What is the Lewis Structure of H₂CO?

What is the Lewis Structure of H₂CO? The Lewis structure has carbon as the central atom with single bonds to both chlorine atoms and a double bond to the oxygen atom. There are two lone pairs of electrons on the oxygen atom and three lone pairs on each chlorine atom.

What is this molecule and what is it used for?

H₂CO is a molecule that you probably already know from biology class as formaldehyde. it has a number of applications, outside of the most common use, which is embalming tissue for preservation. It is also used as a resin in construction, a preservative/disinfectant, as a textile treatment and in manufacturing chemicals and pharmaceuticals.

Method 1: Step method to draw the Lewis structure of H₂CO.

In this method, we find the bonds and lone pairs for the whole molecule, then plug it in to the atoms that we have to get the answer. Here is a little flow chart of how we are going to do this:

We will go through the steps below, but one thing to note here is that all the valence electrons (step 1) are either lone pairs OR bonding electrons. In other words…. Lone Pairs (Step 5) + Bonding electrons (Step 3) = Valence electrons (Step 1) . Let’s go through this example so we can see this a little more clearly.

Step 1: Find valence electrons for all atoms. This is determined by looking at which column on the periodic table the atom is in, ignoring the transition metals in the middle. Add the valence electrons for each atom together.
C : 1×4 = 4
H : 2×1 = 2
O : 1×6 = 6
Total = 12 valence electrons

Step 2: Find octet electrons for each atom and add them together. Most atoms like 8 electrons to form an octet, however hydrogen is one of the exceptions, as it only wants two electrons to form an octet.

C: 1×8 = 8
H: 2×2 = 4
O: 1×8 = 8
Total = 20 “octet” electrons

Step 3: Find the number of bonding electrons. Subtract the valence electrons (step 1) from the octet electrons (step 2). This gives the number of bonding electrons.
20-12=8 bonding electrons.

Step 4: Find number of bonds by diving the number of bonding electrons (step 3) by 2 because each bond is made of 2 e-
8 bonding electrons/2 = 4 bond pairs

Step 5: Find the number of nonbonding (lone pairs) electrons. Subtract bonding electrons (step 3) number from valence electrons (step 1).
12 valence -8 bonding = 4 electrons = 2 lone pair

Now, use the information from step 4 and 5 to draw the Lewis structures. Remembering too (this is important):

Neutral carbon has four bonds and no lone pairs

Hydrogen has one bond and no lone pairs

Neutral oxygen has two bonds and two lone pairs.

[Note: For more information on the natural state of common atoms, see the linked post here.]

The carbon atom will be our central atom because it is the most electropositive. We then place the hydrogen and oxygen atoms around it, creating enough bonds to each to ensure an octet:

Another (easier) method to determine the Lewis structure of H₂CO:

Alternatively a dot method can be used to draw the Lewis structure.
Calculate the total valence electrons in the molecule.
C : 1×4 = 4
H : 2×1 = 2
O : 1×6 = 6
Total = 12 valence electrons

Now, treat the atoms and electrons like puzzle pieces. Put carbons in the center and arrange hydrogen and oxygen atoms on the sides. Remember the natural state of each atom, as discussed above. [Oxygen: 2 bonds, 2 LP; Hydrogen: 1 bonds, 0 LP; Carbon: 4 bonds, 0 LP] Also, make sure sure each atom has an octet, which in the case of hydrogen is only two electrons.

Frequently asked questions:

Q: So what is the difference between the two methods?

A: In the first method, we are figuring out all of the lone pairs and bonds first, then placing those electrons and bonds on the atoms to form a molecule. In the puzzle method, we already have lone pairs and bonding electrons assigned to each atom, so all we need to do is push puzzle pieces together to get a molecule. In each method, we need to remember the “happy state” of each atom, ei hydrogen likes 1 bond and no lone pairs, uncharged carbon likes four bonds and no lone pairs ect.

Q: What is going on with the double bond between oxygen and carbon?

A: A double bond between oxygen and carbon is called a carbonyl. It is electrophilic (meaning it has a small, partial positive charge) and reactive at the carbon atom. With respect to Lewis dot structures, the four electrons between the carbon and the oxygen count as four electrons for both O and C in satisfying the octet rule.

And now some video:

This is a quick video we put together that visually demonstrates the two methods for Lewis structure and Lewis dot problems.

And finally, the Lewis structure study guide:

Here it is, this is our one-page guide to Lewis Dot and Lewis Structures:

Lewis structure study guide

The post Lewis Structure of H2CO [with video and free study guide] appeared first on Organic Chemistry Made Easy by AceOrganicChem.

What is the Lewis Structure of N₂F₄?

What is the Lewis Structure of N₂F₄? The Lewis structure has nitrogen as the central atoms, with two fluorine atoms off each nitrogen. Each fluorine atom has three lone pairs while the nitrogen atoms have one lone pair each.

What is this molecule and what is it used for??

