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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.)

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.

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.

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, N_av = 6.023 × 10²³. Note that N_av = 1 gram/1 amu.

Calculating n, the number of atoms per cm³ in a piece of material of density d (g/cm³).

n = N_av × 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 (X_A – X_B)²], where X_A and X_B 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.
 
 

3.2 Fundamental Concepts

To discuss crystalline structures it is useful to consider atoms as being hard spheres, with well-defined radii. In this scheme, the shortest distance between two like atoms is one diameter.

3.3 Unit Cells

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.

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

3.6 Polymorphism and Allotropy

3.7 – 3.10 Crystallography – Not Covered

3.11 Close-Packed Crystal Structures

Crystalline and Non-Crystalline Materials

3.12 Single Crystals

3.13 Polycrystalline Materials

3.14 Anisotropy

3.16 Non-Crystalline Solids

4.1 Introduction

0D (zero dimension) – point defects: vacancies and interstitials. Impurities.

1D – linear defects: dislocations (edge, screw, mixed)

2D – grain boundaries, surfaces.

3D – extended defects: pores, cracks.

Point Defects

4.2 Vacancies and Self-Interstitials

An interstitial is an atom that occupies a place outside the normal lattice position. It may be the same type of atom as the others (self interstitial) or an impurity atom.

In the case of vacancies and interstitials, there is a change in the coordination of atoms around the defect. This means that the forces are not balanced in the same way as for other atoms in the solid, which results in lattice distortion around the defect.

The number of vacancies formed by thermal agitation follows the law:

N_V = N_A × exp(-Q_V/kT)

where N_A is the total number of atoms in the solid, Q_V is the energy required to form a vacancy, k is Boltzmann constant, and T the temperature in Kelvin (note, not in ^oC or ^oF).

When Q_V is given in joules, k = 1.38 × 10^-23 J/atom-K. When using eV as the unit of energy, k = 8.62 × 10^-5 eV/atom-K.

Note that kT(300 K) = 0.025 eV (room temperature) is much smaller than typical vacancy formation energies. For instance, Q_V(Cu) = 0.9 eV/atom. This means that N_V/N_A at room temperature is exp(-36) = 2.3 × 10^-16, an insignificant number. Thus, a high temperature is needed to have a high thermal concentration of vacancies. Even so, N_V/N_A is typically only about 0.0001 at the melting point.

4.3 Impurities in Solids

Impurities are often added to materials to improve the properties. For instance, carbon added in small amounts to iron makes steel, which is stronger than iron. Boron impurities added to silicon drastically change its electrical properties.

Solid solutions are made of a host, the solvent or matrix) which dissolves the solute (minor component). The ability to dissolve is called solubility. Solid solutions are:

• homogeneous
• maintain crystal structure
• contain randomly dispersed impurities (substitutional or interstitial)

• Similar atomic size (to within 15%)
• Similar crystal structure
• Similar electronegativity (otherwise a compound is formed)
• Similar valence

Miscellaneous Imperfections

4.4 Dislocations—Linear Defects

Edge dislocations occur when an extra plane is inserted. The dislocation line is at the end of the plane. In an edge dislocation, the Burgers vector is perpendicular to the dislocation line.

Screw dislocations result when displacing planes relative to each other through shear. In this case, the Burgers vector is parallel to the dislocation line.

4.5 Interfacial Defects

The density of atoms in the region including the grain boundary is smaller than the bulk value, since void space occurs in the interface.

Surfaces and interfaces are very reactive and it is usual that impurities segregate there. Since energy is required to form a surface, grains tend to grow in size at the expense of smaller grains to minimize energy. This occurs by diffusion, which is accelerated at high temperatures.

Twin boundaries: not covered

4.6 Bulk or Volume Defects

4.7 Atomic Vibrations

5.1 Introduction

Inhomogeneous materials can become homogeneous by diffusion, if the temperature is high enough (temperature is needed to overcome energy barriers to atomic motion.

5.2 Diffusion Mechanisms

There is a difference between diffusion and net diffusion. In a homogeneous material, atoms also diffuse but this motion is hard to detect. This is because atoms move randomly and there will be an equal number of atoms moving in one direction than in another. In inhomogeneous materials, the effect of diffusion is readily seen by a change in concentration with time. In this case there is a net diffusion. Net diffusion occurs because, although all atoms are moving randomly, there are more atoms moving in regions where their concentration is higher.

5.3 Steady-State Diffusion

Steady state diffusion means that J does not depend on time. In this case, Fick’s first law holds that the flux along direction x is:

J = – D dC/dx

Where dC/dx is the gradient of the concentration C, and D is the diffusion constant. The concentration gradient is often called the driving force in diffusion (but it is not a force in the mechanistic sense). The minus sign in the equation means that diffusion is down the concentration gradient.

5.4 Nonsteady-State Diffusion

¶ C/¶ t = D ¶ ²C/¶ x²

The solution depends on the particular value of the boundary conditions (initial values of the concentration and its derivatives.)

5.5 Factors That Influence Diffusion

5.6 Other Diffusion Paths

Mechanical Properties of Metals

1. Introduction

 

 

Often materials are subject to forces (loads) when they are used. Mechanical engineers calculate those forces and material scientists how materials deform (elongate, compress, twist) or break as a function of applied load, time, temperature, and other conditions.

Materials scientists learn about these mechanical properties by testing materials. Results from the tests depend on the size and shape of material to be tested (specimen), how it is held, and the way of performing the test. That is why we use common procedures, or standards, which are published by the ASTM.

2. Concepts of Stress and Strain

 

 

To compare specimens of different sizes, the load is calculated per unit area, also called normalization to the area. Force divided by area is called stress. In tension and compression tests, the relevant area is that perpendicular to the force. In shear or torsion tests, the area is perpendicular to the axis of rotation.

s = F/A₀ tensile or compressive stress

t = F/A₀ shear stress

The unit is the Megapascal = 10⁶ Newtons/m².

There is a change in dimensions, or deformation elongation, DL as a result of a tensile or compressive stress. To enable comparison with specimens of different length, the elongation is also normalized, this time to the length L. This is called strain, e.

e = DL/L

The change in dimensions is the reason we use A₀ to indicate the initial area since it changes during deformation. One could divide force by the actual area, this is called true stress (see Sec. 6.7).

For torsional or shear stresses, the deformation is the angle of twist, q (Fig. 6.1) and the shear strain is given by:

g = tg q

3. Stress—Strain Behavior

 

 

Elastic deformation. When the stress is removed, the material returns to the dimension it had before the load was applied. Valid for small strains (except the case of rubbers).

