AMIE Material Science and Engineering Ebook

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free download amie section a and b question papers 2005 to 2012 and unlimited study materials

Here I am giving you notes for material science engineering course offered by AMIE in PDF file attached with it so you can get it easily..

Chapter 1. INRODUCTION

1 .1 Historical Perspective

Materials are so important in the development of civilization that we associate Ages with them.
In the origin of human life on Earth, the Stone Age, people used only natural materials, like
stone, clay, skins, and wood. When people found copper and how to make it harder by alloying,
the Bronze Age started about 3000 BC. The use of iron and steel, a stronger material that gave
advantage in wars started at about 1200 BC. The next big step was the discovery of a cheap
process to make steel around 1850, which enabled the railroads and the building of the modern
infrastructure of the industrial world.

1.2 Materials Science and Engineering
Understanding of how materials behave like they do, and why they differ in properties was only
possible with the atomistic understanding allowed by quantum mechanics, that first explained
atoms and then solids starting in the 1930s. The combination of physics, chemistry, and the
focus on the relationship between the properties of a material and its microstructure is the
domain of Materials Science. The development of this science allowed designing materials and
provided a knowledge base for the engineering applications (Materials Engineering).

Structure:
•At the atomic level: arrangement of atoms in different ways. (Gives different properties
for graphite than diamond both forms of carbon.)
•At the microscopic level: arrangement of small grains of material that can be identified
by microscopy. (Gives different optical properties to transparent vs. frosted glass.)

Properties are the way the material responds to the environment. For instance, the mechanical,
electrical and magnetic properties are the responses to mechanical, electrical and magnetic
forces, respectively. Other important properties are thermal (transmission of heat, heat
capacity), optical (absorption, transmission and scattering of light), and the chemical stability
in contact with the environment (like corrosion resistance).

Processing of materials is the application of heat (heat treatment), mechanical forces, etc. to
affect their microstructure and, therefore, their properties

1.3 Why Study Materials Science and Engineering?
•To be able to select a material for a given use based on considerations of cost and
performance.
•To understand the limits of materials and the change of their properties with use.
•To be able to create a new material that will have some desirable properties.
All engineering disciplines need to know about materials. Even the most "immaterial", like
software or system engineering depend on the development of new materials, which in turn
alter the economics, like software-hardware trade-offs. Increasing applications of system

engineering are in materials manufacturing (industrial engineering) and complex environmental
systems.

1.4 Classification of Materials
Like many other things, materials are classified in groups, so that our brain can handle the
complexity. One could classify them according to structure, or properties, or use. The one that
we will use is according to the way the atoms are bound together:

Metals: valence electrons are detached from atoms, and spread in an 'electron sea' that "glues"
the ions together. Metals are usually strong, conduct electricity and heat well and are opaque to
light (shiny if polished). Examples: aluminum, steel, brass, gold.

Semiconductors: the bonding is covalent (electrons are shared between atoms). Their electrical
properties depend extremely strongly on minute proportions of contaminants. They are opaque
to visible light but transparent to the infrared. Examples: Si, Ge, GaAs.

Ceramics: atoms behave mostly like either positive or negative ions, and are bound by
Coulomb forces between them. They are usually combinations of metals or semiconductors
with oxygen, nitrogen or carbon (oxides, nitrides, and carbides). Examples: glass, porcelain,
many minerals.

Polymers: are bound by covalent forces and also by weak van der Waals forces, and usually
based on H, C and other non-metallic elements. They decompose at moderate temperatures
(100 – 400 C), and are lightweight. Other properties vary greatly. Examples: plastics (nylon,
Teflon, polyester) and rubber.
Other categories are not based on bonding. A particular microstructure identifies

composites,
made of different materials in intimate contact (example: fiberglass, concrete, wood) to achieve
specific properties. Biomaterials can be any type of material that is biocompatible and used,
for instance, to replace human body parts.

1.5 Advanced Materials
Materials used in "High-Tec" applications, usually designed for maximum performance, and
normally expensive. Examples are titanium alloys for supersonic airplanes, magnetic alloys for
computer disks, special ceramics for the heat shield of the space shuttle, etc.

1.6 Modern Material's Needs
•Engine efficiency increases at high temperatures: requires high temperature structural
materials
•Use of nuclear energy requires solving problem with residues, or advances in nuclear
waste processing.
•Hypersonic flight requires materials that are light, strong and resist high temperatures.
•Optical communications require optical fibers that absorb light negligibly.
•Civil construction – materials for unbreakable windows.
•Structures: materials that are strong like metals and resist corrosion like plastics.