N₂F₄, also known as dinitrogen tetrafluoride or nitrogen tetrafluoride, is a highly toxic and reactive chemical compound. It has several applications including: Rocket propellant, cleaning agent (particularly semiconductors and electronics manufacturing), fluorinating agent (it can be a source of fluorine for chemical reactions), and polymerization initiator (it can start/catalyze polymerization reactions which are often used in plastics).

Method 1: Step method to draw the Lewis structure of N₂F₄.

In this method, we find the bonds and lone pairs for the whole molecule, then plug it in to the atoms that we have to get the answer. Here is a little flow chart of how we are going to do this:

We will go through the steps below, but one thing to note here is that all the valence electrons (step 1) are either lone pairs OR bonding electrons. In other words…. Lone Pairs (Step 5) + Bonding electrons (Step 3) = Valence electrons (Step 1) . Let’s go through this example so we can see this a little more clearly.

Step 1: Find valence electrons for all atoms. This is determined by looking at which column on the periodic table the atom is in, ignoring the transition metals in the middle. Add the valence electrons for each atom together.
2N : 2×5 = 10
4F : 4×7 = 28
Total = 38 valence electrons

Step 2: Find octet electrons for each atom and add them together. Most atoms like 8 electrons to form an octet, but there are exceptions.

2N: 2×8 = 16
4F: 4×8 = 32
Total = 48 “octet” electrons

Step 3: Find the number of bonding electrons. Subtract the valence electrons (step 1) from the octet electrons (step 2). This gives the number of bonding electrons.
48-38=10 bonding electrons.

Step 4: Find number of bonds by diving the number of bonding electrons (step 3) by 2 because each bond is made of 2 e-
10 bonding electrons/2 = 5 bond pairs

Step 5: Find the number of nonbonding (lone pairs) electrons. Subtract bonding electrons (step 3) number from valence electrons (step 1).
38 valence -10 bonding = 28 electrons = 14 lone pair

Now, use the information from step 4 and 5 to draw the Lewis structures. Remembering too (this is important):

Uncharged nitrogen likes 3 bonds and 1 lone pair

Fluorine has one bond and three lone pairs

[Note: For more information on the natural state of common atoms, see the linked post here.]

The N atoms will be our central atoms because it is the most electropositive. We then place the fluorine atoms around them, creating one bond to each F atom. [Hint: Nature loves symmetry, so it is always a good idea to start with a symmetrical molecule and see if that works first!] Behold, N₂F₄ Lewis dot structure.

Another (easier) method to determine the Lewis structure of N₂F₄:

Alternatively a dot method can be used to draw the Lewis structure.
Calculate the total valence electrons in the molecule.
2N : 2×5 = 10
4F : 4×7 = 28
Total = 38 valence electrons

Now, treat the atoms and electrons like puzzle pieces. Put the two nitrogen atoms in the center and arrange the fluorine atoms on the sides. Remember the natural state of each atom, as discussed above. [F: 1 bonds, 3 LP; N: 3 bonds, 1 LP; ] Also, make sure sure each atom has an octet. Behold, the Lewis Structure of N₂F₄.

Frequently asked questions:

Q: So what is the difference between the two methods?

A: In the first method, we are figuring out all of the lone pairs and bonds first, then placing those electrons and bonds on the atoms to form a molecule. In the puzzle method, we already have lone pairs and bonding electrons assigned to each atom, so all we need to do is push puzzle pieces together to get a molecule. In each method, we need to remember the “happy state” of each atom, ei hydrogen likes 1 bond and no lone pairs, uncharged carbon likes four bonds and no lone pairs ect.

Q: Just by looking at the molecule, are there clues it might be unstable?

A: Excellent question! Yes, there are. Your first clue is nitrogen single-bonded to another nitrogen. This is highly unstable because the lone pairs on the N are so close to each other, and because there are so many other lone pairs on the molecule, it can rotate to a lower energy configuration.

And now some video:

This is a quick video we put together that visually demonstrates the two methods for Lewis structure and Lewis dot problems.

And finally, the Lewis structure study guide:

Here it is, this is our one-page guide to Lewis Dot and Lewis Structures:

Lewis structure study guide

The post Lewis Structure of N2F4 [with video and free study guide] appeared first on Organic Chemistry Made Easy by AceOrganicChem.

What is the Lewis Structure of BeF₂?

What is the Lewis Structure of BeF₂? The Lewis structure has beryllium as the central atom, with a fluorine atom on each side. Each fluorine atom has three lone pairs while the beryllium atom has no lone pairs.

What is this molecule?

BeF₂, also known as beryllium fluoride, is a chemical compound primarily used in various industrial applications, to include use as a fluxing agent in the production of ceramics and glass, use as an electrolyte in batteries, and also as a chemical catalyst in certain chemical reactions.

Method 1: Step method to draw the Lewis structure of BeF₂.

In this method, we find the bonds and lone pairs for the whole molecule, then plug it in to the atoms that we have to get the answer. Here is a little flow chart of how we are going to do this:

We will go through the steps below, but one thing to note here is that all the valence electrons (step 1) are either lone pairs OR bonding electrons. In other words…. Lone Pairs (Step 5) + Bonding electrons (Step 3) = Valence electrons (Step 1) . Let’s go through this example so we can see this a little more clearly.