Deformation is reversible, non permanent

Plastic deformation. When the stress is removed, the material does not return to its previous dimension but there is a permanent, irreversible deformation.

In tensile tests, if the deformation is elastic, the stress-strain relationship is called Hooke's law:

s = E e

That is, E is the slope of the stress-strain curve. E is Young's modulus or modulus of elasticity. In some cases, the relationship is not linear so that E can be defined alternatively as the local slope:

E = ds/de

Shear stresses produce strains according to:

t = G g

where G is the shear modulus.

Elastic moduli measure the stiffness of the material. They are related to the second derivative of the interatomic potential, or the first derivative of the force vs. internuclear distance (Fig. 6.6). By examining these curves we can tell which material has a higher modulus. Due to thermal vibrations the elastic modulus decreases with temperature. E is large for ceramics (stronger ionic bond) and small for polymers (weak covalent bond). Since the interatomic distances depend on direction in the crystal, E depends on direction (i.e., it is anisotropic) for single crystals. For randomly oriented policrystals, E is isotropic.

4. Anelasticity

 

 

Here the behavior is elastic but not the stress-strain curve is not immediately reversible. It takes a while for the strain to return to zero. The effect is normally small for metals but can be significant for polymers.

5. Elastic Properties of Materials

 

 

Materials subject to tension shrink laterally. Those subject to compression, bulge. The ratio of lateral and axial strains is called the Poisson's ratio n.

n = e_lateral/e_axial

The elastic modulus, shear modulus and Poisson's ratio are related by E = 2G(1+n)
6. Tensile Properties

 

 

Yield point. If the stress is too large, the strain deviates from being proportional to the stress. The point at which this happens is the yield point because there the material yields, deforming permanently (plastically).

Yield stress. Hooke's law is not valid beyond the yield point. The stress at the yield point is called yield stress, and is an important measure of the mechanical properties of materials. In practice, the yield stress is chosen as that causing a permanent strain of 0.002 (strain offset, Fig. 6.9.)

The yield stress measures the resistance to plastic deformation.

The reason for plastic deformation, in normal materials, is not that the atomic bond is stretched beyond repair, but the motion of dislocations, which involves breaking and reforming bonds.

Plastic deformation is caused by the motion of dislocations.

Tensile strength. When stress continues in the plastic regime, the stress-strain passes through a maximum, called the tensile strength (s_{TS) , and then falls as the material starts to develop a neck and it finally breaks at the fracture point (Fig. 6.10).}

_{Note that it is called strength, not stress, but the units are the same, MPa.}

_{For structural applications, the yield stress is usually a more important property than the tensile strength, since once the it is passed, the structure has deformed beyond acceptable limits.}

_{Ductility. The ability to deform before braking. It is the opposite of brittleness. Ductility can be given either as percent maximum elongation} e_{max or maximum area reduction.}

%EL = e_max x 100 %

%AR = (A₀ - A_f)/A₀

These are measured after fracture (repositioning the two pieces back together).

Resilience. Capacity to absorb energy elastically. The energy per unit volume is the

area under the strain-stress curve in the elastic region.

Toughness. Ability to absorb energy up to fracture. The energy per unit volume is the total area under the strain-stress curve. It is measured by an impact test (Ch. 8).

7. True Stress and Strain

 

 

When one applies a constant tensile force the material will break after reaching the tensile strength. The material starts necking (the transverse area decreases) but the stress cannot increase beyond s_{TS. The ratio of the force to the initial area, what we normally do, is called the engineering stress. If the ratio is to the actual area (that changes with stress) one obtains the true stress.}

8. _{Elastic Recovery During Plastic Deformation}

 

 

_{If a material is taken beyond the yield point (it is deformed plastically) and the stress is then released, the material ends up with a permanent strain. If the stress is reapplied, the material again responds elastically at the beginning up to a new yield point that is higher than the original yield point (strain hardening, Ch. 7.10). The amount of elastic strain that it will take before reaching the yield point is called elastic strain recovery (Fig. 6. 16).}

9. _{Compressive, Shear, and Torsional Deformation}

 

 

_{Compressive and shear stresses give similar behavior to tensile stresses, but in the case of compressive stresses there is no maximum in the} s-e curve, since no necking occurs.

10. Hardness

 

 

Hardness is the resistance to plastic deformation (e.g., a local dent or scratch). Thus, it is a measure of plastic deformation, as is the tensile strength, so they are well correlated. Historically, it was measured on an empirically scale, determined by the ability of a material to scratch another, diamond being the hardest and talc the softer. Now we use standard tests, where a ball, or point is pressed into a material and the size of the dent is measured. There are a few different hardness tests: Rockwell, Brinell, Vickers, etc. They are popular because they are easy and non-destructive (except for the small dent).

11. Variability of Material Properties

 

 

Tests do not produce exactly the same result because of variations in the test equipment, procedures, operator bias, specimen fabrication, etc. But, even if all those parameters are controlled within strict limits, a variation remains in the materials, due to uncontrolled variations during fabrication, non homogenous composition and structure, etc. The measured mechanical properties will show scatter, which is often distributed in a Gaussian curve (bell-shaped), that is characterized by the mean value and the standard deviation (width).

12. Design/Safety Factors

 

 

To take into account variability of properties, designers use, instead of an average value of, say, the tensile strength, the probability that the yield strength is above the minimum value tolerable. This leads to the use of a safety factor N > 1 (typ. 1.2 - 4). Thus, a working value for the tensile strength would be s_{W =} s_{TS / N.}

Chapter 7. Dislocations and Strengthening Mechanisms

1. Introduction

 

 

The key idea of the chapter is that plastic deformation is due to the motion of a large number of dislocations. The motion is called slip. Thus, the strength (resistance to deformation) can be improved by putting obstacles to slip.

2. Basic Concepts

 

 

Dislocations can be edge dislocations, screw dislocations and exist in combination of the two (Ch. 4.4). Their motion (slip) occurs by sequential bond breaking and bond reforming (Fig. 7.1). The number of dislocations per unit volume is the dislocation density, in a plane they are measured per unit area.

3. Characteristics of Dislocations

 

 

There is strain around a dislocation which influences how they interact with other dislocations, impurities, etc. There is compression near the extra plane (higher atomic density) and tension following the dislocation line (Fig. 7.4)

Dislocations interact among themselves (Fig. 7.5). When they are in the same plane, they repel if they have the same sign and annihilate if they have opposite signs (leaving behind a perfect crystal). In general, when dislocations are close and their strain fields add to a larger value, they repel, because being close increases the potential energy (it takes energy to strain a region of the material).

