ENGINEERING MATERIALS SCIENCE (ME 370)

Transcription

1 ENGINEERING MATERIALS SCIENCE (ME 370) Chapter 1 Why Materials? Well, everything that surrounds us is made of some type of materials. After all, we live in a Materialistic Society. This synonym with Capitalistic Society. Materials science has the goal of improving materials performance (e.g. designing a car body that is both light in weight and can better withstand impact or collision). In the last car sample, a light body means a relatively small mass density, and a collision withstanding body means a tough or strong body. Toughness and density of a material are part of its properties. Material scientists try to obtain desirable properties by controlling the structure of a material (e.g. by arranging its composing atoms in a certain way). This arrangement of the structure is facilitated by processing, which is simply the way we treat a raw material (s) to achieve a certain structure. The four components above relate as follows: Processing Structure Properties Performance Examples of material properties: mass density color optical transmittance (e.g. transparent, translucent, or opaque) hardness electrical conductivity heat conductivity toughness strength melting temperature wear resistance corrosion resistance ductility elasticity 1

2 Example: Cake Making (Ingredients: Flour + sugar + eggs + butter + water + baking powder) If you don t mix ingredients very well (processing), you get an ugly globular-like texture (structure), and dense material (property) Taste (performance) is not good. If you don t bake (heat) long enough (processing) your cake will be liquid-like (structure) and dense (property) Taste (performance) is not good. If you do mix well and heat (bake) appropriately, you will get a fluffy texture (structure) which is light or not dense (property) Taste (performance) is good. Classification of Materials METALS: Examples are Aluminum, Copper, Steel, Iron, Solder material, etc. Properties: good electrical and heat conductors. Also strong but deformable (used thus in many structures). CERAMICS: Things like: cement, glass, china (made of clay), etc. Properties: insulative to heat and electricity. Also high temperature resistance and tolerance to harsh environments. POLYMERS: Examples: plastic and rubber. Properties: relatively low density and may be extremely flexible. COMPOSITE: Example: fiberglass (glass fibers mixed with a polymer). Properties: a combination of one or all of the above materials. SEMICONDUCTORS: Materials with electrical conductivity in between conductors and insulators. BIOMATERIALS: Things that can be implanted in the human body (e.g. hips and knee joint replacement). These materials are compatible with body tissues and do not produce toxic substances. NATURAL MATERIALS: Things like wood and bones. ADVANCED MATERIALS: Things employed in high technology (high-tech) applications, e.g. CD s, spacecraft materials, fiber optics, etc. 2

3 FUTURE MATERIALS: Smart materials: materials that can respond on their own to external stimuli. For example, some sunglasses dim with increased sunshine or light intensity. Nanotechnology: nowadays, it is possible to build new materials one atom a time. Each atom has a size in the nanometer range. Class Exercise: In nominal groups of 3 (from neighbors), identify three different types of materials existing in the classroom and its entities, classify them, and describe two properties for each of these materials. Chapter 2 Atomic Structure and Interatomic Bonding Fundamental Concepts Every matter is composed, at a small scale, of a collection of atoms joined together somehow. Each atom is itself composed of tinier objects: a nucleus (made up of proton and neutrons) which is surrounded or encircled by electrons. Electrons are negatively charged: -1.6E-19 C/electron Protons are positively charged: +1.6E-19 C/proton Proton mass = 1.67E-27 kg Electron mass = 9.11E-31 kg Each element in nature (e.g. ydrogen, Oxygen, Argon, Aluminum, etc.) has: - An atomic number (Z) = number of protons in the nucleus - Neutron number (N) = number of neutrons for an element (can vary from the number of protons. This is called isotopes) - The atomic mass (A) = Z+N, Q: why not consider the electron mass as well? A unit of mass for atoms is called an atomic mass unit (amu) = 1/12 of the Carbon 12 ( 12 C) isotope which is the most common isotope of carbon. ere A= As it turns out, 1 amu/atom (or molecule; a molecule being a collection of bonded atoms) = 1 g/mol, where one mole of substance = 6.023E23 atoms or molecules (this is Avogadro s number). Atomic Models The set of laws that govern atomic and subatomic entities (like electrons, for example) are called Quantum Mechanics. Bohr Atomic Model: electrons revolve in discrete circular orbitals around the atomic nucleus. The orbits are also called shells. Each orbit represents a fixed amount of energy that the 3

4 electrons in the orbit have. This energy is quantized, i.e. it changes by discrete amounts from one orbit to another. An electron can move to a higher orbit or state of energy by absorbing energy, or down to a lower orbit or state of energy by emitting energy. The first shell or state of energy can take up to 2 atoms, the second one up to 8, the third one up to 18 and the fourth one up to 32. The simplest atomic model is that for ydrogen which has only one electron. This electron prefers to stay in the lowest energy state or orbit close to the nucleus (this is called ground state for an atom). owever, if excited by gaining energy, it can make discrete jumps to one of the higher orbits or states (this is called excited atomic state ). Bohr s atomic model Another well-known atomic model is called the wave-mechanical model which is more complex than Bohr s model and basically assumes that an electron does not have a welldefined orbit around the nucleus and instead could probably exists at any position on, below or above this orbit. We are not going to discuss this model further. Valence electrons are those in the outermost nonempty shell of an atom. These electrons are important since they participate in bonding between different atoms to form atomic and molecular aggregates. The best example is hydrogen, which is found in nature as 2 (two ydrogen atoms making one molecule. This is attained by sharing their valence electrons). 2 molecule 4

