Write Lewis structures for the following: (a) O2 (b) H2CO (c) AsF3 (d) ClNO (e) SiCl4 (f) H3O (g

Write Lewis structures for the following: (a) O2 (b) H2CO...

Key Concepts

Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms in a molecule and the lone pairs of electrons that may exist in the molecule. They are created by using the valence electrons of the atoms to form bonds and depict the electrons involved in those bonds as well as any non-bonding electrons. These diagrams are essential for visualizing the arrangement of electrons which influences the properties and reactivity of the molecule.

Valence Electrons

Counting valence electrons is a crucial step in constructing Lewis structures, as it determines the total number of electrons available for bonding and for filling lone pairs. Each atom contributes a specific number of electrons based on its position in the periodic table, and the sum of these electrons is used to distribute between bonds and non-bonded regions in such a way that the octet rule (or duet rule for hydrogen) is satisfied where applicable.

Octet Rule

The octet rule is a fundamental guideline used when drawing Lewis structures, stating that atoms tend to form bonds in such a way that they attain a stable configuration that resembles that of the noble gases, usually having eight electrons in their outer shell. There are exceptions to this rule for atoms in lower periods or those capable of expanding their octet by using available d-orbitals.

Formal Charge

Formal charge is a bookkeeping method used to estimate the distribution of electric charge within a molecule, helping to determine the most stable Lewis structure. It is calculated by comparing the number of valence electrons in an isolated atom with the number of electrons assigned to it in the Lewis structure, with lower or zero formal charges generally indicating a more favorable structure.

Bonding and Molecular Geometry

Understanding bonding and molecular geometry is key to interpreting Lewis structures because the arrangement of atoms and electrons influences the shape of the molecule. While Lewis structures provide a two-dimensional representation of electron distribution, the three-dimensional geometry, which can be predicted using theories like VSEPR, determines the actual spatial arrangement of the atoms, affecting properties such as polarity and reactivity.

Transcript

00:02 Hi there.

00:03 In this problem, we are doing some loose dot structures for some molecules and some polyatomic ions.

00:10 When doing dot structures, there are just some basic steps you want to follow.

00:14 The first thing you want to do is determine the total number of valence electrons that you have to work with because that will determine how many dots are going to be in your finished structure.

00:28 Second thing you will do then is you know that these atoms have to be bonded together through at least one pair of shared electrons.

00:39 So you would add a pair of electrons for each bond.

00:53 Right.

00:53 And then what you want to do is complete the octets for each of the atoms.

01:00 Remember an octet means that each atom needs to have eight dots around it.

01:06 With hydrogen being an exception, hydrogen only needs two.

01:10 And if you run out of dots before you complete that octet, then you will have to make double or triple bonds as necessary.

01:17 So let's go ahead and get started.

01:18 Letter a, we have oxygen, o2.

01:23 Oxygen's in group six on the periodic table, so each oxygen has six valence electrons.

01:28 So we have a total of 12 electrons to place in this structure or 12 dots.

01:35 So if we have oxygen bonding to oxygen, they said we need to at least have a pair to bond them together, and then we start adding more dots.

01:50 At this point, i have added 10 dots, and each oxygen still needs two more electrons.

01:57 So if i put them on this one, then the oxygen on the right doesn't have enough.

02:02 And the same would be true if i reverse that, put those two extra dots on the oxygen on the right.

02:07 What this tells us, and we can think of these kind of like a venn diagram, what that tells us is those dots need to be in the intersection.

02:16 In other words, there is a double bond.

02:19 And a moment ago, i mentioned a venn diagram.

02:22 Remember in a venn diagram, the things that are in the intersection of the two circles are held in common.

02:32 So if we look here in this first circle, there are two, four, six, eight dots.

02:38 In the second circle, there are two, four, six, eight dots.

02:42 So we have satisfied the octet rule.

02:45 And this is the dot structure.

02:50 All right, let's move on to letter b.

02:53 In letter b, we have h2 co3, or co, rather, just h2co, my bad.

03:03 So, number of dots.

03:04 Each hydrogen has one, because it's in group one, carbon has four, oxygen has six.

03:11 Adding these together gives us a total of 12 dots in our finished structure.

03:17 Carbon.

03:18 When we have carbon in a compound, it's understood that it's typically in the center.

03:23 So i'm putting the carbon in the center and the other three elements around it.

03:29 We know the hydrogens have to be bonded.

03:31 And now the hydrogens have as many dots around them as they're allowed.

03:36 All right.

03:36 So i know i need a bond there to attach the carbon to the oxygen.

03:41 So so far i've used six of my dots.

03:43 I have six more remaining.

03:46 So if i start putting these on the oxygen, i i end up where i only have two dots remaining, and both carbon and oxygen want these dots.

03:55 So again, that tells me i must have a double bond.

03:59 And this is the completed dot structure.

04:03 Go back and circle the oxygen here.

04:06 All right, let's move on to letter c.

04:11 Letter c is asf3.

04:16 First thing we want to do is determine the total number of dots allowed to us.

04:20 As is in group 5 on the periodic table.

04:24 Fluorine is in group 7, but there are three of those.

04:26 So three times seven is 21.

04:29 Adding these together, i have 26 dots for this structure.

04:35 The single element is in the center with the fluorine surrounding it.

04:42 We know that we need to have at least one pair of electrons to form the bond.

04:47 And then we go about placing in the rest of our electrons to complete the octets.

04:56 So far i have added 24.

04:59 I have two dots remaining, and they are just going to go right there on the as.

05:03 So this is my final structure.

05:06 Everything has eight dots around it, and we have all single bonds.

05:12 In letter d, we have cl, n.

05:21 Cl has seven dots, n has five, o has six.

05:26 And again, this is just the group that they're in on the periodic table.

05:30 Adding these together gives me 18 dots to work with.

05:36 All right, so let's go ahead and get started here.

05:41 C -l -o, o, the bomb there.

05:46 I am going to complete the octet around chlorine.

05:49 Our chlorine has seven dots of its own, so it is only going to share.

05:54 It's typically only going to form single bonds.

05:56 It's only going to share one more electron.

05:59 All right, so i have used 10 dots already.

06:03 12, 13, 14, 15, 16.

06:07 I only have two dots remaining.

06:10 We're going to need another.

06:11 Double bond, it's going to be between the nitrogen and the oxygen.

06:15 Remember, oxygen has six valence electrons of its own.

06:19 So if it shares two dots or two electrons, it forms that double bond and ends up with eight valence electrons.

06:30 Moving on to letter e.

06:34 Letter e is s -i -c -l -4.

06:40 S -i is in the same group as carbon.

06:43 So it has four valence electrons.

06:45 Each cl has seven...

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