Chapter 2. ATOMIC STRUCTURE AND BONDING
2.2 Fundamental Concepts
Atoms are composed of electrons, protons, and neutrons. Electron and protons are negative and
positive charges of the same magnitude, 1.6 × 10-19 Coulombs.
The mass of the electron is negligible with respect to those of the proton and the neutron, which
form the nucleus of the atom. The unit of mass is an atomic mass unit (amu) = 1.66 × 10-27 kg,
and equals 1/12 the mass of a carbon atom. The Carbon nucleus has Z=6, and A=6, where Z is
the number of protons, and A the number of neutrons. Neutrons and protons have very similar
masses, roughly equal to 1 amu. A neutral atom has the same number of electrons and protons,
Z.
A mole is the amount of matter that has a mass in grams equal to the atomic mass in amu of the
atoms. Thus, a mole of carbon has a mass of 12 grams. The number of atoms in a mole is called
the Avogadro number, Nav = 6.023 × 1023. Note that Nav = 1 gram/1 amu.
Calculating n, the number of atoms per cm3 in a piece of material of density d (g/cm3).
n = Nav × d / M
where M is the atomic mass in amu (grams per mol). Thus, for graphite (carbon) with a density
d = 1.8 g/cm3, M =12, we get 6 × 1023 atoms/mol × 1.8 g/cm3 / 12 g/mol) = 9 × 1022 C/cm3.
For a molecular solid like ice, one uses the molecular mass, M(H2O) = 18. With a density of 1

g/cm3, one obtains n = 3.3 × 1022 H2O/cm3. Note that since the water molecule contains 3
atoms, this is equivalent to 9.9 × 1022 atoms/cm3.
Most solids have atomic densities around 6 × 1022 atoms/cm3. The cube root of that number
gives the number of atoms per centimeter, about 39 million. The mean distance between atoms
is the inverse of that, or 0.25 nm. This is an important number that gives the scale of atomic
structures in solids.

2.3 Electrons in Atoms
The forces in the atom are repulsions between electrons and attraction between electrons and
protons. The neutrons play no significant role. Thus, Z is what characterizes the atom.
The electrons form a cloud around the neutron, of radius of 0.05 – 2 nanometers. Electrons do
not move in circular orbits, as in popular drawings, but in 'fuzzy' orbits. We cannot tell how it

moves, but only say what is the probability of finding it at some distance from the nucleus.
According to quantum mechanics, only certain orbits are allowed (thus, the idea of a mini
planetary system is not correct). The orbits are identified by a principal quantum number n,
which can be related to the size, n = 0 is the smallest; n = 1, 2 .. are larger. (They are
"quantized" or discrete, being specified by integers). The angular momentum l is quantized, and
so is the projection in a specific direction m. The structure of the atom is determined by the
Pauli exclusion principle, only two electrons can be placed in an orbit with a given n, l, m – one

for each spin. Table 2.1 in the textbook gives the number of electrons in each shell (given by n)
and subshells (given by l).

2.4 The Periodic Table
Elements are categorized by placing them in the periodic table. Elements in a column share
similar properties. The noble gases have closed shells, and so they do not gain or lose electrons
near another atom. Alkalis can easily lose an electron and become a closed shell; halogens can
easily gain one to form a negative ion, again with a closed shell. The propensity to form closed
shells occurs in molecules, when they share electrons to close a molecular shell. Examples are
H2, N2, and NaCl.
The ability to gain or lose electrons is termed electronegativity or electropositivity, an
important factor in ionic bonds.

2.5 Bonding Forces and Energies
The Coulomb forces are simple: attractive between electrons and nuclei, repulsive between
electrons and between nuclei. The force between atoms is given by a sum of all the individual
forces, and the fact that the electrons are located outside the atom and the nucleus in the center.
When two atoms come very close, the force between them is always repulsive, because the
electrons stay outside and the nuclei repel each other. Unless both atoms are ions of the same

charge (e.g., both negative) the forces between atoms is always attractive at large internuclear
distances r. Since the force is repulsive at small r, and attractive at small r, there is a distance at
which the force is zero. This is the equilibrium distance at which the atoms prefer to stay.
The interaction energy is the potential energy between the atoms. It is negative if the atoms are
bound and positive if they can move away from each other. The interaction energy is the
integral of the force over the separation distance, so these two quantities are directly related.
The interaction energy is a minimum at the equilibrium position. This value of the energy is
called the bond energy, and is the energy needed to separate completely to infinity (the work
that needs to be done to overcome the attractive force.) The strongest the bond energy, the
hardest is to move the atoms, for instance the hardest it is to melt the solid, or to evaporate its
atoms.