Step 1: Find valence electrons for all atoms. This is determined by looking at which column on the periodic table the atom is in, ignoring the transition metals in the middle. Add the valence electrons for each atom together.
Be : 1×2 = 2
F : 2×7 = 14
Total = 16 valence electrons

Step 2: Find octet electrons for each atom and add them together. Most atoms like 8 electrons to form an octet, but there are exceptions and Be is one of those exceptions, as it only needs four electrons for an octet.

Be: 1×4 = 4
F: 2×8 = 16
Total = 20 “octet” electrons

Step 3: Find the number of bonding electrons. Subtract the valence electrons (step 1) from the octet electrons (step 2). This gives the number of bonding electrons.
20-16=4 bonding electrons.

Step 4: Find number of bonds by diving the number of bonding electrons (step 3) by 2 because each bond is made of 2 e-
4 bonding electrons/2 = 2 bond pairs

Step 5: Find the number of nonbonding (lone pairs) electrons. Subtract bonding electrons (step 3) number from valence electrons (step 1).
16 valence -4 bonding = 12 electrons = 6 lone pair

Now, use the information from step 4 and 5 to draw the Lewis structures. Remembering too (this is important):

Beryllium likes 2 bonds and no lone pairs

Fluorine has one bond and three lone pairs

[Note: For more information on the natural state of common atoms, see the linked post here.]

The Be atom will be our central atom because it is the most electropositive. We then place the fluorine atoms around it, creating one bond to each F atom. Behold, BF₂ Lewis dot structure.

Another (easier) method to determine the Lewis structure of BeF₂:

Alternatively a dot method can be used to draw the Lewis structure.
Calculate the total valence electrons in the molecule.
Be : 1×2 = 2
F : 2×7 = 14
Total = 16 valence electrons

Now, treat the atoms and electrons like puzzle pieces. Put beryllium in the center and arrange the fluorine atoms on the sides. Remember the natural state of each atom, as discussed above. [F: 1 bonds, 3 LP; Be: 2 bonds, 0 LP; ] Also, make sure sure each atom has an octet, which in the case of beryllium is only four electrons.

Frequently asked questions:

Q: So what is the difference between the two methods?

A: In the first method, we are figuring out all of the lone pairs and bonds first, then placing those electrons and bonds on the atoms to form a molecule. In the puzzle method, we already have lone pairs and bonding electrons assigned to each atom, so all we need to do is push puzzle pieces together to get a molecule. In each method, we need to remember the “happy state” of each atom, ei hydrogen likes 1 bond and no lone pairs, uncharged carbon likes four bonds and no lone pairs ect.

Q: Why does beryllium only need four electrons for an octet?

A: Beryllium, with an atomic number of 4, belongs to Group 2 of the periodic table. In terms of its electronic configuration, it only has four total electrons, with two electrons in its innermost shell (1s²) and two electrons in its outermost shell (2s²). Atoms become stable when the outermost shell is full, which is usually 8 electrons. Yet, in the case of beryllium, it is an exception to the octet rule because it is stable with only four electrons in its outermost shell.

Because Be has a small atomic radius, the 2s orbital, which is the outermost orbital for beryllium, is closer to the nucleus compared to higher energy orbitals from in other atoms. As a result, the outermost electrons are strongly attracted to the nucleus and held tightly. The energy required to either gain four more electrons or lose the existing two electrons from the 2s orbital is significantly high, making it energetically unfavorable for beryllium to achieve an octet. Instead, beryllium tends to form compounds by sharing its two valence electrons with other elements, resulting in a stable electron configuration with only four electrons in its outermost shell. Those four electrons are the two it brings to the party itself, from its 2s orbital, and the two electrons it picks up from other atoms to form two bonds.

And now some video:

This is a quick video we put together that visually demonstrates the two methods for Lewis structure and Lewis dot problems.

And finally, the Lewis structure study guide:

Here it is, this is our one-page guide to Lewis Dot and Lewis Structures:

Lewis structure study guide

The post Lewis Structure of BeF2 [with video and free study guide] appeared first on Organic Chemistry Made Easy by AceOrganicChem.

What is the Lewis Structure of COCl₂?

What is the Lewis Structure of COCl₂? The Lewis structure has carbon as the central atom with single bonds to both chlorine atoms and a double bond to the oxygen atom. There are two lone pairs of electrons on the oxygen atom and three lone pairs on each chlorine atom.

What is this molecule?

COCl₂ is the chemical formula for carbonyl chloride, also known as phosgene. It is a colorless gas at room temperature with a pungent odor. Phosgene is a highly toxic compound that was historically used as a chemical warfare agent during World War I. Today, it finds application in the production of various chemicals, such as plastics, pesticides, and pharmaceuticals.

Method 1: Step method to draw the Lewis structure of COCl₂.

In this method, we find the bonds and lone pairs for the whole molecule, then plug it in to the atoms that we have to get the answer. Here is a little flow chart of how we are going to do this:

We will go through the steps below, but one thing to note here is that all the valence electrons (step 1) are either lone pairs OR bonding electrons. In other words…. Lone Pairs (Step 5) + Bonding electrons (Step 3) = Valence electrons (Step 1) . Let’s go through this example so we can see this a little more clearly.

