Reading Notes for Chapter 7


These are Dr. Bodwin's reading notes for Chapter 7 of "Chemistry 2e" from OpenStax. I am using a local .pdf copy that was downloaded in May 2020.

Chapter Summary:

The vast majority of what we call "chemistry' is taking place in the electron cloud of atoms, and in most cases it's just the outermost part of that electron cloud, the "valence" electrons. This chapter looks at how atoms and ions interact with one another to form "bonds". This is also one of the most "visual" chapters in that we will be looking at 3-dimensional structure of molecules and ions, so for those of you who are very visual, this should be an exciting journey into that realm and away from some of the math-heavy topics we've looked at recently.

Ionic Interactions ("Ionic Bonding"):

A lot of the material in this short section of your text is review of some of the concepts we looked at back in Chapter 2. Read through closely, though... now that you've been working with this material for a while you will likely catch a few details that you might have missed before.
One thing I definitely want to point out here may seem like a very picky detail, but let me explain... Your textbook (and most textbooks) use the term "Ionic Bonding". When most people see the word "bond", they picture a little string that holds 2 atoms together. For compounds that are purely ionic there is no little string, the positive and negative ions are attracted to each other because "opposites attract". For that reason, I prefer to use the term "ionic interaction". Another reason I like "ionic interaction" is that it's a better description over a wider range of distances. It's hard to call anything longer than a few Angstroms a "bond", but charges can interact (although weakly) over longer distances.
When talking about the interaction between charged particles, we'll often describe the interaction using Coulomb's Law or describe the interaction as a "Coulombic interaction". There are 3 main points to Coulomb's Law when it comes to atomic-scale interactions:
  1. Opposite charges attract, like charges repel
  2. Bigger charges attract or repel more than smaller charges
  3. The attraction or repulsion is inversely proportional to distance; the farther apart the charges are, the smaller the attraction or repulsion.
Although we talk about bonding as being either ionic or covalent, it's actually a spectrum to some extent. Additionally, it is very common to have both covalent bonds and ionic interactions in the same "molecule", so although it is convenient to think of these as very distinct anf discrete ideas, they tend to blend together fairly quickly.

For a very "quick and dirty" visualization of ionic vs. covalent bonding, check out this video:
https://www.youtube.com/watch?v=QqjcCvzWwww

When elements that are far apart on the Periodic Table (like sodium and chlorine) react to form a compound, they are more likely to form ionic compounds. When elements are far apart on the P.P., it usually means that one of them has a (relatively) low ionization energy and the other has a (relatively) strong electron affinity.

Covalent Bonding:

Electronegativity is similar to electron affinity, but read the definitions closely. Electron affinity is the change in energy when an electron is added to an atom or ion; electronegativity is a measure of how strongly an atom attracts electrons. Electron affinity is an energy while electronegativity is an arbitrarily defined scale. That means electronegativity is really only useful or meaningful when we use it compare two elements. Because of this, electronegativity gives us one way to differentiate between ionic interactions and covalent bonds.
Polar covalent bonds describe the interactions that taek place betweeen the ionic and covalent extremes.

Electron Accounting:

Since most chemistry takes place in the electrons, and most of that in the valence electrons, it's good to have a quick, visual way to keep track of these electrons.
G.N.Lewis was a big fan of electrons. We'll mention him a few times when electrons are the focus.
Lewis Symbols of individual atoms and ions are a good starting point, but they're a bit limited. Check out Table 7.9 and the material around that Table. It's a really good start before we dive in to...
Lewis Structures of Covalently Bound Substances - There are a couple different ways to approach Lewis Structures. I prefer a modified version of the "electron counting" approach. The textbook uses a more classical "electron counting" approach that relies upon the octet rule.
The reason I prefer a modified version is because it gives a few different ways to check your work and catch mistakes before going through the whole process. The steps I use are described here:
https://chemistryingeneral.blogspot.com/2010/11/lewis-structures.html
Formal Charge is a way to "assign" electrons to different nuclei in a molecule or polyatomic ion. The textbook provides a "formula" for calcualting formal charge. If you're a more visual person, you can also determine formal charge more visually using this process: Determining Formal Charge
Lewis Structures take practice. Lots and lots of practice...

