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Worked example: Lewis diagram of formaldehyde (CH₂O)

The Lewis diagram of a molecule can be constructed using the following stepwise procedure: (1) Find the number of valence electrons in the molecule. (2) Draw single bonds between bonded atoms. (3) Distribute the remaining electrons throughout the molecule, keeping in mind the duet and octet rules. In this video, we'll use these steps to construct the Lewis diagram of formaldehyde (CH₂O). Created by Sal Khan.

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  • boggle yellow style avatar for user Anuja
    Why is the least electronegative atom the central atom?
    (7 votes)
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    • leaf red style avatar for user Richard
      Atoms which have lower electronegativities hold onto their electrons less tightly and therefore are more prone to share their electrons. The central atom of a molecule needs to be sharing its electrons with multiple atoms which is easier to do so with a less electronegative atom which isn't as reluctant to share its electrons.

      Hope that helps.
      (26 votes)
  • blobby green style avatar for user Fadi Jalal
    hello, in the beginning sal said to try drawing the lewis diagram before he did. so i did draw it and i made sure all of the atoms were Happy with their octet i also made the double bond and also i had the hydrogens the furthest but my drawing wasnt the same? i had the C atom bonded to one H atom and 1 O Atom and the O atom was bonded to another H atom, so basically i connecected the C and O atoms and gave each 1 H atom
    would it still be correct?
    (2 votes)
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    • leaf red style avatar for user Richard
      For a molecular formula of CH2O, we would have 12 valence electrons to distribute around the molecule, or 6 pairs of electrons. If you connect them in the way you’ve described, H-C-O-H, then we will have used 3 pairs to connect them together with single bonds. If we add lone pairs to the oxygen and a double bond between the carbon and oxygen we get, H-C=O-H, then that accounts for all 6 pairs.

      Two problems with this drawing. First the carbon is deficient of electrons with only 6 and does not have an octet. Also the oxygen will have a -1 formal charge and make this molecule an ion, which it is not.

      Sal’s drawing of formaldehyde is correct. Where there is a central carbon double bonded to an oxygen and single bonded to the two hydrogens. And the oxygen has two lone pairs. This gives carbon and oxygen an octet, and have no formal charges.

      Hope that helps.
      (7 votes)
  • stelly orange style avatar for user BootesVoidPointer
    Why are the hydrogen atoms drawn with a 45 degrees angle?
    (3 votes)
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  • male robot donald style avatar for user sylvainfarrel
    How can I know that I must draw the angle diagonally? I need to know the electronegative value right? (Polar and nonpolar covalent bond)
    (2 votes)
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    • leaf red style avatar for user Richard
      The shape of the molecule, and the bond angles, are determined by VSEPR theory, not electronegativity. Electronegativity tells us the dipole moment of a bond and if it is polar.

      VSEPR theory considers the amount of electron groups around a central atom, both bonding and lone pairs, and determines the shape from number of those groups. The idea is that electron groups want to repel each other as much as possible because of the negative charge of the electrons. So they adopt geometries and bond angles that maximize the distance between the electron groups.

      For formaldehyde, the central atom is the carbon atom which has three electron groups around it, two single bonds and a double bond. This would indicate that the geometry is trigonal planar, or a flat triangle where the groups bonded to the central carbon are the vertices of an equilateral triangle with the carbon at the center of that triangle. And the bond angle because the electron groups would be 120° since we’re dividing a full rotation into three equal parts (360°/3 = 120°).

      Now that’s what the molecule looks like in reality, but when we’re doing Lewis structures it’s not always crucial to draw the correct 3D geometry. We’re primarily interested in what atom is bonded to which and where all the valence electrons are. So it would still be a mostly correct structure if we drew it trigonal planar or in a T-shape. It’s just as you draw these often enough and familiarize yourself with the geometries you begin to instinctively draw certain Lewis structure with their correct bond angles.