The number of dislocations increases dramatically during plastic deformation. Dislocations spawn from existing dislocations, and from defects, grain boundaries and surface irregularities.

4. Slip Systems

 

 

In single crystals there are preferred planes where dislocations move (slip planes). There they do not move in any direction, but in preferred crystallographic directions (slip direction). The set of slip planes and directions constitute slip systems.

The slip planes are those of highest packing density. How do we explain this? Since the distance between atoms is shorter than the average, the distance perpendicular to the plane has to be longer than average. Being relatively far apart, the atoms can move more easily with respect to the atoms of the adjacent plane. (We did not discuss direction and plane nomenclature for slip systems.)

BCC and FCC crystals have more slip systems, that is more ways for dislocation to propagate. Thus, those crystals are more ductile than HCP crystals (HCP crystals are more brittle).

5. Slip in Single Crystals

6. Plastic Deformation of Polycrystalline Materials

 

 

Slip directions vary from crystal to crystal. When plastic deformation occurs in a grain, it will be constrained by its neighbors which may be less favorably oriented. As a result, polycrystalline metals are stronger than single crystals (the exception is the perfect single crystal, as in whiskers.)

7. Deformation by Twinning

 

 

This topic is not included.

Mechanisms of Strengthening in Metals

General principles. Ability to deform plastically depends on ability of dislocations to move. Strengthening consists in hindering dislocation motion. We discuss the methods of grain-size reduction, solid-solution alloying and strain hardening. These are for single-phase metals. We discuss others when treating alloys. Ordinarily, strengthening reduces ductility.

8. Strengthening by Grain Size Reduction

 

 

This is based on the fact that it is difficult for a dislocation to pass into another grain, especially if it is very misaligned. Atomic disorder at the boundary causes discontinuity in slip planes. For high-angle grain boundaries, stress at end of slip plane may trigger new dislocations in adjacent grains. Small angle grain boundaries are not effective in blocking dislocations.

The finer the grains, the larger the area of grain boundaries that impedes dislocation motion. Grain-size reduction usually improves toughness as well. Usually, the yield strength varies with grain size d according to:

s_y = s₀ + k_y/ d²

Grain size can be controlled by the rate of solidification and by plastic deformation.
9. Solid-Solution Strengthening

 

 

Adding another element that goes into interstitial or substitutional positions in a solution increases strength. The impurity atoms cause lattice strain (Figs. 7.17 and 7.18) which can "anchor" dislocations. This occurs when the strain caused by the alloying element compensates that of the dislocation, thus achieving a state of low potential energy. It costs strain energy for the dislocation to move away from this state (which is like a potential well). The scarcity of energy at low temperatures is why slip is hindered.

Pure metals are almost always softer than their alloys.

10. . Strain Hardening

 

 

Ductile metals become stronger when they are deformed plastically at temperatures well below the melting point (cold working). (This is different from hot working is the shaping of materials at high temperatures where large deformation is possible.) Strain hardening (work hardening) is the reason for the elastic recovery discussed in Ch. 6.8.

The reason for strain hardening is that the dislocation density increases with plastic deformation (cold work) due to multiplication. The average distance between dislocations then decreases and dislocations start blocking the motion of each one.

The measure of strain hardening is the percent cold work (%CW), given by the relative reduction of the original area, A₀ to the final value A_d :

%CW = 100 (A₀–A_d)/A₀

Recovery, recrystallization and Grain Growth

Plastic deformation causes 1) change in grain size, 2) strain hardening, 3) increase in the dislocation density. Restoration to the state before cold-work is done by heating through two processes: recovery and recrystallization. These may be followed by grain growth.

11. . Recovery

 

 

Heating à increased diffusion à enhanced dislocation motion à relieves internal strain energy and reduces the number of dislocation. The electrical and thermal conductivity are restored to the values existing before cold working.

12. . Recrystallization

 

 

Strained grains of cold-worked metal are replaced, upon heating, by more regularly-spaced grains. This occurs through short-range diffusion enabled by the high temperature. Since recrystallization occurs by diffusion, the important parameters are both temperature and time.

The material becomes softer, weaker, but more ductile (Fig. 7.22).

Recrystallization temperature: is that at which the process is complete in one hour. It is typically 1/3 to 1/2 of the melting temperature. It falls as the %CW is increased. Below a "critical deformation", recrystallization does not occur.

13. . Grain Growth

1. Introduction

 

 

Failure of materials may have huge costs. Causes included improper materials selection or processing, the improper design of components, and improper use.

2. Fundamentals of Fracture

Fracture is a form of failure where the material separates in pieces due to stress, at temperatures below the melting point. The fracture is termed ductile or brittle depending on whether the elongation is large or small.

Steps in fracture (response to stress):

track formation

track propagation

Ductile vs. brittle fracture

Ductile Fracture

Initial necking

small cavity formation (microvoids)

void growth (elipsoid) by coalescence into a crack

fast crack propagation around neck. Shear strain at 45^o

final shear fracture (cup and cone)

Brittle Fracture

5. Principles of Fracture Mechanics

 

 

Fracture occurs due to stress concentration at flaws, like surface scratches, voids, etc. If a is the length of the void and r the radius of curvature, the enhanced stress near the flaw is:

s_m » 2 s_{0 (a/}r)^1/2

where s_{0 is the applied macroscopic stress. Note that a is 1/2 the length of the flaw, not the full length for an internal flaw, but the full length for a surface flaw. The stress concentration factor is:}

_{Kt =} s_m/s_{0 » 2 (a/}r)^1/2

Because of this enhancement, flaws with small radius of curvature are called stress raisers.

6. Impact Fracture Testing

long period of cyclic strain

the most usual (90%) of metallic failures (happens also in ceramics and polymers)

is brittle-like even in ductile metals, with little plastic deformation

it occurs in stages involving the initiation and propagation of cracks.

7. Cyclic Stresses

8. The S—N Curve

9. Crack Initiation and Propagation

slow

along crystallographic planes of high shear stress

flat and featureless fatigue surface

10. . Crack Propagation Rate (not covered)

11. . Factors That Affect Fatigue Life

Mean stress (lower fatigue life with increasing s_mean).