Step 1: Find valence electrons for all atoms. This is determined by looking at which column on the periodic table the atom is in, ignoring the transition metals in the middle. Add the valence electrons for each atom together.
C : 1×4 = 4
Cl : 2×7 = 14
O : 1×6 = 6
Total = 24 valence electrons

Step 2: Find octet electrons for each atom and add them together. Most atoms like 8 electrons to form an octet.

C: 1×8 = 8
Cl: 2×8 = 16
O: 1×8 = 8
Total = 32 “octet” electrons

Step 3: Find the number of bonding electrons. Subtract the valence electrons (step 1) from the octet electrons (step 2). This gives the number of bonding electrons.
32-24=8 bonding electrons.

Step 4: Find number of bonds by diving the number of bonding electrons (step 3) by 2 because each bond is made of 2 e-
8 bonding electrons/2 = 4 bond pairs

Step 5: Find the number of nonbonding (lone pairs) electrons. Subtract bonding electrons (step 3) number from valence electrons (step 1).
24 valence -8 bonding = 16 electrons = 8 lone pair

Now, use the information from step 4 and 5 to draw the Lewis structures. Remembering too (this is important):

Neutral carbon has four bonds and no lone pairs

Chlorine has one bond and three lone pairs

Neutral oxygen has two bonds and two lone pairs.

[Note: For more information on the natural state of common atoms, see the linked post here.]

The carbon atom will be our central atom because it is the most electropositive. We then place the chlorine and oxygen atoms around it, creating enough bonds to each to ensure an octet:

Another (easier) method to determine the Lewis structure of COCl₂:

Alternatively a dot method can be used to draw the Lewis structure.
Calculate the total valence electrons in the molecule.
C : 1×4 = 4
Cl : 2×7 = 14
O : 1×6 = 6
Total = 24 valence electrons

Now, treat the atoms and electrons like puzzle pieces. Put carbons in the center and arrange hydrogen and oxygen atoms on the sides. Remember the natural state of each atom, as discussed above. [Oxygen: 2 bonds, 2 LP; Chlorine: 1 bonds, 3 LP; Carbon: 4 bonds, 0 LP] Also, make sure sure each atom has an octet, which in the case of hydrogen is only two electrons.

Frequently asked questions:

Q: So what is the difference between the two methods?

A: In the first method, we are figuring out all of the lone pairs and bonds first, then placing those electrons and bonds on the atoms to form a molecule. In the puzzle method, we already have lone pairs and bonding electrons assigned to each atom, so all we need to do is push puzzle pieces together to get a molecule. In each method, we need to remember the “happy state” of each atom, ei hydrogen likes 1 bond and no lone pairs, uncharged carbon likes four bonds and no lone pairs ect.

Q: What is going on with the double bond between oxygen and carbon?

A: A double bond between oxygen and carbon is called a carbonyl. It is electrophilic (meaning it has a small, partial positive charge) and reactive at the carbon atom. With respect to Lewis dot structures, the four electrons between the carbon and the oxygen count as four electrons for both O and C in satisfying the octet rule.

And now some video:

This is a quick video we put together that visually demonstrates the two methods for Lewis structure and Lewis dot problems.

And finally, the Lewis structure study guide:

Here it is, this is our one-page guide to Lewis Dot and Lewis Structures:

Lewis structure study guide

The post The Lewis Structure of COCl2 [with free study guide and video] appeared first on Organic Chemistry Made Easy by AceOrganicChem.

What is the Lewis Structure of PCl₅?

What are the Lewis Structure of PCl₅? Phosphorus pentachloride (PCl5) is a chemical compound composed of phosphorus and chlorine. It is a molecule which consists of five chlorine atoms attached to a central phosphorus atom. The P atom is hypervalent in this molecule, and the overall geometry is a trigonal bipyramidal arrangement of chlorine atoms.

What is this molecule?

Phosphorus pentachloride is primarily used as a chlorinating agent in organic synthesis. It can selectively replace hydroxyl groups (-OH) with chlorine atoms (-Cl) in organic compounds, a process known as chlorination. It is particularly useful in the production of various chlorinated compounds, such as acyl chlorides and alkyl chlorides. When phosphorus pentachloride reacts with water, it undergoes a vigorous reaction, producing hydrochloric acid (HCl) and phosphoric acid (H₃PO₄). This reaction is exothermic and releases heat. The compound is also employed as a catalyst in certain chemical reactions, such as the conversion of carboxylic acids to acid chlorides and the preparation of thionyl chloride (SOCl₂).

Method 1: Step method to draw the Lewis structure of PCl₅.

In this method, we find the bonds and lone pairs for the whole molecule, then plug it in to the atoms that we have to get the answer. Here is a little flow chart of how we are going to do this:

We will go through the steps below, but one thing to note here is that all the valence electrons (step 1) are either lone pairs OR bonding electrons. In other words…. Lone Pairs (Step 5) + Bonding electrons (Step 3) = Valence electrons (Step 1) . Let’s go through this example so we can see this a little more clearly.

