5 - Covalent Bonding

Why Do Atoms Fall In Love?

We chemistry teachers have gone to great lengths to explain covalent bonding to students.

• "Valency" as a property that atoms "just have."

• Atoms are "satisfied" when they have a "full octet."

• Fluorine "wants" one more electron "attain a full octet."

• Magnesium "wants" to "donate" its two valence electrons

You've tried them all.

• Yes, Eddie, atoms do have desires that can be satisfied.
• Yes, Molly, fluorine can count up to eight.
• No, Sam, we can't explain why sulfur has four different valencies.

Stronger students still ask "why" as weaker students shake their heads.

Meanwhile, we teachers ignore the 50 chemical bottles filled with real exceptions on the shelves of our chemical prep rooms.

Well, now we can just make all of the silliness go away.

1. Chlorine Atoms Strongly Attract Electrons

All of chlorine's valence electrons are strongly attracted. We know why: they are quite close to a 7+ core charge. Kids can easily intuit that being close means stronger influence, and that a big cause has a big effect.

There is one empty space left in chlorine's valence that can accommodate one more electron. This means that chlorine can strongly attract and hold other atoms' electrons, too.

When you think of it, all of the active non-metals are similar: large core charges, small radii, and 1, 2, 3 or 4 vacancies in the valence shell.

Now - consider the unpaired electron, adjacent to the vacancy.

2. What Could Be Better For These Electrons?

Each unpaired electron is very close to a very strong core charge.

These electrons are strongly attracted to the core, and therefore strongly bound to the chlorine atom. This is a favorable position for these electrons. 

What could be better than being THAT CLOSE to a 7+ core charge?  

3. Being THAT CLOSE to Two 7+ Core Charges!!

Chlorine A strongly attracts the unpaired electron from Chlorine B into the vacancy in its valence shell, and vice versa. Both unpaired electrons are now attracted to both 7+ core charges.

Note that this "shared pair" of electrons is privileged: they  are attracted to two core charges. They are therefore even more strongly bound than the other valence electrons. And, they are even more localized than other valence electrons.

More advanced theories would tell us that the shared pair is in a lower energy state. 

3. Non-metals Form Covalent Bonds 

All non-metals have larger core charges, and smaller radii, than metals and metalloids.

Large core charge and small radius cause high electronegativity, and this in turn causes strongly held, localized bonding pairs of electrons.

Note that this arises naturally from the model. Teachers do not have to pretend that the atoms have wants, needs, or the ability to count to eight.

In the example of nitric acid, to the left, you can see the possibilities of bond polarity, formal oxidation states, and even acid-base behavior. 

 Download this exercise to get a better idea of the scope of this model.

Try this exercise yourself. Then - Let your students do it. You will be surprised how much they can do!

Since ChemStudy 1963

ionic bonding was always first.

You know the drill.....

1. Memorize main block elements
2. Lewis Dot diagrams.
3. Believe that fluorine "wants" this, sodium "wants" that.
4. Na and F do opposite things to obtain the holy grail: the "full octet."
5. Memorize a list of "ionic valencies," some of which violate (4).
6. Memorize math that looks like lowest common denominator. Or...
7. Follow the Criss-Cross rule. And stop thinking.

We can do better.

Lesson 1 - Three salient features of each atom across the periodic table.

Lesson 2 - Students can predict electron attraction, with nearly no instruction!

Lesson 3 - Take a closer look at the elements of the second row.

Lesson 4 - Origins of the Ross model of the atom

Lesson 5 - Covalent bonding

Lesson 6 - Ionic Bonding

Lesson 7 - A post-modern model of science.

Lesson 8 - Learning with IntuitivScience.