Surface defects (scratches, sharp transitions and edges). Solution:

polish to remove machining flaws

add residual compressive stress (e.g., by shot peening.)

case harden, by carburizing, nitriding (exposing to appropriate gas at high temperature)

12. . Environmental Effects

Thermal cycling causes expansion and contraction, hence thermal stress, if component is restrained. Solution:

eliminate restraint by design

use materials with low thermal expansion coefficients.

Corrosion fatigue. Chemical reactions induced pits which act as stress raisers. Corrosion also enhances crack propagation. Solutions:

decrease corrosiveness of medium, if possible.

add protective surface coating.

add residual compressive stresses.

13. . Generalized Creep Behavior

Creep rate: de/dt

Time to failure.

14. . Stress and Temperature Effects

15. . Data Extrapolation Methods (not covered.)

16. . Alloys for High-Temperature Use

9.2 Solubility Limit

9.3 Phases

9.4 Microstructure

The properties of an alloy do not depend only on concentration of the phases but how they are arranged structurally at the microscopy level. Thus, the microstructure is specified by the number of phases, their proportions, and their arrangement in space.

A binary alloy may be

a single solid solution

two separated, essentially pure components.

two separated solid solutions.

a chemical compound, together with a solid solution.

The way to tell is to cut the material, polish it to a mirror finish, etch it a weak acid (components etch at a different rate) and observe the surface under a microscope.

9.5 Phase Equilibriam

9.6 Binary Isomorphous Systems

This very simple case is one complete liquid and solid solubility, an isomorphous system. The example is the Cu-Ni alloy of Fig. 9.2a. The complete solubility occurs because both Cu and Ni have the same crystal structure (FCC), near the same radii, electronegativity and valence.

The liquidus line separates the liquid phase from solid or solid + liquid phases. That is, the solution is liquid above the liquidus line.

The solidus line is that below which the solution is completely solid (does not contain a liquid phase.)

Concentrations: Tie-line method

1. locate composition and temperature in diagram
2. In two phase region draw tie line or isotherm
3. note intersection with phase boundaries. Read compositions.

Fractions: lever rule

construct tie line (isotherm)

obtain ratios of line segments lengths.

Note: the fractions are inversely proportional to the length to the boundary for the particular phase. If the point in the diagram is close to the phase line, the fraction of that phase is large.

Development of microstructure in isomorphous alloys

a) Equilibrium cooling

Solidification in the solid + liquid phase occurs gradually upon cooling from the liquidus line. The composition of the solid and the liquid change gradually during cooling (as can be determined by the tie-line method.) Nuclei of the solid phase form and they grow to consume all the liquid at the solidus line.

b) Non-equilibrium cooling

Solidification in the solid + liquid phase also occurs gradually. The composition of the liquid phase evolves by diffusion, following the equilibrium values that can be derived from the tie-line method. However, diffusion in the solid state is very slow. Hence, the new layers that solidify on top of the grains have the equilibrium composition at that temperature but once they are solid their composition does not change. This lead to the formation of layered (cored) grains (Fig. 9.14) and to the invalidity of the tie-line method to determine the composition of the solid phase (it still works for the liquid phase, where diffusion is fast.)

9.7 Binary Eutectic Systems

Interpretation: Obtain phases present, concentration of phases and their fraction (%).

Solvus line: limit of solubility

Eutectic or invariant point. Liquid and two solid phases exist in equilibrium at the eutectic composition and the eutectic temperature.

Note:

the melting point of the eutectic alloy is lower than that of the components (eutectic = easy to melt in Greek).

At most two phases can be in equilibrium within a phase field.

Single-phase regions are separated by 2-phase regions.

Development of microstructure in eutectic alloys

Case of lead-tin alloys, figures 9.9–9.14. A layered, eutectic structure develops when cooling below the eutectic temperature. Alloys which are to the left of the eutectic concentration (hipoeutectic) or to the right (hypereutectic) form a proeutectic phase before reaching the eutectic temperature, while in the solid + liquid region. The eutectic structure then adds when the remaining liquid is solidified when cooling further. The eutectic microstructure is lamellar (layered) due to the reduced diffusion distances in the solid state.

To obtain the concentration of the eutectic microstructure in the final solid solution, one draws a vertical line at the eutectic concentration and applies the lever rule treating the eutectic as a separate phase (Fig. 9.16).

9.8 Equilibrium Diagrams Having Intermediate Phases or Compounds

9.9 Eutectoid and Peritectic Reactions

9.10 Congruent Phase Transformations

9.11 Ceramic and Ternary Phase Diagrams

9.12 The Gibbs Phase Rule (not discussed)

9.13 The Iron–Iron Carbide (Fe–Fe₃C) Phase Diagram

This is one of the most important alloys for structural applications. The diagram Fe—C is simplified at low carbon concentrations by assuming it is the Fe—Fe₃C diagram. Concentrations are usually given in weight percent. The possible phases are:

a-ferrite (BCC) Fe-C solution

g-austenite (FCC) Fe-C solution

d-ferrite (BCC) Fe-C solution

liquid Fe-C solution

Fe₃C (iron carbide) or cementite. An intermetallic compound.

The maximum solubility of C in a- ferrite is 0.022 wt%. d-ferrite is only stable at high temperatures. It is not important in practice. Austenite has a maximum C concentration of 2.14 wt %. It is not stable below the eutectic temperature (727 C) unless cooled rapidly (Chapter 10). Cementite is in reality metastable, decomposing into a-Fe and C when heated for several years between 650 and 770 C.

For their role in mechanical properties of the alloy, it is important to note that:

9.14 Development of Microstructures in Iron—Carbon Alloys

9.15 The Influence of Other Alloying Elements

As mentioned in section 7.9, alloying strengthens metals by hindering the motion of dislocations. Thus, the strength of Fe–C alloys increase with C content and also with the addition of other elements.

10.2 Basic Concepts

• Diffusion
• Maintaining the type and number of phases (e.g., solidification of a pure metal, allotropic transformation, recrystallization, grain growth.
• Alteration of phase composition (e.g., eutectoid reactions, see 10.5)
• Diffusionless
• Production of metastable phases (e.g., martensitic transformation, see 10.5)

10.3 The Kinetics of Solid-State Reactions

D is called a constant because it does not depend on time, but it depends on temperature as we have seen in Ch. 5. Diffusion occurs faster at high temperatures.

Phase transformation requires two processes: nucleation and growth. Nucleation involves the formation of very small particles, or nuclei (e.g., grain boundaries, defects). This is similar to rain happening when water molecules condensed around dust particles. During growth, the nuclei grow in size at the expense of the surrounding material.

