (π) electrons between the nitrogen and carbon atoms and between the carbon and oxygen atoms. However, the resonance form shown above is the most stable due to the electronegativity of oxygen, which makes it more favorable for oxygen to have a double bond with carbon than nitrogen.
Reactivity of CH_3NCO is influenced by the presence of the isocyanate group (NCO). The carbon in the isocyanate group is electrophilic, meaning it is susceptible to attack by nucleophiles. This reactivity can lead to various chemical reactions, such as nucleophilic addition or substitution, which are important in the synthesis of urethanes and other organic compounds.
Lewis Theory and Covalent Bonds
Strategy for Drawing Lewis Structures
Exceptions to octet rule
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Lewis Theory and Covalent Bonds
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In the previous video,
we talked about Lewis theory and ionic bonds.
In this video, we'll talk about covalent bonds.
We're going to describe the Lewis structure of covalent compounds.
Covalent compounds take place between nonmetals.
We're going to describe 2 examples,
water H_2O and oxygen difluoride, which is OF_2.
Let's begin with water.
Hydrogen has 1 valence electron,
and oxygen has 6 valence electrons,
2 pairs and 2 single electrons that are unpaired.
When the molecule is formed,
1 electron from hydrogen shares with
1 electron from the oxygen to form what we call a bond pair.
Now we have a bond pair between hydrogen and oxygen,
and the other one between oxygen and hydrogen.
I've indicated 1 electron in red and 1 in blue,
but it doesn't really matter because electrons are indistinguishable.
Another way of writing this is to indicate a bond pair by a single line,
that's a single covalent bond.
A bond pair to form
is single covalent bond.
We have 2 bond pairs.
In addition, we can add the pairs of electrons on the oxygen.
These pairs of electrons are called lone pairs.
We have lone pairs and bond pairs.
Now if we look at H_2O,
we can count the number of electrons around each atom.
Around hydrogen, there are 2 electrons,
that we call it duplet.
Around oxygen, there are 8 electrons,
we call that an octet,
another duplet, and we can do the same thing in this picture.
If we look at the oxygen,
2 electrons for each bond pair,
and 2 electrons for each lone pair giving a total of 8.
Now let's look at OF_2.
Fluorine has 7 valence electrons,
3 pairs and 1 unpaired electron.
Oxygen has 6 valence electrons as we saw before.
You'll see electrons in the other fluorine.
The unpaired electrons, fluorine can share with
the unpaired electron on the oxygen to form a bond pair,
and each bond pair we saw before,
can be written as a single straight line.
Once again, we can add all the lone pairs.
We have 3 lone pairs around fluorine and 2 lone pairs around the oxygen.
Once again, we can count the number of electrons around each atom,
around fluorine there's an octet,
around oxygen another octet,
around fluorine another octet.
One thing we should notice is that we've counted some of the electrons more than once.
For example, fluorine has 7 electrons,
oxygen has 6,
and fluorine has 7.
That's a total of 20,
and if we counted an octet,
an octet and an octet that means we've counted 24 electrons.
Some electrons are being counted more than once.
Now we're going to describe
a coordinate covalent bond which is a variant on the covalent bond.
Now in our previous examples,
H_2O and OF_2,
each atom contributed 1 electron to the bond pair.
However, sometimes 1 atom contributes both electrons to the bond pair.
This is called a coordinate covalent bond.
Once it's formed, it seemed distinguishable from a covalent bond.
But when it's formed,
it's a coordinate covalent bond.
Let's take an example.
Going to discuss the reaction of ammonia with hydrogen chloride.
Let's look at the Lewis structure of ammonia.
Ammonia has 3 single bonds connecting the nitrogen with each of the hydrogens.
In addition, it has a lone pair,
so that makes an octet,
2 electrons in each of the bond pairs,
and 2 electrons from the lone pair,
a total of 8 electrons.
HCl has a single bond,
so hydrogen has a duplet around it,
and chlorine has a bond pair,
and another 3 lone pairs.
Now what happens when they combine is that
all the electrons in this bond pair go to the chlorine.
Now, chlorine has 1 electron more than it started with.
It started off with 7 and now it has 8,
so it has an additional electrons,
so it's chlorine minus with an octet around it.
Hydrogen is left without any electrons at all.
Hydrogen is H plus,
so we've gone to H plus plus Cl minus.
Now the lone pair on the nitrogen combines with the hydrogen plus.
The nitrogen donates its lone pair to hydrogen plus,
and we get NH_4 plus.
Here are the single bonds.
We have 4 single bonds between the nitrogen and 4 hydrogens,
and the whole species has a positive charge, it's called ammonium.
Now we have ammonium plus and chlorine minus.
We have ammonium chloride.
Now NH_3 ammonia is called a Lewis base because it donates a pair of
electrons and HCL is called a Lewis acid because it accepts a pair of electrons.
In this video, we discussed single covalent bonds,
and also coordinate covalent bonds.
This video discusses covalent bonds, which take place between nonmetals. It explains the Lewis structure of covalent compounds, such as water (H_2O) and oxygen difluoride (OF_2). It also explains the concept of a coordinate covalent bond, which is a variant of a covalent bond where one atom contributes both electrons to the bond pair. This is demonstrated through the reaction of ammonia (NH_3) with hydrogen chloride (HCl), which forms ammonium chloride (NH_4Cl). The video also explains the concepts of Lewis bases and Lewis acids, which donate and accept electrons, respectively.
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