Materials Science and Engineering/Glossary of Terms

Thermodynamics
Adiabatic Walls: Adiabatic walls don’t allow any exchange of heat with the surroundings.

Adiabatic Transformation: Up to a constant, $$E(X)$$ can be obtained from the amount of work $$\Delta W$$ needed for an adiabatic transformation from an initial state $$X_i$$ to a final state $$X_f$$, using $$\Delta W=E(X_f)-E(X_i)$$.

Calorie: A calorie is a unit of measurement for energy. * The small calorie or gram calorie approximates the energy needed to increase the temperature of 1 gram of water by 1 degree°C. This is about 4.184 joules.
 * The large calorie or kilogram calorie approximates the energy needed to increase the temperature of 1 kg of water by 1 °C. This is about 4.184 kJ, and exactly 1000 small calories.

In some scientific contexts such as physics and chemistry, the name "calorie" refers strictly to the gram calorie, and this unit has the symbol cal (a symbol also used by many for the large calorie).

Celsius Scale: Celsius is, or relates to, the Celsius temperature scale (previously known as the centigrade scale). Until 1954, 0 °C on the Celsius scale was defined as the melting point of ice and 100 °C was defined as the boiling point of water under a pressure of one standard atmosphere; this close equivalency is taught in schools today. However, the unit “degree Celsius” and the Celsius scale are currently, by international agreement, defined by two different points: absolute zero, and the triple point of specially prepared water. This definition also precisely relates the Celsius scale to the Kelvin scale, which is the SI base unit of temperature (symbol: K). Absolute zero—the temperature at which no energy remains in a substance—is defined as being precisely 0 K and −273.15 °C. The triple point of water is defined as being precisely 273.16 K and 0.01 °C.

Classical Thermodynamics: Classical thermodynamics is a branch of physics developed in the nineteenth century, by Sadi Carnot (1824), Emile Clapeyron (1834), Rudolf Clausius (1850), Willard Gibbs (1876), Hermann von Helmholtz (1882), and others that studied heat and work and their relation to the collision and interaction of particles in large, near-equilibrium systems.

The term classical thermodynamics is used in distinction to statistical thermodynamics, which came to be pioneered from the 1860s onwards. Statistical thermodynamics analyses thermodynamic properties by relating them to molecular-level models of microscopic behaviour in the thermodynamic system. In contrast, classical thermodynamics analyses what can be deduced solely from the macroscopic properties of the system and the laws of thermodynamics, regardless of microscopic interpretation.

Closed System: A closed system is an idealization similar to a point particle in mechanics in that it is assumed to be completely isolated by adiabatic walls that don’t allow any exchange of heat with the surroundings.

Continuum: The continuum concept ignores the fact that matter is made of atoms, is not continuous, and that it commonly has some sort of heterogeneous microstructure. It assumes that the substance of the body is distributed uniformly throughout, and completely fills the space it occupies, allowing the approximation of physical quantities, such as energy and momentum, at the infinitesimal limit. Differential equations can thus be employed in solving problems in continuum mechanics. Some of these differential equations are specific to the materials being investigated and are called constitutive equations, while others capture fundamental physical laws, such as conservation of mass or conservation of momentum and energy.

Diathermal Walls: Diathermic walls allow heat exchange for an open system.

Density: In physics, density is mass (m) per unit volume (V) — the ratio of the amount of matter in an object compared to its volume. A small, heavy object, such as a rock or a lump of lead, is denser than a larger object of the same mass, such as a piece of cork or foam.

In the common case of a homogeneous substance, density is expressed as:


 * $$\rho = \frac {m}{V}$$

where, in SI Units:
 * ρ (rho) is the density of the substance, measured in kg·m–3
 * m is the mass of the substance, measured in kg
 * V is the volume of the substance, measured in m3

Empirical Temperature: Despite its apparent simplicity, the zeroth law has the consequence of implying the existence of an important state function, the empirical temperature Θ, such that systems in equilibrium are at the same temperature.