The kinetic behavior often has the S-shape form of Fig. 10.1, when plotting percent of material transformed vs. the logarithm of time. The nucleation phase is seen as an incubation period, where nothing seems to happen. Usually the transformation rate has the form r = A e^-Q/RT (similar to the temperature dependence of the diffusion constant), in which case it is said to be thermally activated.

10.4 Multiphase Transformations

10.5 Isothermal Transformation Diagrams

The family of S-shaped curves at different temperatures can be used to construct the TTT (Time-Temperature-Transformation) diagrams (e.g., fig. 10.4.) For these diagrams to apply, one needs to cool the material quickly to a given temperature T_o before the transformation occurs, and keep it at that temperature over time. The horizontal line that indicates constant temperature T_o intercepts the TTT curves on the left (beginning of the transformation) and the right (end of the transformation); thus one can read from the diagrams when the transformation occurs. The formation of perlite shown in fig. 10.4 also indicates that the transformation occurs sooner at low temperatures, which is an indication that it is controlled by the rate of nucleation. At low temperatures, nucleation occurs fast and grain growth is reduced (since it occurs by diffusion, which is hindered at low temperatures). This reduced grain growth leads to fine-grained microstructure (fine perlite). At higher temperatures, diffusion allows for larger grain growth, thus leading to coarse perlite.

At lower temperatures nucleation starts to become slower, and a new phase is formed, bainite. Since diffusion is low at low temperatures, this phase has a very fine (microscopic) microstructure.

Spheroidite is a coarse phase that forms at temperatures close to the eutectoid temperature. The relatively high temperatures caused a slow nucleation but enhances the growth of the nuclei leading to large grains.

A very important structure is martensite, which forms when cooling austenite very fast (quenching) to below a maximum temperature that is required for the transformation. It forms nearly instantaneously when the required low temperature is reached; since no thermal activation is needed, this is called an athermal transformation. Martensite is a different phase, a body-centered tetragonal (BCT) structure with interstitial C atoms. Martensite is metastable and decomposes into ferrite and perlite but this is extremely slow (and not noticeable) at room temperature.

In the examples, we used an eutectoid composition. For hypo- and hypereutectoid alloys, the analysis is the same, but the proeutectoid phase that forms before cooling through the eutectoid temperature is also part of the final microstructure.

10.7 Mechanical Behavior of Fe-C Alloys

10.8 Tempered Martensite

11.1 Introduction

• relieve stresses
• increase softness, ductility and toughness
• produce a specific microstructure

11.2 Process Annealing

11.3 Stress Relief

11.4 Annealing of Ferrous Alloys

Full anneal involves taking hypoeutectoid alloys to the austenite phase and hypereutectoid alloys over the eutectoid temperature (Fig. 11.1) to soften pieces which have been hardened by plastic deformation, and which need to be machined.

Spheroidizing consists in prolongued heating just below the eutectoid temperature, which results in the soft spheroidite structure discussed in Sect. 10.5. This achieves maximum softness that minimizes the energy needed in subsequent forming operations.

11.5 Hardenability

Hardenability is measured by the Jominy end-quench test (Fig. 11.2). Hardenability is then given as the dependence of hardness on distance from the quenched end. High hardenability means that the hardness curve is relatively flat.
 

11.6 Influence of Quenching Medium, Specimen Size, and Geometry

The shape and size of the piece, together with the heat capacity and heat conductivity are important in determining the cooling rate for different parts of the metal piece. Heat capacity is the energy content of a heated mass, which needs to be removed for cooling. Heat conductivity measures how fast this energy is transported to the colder regions of the piece.
 
 

Precipitation Hardening

Hardening can be enhanced by extremely small precipitates that hinder dislocation motion. The precipitates form when the solubility limit is exceeded. Precipitation hardening is also called age hardening because it involves the hardening of the material over a prolonged time.

11.7 Heat Treatments

a) solution heat treatment where all the solute atoms are dissolved to form a single-phase solution.

b) rapid cooling across the solvus line to exceed the solubility limit. This leads to a supersaturated solid solution that remains stable (metastable) due to the low temperatures, which prevent diffusion.

c) precipitation heat treatment where the supersaturated solution is heated to an intermediate temperature to induce precipitation and kept there for some time (aging).

If the process is continued for a very long time, eventually the hardness decreases. This is called overaging.

The requirements for precipitation hardening are:

appreciable maximum solubility

solubility curve that falls fast with temperature

composition of the alloy that is less than the maximum solubility

11.8 Mechanism of Hardening

11.9 Miscellaneous Considerations

• solution heat treat and quench
• do needed cold working before hardening
• do precipitation hardening

Gold can be agglomerated into larger pieces by cold hammering, but native copper cannot, and an essential step toward the Metal Age was the discovery that metals such as copper could be fashioned into shapes by melting and casting in molds; among the earliest known products of this type are copper axes cast in the Balkans in the 4th millennium BC. Another step was the discovery that metals could be recovered from metal-bearing minerals. These had been collected and could be distinguished on the basis of colour, texture, weight, and flame colour and smell when heated. The notably greater yield obtained by heating native copper with associated oxide minerals may have led to the smelting process, since these oxides are easily reduced to metal in a charcoal bed at temperatures in excess of 700 C (1,300 F), as the reductant, carbon monoxide, becomes increasingly stable. In order to effect the agglomeration and separation of melted or smelted copper from its associated minerals, it was necessary to introduce iron oxide as a flux. This further step forward can be attributed to the presence of iron oxide gossan minerals in the weathered upper zones of copper sulfide deposits.

Bronze

In many regions, copper-arsenic alloys, of superior properties to copper in both cast and wrought form, were produced in the next period. This may have been accidental at first, owing to the similarity in colour and flame colour between the bright green copper carbonate mineral malachite and the weathered products of such copper-arsenic sulfide minerals as enargite, and it may have been followed later by the purposeful selection of arsenic compounds based on their garlic odour when heated.

Arsenic contents varied from 1 to 7 percent, with up to 3 percent tin. Essentially arsenic-free copper alloys with higher tin content--in other words, true bronze--seem to have appeared between 3000 and 2500 BC, beginning in the Tigris-Euphrates delta. The discovery of the value of tin may have occurred through the use of stannite, a mixed sulfide of copper, iron, and tin, although this mineral is not as widely available as the principal tin mineral, cassiterite, which must have been the eventual source of the metal. Cassiterite is strikingly dense and occurs as pebbles in alluvial deposits together with arsenopyrite and gold; it also occurs to a degree in the iron oxide gossans mentioned above.

