Material Science and Engineering PDF

BICOL STATE COLLEGE OF APPLIED SCIENCES AND TECHNOLOGY College of Engineering and Architecture Material Science and Engineering...

BICOL STATE COLLEGE OF APPLIED SCIENCES AND TECHNOLOGY College of Engineering and Architecture Material Science and Engineering Instructor: Engr. Dominic P. Bolima WEEK 1-4 (Part 2) I. INTRODUCTION TO MATERIAL SCIENCE AND ENGINEERING A. Atomic Structure and Interatomic Bonding Atomic Structure Concepts Atom consists of (subatomic particles) a very small nucleus composed of protons and neutrons, which is encircled by moving electrons The magnitude of the charge of electrons and protons is 1.60 x 10-19 C The mass of protons and neutrons is approximately 1.67 x 10-27 kg each while the mass of electron is 9.11 x 10-31 kg. Each chemical element is characterized by the number of protons in the nucleus, or the atomic number (Z). The atomic number also equals the number of electrons. The atomic mass (A) of a specific atom may be expressed as the sum of the masses of protons and neutrons within the nucleus. Isotopes are elements with the same number of protons but different numbers of neutrons and atomic masses. The atomic weight of an element corresponds to the weighted average of the atomic masses of the atom’s naturally occurring isotopes. The atomic weight of an element or the molecular weight of a compound may be specified on the basis of amu per atom (molecule) or mass per mole of material. In one mole of a substance there are 6.023 x 1023 (Avogadro’s number) atoms or molecules. These two atomic weight schemes are related through the following equation: amu 1 = 1g/mol atom(molecule) 1 1 amu = 𝑎𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 𝑜𝑓 𝐶𝑎𝑟𝑏𝑜𝑛 12 12 Example: Carbon 12 : + 6 neutrons Carbon 14: + 8 neutrons Carbon has 15 known isotopes, of which only two are stable. Others exist only in trace amounts, so it makes no measurable contribution to the average atomic mass. Electrons in Atoms Quantum Mechanics – the branch of mechanics that deals with the mathematical description of the motion and interaction of subatomic particles, incorporating the concepts of quantization of energy, wave-particle duality, the uncertainty principle, and the correspondence principle. Bohr Atomic Model – electrons are assumed to revolve around the atomic nucleus in discrete orbitals, and the position of any particular electron is more or less well defined in terms of its orbital. – the energies of electrons are quantized; that is, electrons are permitted to have only specific values of energy. Page 1 of 8  BICOL STATE COLLEGE OF APPLIED SCIENCES AND TECHNOLOGY College of Engineering and Architecture Material Science and Engineering Instructor: Engr. Dominic P. Bolima Figure 1. Bohr’s Atomic Model Wave-mechanical model – the electron is considered to exhibit both wave-like and particle-like characteristics. – electron’s position is considered to be the probability of an electron’s being at various locations around the nucleus. – an electron does not have a well-defined orbit around the nucleus and instead could probably exists at any position on, below or above this orbit. Figure 2. Wave-Mechanical Atomic Model Valence Electrons – are those in the outermost nonempty shell of an atom. Quantum Numbers - the four parameters that characterized electron in an atom. 1. Principal Quantum Number, n – designated by letters K, L, M, N, O… – related to the distance of an electron from the nucleus, or its position 2. Second Quantum Number, l – signifies the subshell denoted by s, p, d, or f. – related to the shape of the electron subshell 3. Third Quantum Number, ml – the number of energy states. 4. Spin Moment, ms – the orientation of the electron, which is either up (+1/2) or down (-1/2). Page 2 of 8  BICOL STATE COLLEGE OF APPLIED SCIENCES AND TECHNOLOGY College of Engineering and Architecture Material Science and Engineering Instructor: Engr. Dominic P. Bolima Figure 3. Quantum Numbers Electron State – values of energy that are permitted for electrons. Pauli Exclusion Principle – states that each electron state can hold no more than two electrons, which must have two opposite spins. Electron Configuration – or the structure of an atom represents the manner in which energy states are occupied. Valence Electrons – electrons that occupy the outermost shell. Stable Electron Configurations – the states within the outermost shell are completely filled. Figure 4. Electron Configuration All the elements have been classified according to electron configuration in the periodic table. Here, the elements are situated, with increasing atomic number, in seven horizontal rows called periods. The arrangement is such that all elements arrayed in a given column or group have similar valence electron structures, as well as chemical and physical properties. These properties change gradually, moving horizontally across each period and vertically down each column. Some of the groups of elements and their characteristics. The elements positioned in Group 0, the rightmost group, are the inert gases, which have filled electron shells and stable electron configurations. Group VIIA and VIA elements are one and