Boiling is the rapid vaporization of a liquid, which occurs when a liquid is heated to its boiling point, the temperature at which the vapour pressure of the liquid is equal to the pressure exerted on the liquid by the surrounding atmosphere. There are two main types of boiling: nucleate boiling where small bubbles of vapour form at discrete points, critical heat flux boiling where the boiling surface is heated above a certain critical temperature and a film of vapor forms on the surface. Transition boiling is an unstable form of boiling with elements of both types; the boiling point of water is 100 °C or 212 °F but is lower with the decreased atmospheric pressure found at higher altitudes. Boiling water is used as a method of making it potable by killing microbes; the sensitivity of different micro-organisms to heat varies, but if water is held at 70 °C for ten minutes, many organisms are killed, but some are more resistant to heat and require one minute at the boiling point of water. Boiling is used in cooking.
Foods suitable for boiling include vegetables, starchy foods such as rice and potatoes, eggs, "meats", sauces and soups. As a cooking method, it is suitable for large-scale cookery. Tough meats or poultry can be given a long, slow cooking and a nutritious stock is produced. Disadvantages include loss of water-soluble minerals. Commercially prepared foodstuffs are sometimes packed in polythene sachets and sold as "boil-in-the-bag" products. Nucleate boiling is characterized by the growth of bubbles or pops on a heated surface, which rises from discrete points on a surface, whose temperature is only above the liquids. In general, the number of nucleation sites are increased by an increasing surface temperature. An irregular surface of the boiling vessel or additives to the fluid can create additional nucleation sites, while an exceptionally smooth surface, such as plastic, lends itself to superheating. Under these conditions, a heated liquid may show boiling delay and the temperature may go somewhat above the boiling point without boiling.
As the boiling surface is heated above a critical temperature, a film of vapor forms on the surface. Since this vapor film is much less capable of carrying heat away from the surface, the temperature rises rapidly beyond this point into the transition boiling regime; the point at which this occurs is dependent on the characteristics of boiling fluid and the heating surface in question. Transition boiling may be defined as the unstable boiling, which occurs at surface temperatures between the maximum attainable in nucleate and the minimum attainable in film boiling; the formation of bubbles in a heated liquid is a complex physical process which involves cavitation and acoustic effects, such as the broad-spectrum hiss one hears in a kettle not yet heated to the point where bubbles boil to the surface. If a surface heating the liquid is hotter than the liquid film boiling will occur, where a thin layer of vapor, which has low thermal conductivity, insulates the surface; this condition of a vapor film insulating the surface from the liquid characterizes film boiling.
As a method of disinfecting water, bringing it to its boiling point at 100 °C, is the oldest and most effective way since it does not affect the taste, it is effective despite contaminants or particles present in it, is a single step process which eliminates most microbes responsible for causing intestine related diseases. Water's boiling point rests at around 100.0 degrees Celsius, when at an elevation of 0. In places having a proper water purification system, it is recommended only as an emergency treatment method or for obtaining potable water in the wilderness or in rural areas, as it cannot remove chemical toxins or impurities; the elimination of micro-organisms by boiling follows first-order kinetics—at high temperatures, it is achieved in less time and at lower temperatures, in more time. The heat sensitivity of micro-organisms varies, at 70 °C, Giardia species can take ten minutes for complete inactivation, most intestine affecting microbes and E. coli take less than a minute. Boiling does not ensure the elimination of all micro-organisms.
Thus for human health, complete sterilization of water is not required. The traditional advice of boiling water for ten minutes is for additional safety, since microbes start getting eliminated at temperatures greater than 60 °C and bringing it to its boiling point is a useful indication that can be seen without the help of a thermometer, by this time, the water is disinfected. Though the boiling point decreases with increasing altitude, it is not enough to affect the disinfecting process. Boiling is the method of cooking food in boiling water or other water-based liquids such as stock or milk. Simmering is gentle boiling; the boiling point of water is considered to be 100 °C or 212 °F. Pressure and a change in the composition of the liquid may alter the boiling point of the liquid. For this reason, high elevation cooking takes longer since boiling point is a function of atmospheric pressure. In Denver, Colorado, USA, at an elevation of about one mile, water boils at 95 °C or 203 °F. Depending on the type of food and the elevation, the boiling water may not be hot enough to cook the food properly.
Enthalpy of vaporization
The enthalpy of vaporization known as the heat of vaporization or heat of evaporation, is the amount of energy that must be added to a liquid substance, to transform a quantity of that substance into a gas. The enthalpy of vaporization is a function of the pressure; the enthalpy of vaporization is quoted for the normal boiling temperature of the substance. The heat of vaporization is temperature-dependent, though a constant heat of vaporization can be assumed for small temperature ranges and for reduced temperature T r ≪ 1; the heat of vaporization diminishes with increasing temperature and it vanishes at a certain point called the critical temperature. Above the critical temperature, the liquid and vapor phases are indistinguishable, the substance is called a supercritical fluid. Values are quoted in J/mol or kJ/mol, although kJ/kg or J/g, older units like kcal/mol, cal/g and Btu/lb are sometimes still used, among others; the enthalpy of condensation is by definition equal to the enthalpy of vaporization with the opposite sign: enthalpy changes of vaporization are always positive, whereas enthalpy changes of condensation are always negative.