Step 1: Find valence electrons for all atoms. This is determined by looking at which column on the periodic table the atom is in, ignoring the transition metals in the middle. Add the valence electrons for each atom together.
P: 1×5 = 5
Cl: 5×7 = 35
Total = 40 valence electrons

Step 2: Find octet electrons for each atom and add them together. Most atoms like 8 electrons to form an octet, however H is an exception to this.

P: 1x 10 = 10**
Cl: 5 x 8 = 40
Total = 50 “octet” electrons

NOTE: Phosphorous gets 10 valence electrons in this case since it must make 5 bonds with surrounding Cl atoms.

Step 3: Find the number of bonding electrons. Subtract the valence electrons (step 1) from the octet electrons (step 2). This gives the number of bonding electrons.
50-40=10 bonding electrons.

Step 4: Find number of bonds by diving the number of bonding electrons (step 3) by 2 because each bond is made of 2 e-
10 bonding electrons/2 = 5 bond pairs

Step 5: Find the number of nonbonding (lone pairs) electrons. Subtract bonding electrons (step 3) number from valence electrons (step 1).
40 valence – 10 bonding = 30 electrons = 15 lone pairs

Now, use the information from step 4 and 5 to draw the Lewis structures. Remembering too (this is important):

Phosphorus will have five bonds

Chlorine will have 1 bond and three lone pairs

[Note: For more information on the natural state of common atoms, see the linked post here.]

The phosphorus atom will be the central atom, and will connect to the five chlorine atoms. Each chlorine atom will have three lone pairs to complete the octet. Therefore the Lewis structure of PCl₅ looks like this:

Another (easier) method to determine the Lewis structure of PCl₅:

Alternatively a dot method can be used to draw the Lewis structure.
Calculate the total valence electrons in the molecule.
P: 1×5 = 5
Cl: 5×7 = 35
Total = 40 valence electrons

Now, treat the atoms and electrons like puzzle pieces. Put phosphorus in the center and arrange the chlorine atoms on the sides. Remember the natural state of each atom, as discussed above. [Phosphorus: 5 bonds, 0 LP; Chlorine: 1 bonds, 3 LP]. We will confirm each atom has their desired octet before we say this is our final answer.

Frequently asked questions:

Q: So what is the difference between the two methods?

A: In the first method, we are figuring out all of the lone pairs and bonds first, then placing those electrons and bonds on the atoms to form a molecule. In the puzzle method, we already have lone pairs and bonding electrons assigned to each atom, so all we need to do is push puzzle pieces together to get a molecule. In each method, we need to remember the “happy state” of each atom, ei hydrogen likes 1 bond and no lone pairs, uncharged carbon likes four bonds and no lone pairs ect.

Q: How can phosphorus have five bonds?

A: Overall, the ability of phosphorus to form compounds with five bonds is due to its ability to expand its valence shell by utilizing its vacant d orbitals, allowing for additional bonding opportunities. This is called hypervalency.
Now the super technical explanation: Phosphorus typically forms compounds with five bonds by utilizing its vacant d orbitals, which can accommodate additional electrons. In its ground state, phosphorus has three occupied orbitals in its valence shell (3s²3p³), leaving two vacant d orbitals available for bonding. The electronic configuration of phosphorus in PCl₅ involves hybridization of the 3s, 3p, and 3d orbitals to form five sp³d hybrid orbitals. Each of these hybrid orbitals overlaps with a 3p orbital of a chlorine atom, resulting in the formation of five sigma bonds. It’s important to note that the concept of five covalent bonds for phosphorus is a simplification used to describe the bonding in certain compounds. In reality, the electron density is spread out over the bonding regions, and the electron pairs are delocalized to some extent. Additionally, the actual nature of bonding can be described using molecular orbital theory, where the orbitals from multiple atoms combine to form molecular orbitals.

And now some video:

This is a quick video we put together that visually demonstrates the two methods for Lewis structure and Lewis dot problems.

And finally, the Lewis structure study guide:

Here it is, this is our one-page guide to Lewis Dot and Lewis Structures:

Lewis structure study guide

The post Lewis Structure of PCl5? [with free study guide and video] appeared first on Organic Chemistry Made Easy by AceOrganicChem.

What is the Lewis Structure of ClO₄^–?

The Lewis Structure of ClO₄^– is an ionic structure which has chlorine as the central atom connected to four oxygen atoms. Three of the four oxygen atoms are double bonded to chlorine and are neutral, while one chlorine atom is single bonded to chlorine and is negatively charged. It has a tetrahedral geometry.

ClO₄^–, also known as the perchlorate ion, is a highly oxidizing and reactive ion that can be synthesized through various methods, including the reaction of chlorine with sodium hydroxide. Perchlorate ions are used in a variety of industrial processes, including as rocket propellants, fireworks and explosives. They are considered to be environmental pollutants as they are quite persistent after dispersion. Unfortunately, they can be found in drinking water, soil, and various food products. Perchlorate ions are believed to have harmful effects on human beings, as they can interfere with the proper functioning of the thyroid gland by inhibiting the uptake of iodide, which is necessary for the production of thyroid hormones. This can lead to a condition called hypothyroidism, which can cause developmental delays in children and other health problems in adults.