two electrons deficient, respectively, from having stable structures. The Group VIIA elements (F, Cl, Br, I, and At) are sometimes termed the halogens. The alkali and the alkaline earth metals (Li, Na, K, Be, Mg, Ca, etc.) are labeled as Groups IA and IIA, having, respectively, one and two electrons in excess of stable structures. Page 3 of 8  BICOL STATE COLLEGE OF APPLIED SCIENCES AND TECHNOLOGY College of Engineering and Architecture Material Science and Engineering Instructor: Engr. Dominic P. Bolima The elements in the three long periods, Groups IIIB through IIB, are termed the transition metals, which have partially filled d electron states and in some cases one or two electrons in the next higher energy shell. Groups IIIA, IVA, and VA (B, Si, Ge, As, etc.) display characteristics that are intermediate between the metals and nonmetals by virtue of their valence electron structures. Figure 5. The Periodic Table of Elements Alkali metals. The alkali metals make up group 1 of the Table, and comprise Li through Fr. They have very similar behavior and characteristics. Hydrogen is group 1 but exhibits few characteristics of a metal and is often categorized with the nonmetals. Alkaline earth metals. The alkaline earth metals make up group 2 of the periodic table, from Be through Ra. The alkaline earth metals have very high melting points and oxides that have basic alkaline solutions. Their characteristics are well described and consistent down the group. Transition metals. The transition elements are metals that have a partially filled d subshell (CRC Handbook of Chemistry and Physics) and comprise groups 3 through 12 and the lanthanides and actinides. Post-transition metals. The post-transition elements are Al, Ga, In, Tl, Sn, Pb and Bi. As their name implies, they have some of the characteristics of the transition elements. They tend to be softer and conduct more poorly than the transition metals. Metalloid (or "semi-metal" or "poor metal"). The metalloids are B, Si, Ge, As, Sb, Te, and Po. They sometimes behave as semiconductors (B, Si, Ge) rather than as conductors. Lanthanides. The lanthanides comprise elements 57 (lanthanum, hence the name of the set) through 71. They are grouped together because they have similar chemical properties. They, along with the actinides, are often called "the f-elements" because they have valence electrons in the f shell. Actinides. The actinides comprise elements 89 through 103. They, along with the lanthanides, are often called "the f- elements" because they have valence electrons in the f shell. Only thorium and uranium are naturally occurring actinides with significant abundance. They are all radioactive. Nonmetals. The term "nonmetals" is used to classify the elements H, C, N, P, O, S, and Se. Page 4 of 8  BICOL STATE COLLEGE OF APPLIED SCIENCES AND TECHNOLOGY College of Engineering and Architecture Material Science and Engineering Instructor: Engr. Dominic P. Bolima Halogens. The halogen elements are a subset of the nonmetals. They comprise group 17 of the periodic table, from F through At. They generally very chemically reactive and are present in the environment as compounds rather than as pure elements. Noble gases. The noble gases comprise group 18. They are generally very stable chemically and exhibit similar properties of being colorless and odorless. B. Atomic Bonding in Solids There are attractive and repulsive forces between two atoms that vary according to the distance between the atoms. 𝐅𝐍 = 𝐅𝐀 + 𝐅𝐑 ≫ 𝒏𝒆𝒕 𝒇𝒐𝒓𝒄𝒆 For FA = FR, an equilibrium state exists. Thus, the center of the two atoms will remain separated by a constant spacing. Bonding Energy – represents the energy that would be required to separate two atoms to an infinite separation. The magnitude of bonding energies vary from material to material and they both depend on the type of atomic bonding. Classifications of Bonds 1. Primary Bonds or Chemical Bonds – depends on the electron structures Ionic Bonding Covalent Bonding Metallic Bonding 2. Secondary Bonds or Physical Bonds Induced Dipole Polar Molecule-Induced Dipole Bond Permanent Dipole Bonds Primary Bonds or Chemical Bonds Ionic Bonding - found in compounds that are composed of metallic and nonmetallic elements. - atoms of a metallic element easily give up their valence electrons to the nonmetallic atoms. - in the process all the atoms acquire stable or inert gas configurations and, in addition, an electrical charge; that is, they become ions Example: Sodium Chloride (NaCl) Figure 6. NaCl Bonding Page 5 of 8  BICOL STATE COLLEGE OF APPLIED SCIENCES AND TECHNOLOGY College of Engineering and Architecture Material Science and Engineering Instructor: Engr. Dominic P. Bolima - Sodium and Chlorine will have the same electron configurations with Neon and Argon respectively. - Coulombic bonding force – attractive forces between positive and negative ions due by virtue of their net electrical charges. - Ionic materials are mostly hard and brittle, and electrically and thermally insulative. Other Examples: NaBr - sodium bromide KBr - potassium bromide NaCl - sodium chloride NaF - sodium fluoride KI - potassium iodide KCl - potassium chloride CaCl2 - calcium chloride K2O - potassium oxide MgO - magnesium oxide Covalent Bonding - stable electron configurations are assumed by the sharing of electrons between adjacent atoms - two atoms that are covalently bonded will each contribute at least one electron to the bond - the shared electrons may be considered to belong to both atoms E.g. H2, Cl2, F2, etc. CH4, H20, HNO3, HF, Diamond(carbon) Figure 7. CH4 (Methane) Bonding Types of Covalent Bonds Nonpolar covalent bonds are a type of chemical bond where two atoms share a pair of electrons with each other. Non-polar Molecule – when the electrons in a bond are perfectly shared, then there is no dipole and neither of end of the band carries any partial charge. example: H2, CO2 Polar covalent bond is a type of chemical bond where a pair of electrons is unequally shared between two atoms. For atoms with different electronegativity, the electrons tend to spend more time at one end of the bond than the other. This sets up what is known as a dipole. example: H2O Page 6 of 8  BICOL STATE COLLEGE OF APPLIED SCIENCES AND TECHNOLOGY College of Engineering and Architecture Material Science and Engineering Instructor: Engr. Dominic P. Bolima Electronegativity - measure of the tendency of an atom to attract a bonding pair of electrons. - the wider the separation (both horizontally—relative to Group IVA—and vertically) from the lower left to the upper- right-hand corner (i.e., the greater the difference in electronegativity), the more ionic the bond. Conversely, the closer the atoms are together (i.e., the smaller the difference in electronegativity), the greater the degree of covalency. - 1.7 and less of electronegativity difference between two atoms means that the compound has covalent bond. Examples: PCl3 - phosphorus trichloride CH3CH2OH - ethanol O3 - ozone H2 - hydrogen H2O - water HCl - hydrogen chloride CH4 - methane NH3 - ammonia CO2 - carbon dioxide Metallic Bond - is the force of attraction between valence electrons and the metal atoms. It is the sharing of many detached electrons between many positive ions, where the electrons act as a "glue" giving the substance a definite structure. - The remaining non-valence electrons and atomic nuclei form what are called ion cores, which possess a net positive charge equal in magnitude to the total valence electron charge per atom. Figure 8. Metallic Bond - metallic bonding is found in the periodic table for Group IA and IIA elements and, in fact, for all elemental metals. - metals tend to have high melting points and boiling points suggesting strong bonds between the atoms. Some general behaviors of the various material types (i.e., metals, ceramics, polymers) may be explained by bonding type. For example, metals are good conductors of both electricity and heat, as a consequence of their free electrons. By way of contrast, ionically and covalently bonded materials are typically electrical and thermal insulators, due to the absence of large numbers of free electrons. At room temperature, most metals and their alloys fail in a ductile manner; that is, fracture occurs after the materials have experienced significant degrees of permanent deformation. This behavior is explained in terms of deformation mechanism, which is implicitly related to the characteristics of the metallic bond. Conversely, at room temperature ionically bonded materials are intrinsically brittle as a consequence of the electrically charged nature of their component ions Page 7 of 8  BICOL STATE COLLEGE OF APPLIED SCIENCES AND TECHNOLOGY College of Engineering and Architecture Material Science and Engineering Instructor: Engr. Dominic P. Bolima Secondary Bonds or Physical Bonds - Van der Waals Bond - mostly in solids, weaker than primary bonds. - exists between virtually all atoms or molecules but hard to detect due to primary bonds - result from forces between electric dipoles can exist between the following: 1. induced dipoles 2. induced dipoles + polar molecule 3. polar molecules Induced Dipoles Bond - fluctuating induced dipole in electrically symmetric molecule - atoms experienced constant vibrational motions that cause instantaneous and short-lived distortions on electrical symmetry for some of the atoms or molecules, and the create small electric dipole - induced polarity in neighboring atom/molecule - very weak force, low bond energy and low melting point. Figure 9. Induced Dipole Bond Polar Molecule – Induced Dipole Bonds - asymmetrical electrical distribution – polar molecules - polar molecules have permanent dipole moment - induced dipole moment in other electrically symmetrical atom/molecule or nonpolar molecules - attractive forces develop stronger than fluctuating induced dipoles Permanent Dipole Bond - attractive force between two polar molecules - stronger force than previous two configurations - strongest among all the secondary bonds and high melting point - special case is hydrogen bonding: example HF Hydrogen Bond – special type of secondary bonding that exists between molecules that have hydrogen as one of its constituents. - Hydrogen shares its only electron, strong positive end - very strong attraction with adjacent F of HF which is negative examples: HF, HCl, H2O, HN Page 8 of 8