      Hope that helps.
      (5 votes)
  • blobby green style avatar for user 1654jduan
    Q: How can you be sure of the amount of covalent electrons, and what if there is 1 extra or 1 less electron in the molecule structure?
    (3 votes)
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    • leaf red style avatar for user Richard
      Valid Lewis structures are ones where the atoms have their octets filled and their formal charges are as low as possible. This will dictate how many covalent bonds exist between atoms.

      When you say one extra or less electron, what exactly do you mean by this? Because this can be interpreted as simply having more electrons than protons (or more protons than electrons), and thus would just be an ion. Or it could mean that we have an odd number of electrons where one electron is no paired, either as a bonding pair or a lone pair. And in this case that would be called a radical.

      Hope that helps.
      (2 votes)
  • blobby green style avatar for user Amin Khateeb
    Towards the end at sal said that carbon would love to have a full outer shell, hence 8 electrons. But the outer shell of carbon is a p type shell and can have up to 6 electrons. So why did he add other 2 electrons?
    (2 votes)
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  • stelly green style avatar for user burrowwitch
    Why are the four electrons around oxygen in pairs, rather than having one pair and two single electrons?
    (1 vote)
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    • leaf red style avatar for user Richard
      Electrons are more stable existing in pairs as opposed to being unpaired. Electrons which are not paired are called radicals and, as the name suggests, are unstable and reactive. So the electrons would prefer to be more stable and exist as lone pairs than exist as radicals.

      Hope that helps.
      (3 votes)
  • blobby green style avatar for user Steven Chelney
    Why is hydrogen never the central atom?
    (1 vote)
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    • leaf red style avatar for user Richard
      Hydrogen only forms one single bond in these sort of problems. And this is a result of hydrogen's small valence shell which can only hold a maximum of two electrons. So if we wanted to draw the Lewis structure of this molecule with hydrogen as the central atom, then we would only be able have the central hydrogen with a single bond to one of the atoms and wouldn't have room for additional bonds to the rest of the atoms. Hence our proposed Lewis structure with a central hydrogen wouldn't match up with the chemical formula CH2O and we wouldn't consider it correct.

      The only time we can consider hydrogen as a central atom in some capacity is for small diatomic molecules like hydrogen gas, H2, or hydrochloric acid, HCl. Even then it's a stretch saying a molecule with only two atoms has a central atom.

      So hydrogen's limit in the amount of bonds it can form resigns it from being able to act as the central atom.

      Hope that helps.
      (3 votes)
  • blobby green style avatar for user Ev
    how do you know The placement of the remaining electrons?
    Why couldn't you take the pair at the top of O and turn that into a bond?
    (1 vote)
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  • starky sapling style avatar for user Oleander
    Is there anyway I can extract/make Formaldehyde at home?
    I understand it is naturally occurring and wondering if there would be a way to safely extract it for use with wet specimens.

    Sorry if this is a silly question but I cannot buy where I live.
    (1 vote)
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    • starky ultimate style avatar for user ++§ Αλεκσανδαρ
      For what exactly do you need formaldehyde?
      It's a fairly reactive specie that can be easily oxidized, and in its pure form it's not chemically stable. Commercial formaldehyde solutions always have some methanol added to increase their stability. Knowing that, formaldehyde can not be extracted as such, but it can be prepared from appropriate source. Probably the easiest way to prepare formaldehyde is by oxidation of methanol.

      That being said, if you don't know how to prepare formaldehyde, you are officially not competent enough to do it. In NFPA 704 standard, formaldehyde has the highest score in the health hazard category. It is harmful even at low concentrations and short exposure. Without appropriate safety measures and experience, i'd say your chances of killing yourself are pretty high.