Equilibrium: A system under study is said to be in equilibrium when its properties do not change appreciably with time over the intervals of interest (observation times). The dependence on the observation time makes the concept of equilibrium subjective. For example, window glass is in equilibrium as a solid over many decades, but flows like a fluid over time scales of millennia. At the other extreme, it is perfectly legitimate to consider the equilibrium between matter and radiation in the early universe during the first minutes of the big bang.

Extensive Quantity: Quantity proportional to the system size

First Law: The amount of work required to change the state of an otherwise adiabatically isolated system depends only on the initial and final states, and not on the means by which the work is performed, or on the intermediate stages through which the system passes.

Force Constant: The (infinitesimal) ratio of displacement to force and are generalizations of the spring constant.

Heat Capacity: The quantity of heat required (usually in Joules) to change the temperature of an object or system by one degree celsius or Kelvin. Heat capacity is a proportionality constant that depends on the size and composition of the sample. Heat is a state function and does not depend on the path to its current value, only on the objects current state. Pressure, temperature and volume are state functions. A system's current temperature, for example, does not depend on what is was yesterday.

Ideal Gas: The ideal gas occupies an important place in thermodynamics. Empirical observations indicate that the product of pressure and volume is constant along the isotherms of any gas that is sufficiently dilute. The ideal gas refers to this dilute limit of real gases, and the ideal gas temperature is proportional to the product.

Incompressible Substances:

Independent Properties:

Intensive Quantity: Quantity independent of size

Isobaric Process: An isobaric process is a thermodynamic process in which the pressure stays constant: Δp = 0. The heat transferred to the system does work but also changes the internal energy of the system. An isobaric process is shown on a P-V diagram as a straight horizontal line, connecting the initial and final thermostatic states. If the process moves towards the right, then it is an expansion. If the process moves towards the left, then it is a compression.

Isochoric Process: An isochoric process, also called an isometric process or an isovolumetric process, is a process during which volume remains constant. The name is derived from the Greek "iso" meaning "equal", and "chor" meaning "place".

If an ideal gas is used in an isochoric process, and the quantity of gas stays constant, then the increase in energy is proportional to an increase in temperature and pressure. Take for example a gas heated in a rigid container: the pressure and temperature of the gas will increase, but the volume will remain the same.

A real life example of an isochoric process is burning of a petrol-air mixture in an internal combustion engine, of any petrol/gasoline car.

Isolated System: In the natural sciences an isolated system, as contrasted with a open system, is a physical system that does not interact with its surroundings. It obeys a number of conservation laws: its total energy and mass stay constant. They cannot enter or exit, but can only move around inside.

Isothermal Process: An isothermal process is a thermodynamic process in which the temperature of the system stays constant: ΔT = 0. This typically occurs when a system is in contact with an outside thermal reservoir (heat bath), and processes occur slowly enough to allow the system to continually adjust to the temperature of the reservoir through heat exchange. An alternative special case in which a system exchanges no heat with its surroundings (Q = 0) is called an adiabatic process. During an isothermal process, all heat accepted by the system from its surroundings must have its energy entirely converted to work which it performs on the surroundings. That is, all the energy which comes into the system comes back out; the internal energy and thus the temperature of the system remain constant.

Joule: The joule is the SI unit of energy. It was named after James Prescott Joule for his work on the relationship between heat, electricity and mechanical work. One joule is the work done, or energy expended, by a force of one newton moving an object one meter along the direction of the force. This quantity is also denoted as a Newton-meter with the symbol N·m.

Kelvin Scale: The Kelvin scale is a thermodynamic (absolute) temperature scale where absolute zero — the coldest possible temperature — is zero kelvins (0 K).

Kinetic Energy: The kinetic energy of an object is the extra energy which it possesses due to its motion. It is defined as the work needed to accelerate a body of a given mass from rest to its current velocity. Having gained this energy during its acceleration, the body maintains this kinetic energy unless its speed changes. Negative work of the same magnitude would be required to return the body to a state of rest from that velocity.

Latent Energy: The latent internal energy of a system is the internal energy a system has due to phase change. Latent internal energy is the energy required for 1kg of a substance to change phase.