Method to draw the Lewis dot structure of perchlorate ion, ClO₄^–.

Here is a little flow chart of how we are going to do this:

We will go through the steps below, but one thing to note here is that the valence electrons (step 1) are either lone pairs OR bonding electrons. In other words…. Lone Pairs (Step 5) + Bonding electrons (Step 3) = Valence electrons (Step 1) . Let’s go through this example so we can see this a little more clearly.

Step 1: Find valence electrons for all atoms. This is determined by looking at which column on the periodic table the atom is in, ignoring the transition metals in the middle. Add the valence electrons for each atom together. (we will use the shorthand of “e-” to symbolize the word electron)

Cl has 7 valence electrons. When we picture it like this, it becomes easier to see that stable and uncharged chlorine likes to have one bond and three lone pairs of electrons.

O: 6 x 4 atoms =24 total valance electrons. When we picture it like this, it becomes easier to see that stable and uncharged oxygen likes to have two bonds and two lone pairs of electrons.

Total = 31+1 = 32*

* add an electron to the total since the molecule has -1 charge on it.

Step 2: Find octet e- for each atom and add them together.

Cl : 8
O = 8×4 = 32

Total = 40

Step 3: Find the bonding e-. These are the electrons that comprise the bonds of the atom. Subtract step 1 total from step 2.

40 – 32 = 8e- bonding electrons

Step 4: Find number of bonds by diving the number in step 3 by 2 (because each bond is made of 2 e-)

8e- / 2= 4 bond pairs

Step 5: Find the number of nonbonding (lone pairs) e-. Subtract step 3 number from step 1.

32 – 8 = 24e- = 12 lone pair

Use information from step 4 and 5 to draw the Lewis structure.

HOWEVER, this structure above is not the right one! While every atom has an octet, and it follows our steps properly, it does not give you the right answer because this molecule has a charge of -4, which is not what perchlorate has. Therefore, we have to move some lone pairs around to get to a charge of -1.

In the structure above, we have moved three lone pairs and turned them into bonding electrons, which puts double bonds on three of the four oxygen atoms. While this may not look right, because Cl has more than an octet, it is the correct structure of ClO₄^–.

Another (easier) method to determine the structure of ClO₄^–

Alternatively a dot method can be used to draw the Lewis structure.
Calculate the total valence electrons in the molecule.
Cl : 7
O : 6×4 = 24

Total = 31+1 = 32*

Treat the atoms and electrons like puzzle pieces. Put chlorine in center and arrange oxygen atoms on the sides. put a pair of electrons connecting the chlorine atom and the remaining electrons on oxygen atoms.

Frequently Asked Questions:

Q: So what is the difference between the two methods to find the Lewis Structure of ClO₄^–?

A: In the first method, we are figuring out all of the lone pairs and bonds first, then placing those electrons and bonds on the atoms to form a molecule. In the puzzle method, we already have lone pairs and bonding electrons assigned to each atom, so all we need to do is push puzzle pieces together to get a molecule. In each method, we need to remember the “happy state” of each atom, ei hydrogen likes 1 bond and no lone pairs, uncharged carbon likes four bonds and no lone pairs ect.

Q: Why does chlorine appear to have 14 electrons around it?

A: In some cases, chlorine can take on more valence electrons than in an octet, this is referred to as hypervalence. Generally, this indicates the ion will be highly reactive.

Q: How does resonance affect this ion?

A: Any one of the oxygen atoms can be negatively charged at any one time. Another way to think about this is that the negative charge/single bond rotates quickly between the four oxygen atoms, essentially equally distributing the negative charge over all four.

And now some video:

This is a quick video we put together that visually demonstrates the two methods for Lewis structure and Lewis dot problems.

And finally, the Lewis Structure study guide:

Here it is, this is our one-page guide to Lewis Dot and Lewis Structures:

Lewis structure study guide

The post Lewis Structure of ClO4- [with video and free guide] appeared first on Organic Chemistry Made Easy by AceOrganicChem.

What is the Lewis Structure of C₂H₄O?

What are the Lewis Structures of C₂H₄O? Actually, there are several plausible structures. This molecular formula lead to either acetaldehyde, Ethylene Oxide, or Vinyl Alcohol. All are correct answers. Below, we show two methods to arrive at any one of these answers.

What are these molecules?

Acetaldehyde is colorless liquid with a pungent, fruity odor and is an important intermediate in various chemical reactions and is used in the production of many chemicals and materials. Acetaldehyde is primarily known for its role in alcohol metabolism. When alcohol is consumed, it is first broken down by the liver into acetaldehyde before being further metabolized into acetic acid and eventually carbon dioxide and water. The buildup of acetaldehyde in the body is what leads to some of the more unpleasant effects of alcohol consumption, such as hangovers.

