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Verified Solution by Proprep Tutor
What is SCI2 Lewis structure?
The Lewis structure of SCl2 (sulfur dichloride) can be determined by following a systematic approach.
This process involves understanding the valence electrons of sulfur and chlorine, arranging these electrons to satisfy the octet rule, and ensuring the correct molecular geometry.
This process involves understanding the valence electrons of sulfur and chlorine, arranging these electrons to satisfy the octet rule, and ensuring the correct molecular geometry.
Here's a detailed, step-by-step guide to drawing the Lewis structure of SCl2:
Step 1: Counting Valence Electrons
- Sulfur (S) is in Group 16 of the periodic table and thus has 6 valence electrons.
- Chlorine (Cl) is in Group 17 of the periodic table, so each chlorine atom has 7 valence electrons.
- Since there are two chlorine atoms in SCl2, the total number of valence electrons from chlorine is 2×7=14.
- Adding these together, the total number of valence electrons available for the SCl2 molecule is 6+14=20 electrons.
Step 2: Drawing the Skeleton Structure
- Place the sulfur atom in the center as it is less electronegative than chlorine.
- Connect each chlorine atom to the sulfur atom with a single bond.
- Each single bond accounts for 2 electrons.
Step 3: Distributing Remaining Electrons
- After forming single bonds, 20−2×2=16 valence electrons remain.
- Distribute these electrons around the chlorine atoms first to complete their octets.
- Each chlorine atom needs 6 more electrons, so 12 electrons (or 6 pairs) are used here.
Step 4: Completing the Octet for Sulfur
- After assigning electrons to chlorine, 4 valence electrons remain.
- Place these remaining electrons on the sulfur atom.
Step 5: Checking the Octet Rule
- Each chlorine atom now has 8 electrons around it, satisfying the octet rule.
- Sulfur, having 6 electrons in its valence shell, can expand its octet.
- In SCl2, sulfur has 8 electrons around it (2 from each bond and 4 non-bonding electrons), satisfying the expanded octet rule.
The Lewis Structure for sulfur dichloride (SCl2):
- The final Lewis structure of SCl2 shows sulfur in the center with two single bonds to two chlorine atoms and two lone pairs of electrons on the sulfur atom.
This Lewis structure of SCl2 demonstrates the molecule's valence electron arrangement, satisfying both the octet rule and formal charge considerations, ensuring a stable molecular structure.
Lewis Theory and Covalent Bonds
Strategy for Drawing Lewis Structures
Exercise 3 - Is SCl₂ a polar or nonpolar molecule?
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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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