Reading Notes for Chapter 6

These are Dr. Bodwin's reading notes for Chapter 6 of "Chemistry 2e" from OpenStax. I am using a local .pdf copy that was downloaded in May 2020.

Chapter Summary:

Almost all of what we call "chemistry" takes place in the electrons. Protons in the nucleus are critically important because they define the identity of the element; neutrons are important for a number of reasons, the most basic of which is to define (with protons) the mass of an atom, but electrons are involved in almost all chemical reactions and transformations and they help us make a variety of predictions about physical properties, microscopic shapes, reactivity, and stability. Because electrons are so important to almost all of chemistry, it is important that we look more closely at them and develop ways to "address" the individual electrons in an atom, ion, or molecule.

Electromagnetic Energy:

There are a wide variety of energy forms, but it is important to realize that (for the most part) these are all just different regions of the same "big picture", the electromagnetic spectrum. The microwaves that heat up your lunch, the ultraviolet rays that cause damage to your skin, and the light that lets you read these words are all just different parts of the electromagentic spectrum.
Waves, whether electromagnetic or otherwise, have some defining qualities that describe their energy. Frequency, amplitude, wavelength... all are useful in determining the "energy" of a specific wave of electromagnetic radiation. Frequency and wavelength are related by the speed of light.
The energy of a specific radiation can be calculated from its frequency and Planck's Constant using E=hv
For our purposes, the reason we need to have a good, fundamental understanding of energy is because we study matter by observing how energy changes when it interacts with that matter.
Particle-wave duality - this one is a head-scratcher. Electromagnetic radiation bahves like a wave when we do some experiments, but it behaves like a beam of particles in other experiments. The best explanation for this was that electromagnetic radiation is just something special that can behave like both a particle and a wave, hence the particle-wave duality

Photoelectric Effect:

The photoelectric effect is a great example of "quantized energy". The easiest physical description or analogy of quantized energy levels is comparing a ramp to a flight of stairs - the ramp is continuous, the stairs are quantized.
The photoelectric effect is a clear demonstration that the energy with which electrons are help is quantized. These quantized electron energy levels for the basis for our understaning of the electronic stucture of atoms.

Quantum Mechanics:

Quantum mechanics is a vast and complex field, but there are a number of things we can extract and use to help our understanding of electrons.
The Dr. Quantum video is really quite good (https://www.youtube.com/watch?v=Q1YqgPAtzho)
deBroglie Wavelength - this is a fun little party game... macroscopic particles have wave-like behaviour, and the deBroglie equation can be used to calculate the wavelength. How fast does a golf ball have to be going to develop a measurable wavelength?

Electron Orbitals:

Orbitals are probability surfaces or areas. This is well described in Figure 6.20
Each electron in an atom (or ion) is unique. To study them, we need a consistent way to provide an "address" for each individual electron. This is the whole point of quantum numbers - to provide a unique address system,
If you're a "rule follower" type, quantum numbers are great because they have very specific and well-defined rules as summarized in Table 6.1.
Although quantim numbers are a very good address system, they can sometimes get lost as a sea of numbers, so we use some specialized "shorthand" to help us out. The "angular momentum quantum number" defines the type of orbital, so rather than using numbers, chemists tend to use letters (s, p, d, f), so an orbital with l=0 is referred to as a "s orbital" regardless of the value of n.

Electron Configurations:

A full set of quantim numbers for a given electron provides a LOT of information, often more than we really need to make some useful predictions. Specifically, for many purposes, the value of the magnetic quantum number really isn't all that informative and the value of the spin quantum number is only important in the context that there can be only 2 electrons in each "orbital".
Electron configurations are quantum number shorthand that only includes the most important details, with the finer details left to the reader to define if they become necessary.
"Core" vs "Valence" electrons - "core" electrons are full electronic shells like those present in a noble gas. Valence electrons are those that are outside the core.
Most chemistry takes place in the valence electrons.

Predicting Behaviour & Properties:

The arrangement of the Periodic Table might make a little mroe sense now... The forst 2 columns are often called the "s-block", the transition metals are the "d-block" and the main group elements are the "p-block" because that is the orbital type that houses the valnce electrons in each block.
Whenever predicitng stable electronic configurations, look for full shells, full subshells, and half-full subshells. These tend to be relatively stable. For example, tin atoms have a valence electron of [Kr]5s²4d¹⁰5p². From this electron configuration, we could lose 2 electrons (the "5p" electrons) to form a +2 ion that only has full subshells. It turns out Sn(II) is a relatively stable ion. Similarly, if we lose 2 more electrons (the "5s" electrons), we woul dbe left with a full-subshells-only ion with +4 charge... and Sn(IV) is a relatively stable ion!
There are a number of size-based trends that we can explain with the shell/subshell model.
Ionization Energy - the energy required to remove an electron from an atom or ion.
Electron Affinities - the energy associated with adding an electron to an atom or ion

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