2.6 Primary Interatomic Bonds
Ionic Bonding
This is the bond when one of the atoms is negative (has an extra electron) and another is
positive (has lost an electron). Then there is a strong, direct Coulomb attraction. An example is
NaCl. In the molecule, there are more electrons around Cl, forming Cl- and less around Na,
forming Na+. Ionic bonds are the strongest bonds. In real solids, ionic bonding is usually
combined with covalent bonding. In this case, the fractional ionic bonding is defined as %ionic
= 100 × [1 – exp(-0.25 (XA – XB)2], where XA and XB are the electronegativities of the two
atoms, A and B, forming the molecule.

Covalent Bonding
In covalent bonding, electrons are shared between the molecules, to saturate the valency. The
simplest example is the H2 molecule, where the electrons spend more time in between the
nuclei than outside, thus producing bonding.

Metallic Bonding
In metals, the atoms are ionized, loosing some electrons from the valence band. Those electrons
form a electron sea, which binds the charged nuclei in place, in a similar way that the electrons
in between the H atoms in the H2 molecule bind the protons
.
2.7 Secondary Bonding (Van der Waals)

Fluctuating Induced Dipole Bonds
Since the electrons may be on one side of the atom or the other, a dipole is formed: the +
nucleus at the center, and the electron outside. Since the electron moves, the dipole fluctuates.
This fluctuation in atom A produces a fluctuating electric field that is felt by the electrons of an
adjacent atom, B. Atom B then polarizes so that its outer electrons are on the side of the atom
closest to the + side (or opposite to the – side) of the dipole in A. This bond is called van der
Waals bonding.

Polar Molecule-Induced Dipole Bonds
A polar molecule like H2O (Hs are partially +, O is partially – ), will induce a dipole in a
nearby atom, leading to bonding.

Permanent Dipole Bonds
This is the case of the hydrogen bond in ice. The H end of the molecule is positively charged
and can bond to the negative side of another dipolar molecule, like the O side of the H2O
dipole.

2.8 Molecules
If molecules formed a closed shell due to covalent bonding (like H2, N2) then the interaction
between molecules is weak, of the van der Waals type. Thus, molecular solids usually have
very low melting points.

Chapter-3: STRUCTURE OF CRYSTALS
3.2 Fundamental Concepts
Atoms self-organize in crystals, most of the time. The crystalline lattice, is a periodic array of
the atoms. When the solid is not crystalline, it is called amorphous. Examples of crystalline
solids are metals, diamond and other precious stones, ice, graphite. Examples of amorphous
solids are glass, amorphous carbon (a-C), amorphous Si, most plastics
To discuss crystalline structures it is useful to consider atoms as being hard spheres, with welldefined
radii. In this scheme, the shortest distance between two like atoms is one diameter.

3.3 Unit Cells
The unit cell is the smallest structure that repeats itself by translation through the crystal. We
construct these symmetrical units with the hard spheres. The most common types of unit cells
are the faced-centered cubic (FCC), the body-centered cubic (FCC) and the hexagonal closepacked
(HCP). Other types exist, particularly among minerals. The simple cube (SC) is often
used for didactical purpose, no material has this structure.

3.4 Metallic Crystal Structures

Important properties of the unit cells are
•The type of atoms and their radii R.
•cell dimensions (side a in cubic cells, side of base a and height c in HCP) in terms of R.
•n, number of atoms per unit cell. For an atom that is shared with m adjacent unit cells,
we only count a fraction of the atom, 1/m.
•CN, the coordination number, which is the number of closest neighbors to which an
atom is bonded.

•APF, the atomic packing factor, which is the fraction of the volume of the cell actually
occupied by the hard spheres. APF = Sum of atomic volumes/Volume of cell.

Unit Cell n CN a/R APF
SC 1 6 2 0.52
BCC 2 8 4Ö 3 0.68
FCC 4 12 2Ö 2 0.74
HCP 6 12 0.74
The closest packed direction in a BCC cell is along the diagonal of the cube; in a FCC cell is
along the diagonal of a face of the cube.

3.5 Density Computations

The density of a solid is that of the unit cell, obtained by dividing the mass of the atoms (n
atoms x Matom) and dividing by Vc the volume of the cell (a3 in the case of a cube). If the mass
of the atom is given in amu (A), then we have to divide it by the Avogadro number to get Matom.
Thus, the formula for the density is:

3.6 Polymorphism and Allotropy
Some materials may exist in more than one crystal structure, this is called polymorphism. If the
material is an elemental solid, it is called allotropy. An example of allotropy is carbon, which
can exist as diamond, graphite, and amorphous carbon.

3.11 Close-Packed Crystal Structures
The FCC and HCP are related, and have the same APF. They are built by packing spheres on
top of each other, in the hollow sites (Fig. 3.12 of book). The packing is alternate between two
types of sites, ABABAB.. in the HCP structure, and alternates between three types of positions,
ABCABC… in the FCC crystals.


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