Visualizing Electron Distributions:

Valence Shell Electron Pair Respulsion Theory (VSEPR)
Opposite charges attract, like charges repel. This is always interesting when we think about atoms because all the positive charge is pocked together in the nucleus and all the negative charge is in the "electron cloud".
VSEPR tells us that electrons around a nucleus try to get as far apart as possible to minimize the repulsion between those electrons. The first step to a 3-dimensional VSEPR structure is drawing a good Lewis structure, so practice, practice, practice!
Here's another bit of info about Lewis Structures and VSEPR:
http://www.drbodwin.com/teaching/genchemlab/c150L2012d10vseprslides1.pdf
If you're looking for some lists of structures to practice, here are some good ones:
http://www.drbodwin.com/teaching/genchemlab/c150L2012d10vsepr1.pdf

Bond Strengths, Structures, and Applications:

The Born-Haber Cycle is an excellent example of the power of Hess's Law. The textbook also uses a nice reaction coordinate diagram in Figure 7.13 to visually show how Hess's Law is used to determine an un-measurable (or difficult to measure) quantity.

Resonance structures - in some structures (like CH4 or CO2, for example), all of the bonds are the same. In other structures (like SO4-2 or PO4-3), we typically draw Lewis structures that have some single bonds and some double bonds between the same type of atoms. In SO4-2, there are two single bonds and two double bonds... how do you decide which ones to make double bonds? In a case like this, it doesn't matter, and in fact, a Lewis structure is just an "instant in time" snapshot picture... in reality, all 4 bonds are the same and they are half way between single and double. This is resonance; two or more reasonable Lewis structures that have the same arrangement of nuclei, but different arrangements of electrons.
Resonance is why we talk about "regions of electron density" when thinking about VSEPR... in many cases, the bond order doesn't matter, all that matters is that there is some type of a bond there.
EXCEPTION: just because resonance exists, don't fall too head-over-heels in love with it. A molecule like COF2 (carbon is the central atom, double bond to the oxygen) does not exibit resonance structures with a double bond to the fluorine and a single bond to the oxygen... check the formal charges for some good reasons why this is not a favorable structure. It's only when similarly "good" Lewsi structures are possible that resonance is important.

Deviations from "ideal" VSEPR geometry - The bond angles in "ideal" VSEPR shapes are a great starting point, but there are many things that can nudge them around a bit. Lone pairs "take up more space" than bonding pairs. "Big" atoms take up more space around a central atomn. Shorter bonds take up more space than longer bonds. These are all subtle effects that come up pretty regularly, but focus on drawing good Lewis and VSEPR structures before you dig too deeply into the mroe subtle details.

Polarity - Between pure covalent bonds and pure ionic interactions, there's avast spectrum of "polar covalent" bonds in which the electrons are shared, but not shared equally. These bond polarities become ver important when looking at molecular structure and physical properties. A polar molecule or polyatomic ion is one in which the individual bon polarities give rise to an overall polarity of the molecule. For example, F2 is a non-polar molecule because the fluorine atoms both attract the bonding electrons equally - this is a pure covalent bond. What about H-F? The fluorine attracts electrons mroe strongly than the H, so the electron cloud is pulled toward the F, leading to a slightly negative F end of the molecule and a slightly positive H end of the molecule. A polar covalent bond leading to a polar molecule!
But what happens when we go to more than diatomic molecules? Try CO2. Both bonds are polar, pulling electron clouds toward the oxygen, BUT the two oxygens pull in exactly opposite directions with exactly the same force, so although the bonds are polar, the molecule is not Next up: water... Water is "bent", and both bonds are polar, pulling electron clouds toward the oxygen. But now, these forces don't cancel out, so the water molecule is polar with the oxygen side being slightly negative and the hydrogen side being slightly positive.
If you have experience with vectors, you can think of "polar bonds" as a dipole vector with the arrow pointing toward the negative end of the bond.

Another nice application of Lewis structures is when we want to find the oxidation number of an atom in a molecule or polyatomic ion. Finding oxidation numbers is very similar to finding formal charge, the only difference lies in how we assign electrons. Formal charge treats all bonds as if they were purely covalent and the bonding electrons are equally shared between the bonding atoms. Oxidation number treats all bonds as if they are purely ionic and all electrons are assigned to the more electronegative side of the bond. In both cases, these are just ways for us to think about the electrons, they are not concepts that rigidly define the actual physical structure of the moelcule or polyatomic ion.


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