97 - Atomic Orbitals (w/ Adele Musicant!)
97. Atomic Orbitals
A common misconception is that electrons orbit the nucleus of an atom, like a planet around the sun. Why is this wrong? What are atomic orbitals? Let’s learn to be scientifically conversational.
General Learning Concepts
1) Background knowledge
a. Previous particle knowledge: We previously discussed some preliminary information about atoms in episode 6 ‘Batteries: Zap, but How?’. An atom a fundamental piece of matter. (Matter is anything that can be touched physically.) Everything in the universe (except energy) is made of matter, and, so, everything in the universe is made of atoms.
i. What is a proton? A particle. Protons reside in the nucleus of atoms, containing a positive charge. They are responsible for most of the mass along with neutrons. (Protons and neutrons are responsible for most of the atomic mass e.g in a 150 person 149 lbs, 15 oz are protons and neutrons while only 1 oz. is electrons.)
ii. What is an electron? A particle. Electrons reside in orbitals around the nucleus. They contain a negative charge.
b. What is the periodic table? A list of chemical elements organized so that scientists can quickly discern the properties of individual elements such as their mass, electron number, electron configuration and their unique chemical properties.
2) What is an atomic orbital?
a. Atomic Orbital: An orbital is a three dimensional description of the most likely location of an electron around an atom. This is where we need to be very careful about defining the position of an electron: it is very likely NOT orbiting the nucleus like the popularization of the Bohr model. Instead, we are talking about energy levels of electrons. (Chem.libretexts: an orbital is the quantum mechanical refinement of Bohr’s orbit. In contrast to his concept of a simple circular orbit with a fixed radius, orbitals are mathematically derived regions of space with different probabilities of having an electron.) [2]
i. Heisenberg Uncertainty Principle: It is impossible to define with absolute precision (at the same time) both the position and the momentum of the electron.
ii. Suppose you had a single hydrogen atom and at a particular instant plotted the position of the one electron. Soon afterwards, you do the same thing, and find that it is in a new position. You have no idea how it got from the first place to the second. You keep on doing this over and over again, and gradually build up a sort of 3D map of the places that the electron is likely to be found. In the hydrogen case, the electron can be found anywhere within a spherical space surrounding the nucleus. The diagram shows a cross-section through this spherical space. 95% of the time (or any other percentage you choose), the electron will be found within a fairly easily defined region of space quite close to the nucleus. Such a region of space is called an orbital. You can think of an orbital as being the region of space in which the electron lives.
b. Types of orbitals: Four orbital types are common: s (sharp), p (principle), d (diffuse), and f (fundamental). These orbitals attempt to localize the distribution of electrons within the available shells; that being, at any point in time an electron is likely contained within the volume of the orbital shape.
i. How many electrons can be held within certain energy levels? S can hold 2 electrons; p = 6, d = 10, and f = 14.
ii. Shapes of orbitals: The s orbital is spherical. The p orbitals are dumbbell shaped along different 3D axis. Four of the five d orbitals are cloverleaf shaped. The final d orbital is shaped like a dumbbell with a doughnut around its center. The f orbitals are far more complex to describe.
c. The Aufbau Principle: From the German word “Aufbauen” which means “to build”. Orbitals are filled in order of increasing atomic number, meaning that electrons pile up into higher energy levels (starting from the lowest available energy states). However, the energy of orbitals is dictated by both the principle quantum number (n, describes the energy of an electron and the most probable distance of the electron from the nucleus) and the azimuthal quantum number (l, a quantum number for an atomic orbital that determines its orbital angular momentum and describes the shape of the orbital). Therefore, the energy level of 4s (n=4, l=0, n+l=4) is actually lower than the 3d orbital (n=3, l=2, n+l=5) and will preferably fill faster than 3d.
3) Fun Tidbits
a. Exceptions to filling: Full or half-full subshells tend to have enhanced stability. When examining the first 30 elements, only copper (atomic number 24, expected 1s2 2s2 2p6 3s2 3p6 3d9 4s2, actual 1s2 2s2 2p6 3s2 3p6 3d10 4s1) and chrome (atomic number 29, expected 1s2 2s2 2p6 3s2 3p6 3d4 4s2, actual 1s2 2s2 2p6 3s2 3p6 3d5 4s1) are exceptions to the Aufbau principle.
4) Solicited Questions
a. Why are there two electrons per orbital? The Pauli Exclusion Principle states that, in an atom or molecule, no two electrons can have the same four electronic quantum numbers. As an orbital can contain a maximum of only two electrons, the two electrons must have opposing spins. This means if one is assigned an up-spin ( +1/2), the other must be down-spin (-1/2).