While there may have been some independent development of bronze in varying localities, it is most likely that the bronze culture spread through trade and the migration of peoples from the Middle East to Egypt, Europe, and possibly China. In many civilizations the production of copper, arsenical copper, and tin bronze continued together for some time. The eventual disappearance of copper-arsenic alloys is difficult to explain. Production may have been based on minerals that were not widely available and became scarce, but the relative scarcity of tin minerals did not prevent a substantial trade in that metal over considerable distances. It may be that tin bronzes were eventually preferred owing to the chance of contracting arsenic poisoning from fumes produced by the oxidation of arsenic-containing minerals.

As the weathered copper ores in given localities were worked out, the harder sulfide ores beneath were mined and smelted. The minerals involved, such as chalcopyrite, a copper-iron sulfide, needed an oxidizing roast to remove sulfur as sulfur dioxide and yield copper oxide. This not only required greater metallurgical skill but also oxidized the intimately associated iron, which, combined with the use of iron oxide fluxes and the stronger reducing conditions produced by improved smelting furnaces, led to higher iron contents in the bronze.

Iron

It is not possible to mark a sharp division between the Bronze Age and the Iron Age. Small pieces of iron would have been produced in copper smelting furnaces as iron oxide fluxes and iron-bearing copper sulfide ores were used. In addition, higher furnace temperatures would have created more strongly reducing conditions (that is to say, a higher carbon monoxide content in the furnace gases). An early piece of iron from a trackway in the province of Drenthe, Neth., has been dated from 1350 BC, a date normally taken as the Middle Bronze Age for this area. In Anatolia, on the other hand, iron was in use as early as 2000 BC. There are also occasional references to iron in even earlier periods, but this material was of meteoric origin.

Once a relationship had been established between the new metal found in copper smelts and the ore added as flux, the operation of furnaces for the production of iron alone naturally followed. Certainly by 1400 BC in Anatolia, iron was assuming considerable importance, and by 1200-1000 BC it was being fashioned on quite a large scale into weapons, initially dagger blades. For this reason, 1200 BC has been taken as the beginning of the Iron Age. Evidence from excavations indicates that the art of iron making originated in the mountainous country to the south of the Black Sea, an area dominated by the Hittites. Later the art apparently spread to the Palestinians, for crude furnaces dating from 1200 BC have been unearthed at Gerar, together with a number of iron objects.

Smelting of iron oxide with charcoal demanded a high temperature, and, since the melting temperature of iron at 1,540 C (2,800 F) was not attainable then, the product was merely a spongy mass of pasty globules of metal intermingled with a semiliquid slag. This product, later known as bloom, was hardly usable as it stood, but repeated reheating and hot hammering eliminated much of the slag, creating wrought iron, a much better product.

The properties of iron are much affected by the presence of small amounts of carbon, with large increases in strength associated with contents of less than 0.5 percent. At the temperatures then attainable--about 1,200 C (2,200 F)--reduction by charcoal produced an almost pure iron, which was soft and of limited use for weapons and tools, but when the ratio of fuel to ore was increased and furnace drafting improved with the invention of better bellows, more carbon was absorbed by the iron. This resulted in blooms and iron products with a range of carbon contents, making it difficult to determine the period in which iron may have been purposely strengthened by carburizing, or reheating the metal in contact with excess charcoal.

Carbon-containing iron had the further great advantage that, unlike bronze and carbon-free iron, it could be made still harder by quenching--i.e., rapid cooling by immersion in water. There is no evidence for the use of this hardening process during the early Iron Age, so that it must have been either unknown then or not considered advantageous, in that quenching renders iron very brittle and has to be followed by tempering, or reheating at a lower temperature, to restore toughness. What seems to have been established early on was a practice of repeated cold forging and annealing at 600-700 C (1,100-1,300 F), a temperature naturally achieved in a simple fire. This practice is common in parts of Africa even today.

By 1000 BC iron was beginning to be known in central Europe. Its use spread slowly westward; iron making was fairly widespread in Great Britain at the time of the Roman invasion in 55 BC. In Asia iron was also known in ancient times, in China by about 700 BC.

To make iron you start with iron ore. Iron ore is nothing more than rock that happens to contain a high concentration of iron. One thing that gave certain countries an edge between the 15th and 20th centuries was the availability of iron ore deposits. For example, England, the U.S., France, Germany, Spain and Russia all have good iron ore deposits. When you think of the historical importance of all of these countries, you can see the correlation!

Common iron ores include:

• Hematite - Fe2O3 - 70% iron
• Magnetite - Fe3O4 - 72% iron
• Limonite - Fe2O3 + H2O - 50% to 66% iron
• Siderite - FeCO3 - 48% iron

Creating Iron

You can see in the previous section that all of the iron ores contain iron combined with oxygen. To make iron from iron ore, you need to eliminate the oxygen to create pure iron.

The most primitive facility used to refine iron from iron ore is called a bloomery. In a bloomery you burn charcoal with iron ore and a good supply of oxygen (provided by a bellows or blower). Charcoal is essentially pure carbon. The carbon combines with oxygen to create carbon dioxide and carbon monoxide (releasing lots of heat in the process). Carbon and carbon monoxide combine with the oxygen in the iron ore and carry it away, leaving iron metal.

In a bloomery the fire does not get hot enough to melt the iron completely, so you are left with a spongy mass containing iron and silicates from the ore (the bloom). By heating and hammering the bloom, the glassy silicates mix into the iron metal to create wrought iron. Wrought iron is tough and easy to work, making it perfect for creating tools in a blacksmith shop.

The more advanced way to smelt iron is in a blast furnace. A blast furnace is charged with iron ore, charcoal or coke (coke being charcoal made from coal) and limestone (CaCO4). Huge quantities of air blast in at the bottom of the furnace. The calcium in the limestone combines with the silicates to form slag. At the bottom of the blast furnace, liquid iron collects along with a layer of slag on top. Periodically you let the liquid iron flow out and cool. Typically the liquid iron flows into a channel and indentations in a bed of sand. Once it cools, this metal is known as pig iron.

To create a ton of pig iron you start with 2 tons of ore, 1 ton of coke and half a ton of limestone, and the fire consumes 5 tons of air. The temperature reaches almost 3000 degrees F (1600 or so degrees C) at the core of the blast furnace!