Ethylene oxide is a flammable and colorless gas at room temperature, but it is commonly stored and transported as a liquid due to its low boiling point. Ethylene oxide has a sweet, ether-like odor. Ethylene oxide is an extremely versatile chemical and is primarily used in the production of other chemicals. It is a key building block for the synthesis of various products, including plastics, resins, solvents, detergents, textiles, and pharmaceuticals. It is also used as a sterilizing agent for medical equipment and supplies. One of the major applications of ethylene oxide is in sterilization processes. Due to its ability to penetrate materials and kill a wide range of microorganisms, it is commonly used to sterilize heat- or moisture-sensitive medical devices, such as surgical instruments, syringes, and catheters.

Vinyl alcohol, also known as ethenol, is an unstable molecule that is rarely isolated in pure form due to its tendency to undergo tautomerization, converting into its more stable tautomer, acetaldehyde. Vinyl alcohol is an important intermediate in various chemical reactions and is of interest in the field of organic chemistry. It is formed as an intermediate during the hydrolysis of vinyl acetate, a common chemical used in the production of polyvinyl acetate (PVA) and polyvinyl alcohol (PVOH).

Method 1: Step method to draw the Lewis dot structure of C₂H₄O.

In this method, we find the bonds and lone pairs for the whole molecule, then plug it in to the atoms that we have to get the answer. Here is a little flow chart of how we are going to do this:

We will go through the steps below, but one thing to note here is that all the valence electrons (step 1) are either lone pairs OR bonding electrons. In other words…. Lone Pairs (Step 5) + Bonding electrons (Step 3) = Valence electrons (Step 1) . Let’s go through this example so we can see this a little more clearly.

Step 1: Find valence electrons for all atoms. This is determined by looking at which column on the periodic table the atom is in, ignoring the transition metals in the middle. Add the valence electrons for each atom together.
C : 2×4 = 8
H : 4×1 = 4
O : 1×6 = 6
Total = 18 valence electrons

Step 2: Find octet electrons for each atom and add them together. Most atoms like 8 electrons to form an octet, however H is an exception to this.

C: 2×8 =16
H: 4×2 = 8
O: 1×8 = 8
Total = 32 “octet” electrons

Step 3: Find the number of bonding electrons. Subtract the valence electrons (step 1) from the octet electrons (step 2). This gives the number of bonding electrons.
32-18=14 bonding electrons.

Step 4: Find number of bonds by diving the number of bonding electrons (step 3) by 2 because each bond is made of 2 e-
14 bonding electrons/2 = 7 bond pairs

Step 5: Find the number of nonbonding (lone pairs) electrons. Subtract bonding electrons (step 3) number from valence electrons (step 1).
18 valence -14 bonding = 4 electrons = 2 lone pair

Now, use the information from step 4 and 5 to draw the Lewis structures. Remembering too (this is important):

Neutral carbon has four bonds and no lone pairs

Hydrogen has one bond and no lone pairs

Neutral oxygen has two bonds and two lone pairs.

[Note: For more information on the natural state of common atoms, see the linked post here.]

The two carbon atoms are the central atoms and go side by side. Now, arrange hydrogens and the oxygen atom on the sides of carbon. The following three Lewis structures are possible for C₂H₄O:

Another (easier) method to determine the structure of C₂H₄O:

Alternatively a dot method can be used to draw the Lewis structure.
Calculate the total valence electrons in the molecule.
C : 2×4 = 8
H : 4×1 = 4
O : 1×6 = 6
Total = 18 valence electrons

Now, treat the atoms and electrons like puzzle pieces. Put carbons in the center and arrange hydrogen and oxygen atoms on the sides. Remember the natural state of each atom, as discussed above. [Oxygen: 2 bonds, 2 LP; Hydrogen: 1 bonds, 0 LP; Carbon: 4 bonds, 0 LP] Also, make sure sure each atom has an octet, which in the case of hydrogen is only two electrons.

Frequently asked questions:

Q: So what is the difference between the two methods?

A: In the first method, we are figuring out all of the lone pairs and bonds first, then placing those electrons and bonds on the atoms to form a molecule. In the puzzle method, we already have lone pairs and bonding electrons assigned to each atom, so all we need to do is push puzzle pieces together to get a molecule. In each method, we need to remember the “happy state” of each atom, ei hydrogen likes 1 bond and no lone pairs, uncharged carbon likes four bonds and no lone pairs ect.

Q: How is it that we can get three different molecules from the same formula?

A: It is all in how you arrange the atoms. As long as you are following the rules and placing the appropriate number of electrons on each atom, there may be several correct answers. BTW, when you have different molecules with the same formula, they are called constitutional isomers. Here is a good post about constitutional isomers.

Q: OK, so which one is right?

A: That depends on which Lewis Structure of C₂H₄O your professor is looking for. Any of the three are technically correct, so if your professor did not give any guidance (ei a name of a compound), all should get you full credit.

And now some video:

This is a quick video we put together that visually demonstrates the two methods for Lewis structure and Lewis dot problems.