The enthalpy of vaporization can be written as Δ H v a p = Δ U v a p + p Δ V It is equal to the increased internal energy of the vapor phase compared with the liquid phase, plus the work done against ambient pressure. The increase in the internal energy can be viewed as the energy required to overcome the intermolecular interactions in the liquid. Hence helium has a low enthalpy of vaporization, 0.0845 kJ/mol, as the van der Waals forces between helium atoms are weak. On the other hand, the molecules in liquid water are held together by strong hydrogen bonds, its enthalpy of vaporization, 40.65 kJ/mol, is more than five times the energy required to heat the same quantity of water from 0 °C to 100 °C. Care must be taken, when using enthalpies of vaporization to measure the strength of intermolecular forces, as these forces may persist to an extent in the gas phase, so the calculated value of the bond strength will be too low; this is true of metals, which form covalently bonded molecules in the gas phase: in these cases, the enthalpy of atomization must be used to obtain a true value of the bond energy.
An alternative description is to view the enthalpy of condensation as the heat which must be released to the surroundings to compensate for the drop in entropy when a gas condenses to a liquid. As the liquid and gas are in equilibrium at the boiling point, ΔvG = 0, which leads to: Δ v S = S g a s − S l i q u i d = Δ v H / T b As neither entropy nor enthalpy vary with temperature, it is normal to use the tabulated standard values without any correction for the difference in temperature from 298 K. A correction must be made if the pressure is different from 100 kPa, as the entropy of a gas is proportional to its pressure: the entropies of liquids vary little with pressure, as the compressibility of a liquid is small; these two definitions are equivalent: the boiling point is the temperature at which the increased entropy of the gas phase overcomes the intermolecular forces. As a given quantity of matter always has a higher entropy in the gas phase than in a condensed phase, from Δ G = Δ H − T Δ S,the Gibbs free energy change falls with increasing temperature: gases are favored at higher temperatures, as is observed in practice.
Estimation of the enthalpy of vaporization of electrolyte solutions can be carried out using equations based on the chemical thermodynamic models, such as Pitzer model or TCPC model. The vaporization of metals is a key step in metal vapor synthesis, which exploits the increased reactivity of metal atoms or small particles relative to the bulk elements. Enthalpies of vaporization of common substances, measured at their respective standard boiling points: Enthalpy of fusion Enthalpy of sublimation Joback method CODATA Key Values for Thermodynamics Gmelin, Leopold. Gmelin-Handbuch der anorganischen Chemie / 08 a. Berlin: Springer. Pp. 116–117. ISBN 978-3-540-93516-2. NIST Chemistry WebBook Young, Francis W. Sears, Mark W. Zemansky, Hugh D.. University physics. Read
Freeze drying known as lyophilisation or cryodesiccation, is a low temperature dehydration process which involves freezing the product, lowering pressure removing the ice by sublimation. This is in contrast to dehydration by most conventional methods that evaporate water using heat. Freeze drying results in a high quality product because of the low temperature used in processing; the original shape of the product is maintained and quality of the rehydrated product is excellent. Primary applications of freeze drying include biological and food processing and preservation; the first application of the freeze-drying process were in the Andes where Indigenous people would take low-land tubers up to high altitudes and leave them to freeze. The potatoes were squashed to facilitate the water loss but when left to freeze at night and exposed through the days, the mash would lose its water while the nutrients were preserved through the lyophilization process; this resulted in a product called chuño, a long shelf life food eaten well past its seasonal availability and stored for emergency rations, if needed.
Freeze drying as an industrial process began in as early as 1890 by Richard Altmann who devised a method to freeze dry tissues, but went unnoticed until the 1930s. In 1909, Shackell independently created the vacuum chamber by using an electrical pump. However, no further data on freeze drying was documented until Tival in 1927 and Elser in 1934 had patented freeze drying systems with improvements to freezing and condenser steps. A significant turning point for freeze drying occurred during World War II. Blood plasma and penicillin were needed to treat the wounded in the field, because of the lack of simultaneous refrigeration and transport, many serum supplies were spoiling before reaching their intended recipients; the freeze-drying process was developed as a commercial technique that enabled blood plasma and penicillin to be rendered chemically stable and viable without having to be refrigerated. In the 1950s-60s, freeze drying began to be viewed for its multi-purpose application to both pharmaceuticals and food processing.
