EEE208 — Electrical Engineering Materials: Lecture Note

About this lecture note

This is the complete lecture note for EEE208 Electrical Engineering Materials. You can read it online below, open it in the full viewer, or download the PDF for offline study. It pairs with the course's video lectures, so you can read and watch together.

Preview: opening of the note

Course code: EEE 208 Course title: Electrical Engineering Materials Course credit load: 3 Units Series 1: Atomic Structure and Bonding in Materials • Part 1.1: The Atom and Its Internal Structure o Sub Part 1.1.1: Introduction to atomic structure o Sub Part 1.1.2: Atomic number, mass number, and isotopes o Sub Part 1.1.3: The periodic table and electron configuration overview • Part 1.2: The Electron Orbital Model o Sub Part 1.2.1: Bohr's model and its limitations o Sub Part 1.2.2: Quantum numbers and atomic orbitals o Sub Part 1.2.3: Electron configuration and the Aufbau principle • Part 1.3: Valence Electrons and the Basis of Bonding o Sub Part 1.3.1: Valence electrons and the octet rule o Sub Part 1.3.2: Ionization energy and electron affinity o Sub Part 1.3.3: Electronegativity and bond character prediction • Part 1.4: Primary (Strong) Bonding in Materials o Sub Part 1.4.1: Ionic bonding o Sub Part 1.4.2: Covalent bonding o Sub Part 1.4.3: Metallic bonding • Part 1.5: Secondary Bonding and Bonding's Effect on Material Properties o Sub Part 1.5.1: Van der Waals (secondary) bonding o Sub Part 1.5.2: Hydrogen bonding o Sub Part 1.5.3: How bonding type influences material properties

Introduction This course examines the materials from which every electrical and electronic device is ultimately made, and why those materials behave as they do. This first Series lays the essential foundation, beginning with the internal structure of the atom and the electron orbital model, before extending this to explain the different types of bonding that hold materials together. The aim of this Series is to equip you with a sound understanding of the internal structure of the atom, the electron orbital model, valence electrons and the periodic trends that govern bonding, and the primary and secondary bonding types found in engineering materials. We shall commence this Series by examining the internal structure of the atom, its atomic and mass numbers, and its place in the periodic table. Next, we will consider the electron orbital model, including Bohr's model, quantum numbers, and electron configuration. Following that, our attention turns to valence electrons and the periodic trends of ionisation energy, electron affinity, and electronegativity. We will then examine the primary bonding types: ionic, covalent, and metallic bonding. Finally, we will conclude by examining secondary bonding and how bonding type influences the properties of a material. Series Learning Outcomes Upon completing this Series, you should be able to: • SLO 1.1: describe the internal structure of an atom, including its nucleus and electrons, • SLO 1.2: distinguish between atomic number, mass number, and isotopes, • SLO 1.3: describe the Bohr model of the atom and explain its limitations, • SLO 1.4: describe the quantum mechanical orbital model and determine electron configurations, • SLO 1.5: explain the significance of valence electrons and the octet rule, • SLO 1.6: explain ionisation energy, electron affinity, and electronegativity, • SLO 1.7: describe ionic and covalent bonding, • SLO 1.8: describe metallic bonding and explain its relationship to electrical conductivity, • SLO 1.9: describe secondary (van der Waals) bonding and hydrogen bonding, • SLO 1.10: explain how bonding type influences the properties of a material.

