The Elemental Flow arc is long past us, but our objective is to make sure you understand what you’ve learned…It’s one thing to say that you’ve been taught. It’s another to say that you can remember – a general flaw brought about by time constraints in school.
Electrons behave in a simple manner. There is only one factor that motivates it – energy. Electrons often hold high energy due to every other particle around them. Conversely to organisms, such as yourself, electrons seek to go from high-energy to low energy. This energy comes from attraction to the positively-charged nucleus, as well as repulsion from other negatively-charged electrons. Motion is one of the methods that electrons use to expel energy. The energy of motion is called kinetic energy; it is one part of a pair of intimately linked types of energy. Its partner, potential energy, represents stored energy. The conventional example is dropping a ball off a cliff. As you hold the ball over a cliff, it has potential energy – the energy imposed upon it by the force of gravity and its position above ground. As the ball drops, that stored energy begins to dissipate. But energy is always conserved. That means the potential energy must not be disappearing, but converting to another energy source. Electrons move toward the positively-charged nucleus, in an attempt to rid themselves of their potential energy, but are flung away by the massive amount of kinetic energy they experience at the center. The cycle continues, creating a particle that constantly moves.
How does that electron move? Not like the orbiting model you know, where electrons all revolve around a center in the exact same orbit every second of every moment. That came from the observation by Niels Bohr that inputting energy into a system, through light, yielded an output of energy, also light. This is evidence of a particle trying to remove energy. But this didn’t work for all wavelengths of light. Only certain lights – certain amounts of energy – yielded an output of light. These specific, discrete energies led Louis de Broglie to look at the photon, a particle of light discovered by the famed Albert Einstein, which Bohr observed, but did not believe in. Both electrons and photons are called quantized – electrons because they exist at only discrete values and photons because they are the lowest discrete magnitude of light (called light’s quanta).
For the electron to emit specific colors of light, energy that exists at specific levels must have influence on the electron’s electromagnetic field, which all moving charges create. de Broglie proved that these electrons are under the influence of their constantly oscillating electromagnetic waves and behave like waves themselves. He kept in line with Bohr’s correct analysis that the electrons must be some distance away from the nucleus at all times, resulting in the theory of electrons having a dual particle and wave, just like Einstein theorized a photon had. These waves, constant and unchanging, are called standing waves.
This wave-particle duality has since been proven in many experiments, explaining even Bohr’s model for the hydrogen, which takes the circular orbit that Bohr incorrectly theorized all atoms have. Schrodinger created a mathematical model based on each of these discoveries, giving us the model that we use to find the probable location of an electron within an atom.
Electrons are incredibly fast. Not only do they behave like both waves and particles, but, sometimes, both at once. This is the magic of quantum mechanics.
Electrons exist in layers. The more electrons there are, the more layers there will be. But the most important layer, and the layer that defines what orbital, or classification, an atom is designated under is the outermost. It is called the valence shell. To find out what shell any atom is in, you need to understand the periodic table. We will discuss this further in the next arc, but right now, just a preliminary understanding.
Take a glance at the periodic table and you will see a number at the top left. This is referred to as the atomic number, which represents the number of protons in any atom. If there is no loss or gain of electrons, and the system has a net charge of 0 (neither positive or negative), then the atomic number also represents the number of electrons.
Experimentation has uncovered that most atoms in ordinary circumstances, desire either two electrons or eight electrons in their valence shell, according to their nearest noble gas. The rule is somewhat shaky when it comes to atoms combining with other atoms.
Hydrogen with its one electron, for example, wishes to bind to something else that will give it one electron. Lithium, on the other hand, with three electrons, wishes to give away one of its electrons to match helium, which is the noble gas with a complete valence shell of two electrons.
There are four quantum numbers that determine how an electron is ordered on a periodic table.
First, the Principle Quantum Number (denoted as n) determines the energy level, and thus, the electron’s shell. It is conventionally shown with a number from 1 to 5. Next, the Orbital Quantum Number (l) determines where the electrons are with respect to the nucleus. It shows the sub-shell and is typically denoted with a letter (s, p, d or f) according to its number (0, 1, 2 and 3). The Magnetic Quantum Number (m) tells how many subshells there are. It is determined by looking at the subshell and taking the range from -l to l. Lastly, the complex Spin Quantum Number (s) determines whether the electron’s “spin” is +1/2 or -1/2.
Each atom fits in the periodic table, going from left to right according to their orbital designation.
How do electrons fill their orbitals? There is a principle of building up that takes place, governed by the Pauli Exclusion Principle, which states that only two electrons can be in any orbital subshell. Generally, the lowest energy orbitals are filled first. But things get dicey when you get to the second principle number. The Hund’s Rule takes charge here, which states that each subshell in a given orbital must be filled with an electron containing the same Spin Quantum Number before the opposite Spin Quantum Number is filled. It’s easier to understand with this diagram.
Note that unlike the transition from Lithium (Le) to Beryllium (Be), there is a different transition taking place from Boron (B) to Carbon (C). This is Hund’s Rule in action.
Fortunately, if you’ve understood all of this, then the “Diagonal Rule” will come naturally to you.
This gives you an easy road map to understanding how orbitals are filled in. There is only one trick here. Did you notice that the order, upon reaching 3p, is 4s then 3d? The reason why this is so unlike the earlier organization (1s, 2s, 2p, 3s…) is because of the Madelung Rule, which helps you quantify the energy level of each of these orbitals. Simply add the principle quantum number, n, and the orbital quantum number, l to obtain the orbital with the lowest energy (remember electrons fill lowest energy orbitals first). 3p has a principle number of 3 and an orbital number of 1, making 4. 4s comes before 3d because the former has an energy level of 4, while the latter has an energy level of 5, following the Madelung Rule. Therefore, 4s comes first, validating the Diagonal Rule.
Lastly, this buildup explains the Noble Gases. With full subshells, it is impossible for a noble gas to take on additional electrons, nor is it willing to part with its electrons. Therefore, it is nonreactive.
Next: The Dance of Chemistry
With orbitals behind us, it is about time to introduce the next topic, where we discuss in further detail how individual atoms combine.
The best part is, if you’ve followed all of this, then the next part will be a breeze. I promise.