Macroscopic System:

Mechanical Equilibrium: A system is in mechanical equilibrium if its position in configuration space is a point at which the gradient of the potential energy is zero.

Microscopic:

Newton: The newton is the amount of force that is required to accelerate a kilogram of mass at a rate of one meter per second squared. Algebraically:


 * $${\rm 1~N = 1~\frac{kg\cdot m}{s^2}}.$$

Non-Adiabatic Transformation: In a generic (non–adiabatic) transformation, the amount of work does not equal to the change in the internal energy. The difference $$\Delta Q=\Delta E-\Delta W$$is defined as the heat intake of the system from its surroundings.

Open System: Diathermic walls allow heat exchange for an open system.

Pascal: The pascal (symbol: Pa) is the SI derived unit of pressure or stress (also: Young's modulus and tensile strength). It is a measure of perpendicular force per unit area i.e. equivalent to one newton per square meter or one Joule per cubic meter.

Path of a Process:

Phase Equilibrium: Gibbs' phase rule, stated by Josiah Willard Gibbs in the 1870s, is the fundamental rule on which phase diagrams are based.


 * F = 2 &minus; π + C

where π is the number of phases present in equilibrium (Types of solid, liquid, gas phases etc). F is the number of degrees of freedom or independent variables taken from temperature, pressure and composition of the phases present. C is the number of chemical components required to describe the system

Piezoelectric Effect: Piezoelectricity is the ability of some materials (notably crystals and certain ceramics) to generate an electric potential[1] in response to applied mechanical stress. This may take the form of a separation of electric charge across the crystal lattice. If the material is not short-circuited, the applied charge induces a voltage across the material. The word is derived from the Greek piezein, which means to squeeze or press.

The piezoelectric effect is reversible in that materials exhibiting the direct piezoelectric effect (the production of electricity when stress is applied) also exhibit the converse piezoelectric effect (the production of stress and/or strain when an electric field is applied). For example, lead zirconate titanate crystals will exhibit a maximum shape change of about 0.1% of the original dimension. The effect finds useful applications such as the production and detection of sound, generation of high voltages, electronic frequency generation, microbalances, and ultra fine focusing of optical assemblies.

Potential Energy: Potential energy can be thought of as energy stored within a physical system. This energy can be released or converted into other forms of energy, including kinetic energy. It is called potential energy because it has the potential to change the states of objects in the system when the energy is released. A formal definition is that potential energy is the energy of position, that is, the energy an object is considered to have due to its position in space.

Pressure: The force per unit area applied on a surface in a direction perpendicular to that surface. Mathematically:

p = \frac{F}{A}\ \mbox{or}\ p = \frac{dF}{dA} $$

where:
 * $$p$$ is the pressure,
 * $$F$$ is the normal force
 * $$A$$ is the area.

Process: In science, a process is every sequence of changes of a real object/body which is observable using scientific method. Therefore, all sciences analyse and model processes.

Property: A property of an object is some intrinsic or extrinsic quality of that object

Quasi-static Transformation: A quasi-static transformation is one that is performed sufficiently slowly so that the system is always in equilibrium. Thus at any stage of the process, the thermodynamic coordinates of the system existandcan in principle be computed.For such transformations, the work done on the system (equal in magnitude but opposite in sign to the work done by the system) can be related to changes in these coordinates. Typically one can divide the state functions $${X}$$ into a set of generalized displacements $${x}$$, and their conjugate generalized forces $${J}$$, such that for an infinitesimal quasi-static transformation $$dW=\sum_{i} J_i dx_i$$.

Secondary Dimensions:

Simple Compressible System:

Specific Gravity: Specific Gravity (SG) is a special case of relative density defined as the ratio of the density of a given substance, to the density of water (H2O). Substances with a specific gravity greater than 1 are heavier than water, and those with a specific gravity of less than 1 are lighter than water.

$$\mbox{SG} = \frac{\rho_\mathrm{substance}}{\rho_{\mathrm{H}_2\mathrm{O}}}$$

Specific Properties: Specific properties of a substance are derived from other intrinsic and extrinsic properties (or intensive and extensive properties) of that substance. For example, the density of steel (a specific and intrinsic property) can be derived from measurements of the mass of a steel bar (an extrinsic property) divided by the volume of the bar (another extrinsic property). Similarly, the specific gravity of a liquid is derived from the density of the liquid divided by the density of water (two intrinsic properties).