      I recommend you do not attempt to synthesize it, or purchase it.
      (1 vote)

Video transcript

- [Instructor] What we're gonna do in this video is get a little bit more practice constructing Lewis diagrams, and in particular, we're going to try to construct the Lewis diagram for formaldehyde. Formaldehyde has one carbon, two hydrogens, and an oxygen, CH2O. So pause this video and have a go at it. Try to construct a valid Lewis structure, or a Lewis diagram for formaldehyde. All right, now let's do this together. Now the first step, and we saw this in a previous video, we want to think about all of the valence electrons for this molecule. So we want to account, account for the valence electrons. Now the reason why we wanna do that is so that while we're trying to create this structure, we are making use of all of the valence electrons. And to figure out how many total valence electrons we have, we can look at a periodic table of elements. We can see that carbon, it's in that second row, in that second period, so its second shell is its outer shell. And in that shell, it has one, two, three, four valence electrons. So, carbon has four valence electrons. A neutral free hydrogen atom is going to have one valence electron, but we have two of them here, so it's gonna be two times one. And then, oxygen, it also is in the second period, and in its second shell it has one, two, three, four, five, six valence electrons. And so the total valence electrons in this molecule are gonna be four plus two, which is six, plus six, which is equal to 12 valence electrons. Now the next step is to try to draw a structure. Try to draw, draw single bonds, I'll say, single bonds. And a key question is, what do we think is going to be our central atom? And the rule of thumb is the least electronegative atom, that is not hydrogen, is a good candidate for our central atom. So we can rule out hydrogen. So between carbon and oxygen, we know that oxygen is one of the most electronegative atoms, well one of the most electronegative elements on the periodic table of elements. It's very close to fluorine. And so carbon is a good candidate for the central atom. So let's put the carbon right over here, and then let's put these other atoms around it. We could call them terminal atoms. So, let's put our oxygen right over there, and then we have two hydrogens. Hydrogen there, a hydrogen there. And let me draw the bonds. So that's a single bond. That accounts for two valence electrons. That accounts for two valence electrons. That accounts for two valence electrons. So I've just used two, four, six valence electrons. So if I subtract six valence electrons, I am now left with six valence electrons, six valence electrons. So the next step is allocate the remaining valence electrons, trying to get to the octet rule for atoms that are not hydrogen, and then for hydrogen, trying to get it to have two valence electrons. So allocate, allocate the remaining, remaining valence electrons. All right, so let's start with this oxygen. This oxygen already has these two electrons that it's sharing hanging around. So in order to get to the octet rule, it needs six more. So let's give it six electrons. So, one, two, three, four, five, six. Well I've just used up the remaining six valence electrons. So I don't really have any more to play with, but let's see how the other atoms are feeling. So hydrogen here, it's able to share these two electrons that are in this covalent bond, so it's feeling good. It can kind of pretend that it has a full outer shell, 'cause its outer shell is just that one, that first shell, that's filled with two electrons. Same thing for this other hydrogen. So at least the terminal atoms, the oxygen and the two hydrogens, are feeling like they have a full outer shell. But then in the fourth step, we're going to look at our central atom. So, let's focus on the central atom, central atom, and do we need more bonds, or do we need to do something interesting here? And what we see is that carbon, it's able to have two, four, six electrons hanging around it, but it would love to have eight. Carbon would love to have a full outer shell, so how could we do that? Well, we could add more bonds. Where could the bonds come from? Well it would come from some lone pair of electrons. Well the only lone pairs of electrons are hanging around this oxygen. So what if we were to take, say, this lone pair of electrons, and then construct another covalent bond with that? Then, our Lewis diagram will look like this. I will actually redraw it. So you have your carbon, you have your three original covalent bonds, you had a hydrogen, a hydrogen, and then you had your oxygen, right over here, and now we've formed a new covalent bond, just like this, and then you have these two other lone pairs around the oxygen. So let me draw that. So, two, then another two around the oxygen. And this is looking pretty good, because the oxygen, it still has eight electrons hanging around, four in lone pairs, and then four, they're in this double bond that it is sharing. The hydrogens still have two electrons hanging around. They're able to share the electrons in each of these covalent bonds. And now the carbon is participating in, you could think of it as four covalent bonds, two single bonds and one double bond, so each of those have two electrons associated with it, so it has eight electrons hanging around. So this is looking really good as a legitimate Lewis structure, or Lewis diagram for formaldehyde.