5 Elements are arranged in the Periodic Table in rows according to their atomic number and in columns according to the number of valence electrons they have (which dictates also their chemical and physical properties). The ones that have relatively very few atoms in their outer most shell are in the left columns (e.g. Na, Li, Mg, etc.). Those with a relatively large number of atoms are in the right columns (e.g. F, Cl). Elements with completely filled outer electron shells are inert or non-reactive (e.g. Ar, Kr, e). Those are in the right-most column. An ion is simply an atom that has gained one or more electrons (called Anion and is negatively charged) or lost one or more (called cation or positively charged). Atomic Bonding in Solids Bonding Forces and Energies Consider a molecule made up of two atoms separated by an equilibrium distance r o. If one tries to pull these atoms apart the atoms will feel an attraction (an attractive force) that wants to pull them back to where they were (in proximity to one another). Left alone the two atoms will go back to their initial configuration (i.e. experience an attractive force, F). This is because of the bond that these atoms have, which can be thought of as a spring attached between the atoms such that it is unstretched in the original configuration before the pull. If one now instead tries to push the atoms closer to one another, a repulsive force (F) will be felt which wants to keep the two atoms apart (so as not to keep them compressed against one another) again back to where they were. Atom A Interatomic separation, r Atom B The work done in separating or pushing the atoms closer is (from mechanics) given by: Work = Potential Energy E = F dr, where F here is the bonding force resisting the change in r. The work is stored here as potential energy (remember the spring analogy). Assuming a force F vs. r curve as below, the E vs. r curve will then look as given below the force curve. 5

6 Attractive + r o Force F Interatomic separation, r Repulsion - Repulsion + Potential Energy E Interatomic separation, r E o Attractive - The energy corresponding to r o, E o, is the minimum energy state and is called the bonding energy There are few implications of the bonding energy as one might anticipate. For example, materials with strong bonding energies exhibit high melting points or temperatures. Also, at room temperatures, these materials are in a solid state. If the bonds were weak, the material is typically a gas. If it is intermediate, then it is typically a liquid. Types of atomic bonds: primary (or chemical) bonds and secondary (van der Waals or physical) bonds. Primary bonds are: Ionic, covalent, and metallic. 6

7 Primary Bonds Ionic Bonding In Ionic Bonding, one atom gives out one valence electron and another accepts that electron. The ones that loose electrons become Cations and the ones that gain them become Anions. The elements (or correspondingly their atoms) that have tendency to become Cations are called electropositive and those who have tendency to become Anions are called electronegative. The best example of such bonding in materials is NaCl (Sodium Chloride or rock salt). See sketch for NaCl. Coulombic bonding force Ionic Bonding in Sodium Chloride (NaCl). The large circles are Cl - and the small ones are Na +. Since in Ionic Bonding, positive and negative ions attract, the nature of the forces is coulombic or electrical and such bonding is nondirectional (equal in all directions) and hence the symmetric distribution of Na + around Cl -. In Ionic Bonding, bonding energies are relatively large (between 600 and 1500 kj/mol) and hence result in high melting temperatures. Other examples of Ionic Bonding is MgO. 7

8 Covalent Bonding In Covalent Bonding none of the atoms involved loose or gain atoms per say but rather they get close to one another such that they can share the valence electrons. See the C 4 example below whereby one Carbon atom shares its 4 valence electrons with 4 ydrogen atoms located symmetrically around it. The sharing is done to achieve a stable electronic structure by filling the outermost atomic shell. Similarly, diamond is simply made of carbon atoms each sharing 4 electrons with 4 neighboring atoms. Other examples of covalent bonding is Cl 2, F 2, 2 O, F, Si, GaAs, and SiC. Some Covalent Bonds are very strong. For example diamond is very hard and has a high melting temperature because of that (6400 o F). Also, many Polymers exhibit this type of bonding. Shared electron from hydrogen Shared electron from carbon C Metallic Bonding In many metals (e.g. Al, Cu, Ag, Pd, etc.), the atoms share their valence electrons but not only or necessarily with neighboring atoms but rather with all the atoms in the matter. This is facilitated by the electrons forming a sea or cloud that engulfs the rest of the atom (i.e. the positively charged nucleus with its remaining non-valence electrons). See the sketch for an explanation. Ion cores Sea of valence electrons 8

9 This type of bonding can result in either week or strong bonds and thus low or large melting points (varying from 39 o C o C for Mercury through Tungsten). The free electron 9

10 ydrogen Bonds A special type of permanent molecule-induced dipole bonds is called ydrogen Bonds. It exists in compounds like F, 2 O, and N 3 (see sketch of this bond). In these atoms the hydrogen proton (which is not shielded by any negative electrons) acts as a strong positive pole that can form bonds with negative poles from other molecules. This is the strongest secondary bond and is responsible for relatively high melting temperature like for water. F F ydrogen bond 10