1.1 The Atom and Its Internal Structure 1.1.1 Introduction to atomic structure An atom is the smallest unit of an element that retains the chemical properties of that element, consisting of a small, dense central nucleus, containing positively charged protons and electrically neutral neutrons, surrounded by a comparatively vast region occupied by negatively charged electrons, with the atom as a whole electrically neutral because the number of protons exactly equals the number of electrons. The nucleus accounts for almost the entire mass of an atom, since a proton or neutron is roughly 1836 times more massive than an electron, yet occupies only a tiny fraction of the atom's total volume, meaning that an atom is, by volume, overwhelmingly empty space through which the comparatively light electrons move. Worked example: A neutral atom of carbon contains 6 protons and 6 electrons; since the number of protons equals the number of electrons, the positive charge of the nucleus is exactly balanced by the negative charge of the surrounding electrons, leaving the atom overall electrically neutral. Application: This basic picture of a small, massive, positively charged nucleus surrounded by light, negatively charged electrons underlies every subsequent topic in this Series, from the electron orbital model through to the bonding that holds materials together, and is the essential starting point for understanding the electrical and mechanical properties of engineering materials. Fill in the blank: The central, positively charged part of an atom is called the ________. An atom is electrically neutral because the number of protons equals the number of ________. 1.1.2 Atomic number, mass number, and isotopes The atomic number (Z) of an element is the number of protons in the nucleus of an atom of that element, a value unique to each element and the property that fundamentally determines its chemical identity, while the mass number (A) is the total number of protons and neutrons together in the nucleus, so that the number of neutrons in a given atom is found from A − Z. Isotopes are atoms of the same element, and therefore having the same atomic number Z, but differing in their number of neutrons, and consequently in their mass number A; because chemical behaviour is governed almost entirely by the number and arrangement of electrons, which in turn depends on the atomic number, isotopes of the same element behave almost identically chemically despite their differing mass. Worked example: Carbon has two naturally common isotopes, carbon-12 (6 protons, 6 neutrons, mass number 12) and carbon-14 (6 protons, 8 neutrons, mass number 14); both have the same atomic number, 6, and hence the same chemical behaviour, but differ in mass and, in the case of carbon-14, in nuclear stability.

Application: Distinguishing atomic number from mass number is essential when identifying elements and their isotopes on the periodic table, and the near-identical chemical behaviour of isotopes, despite their differing mass, is precisely why isotopic labelling is such a useful technique in materials analysis and tracing. Fill in the blank: The atomic number of an element is the number of ________ in its nucleus. Isotopes of an element have the same atomic number but a different ________ number. 1.1.3 The periodic table and electron configuration overview The periodic table arranges all known elements in order of increasing atomic number, organised into rows, called periods, and columns, called groups, such that elements within the same group share broadly similar chemical and physical properties, since they share a similar arrangement of electrons in their outermost shell. This periodic recurrence of similar properties arises directly from the underlying electron structure of the atoms concerned, examined in detail in the next Part of this Series, and is why the periodic table remains an indispensable reference tool for an engineer, allowing the likely bonding behaviour, electrical conductivity, and other key properties of an unfamiliar material to be anticipated from the position of its constituent elements alone. Worked example: Sodium and potassium, both found in the same group of the periodic table, share the characteristic property of readily losing a single outer electron to form a positively charged ion, a similarity in chemical behaviour that follows directly from their sharing the same number of electrons in their outermost shell. Application: An engineer selecting a material can use its constituent elements' positions on the periodic table as an immediate, first-approximation guide to its likely properties, for example anticipating that a material composed of a metal from the left of the table and a non-metal from the right is likely to involve ionic bonding, a connection examined in detail later in this Series. Fill in the blank: The periodic table arranges elements in order of increasing ________ number. Elements in the same group of the periodic table share a similar arrangement of electrons in their outermost ________.

Figure 1.1: Basic structure of an atom showing the nucleus and surrounding electron shells. Pilot questions 1.1 1. Name the three main subatomic particles and state their charges. (Linked to SLO 1.1) 2. Explain why an atom is electrically neutral. (Linked to SLO 1.1) 3. Distinguish between atomic number and mass number. (Linked to SLO 1.2) 4. Define isotopes. (Linked to SLO 1.2) 5. Explain why isotopes of an element behave almost identically chemically. (Linked to SLO 1.2) 1.2 The Electron Orbital Model 1.2.1 Bohr's model and its limitations Bohr's model of the atom, proposed in 1913, describes electrons as orbiting the nucleus in fixed, well-defined circular paths, called energy levels or shells, each corresponding to a specific, discrete energy, with an electron able to jump between shells only by absorbing or emitting a photon of exactly the energy difference between the two shells involved.

Although Bohr's model successfully explained the discrete spectral lines observed in hydrogen's emission spectrum, a considerable achievement at the time, it fails to accurately predict the spectra of atoms with more than one electron, and, more fundamentally,

This is the opening of the note. Read the full note or download the PDF to continue.