The period from the late 18th to the 19th century is referred to as the industrial age, largely due to the rise of mechanization. Similarly, the 20th century marked the beginning of the electronic age, with electronic vacuum tubes dominating the first half of the century and enabling the creation of devices such as radios, televisions, and telephones. However, the technology of the vacuum tube was soon surpassed by solid-state semiconductors, which were first invented in 1947 at Bell Labs by John Bardeen, Walter Brattain, and William Shockley. These semiconductors proved to be smaller, lighter, more reliable, and cheaper to produce than vacuum tubes. The introduction of transistors and related devices enabled the creation of new applications and significantly improved the performance of existing ones. Over time, the development of integrated circuits allowed for the combination of numerous transistors in a single device, resulting in the widespread use of cell phones, GPS devices, laptops, and other electronic devices. Despite the prevalence of these devices in our daily lives, many people have limited knowledge of how they work.
Arthur C. Clarke once said that advanced technology can seem like magic to those who do not understand it. Despite the widespread use of electronic devices in industrialized countries, many people have limited knowledge of how these devices function. However, the operation of these devices is based on scientific principles and human innovation, not magic. Additionally, while more people can use a cell phone than design one, there is a greater need for individuals who can manufacture and maintain semiconductor-based devices rather than designing the semiconductors themselves. As a result, this text will focus on the operation and application of semiconductor devices, rather than their design. This text focuses on the operation and application of semiconductor devices, rather than the design of the semiconductors themselves.
When attempting to comprehend the functioning of semiconductors, a crucial inquiry to make is, “What is the internal structure of an atom?” It is important to note that it is illogical to question what an atom “looks like” since all of its components are smaller than the shortest wavelengths of visible light. Instead, we require a model that can clarify its observed behavior. The most prevalent model in popular culture is the planetary model displayed in Figure 1.1, where the nucleus or core is located at the center, holding positively charged protons and neutral neutrons. Revolving around this core are negatively charged electrons, each following a regular planar path similar to a planet orbiting the sun. Unfortunately, this model is wholly incorrect, although it has been utilized as a symbol for nuclear regulatory agencies and even as the cover art for a DEVO album in the 1970s. Before we devise a more precise and valuable model, let us take a closer look at the sub-components, namely the proton, neutron, and electron. Firstly, most of an atom’s mass arises from its protons and neutrons, which have comparable masses of about 1.67E−24 grams each. An electron’s mass is approximately 2000 times smaller. The radius of a proton is approximately 0.87E−15 meters, and the mean distance to the closest electron is around 5.3E−11 meters. This means that the electron is around 60,000 times further away from the proton than the size of the proton itself. To put this into perspective, this is similar to the ratio between a golf ball and a sphere with a radius of 3/4ths of a mile or 1200 meters. This is applicable for a hydrogen atom as it comprises a single proton and electron. The magnitude of this ratio does not differ significantly for other substances, including materials that incorporate semiconductors.
To comprehend the workings of semiconductors, it is crucial to understand the internal structure of atoms. Atoms consist of protons, neutrons, and electrons. The planetary model is the most widely recognized model in popular culture, where the nucleus or core holds positively charged protons and neutral neutrons, while negatively charged electrons revolve around them. However, this model is not accurate, and we require a better model to explain atom behavior.
Atoms of crystalline carbon (diamond) and quartz (a molecule of silicon and oxygen) are hard and solid. However, solidity is an illusion because, at the atomic level, most of what we consider solid is empty space. For example, when you sit on a chair, the feeling of your buttocks pressed against the chair is just the result of atomic forces between the two objects.
Understanding the concept of solidity and emptiness at an atomic level is essential in the study of semiconductors
In this passage, we learn about the flaws in the popular planetary model of the atom. The author explains that the idea of electrons whirling around the nucleus in stable, planet-like orbits is simply not true. Instead, due to the Heisenberg Uncertainty Principle, we can’t precisely plot the position and trajectory of a given electron. The best we can do is make a plot of where the electron is likely to be, which is called a probability contour.
The author then introduces the concept of an orbital, which is a cloud of dots around the nucleus that represents the likely position of the electron based on multiple measurements. It is important to note that an orbital should not be confused with an orbit (like a planetary orbit), as they are two different things.
The author then explains that only certain orbitals are allowed due to quantum physics. The permissible electron energy levels are first grouped into shells, then subshells and finally orbitals. Shells are denoted by their principal quantum number, n; 1, 2, 3, etc. The higher the number, the more subshells it can contain. Subshells are organized by their orbital shape and are designated by letters, the first four being s, p, d, and f. Shell 1 contains only subshell s while shell 2 contains subshell types s and p. Shell 3 contains subshell types s, p and d, and so on.
Thus, we see designations such as 1s, 2s and 2p. These subshells may also have variations within them, such as one variation on s, three variations on p, five variations on d, etc. Each orbital can hold a maximum of two electrons.
Overall, this passage provides a detailed explanation of the structure of an atom, highlighting the flaws in the popular planetary model and introducing the concept of orbitals, shells, and subshells. The author provides specific examples and designations to help the reader understand the different types of subshells and orbitals, and emphasizes the importance of these concepts in future discussions.
In the structure of an atom, electrons are arranged in shells around the nucleus. The first shell, closest to the nucleus, can hold a maximum of two electrons in the single s subshell orbital (1s). The second shell can hold a maximum of eight electrons, consisting of two in the s subshell (2s) and two in each of the three p subshell orbitals (2p). The third shell can hold a maximum of 18 electrons, consisting of two in 3s, six in 3p, and two in each of the five d subshell orbitals (3d). This can be simplified into the formula 2n^2, where n is the shell number.
The electron probability contour of the innermost orbital, 1s, is shown in Figure 1.2. It is spherical in shape, with the nucleus located at the center. Higher order orbitals can take on a variety of forms. Figure 1.3 shows the electron probability contour for the 2p orbitals, which have three variations oriented along the X, Y, and Z axes. The Bohr model, named after Danish physicist Niels Bohr, is a more functional graphic that represents the energy description of the atom. It consists of a nucleus at the center and concentric rings representing the electron shells, with higher numbers indicating greater energy levels.
Using the Bohr model, we can create diagrams to represent individual elements. For example, copper has an atomic number of 29 and an electron shell configuration of 2-8-18-1, with the first three shells completely filled and a single electron in the fourth shell. The Bohr model for copper would simply show four rings, with a single electron in the fourth ring. Figure 1.5 shows the Bohr model of an atom of Silicon, atomic number 14, with an electron shell configuration of 2-8-4.
To simplify the Bohr model further, the filled inner shells can be omitted, and the atomic number can be replaced by the number of electrons in the outermost, or valence, shell. This is shown in Figure 1.6. The valence shell is particularly important as it gives insight into the general behavior of the material. Alternatively, the Bohr model can be depicted as energy levels graphically as lines or bands, without counting specific electrons, as shown in Figure 1.7.
One response to “Semiconductor Basics: Atomic Structure”
[…] accelerated and quite literally smash into the target, dislodging andreplacing some of the original atoms in the […]