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6.05a Determine the formal charg on atoms in NH4+

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化学

肯塔基大学

4.7（666 個評分） | 67K 名學生已註冊

This course is designed to cover subjects in advanced high school chemistry courses, correlating to the standard topics as established by the American Chemical Society. This course is a precursor to the Advanced Chemistry Coursera course. Areas that are covered include atomic structure, periodic trends, compounds, reactions and stoichiometry, bonding, and thermochemistry.

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從本節課中

Week 6

This unit looks more in depth at molecular compounds to see how they are bonded together by looking at Lewis structures. These structures provide information about the types of bonds (single, double, or triple) as well as the connectivity of atoms. By knowing the Lewis structure, we can also predict the three-dimensional geometry of an individual molecule.

6.01 Lewis symbols 4:36

6.02 Covalent bonds 5:19

6.03 Electronegativity 11:21

6.04 Lewis Structures Part 1 13:11

6.04 Lewis Structures Part 2 18:49

6.04a Draw the Lewis structure for H2O 1:13

6.04b Draw the Lewis structure for H2CO 2:00

6.04c Draw the Lewis structure for CCL4 1:21

6.04d Draw the Lewis structure for NH3 1:16

6.05 Resonance & Formal Charge 16:27

6.05a Determine the formal charg on atoms in NH4+ 2:24

6.05b Determine the formal charge on atoms in H2CO 3:17

6.06 Exceptions to the Octet Rule 8:48

6.06a Draw the Lewis structure for I3- 1:16

6.06b Draw the Lewis structure for ClF4 1:41

6.06c Draw the Lewis structure for XeF4 1:23

6.07 VSEPR Part 1 11:01

6.07 VSEPR Part 217:07

6.07a Determine the electron pair and molecular geometries for I3- 1:42

6.07b Determine the electron pair and molecular geometries for NH3 1:22

6.07c Determine the electron pair and molecular geometries for ICl4- 1:57

6.07d Determine the electron pair and molecular geometries for PF5 1:52

6.07e Determine the electron pair and molecular geometries for XeF2 2:16

教學方

Dr. Allison Soult
Lecturer
Dr. Kim Woodrum
Sr. Lecturer

腳本

Determine the formal change on atoms in NH4 plus.

Before I can determine the formal charge,

I must first have a Lewis structure for my molecule, and

in order to draw the Lewis structure I must have the number of valence electrons.

So, nitrogen has five valence electrons, plus four hydrogens,

each with one valence electron, and because I have a plus charge which

indicates the loss of an electron, I'm going to subtract one from that.

And so what I see is that I have 8 electrons in my Lewis structure.

Because that's the total number of valence electrons.

So I draw my nitrogen in the middle, because hydrogen must be a terminal atom.

And I'm going to draw everything with single bonds.

Now I see that in my structure I've used up eight electrons,

which is all the electrons I had.

Draw my brackets and my plus sign just so I have a correct structure.

And that's okay, because, remember, we don't need any additional electrons around

hydrogen it only wants two, so it's perfectly fine that way.

Now I want to look at my formal charges, so I'm going to look at my hydrogen first.

I see that, in an isolated atom of hydrogen,

that it would have one electron assigned to it.

I'm going to subtract off the number of lone pair or

lon, non-bonding electrons a sign to it in this structure which is zero and

minus half of the bonding electrons.

So I get a formal charge of 0, on each of my hydrogen atoms.

Because they all have the same structure within this molecule.

Now I'm going to look at my nitrogen.

And I know that nitrogen has five electrons in the isolated atom.

Here, it has no non-bonding electrons, and it has 8 electrons around it.

So, half of eight equals four.

So, five minus four is one.

And so our nitrogen has a formal charge of plus one.

And this is okay, because remember that the sum of our formal charges,

must equal the sum of the charge on the particular substance.

Here we have a plus one charge, so

the sum of our formal charges should equal plus one.

And we can't really get in any lower because we have to

have that plus 1 charge.

We do have lots of zeros, which indicates some stability in this structure.

In general we're looking at formal charges,

to compare whether something is a better structure than another.

We have to have two different structures, and

compare the sets of formal charges with each.

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