Pig iron contains 4% to 5% carbon and is so hard and brittle it is almost useless. You do one of two things with pig iron:

1. You melt it, mix it with slag and hammer it to eliminate most of the carbon (down to 0.3%) and create wrought iron. Wrought iron is the stuff a blacksmith works with to create tools, horseshoes, and so on. When you heat wrought iron, it is malleable, bendable, weldable and very easy to work with.
2. You create steel (see the next section).

• Relatively speaking, and especially when compared to wood or copper, iron is extremely strong.
• By heating it, iron is relatively easy to bend and shape using simple tools.
• Unlike wood, iron can handle heat, so you can build things like engines from it.
• Unlike most substances, you can magnetize iron, making it useful in the creation of electric motors and generators.
• Iron is plentiful -- 5% of the earth's crust is iron, and in some areas it concentrates in ores that contain as much as 70% iron.
• It is relatively easy to refine iron using simple tools.

An object like the flintlock rifle would be impossible to create without iron. Fortunately, iron can be created relatively easily with tools available to primitive societies. There will likely come a day when we become so technologically advanced that iron is completely replaced by aluminum, plastics and things like carbon and glass fibers. But right now the economic equation gives inexpensive iron and steel a huge advantage over these much more expensive alternatives.

Brass

While some zinc appears in bronzes dating from the Bronze Age, this was almost certainly an accidental inclusion, although it may foreshadow the complex ternary alloys of the early Iron Age, in which substantial amounts of zinc as well as tin may be found. Brass, as an alloy of copper and zinc without tin, did not appear in Egypt until about 30 BC, but after this it was rapidly adopted throughout the Roman world, for example, for currency. It was made by the calamine process, in which zinc carbonate or zinc oxide were added to copper and melted under a charcoal cover in order to produce reducing conditions. The general establishment of a brass industry was one of the important metallurgical contributions made by the Romans.

Precious metals

Bronze, iron, and brass were, then, the metallic materials on which successive peoples built their civilizations and of which they made their implements for both war and peace. In addition, by 500 BC, rich lead-bearing silver mines had opened in Greece. Reaching depths of several hundred metres, these mines were vented by drafts provided by fires lit at the bottom of the shafts. Ores were hand-sorted, crushed, and washed with streams of water to separate valuable minerals from the barren, lighter materials. Because these minerals were principally sulfides, they were roasted to form oxides and were then smelted to recover a lead-silver alloy.

Lead was removed from the silver by cupellation, a process of great antiquity in which the alloy was melted in a shallow porous clay or bone-ash receptacle called a cupel. A stream of air over the molten mass preferentially oxidized the lead. Its oxide was removed partially by skimming the molten surface; the remainder was absorbed into the porous cupel. Silver metal and any gold were retained on the cupel. The lead from the skimmings and discarded cupels was recovered as metal upon heating with charcoal.

Native gold itself often contained quite considerable quantities of silver. These silver-gold alloys, known as electrum, may be separated in a number of ways, but presumably the earliest was by heating in a crucible with common salt. In time and with repetitive treatments, the silver was converted into silver chloride, which passed into the molten slag, leaving a purified gold. Cupellation was also employed to remove from the gold such contaminates as copper, tin, and lead. Gold, silver, and lead were used for artistic and religious purposes, personal adornment, household utensils, and equipment for the chase.

From 500 BC To AD 1500

In the thousand years between 500 BC and AD 500, a vast number of discoveries of significance to the growth of metallurgy were made. The Greek mathematician and inventor Archimedes, for example, demonstrated that the purity of gold could be measured by determining its weight and the quantity of water displaced upon immersion--that is, by determining its density. In the pre-Christian portion of the period, the first important steel production was started in India, using a process already known to ancient Egyptians. Wootz steel, as it was called, was prepared as sponge (porous) iron in a unit not unlike a bloomery. The product was hammered while hot to expel slag, broken up, then sealed with wood chips in clay containers and heated until the pieces of iron absorbed carbon and melted, converting it to steel of homogeneous composition containing 1 to 1.6 percent carbon. The steel pieces could then be heated and forged to bars for later use in fashioning articles, such as the famous Damascus swords made by medieval Arab armourers.

Arsenic, zinc, antimony, and nickel may well have been known from an early date but only in the alloy state. By 100 BC mercury was known and was produced by heating the sulfide mineral cinnabar and condensing the vapours. Its property of amalgamating (mixing or alloying) with various metals was employed for their recovery and refining. Lead was beaten into sheets and pipes, the pipes being used in early water systems. The metal tin was available and Romans had learned to use it to line food containers. Although the Romans made no extraordinary metallurgical discoveries, they were responsible for, in addition to the establishment of the brass industry, contributing toward improved organization and efficient administration in mining.

Beginning about the 6th century, and for the next thousand years, the most meaningful developments in metallurgy centred on iron making. Great Britain, where iron ore was plentiful, was an important iron-making region. Iron weapons, agricultural implements, domestic articles, and even personal adornments were made. Fine-quality cutlery was made near Sheffield. Monasteries were often centres of learning of the arts of metalworking. Monks became well known for their iron making and bell founding, the products made either being utilized in the monasteries, disposed of locally, or sold to merchants for shipment to more distant markets. In 1408 the bishop of Durham established the first water-powered bloomery in Britain, with the power apparently operating the bellows. Once power of this sort became available, it could be applied to a range of operations and enable the hammering of larger blooms.

In Spain, another iron-making region, the Catalan forge had been invented, and its use later spread to other areas. A hearth type of furnace, it was built of stone and was charged with iron ore, flux, and charcoal. The charcoal was kept ignited with air from a bellows blown through a bottom nozzle, or tuyere (see figure). The bloom that slowly collected at the bottom was removed and upon frequent reheating and forging was hammered into useful shapes. By the 14th century the furnace was greatly enlarged in height and capacity.
 
 
 
 

If the fuel-to-ore ratio in such a furnace was kept high, and if the furnace reached temperatures sufficiently hot for substantial amounts of carbon to be absorbed into the iron, then the melting point of the metal would be lowered and the bloom would melt. This would dissolve even more carbon, producing a liquid cast iron of up to 4 percent carbon and with a relatively low melting temperature of 1,150 C (2,100 F). The cast iron would collect in the base of the furnace, which technically would be a blast furnace rather than a bloomery in that the iron would be withdrawn as a liquid rather than a solid lump.

