From Bohr's to Lewis'
1. Review
Let's start this section by reviewing Bohr's models. To do this, we will work together through the following worksheet:
| a3_electrons_worksheet.pdf |
As you complete this sheet, take a look at the very last layer of the atoms you are building. How many electons are there in each layer? Write it down near the model.
2. Simplifying Bohr's Models
As mentioned before, there isn't a chemist in the world who will create formulas using Bohr's models. It's boring and tedious after a while, specially if you get on the bigger atoms.
We also mentioned that the only layer that really matters for us is the last one (or, the last ones, depending on the atoms we're talking about). The inner layers? Not so much.
On ionic and covalent bonds, it's really easy to know which layers matter to us. Do you remember valences? They are located in the last layer in an atom.
The GROUP indicates how many valence electrons it has!
We also mentioned that the only layer that really matters for us is the last one (or, the last ones, depending on the atoms we're talking about). The inner layers? Not so much.
On ionic and covalent bonds, it's really easy to know which layers matter to us. Do you remember valences? They are located in the last layer in an atom.
The GROUP indicates how many valence electrons it has!
We can tell how many electrons are on the last layer of each atom found on groups 1, 2, 13, 14, 15, 16, 17 and 18, just by looking at the table above. I mean, Mendeleev had to walk over 2000 kilometres on the snow for us to have this table, might as well use it as he intended, right?
Let's think of that last layer, since it is the most important and there isn't a lot of variation regarding charges. Most atoms want to have 8 electrons on their last layer; this is called the octate rule.
Gilbert Lewis, a chemist from 1916, had this idea of simplifying the valence representation so it would actually make sense to people.
Since the goal is to have (usually) eight electrons for each atom, Lewis thought of representing a noble gas (super happy atoms) by adding 8 dots around. The dots represent their last layer, or valence; we can easily see there is no more space for sharing.
Let's think of that last layer, since it is the most important and there isn't a lot of variation regarding charges. Most atoms want to have 8 electrons on their last layer; this is called the octate rule.
Gilbert Lewis, a chemist from 1916, had this idea of simplifying the valence representation so it would actually make sense to people.
Since the goal is to have (usually) eight electrons for each atom, Lewis thought of representing a noble gas (super happy atoms) by adding 8 dots around. The dots represent their last layer, or valence; we can easily see there is no more space for sharing.
Helium only has two dots because -- just like Hydrogen -- it is an exception to the rule, as there are no other electrons on an inner layer.
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Calcium, on the other hand, has two dots just like Helium -- but it is not happy. It is part of group 2 and has electrons on an inner layer. It needs eight electrons to stabilize.
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Let's understand how this works by completing a worksheet!
| explanation_valence_and_lewis.pdf |
3. Covalent Lewis Dot Structures
Bonds: Covalent bonds are represented by lines. Each line represents two shared electrons. If you see two lines between two atoms, that is called a double bond (4 shared electrons). Three lines indicates a triple bond (6 shared electrons).
However, there is no such thing as a quadruple bond (atoms would never need to share 8 electrons between them!)
Non-Bonding Electrons: Electrons that do not participate in bonding are represented by dots. Most of the time, non-bonding electrons exist in pairs. A pair of nonbonding electrons is called a lone pair.
However, there is no such thing as a quadruple bond (atoms would never need to share 8 electrons between them!)
Non-Bonding Electrons: Electrons that do not participate in bonding are represented by dots. Most of the time, non-bonding electrons exist in pairs. A pair of nonbonding electrons is called a lone pair.
Let's imagine a molecule, such as POCl. That is harder to build than a simple pair, and we don't know for sure which atom to put in the centre. Here are some steps you can follow to figure it out:
Step 1: Determine the central atom.
Rule 1. The central atom is usually the minority. In the formula for water, for instance, the oxygen is alone, while there are two hydrogens. So, we place the oxygen in the middle. Since our molecule does not have a minority (we only have one atom of each element), then let's go to...
Rule 2. The central atom is the least electronegative element.
HOLD UP! What the hell is electronegativity?
That is just the ability an atom has to "suck in" another atom's electron. So, obviously something that is very light and has 7 electrons on the last layer would be super efficient, while something that is very heavy and has 1 electron on the last layer would not be so good at it.
This is best seen as a visual thing which we can see on the table below. The lowest and leftest (is that a word?) is the most electronegative. Everything goes up as you move to the right and up.
Rule 2. The central atom is the least electronegative element.
HOLD UP! What the hell is electronegativity?
That is just the ability an atom has to "suck in" another atom's electron. So, obviously something that is very light and has 7 electrons on the last layer would be super efficient, while something that is very heavy and has 1 electron on the last layer would not be so good at it.
This is best seen as a visual thing which we can see on the table below. The lowest and leftest (is that a word?) is the most electronegative. Everything goes up as you move to the right and up.
So phosphorous should be to the middle, while the other atoms should be one on each side.
Step 2: Add dots to represent the valence layer.
Think like this: "Phosphorous (in the middle) needs 3 electrons. He can share two with oxgen who needs two, and one with chlorine who needs one."
Step 3: Make bonds to reach 8 electrons, using lines.
4. Now practice!
| covalent_bonds_practice_packet.pdf |
| ionic_compounds_practice_packet.pdf |