And finally, the Lewis structure study guide:

Here it is, this is our one-page guide to Lewis Dot and Lewis Structures:

https://www.biochemhelp.com/Lewis-Dot-Structure-guide.pdf

The post The Lewis Structures of C2H4O appeared first on Organic Chemistry Made Easy by AceOrganicChem.

Structures of molecules can be difficult to piece together at first when you are just starting in an organic chemistry class. Hopefully you retained some of this knowledge from general chemistry. If not, one of the tricks that can greatly help with this is to know the uncharged or “normal” state for atoms that are commonly found in organic molecules.   Here is a table of the most common of those:

      – C has four bonds and no lone pairs

       – N has three bonds and one lone pair

       – Halogens (F, Cl, Br, I) have one bond and three lone pairs. 

       – O has two bonds and two lone pairs

       – H has one bond and no lone pairs

Three more rules:

–          C, N, O are central atoms, meaning that they will always be in the middle of your molecule.

–          H and halogens are terminal atoms, meaning that they will only have one bond and be at the ends of molecules.

–          With the exception of H, atoms in group I & group II are only counterions (+1 or +2 and not involved in resonance).

Remember, these rules are for when the atom is uncharged; this does not apply to charged atoms.  For example, a carbocation (a positively charged carbon atom) will have only three bonds with no lone pairs while a carbanion (a negatively charged carbon atom) wlll have three bonds with one lone pair, and a carbene will have two bonds with two lone pairs.

Notice that all of these carbons still follow the octet rule.  However, beware of atoms that do not follow the octet rule, as phosphorus is an example of an atom that can have more than an octet of electrons.  Shown below is triphenylphosphine oxide, a byproduct of the Wittig reaction.

Elements with open d-subshells, like phosphorous and sulfur, do not always follow the octet rule.  More examples of this are SF₆ and PCl₅.  However, carbon, nitrogen and oxygen will follow the octet rule.

The post Know the “normal” state for common organic atoms [3 rules to live by] appeared first on Organic Chemistry Made Easy by AceOrganicChem.

Emil Fischer is considered by many to be the greatest organic chemist to ever live.  His problem was that he created a way of looking at organic molecules that is very confusing to undergraduates.  These structures are necessary to learn and are very helpful when looking at certain molecules (such as carbohydrates), but they are also very easy to jumble.  This is because Fischer structures are drawn as crosses, which could lead one to erroneously think that the central carbon is flat, when it is actually still tetrahedral.

The easiest way to look at these is to think of them as bowties that have been strung together:

3-dimensionally speaking, the substituents that are on the sides of the structure are depicted at the end of the bowtie and are represented as “coming out of the paper”.  The backbone is composed of dashed lines, which are meant to represent that those portions “are going into the paper”.  This is now a much easier way to view these structures, as it is more apparent what area each substituent occupies.

The useful part of the bowtie projection is that it is now easier to assess the stereochemistry at each chiral center.   It should be much easier to visualize that the bottom chiral center is “R”.  This was not as obvious when viewing the Fischer projection as a cross

For more help like this, please go to organic chemistry

The post Fischer Projections are a Black Tie Affair appeared first on Organic Chemistry Made Easy by AceOrganicChem.

Hi Everybody–Resonance is one of those issues that you will have to deal with for both semester I & II organic chemistry.  It is much better to have a solid understanding of it now, rather than have to worry about it later.  The basic goal of resonance structures is to show that molecules can move electrons and charges onto different atoms on the molecule.  This makes the molecule generally more stable because the charge is now delocalized and not “forced” on an atom that does not want it.

Below are some handy rules of resonance.  If you learn these and think about them when tackling different resonance problems, you will be able to handle whatever is thrown at you.

1) Know each atom’s “natural state”.  You need to recognize what each atom generally looks like, in an uncharged state.  This will help you to construct the Lewis Dot structure on which you will base your resonance structures.  In most uncharged cases:

       – C has four bonds and no lone pairs

       – N has three bonds and one lone pair

       – Halogens (F, Cl, Br, I) have one bond and three lone pairs. 

       – O has two bonds and two lone pairs

       – H has one bond and no lone pairs

       – With the exception of H, everyone in group I & group II are only counterions (+1 or +2 and not involved in resonance).

Remember that halogens and hydrogens are always terminal, meaning that are at the end of the molecule and only have one bond, and therefore, they will not participate in resonance.

2) Atom positions will not change.  Once you have determined that an atom is bonded to another atom, that will not change in a resonance structure.  If they do change, it is no longer a resonance strucutre, but is now a constitutional isomer.

3) Check the structure you have created to make sure that it follows the octet rule.  This will become much easier once you have a better handle on the “natural state” of atoms.

4) When two or more resonance structures can be drawn, the one with the fewest total charges is the most stable.  In the example below, A is more stable than B.

5) When two or more resonance structures can be drawn, the more stable has the negative charge on the more electronegative atom.  In the example below, A is more stable than B.

6) In the end, each resonance structure should have the same overall charge and total number of electrons (bonds + lone pairs) as when you started.

The post Organic Chemistry Help: Resonance appeared first on Organic Chemistry Made Easy by AceOrganicChem.