Specific Volume: Specific volume (v) is the volume occupied by a unit of mass of a material. It is equal to the inverse of density. Specific volume may be expressed in $$ \frac{m^3}{kg} $$


 * $$ v = \frac{V}{m} = \frac{1}{\rho} $$

Specific Weight: The specific weight (also known as the unit weight) is the weight per unit volume of a material, or:


 * $$\gamma = \rho \, g$$

where
 * $$\gamma$$ is the specific weight of the material (Weight per unit volume, typically N·m−3 units)
 * $$\rho$$ is the density of the material (Mass per unit volume, typically kg·m−3)
 * $$g$$ is acceleration due to gravity (rate of change of velocity, given in m·s−2)

State: In thermodynamics/statistical mechanics, a thermodynamic state, or more precisely, a macrostate, is the specification of a particular combination of physical properties (e.g. temperature, volume, pressure, etc). For example, an equation of state describes the relationship between macrostates of the system. In the same context as above, a microstate is a detailed description of a collection of atoms or other particles. There may be many microstates corresponding to the same macrostate.

State Functions: In thermodynamics, a state function, or state quantity, is a property of a system that depends only on the current state of the system, not on the way in which the system got to that state. A state function describes the equilibrium state of a system. For example, internal energy, enthalpy and entropy are state quantities because they describe quantitatively an equilibrium state of thermodynamic systems. At the same time, mechanical work and heat are process quantities because they describe quantitatively the transition between equilibrium states of thermodynamic systems.

Statistical Thermodynamics: In thermodynamics, statistical thermodynamics is the study of the microscopic behaviors of thermodynamic systems using probability theory. Statistical thermodynamics, generally, provides a molecular level interpretation of thermodynamic quantities such as work, heat, free energy, and entropy. Statistical thermodynamics was born in 1870 with the work of Austrian physicist Ludwig Boltzmann, much of which was collectively published in Boltzmann's 1896 Lectures on Gas Theory.

Thermal Equilibrium: Two systems are in thermal equilibrium when their temperatures are the same.

Thermal Responses: Change in the thermodynamic coordinates with temperature.

Thermodynamics: A phenomenological description of equilibrium properties of macroscopic systems

Thermodynamic Coordinates: Some common examples of such coordinates are pressure and volume (of a fluid), surface tension and area (of a film), tension and length (of a wire), electric field and polarization (of a dielectric)

Thermodynamic System: In thermodynamics, a thermodynamic system, originally called a working substance, is defined as that part of the universe that is under consideration. A real or imaginary boundary separates the system from the rest of the universe, which is referred to as the environment or surroundings (sometimes called a reservoir.) A useful classification of thermodynamic systems is based on the nature of the boundary and the quantities flowing through it, such as matter, energy, work, heat, and entropy. A system can be anything, for example a piston, a solution in a test tube, a living organism, a planet, etc.

Thermodynamic Temperature Scale:

Total Energy: In classical physics, the total energy of an object is the sum of its potential energy and its kinetic energy. Note that since all other forms of energy can be derived from these two types, the total energy is effectively the theoretical maximum amount of energy that could be taken from the object.

In modern physics, the total energy of an object is the sum of its rest energy, its total kinetic energy, and its potential energy.

Triple Point: In physics and chemistry, the triple point of a substance is the temperature and pressure at which three phases (gas, liquid, and solid) of that substance may coexist in thermodynamic equilibrium.

Work: In thermodynamics, work is the quantity of energy transferred from one system to another without an accompanying transfer of entropy. It is a generalization of the concept of mechanical work in mechanics. In the SI system of measurement, work is measured in joules (symbol: J). The rate at which work is performed is power.

Zeroth Law: If two systems, A and B, are separately in equilibrium with a third system C, then they are also in equilibrium with one another.