While the Iron Age peoples of Anatolia and Europe on occasion may have accidently made cast iron, which is chemically the same as blast-furnace iron, the Chinese were the first to realize its advantages. Although brittle and lacking the strength, toughness, and workability of steel, it was useful for making cast bowls and other vessels. In fact, the Chinese, whose Iron Age began about 500 BC, appear to have learned to oxidize the carbon from cast iron in order to produce steel or wrought iron indirectly, rather than through the direct method of starting from low-carbon iron.

After 1500

During the 16th century, metallurgical knowledge was recorded and made available. Two books were especially influential. One, by the Italian Vannoccio Biringuccio, was entitled De la pirotechnia (Eng. trans., The Pirotechnia of Vannoccio Biringuccio, 1943). The other, by the German Georgius Agricola, was entitled De re metallica. Biringuccio was essentially a metalworker, and his book dealt with smelting, refining, and assay methods (methods for determining the metal content of ores) and covered metal casting, molding, core making, and the production of such commodities as cannons and cast-iron cannonballs. His was the first methodical description of foundry practice.

Agricola, on the other hand, was a miner and an extractive metallurgist; his book considered prospecting and surveying in addition to smelting, refining, and assay methods. He also described the processes used for crushing and concentrating the ore and then, in some detail, the methods of assaying to determine whether ores were worth mining and extracting. Some of the metallurgical practices he described are retained in principle today.

Ferrous metals

From 1500 to the 20th century, metallurgical development was still largely concerned with improved technology in the manufacture of iron and steel. In England, the gradual exhaustion of timber led first to prohibitions on cutting of wood for charcoal and eventually to the introduction of coke, derived from coal, as a more efficient fuel. Thereafter the iron industry expanded rapidly in Great Britain, which became the greatest iron producer in the world. The crucible process for making steel, introduced in England in 1740, by which bar iron and added materials were placed in clay crucibles heated by coke fires, resulted in the first reliable steel made by a melting process.

One difficulty with the bloomery process for the production of soft bar iron was that, unless the temperature was kept low (and the output therefore small), it was difficult to keep the carbon content low enough so that the metal remained ductile. This difficulty was overcome by melting high-carbon pig iron from the blast furnace in the puddling process, invented in Great Britain in 1784. In it, melting was accomplished by drawing hot gases over a charge of pig iron and iron ore held on the furnace hearth. During its manufacture the product was stirred with iron rabbles (rakes), and, as it became pasty with loss of carbon, it was worked into balls, which were subsequently forged or rolled to a useful shape. The product, which came to be known as wrought iron, was low in elements that contributed to the brittleness of pig iron and contained enmeshed slag particles that became elongated fibres when the metal was forged. Later, the use of a rolling mill equipped with grooved rolls to make wrought-iron bars was introduced.

The most important development of the 19th century was the large-scale production of cheap steel. Prior to about 1850, the production of wrought iron by puddling and of steel by crucible melting had been conducted in small-scale units without significant mechanization. The first change was the development of the open-hearth furnace by William and Friedrich Siemens in Britain and by Pierre and Émile Martin in France. Employing the regenerative principle, in which outgoing combusted gases are used to heat the next cycle of fuel gas and air, this enabled high temperatures to be achieved while saving on fuel. Pig iron could then be taken through to molten iron or low-carbon steel without solidification, scrap could be added and melted, and iron ore could be melted into the slag above the metal to give a relatively rapid oxidation of carbon and silicon--all on a much enlarged scale. Another major advance was Henry Bessemer's process, patented in 1855 and first operated in 1856, in which air was blown through molten pig iron from tuyeres set into the bottom of a pear-shaped vessel called a converter. Heat released by the oxidation of dissolved silicon, manganese, and carbon was enough to raise the temperature above the melting point of the refined metal (which rose as the carbon content was lowered) and thereby maintain it in the liquid state. Very soon Bessemer had tilting converters producing 5 tons in a heat of one hour, compared with four to six hours for 50 kilograms (110 pounds) of crucible steel and two hours for 250 kilograms of puddled iron.

Neither the open-hearth furnace nor the Bessemer converter could remove phosphorus from the metal, so that low-phosphorus raw materials had to be used. This restricted their use from areas where phosphoric ores, such as those of the Minette range in Lorraine, were a main European source of iron. The problem was solved by Sidney Gilchrist Thomas, who demonstrated in 1876 that a basic furnace lining consisting of calcined dolomite, instead of an acidic lining of siliceous materials, made it possible to use a high-lime slag to dissolve the phosphates formed by the oxidation of phosphorus in the pig iron. This principle was eventually applied to both open-hearth furnaces and Bessemer converters.

As steel was now available at a fraction of its former cost, it saw an enormously increased use for engineering and construction. Soon after the end of the century it replaced wrought iron in virtually every field. Then, with the availability of electric power, electric-arc furnaces were introduced for making special and high-alloy steels. The next significant stage was the introduction of cheap oxygen, made possible by the invention of the Linde-Frankel cycle for the liquefaction and fractional distillation of air. The Linz-Donawitz process, invented in Austria shortly after World War II, used oxygen supplied as a gas from a tonnage oxygen plant, blowing it at supersonic velocity into the top of the molten iron in a converter vessel. As the ultimate development of the Bessemer/Thomas process, oxygen blowing became universally employed in bulk steel production.

There are two process routes for making steel in the UK today: the electric arc furnace and the basic oxygen converter. The latter requires a charge of molten iron, which is produced in blast furnaces.

The raw materials for producing molten iron are iron ore, coking coal and fluxes (materials that help the chemical process) - mainly limestone. The iron ore and coal used in the UK is imported (primarily from the USA, Canada, Brazil, Australia and Scandinavia), because the UK's resources of good quality coking coal and ore are limited and not economically viable.

The coal and ore arrives by sea in very large ships and is off-loaded at deepwater harbours close to the four steelworks that use it. These steelworks are at Teesside and Scunthorpe on the North East coast, and Llanwern and Port Talbot in South Wales. The iron ores arrive in a number of forms: lumps of ore in the form in which they were mined; fine-sized iron ores; and pellets - fine ores which have been processed to stick together to form hard spheres of iron ore. The coals and ores are transported by conveyor belt or rail to stockyards where they are stored and carefully blended.

Blended coal is first heated in coke ovens to produce coke. This process is known as carbonisation. The gas produced during carbonisation is extracted and used for fuel elsewhere in the steelworks. Other by-products (such as tar and benzole) are also extracted for further refining and sale. Once carbonised, the coke is pushed out of the ovens and allowed to cool.

Fine-sized ore is first mixed with coke and fluxes and heated in a sinter plant. This is a continuous moving belt on w