MOLECULES Part 1: LEWIS THEORY & BONDING

Lewis structures show activity of valence electrons in chemical bonding.

LEWIS THEORY

^What to catch from above slide:

Valence electrons are outermost electrons (see image below)
Ionic bond refers to TRANSFER of electrons from one atom to another
Covalent bond refers to sharing of electrons between two atoms
The reason for ionic and covalent bonding is for each atom to achieve the octet (= 8 electrons in the outermost shell)

^What to catch from above slide:

Lewis symbol consists of two parts: 1) chemical symbol (e.g. C, Na, F, etc.) and 2) valence electrons as represented by DOTS
Rule to remember: note that you must draw one dot all around the chemical symbol before you pair them up (this is sort of in relation to Hund's rule)

^What to catch from above slide:

For ionic bonds (transfer of e-), Lewis structure is drawn using square brackets on each atom, and its corresponding charge.

^Lewis structure for sodium chloride (NaCl)

For covalent bonds (sharing of e-), Lewis structure is drawn combining both atoms together by their valence e-

^Lewis structure for hydrogen chloride (HCl)

See more examples of Lewis Structures

^Barium oxide (BaO)

^Magnesium chloride (MgCl2)

^Aluminum oxide (Al2O3)

NOTICE THAT ALL IONIC BONDS WILL INVOLVE METALS WITH NO VALENCE ELECTRONS BECAUSE THEY HAVE GIVEN THEM AWAY TO NON-METALS

^What to catch from above slide:

As we know already, covalent bond is represented by a Lewis structure that combines two or more atoms together by their valence e-
The reason for the bond is so that each atom can achieve the octet (= 8 electrons that fully fill the energy shell to achieve stability)
Bond pair = shared (or bonded) electrons between two atoms
Lone pair = non-bonded electrons

^What to catch from above slide:

Coordinate covalent bond = covalent bond where a single atom gives both electrons to a shared pair (see ammonia and water as examples)
The reason ammonium and hydronium ion both have a positive charge is because the "+" charge was there in hydrogen (H+) to begin with.

^What to catch from above slide:

Double bonds and triple bonds represent multiple covalent bonds
This happens when a single bond cannot fulfill the octet rule for each of the atoms

^diatomic oxygen has a double bond, which represents two covalent bonds (4 shared electrons)

^What to catch from the above slide:

Polar covalent bond = covalent bond where electrons are NOT SHARED EQUALLY between 2 atoms
Now the question is "what do you mean not shared equally?" The answer is on electronegativity. A more electronegative atom will HOG the electron density (or the negative charge) towards itself more than the less electronegative atom (e.g. CO molecule would have oxygen hog more of the electron density; likewise, HCl molecule would have chlorine hog the electron density more)
The analogy I like to give is sitting on a bench with your big bro. Although you might be sharing the bench, your big bro who is clearly bigger than you will take up more of the bench than yourself.

^Now a little gist about electronegativity:

Electronegativity (EN) = the degree to which an atom likes and attracts an electron
As we saw in our previous blog post (See "ATOMS #4: PERIODIC TABLE), electronegativity increases right and up the periodic table. In other words, EN increases with increasing non-metallic properties, and decreases with more metallic properties

^what to catch from the above slide:

The difference in EN between 2 atoms will determine the bond character (whether they are ionic, polar covalent, or covalent).
As we can see in the graph, the higher the difference, the closer it gets to IONIC CHARACTER. The lower the difference, the more COVALENT it becomes.

^What to catch from above slide:

It shows three different molecules, HCl, HBr, and HI
The EN difference is biggest in HCl (with the value, 0.9), thus it exhibits the most IONIC CHARACTER out of these 3 molecules.

PRACTICE DRAWING BASIC LEWIS STRUCTURES

ANSWER KEY

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