Freeze-drying products became a major commodity for military food rations. What began for astronaut crews as tubed meals and freeze-dried snacks that were difficult to rehydrate, they are now able to enjoy warm hot meals while in space by improving the easability for rehydrating freeze-dried meals with water; as technology and food processing improved NASA looked for ways to provide a complete nutrient profile while reducing crumbs, disease-producing bacteria, toxins. The complete nutrient profile was improved with the addition of an algae-based vegetable-like oil to add polyunsaturated fatty acids. Polyunsaturated fatty acids are beneficial in mental and vision development, as it remains stable after space travel can provide astronauts with its added benefits; the crumb problem was solved with the addition of a gelatin coating on the foods to lock in and prevent crumbs. Disease-producing bacteria and toxins were reduced by quality control and the development of the Hazard Analysis Critical Control Point plan, used today to evaluate food material before and after processing.
With the combination of these 3 things, NASA could provide safe and wholesome foods to their crews while in space in a freeze-dried meal source. Military rations have come a long way from being served spoiled pork and corn meal to beefsteak with mushroom gravy. How rations are chosen and developed are based on acceptance, wholesomeness, producibility and sanitation. Additional requirements for rations include a minimum shelf life of 3 years, be deliverable by air, consumable in worldwide environments, provide a complete nutritional profile; the new tray rations, improved upon by increasing acceptable items and provide high quality meals while in the field. Freeze-dried coffee was incorporated by replacing spray-dried coffee within the meal, ready-to-eat category. There are four stages in the complete freeze drying process: pretreatment, primary drying, secondary drying. Pretreatment includes any method of treating the product prior to freezing; this may include concentrating the product, formulation revision, decreasing a high-vapor-pressure solvent, or increasing the surface area.
Food pieces are IQF treated to make it free flowing prior to freeze drying. In many instances the decision to pretreat a product is based on theoretical knowledge of freeze-drying and its requirements, or is demanded by cycle time or product quality considerations. During the freezing stage, the material is cooled below its triple point, the lowest temperature at which the solid and gas phases of the material can coexist; this ensures. To facilitate faster and more efficient freeze drying, larger ice crystals are preferable; the large ice crystals forms a network within the product which promotes faster removal of water vapor during sublimation. To produce larger crystals, the product should be frozen or can be cycled up and down in temperature in a process called annealing; the freezing phase is the most critical in the whole freeze-drying process, as the freezing method can impact the speed of reconstitution, duration of freeze-drying cycle, product stability, appropriate crystallization. Amorphous materials do not have a eutectic point, but they do have a critical point, below which the product must be maintained to prevent melt-back or collapse during primar
Van der Waals force
In molecular physics, the van der Waals force, named after Dutch scientist Johannes Diderik van der Waals, is a distance-dependent interaction between atoms or molecules. Unlike ionic or covalent bonds, these attractions do not result from a chemical electronic bond; the Van der Waals force vanishes at longer distances between interacting molecules. Van der Waals force plays a fundamental role in fields as diverse as supramolecular chemistry, structural biology, polymer science, surface science, condensed matter physics, it underlies many properties of organic compounds and molecular solids, including their solubility in polar and non-polar media. If no other force is present, the distance between atoms at which the force becomes repulsive rather than attractive as the atoms approach one another is called the van der Waals contact distance; the van der Waals force has the same origin as the Casimir effect, arising from quantum interactions with the zero-point field. The term van der; the term always includes the London dispersion force between instantaneously induced dipoles.
It is sometimes applied to the Debye force between a permanent dipole and a corresponding induced dipole or to the Keesom force between permanent molecular dipoles. Van der Waals forces include attraction and repulsions between atoms and surfaces, as well as other intermolecular forces, they differ from covalent and ionic bonding in that they are caused by correlations in the fluctuating polarizations of nearby particles. Being the weakest of the weakest chemical forces, with a strength between 0.4 and 4kJ/mol, they may still support an integral structural load when multitudes of such interactions are present. Such a force results from a transient shift in electron density; the electron density may temporarily shift more to one side of the nucleus. This generates a transient charge to which a nearby atom can be either repelled; when the interatomic distance of two atoms is greater than 0.6 nm the force is not strong enough to be observed. In the same vein, when the interatomic distance is below 0.4 nm the force becomes repulsive.
Intermolecular forces have four major contributions: A repulsive component resulting from the Pauli exclusion principle that prevents the collapse of molecules. Attractive or repulsive electrostatic interactions between permanent charges, quadrupoles, in general between permanent multipoles; the electrostatic interaction is sometimes called the Keesom interaction or Keesom force after Willem Hendrik Keesom. Induction, the attractive interaction between a permanent multipole on one molecule with an induced multipole on another; this interaction is sometimes called Debye force after Peter J. W. Debye. Dispersion, the attractive interaction between any pair of molecules, including non-polar atoms, arising from the interactions of instantaneous multipoles. Returning to nomenclature, different texts refer to different things using the term "van der Waals force"; some texts describe the van der Waals force as the totality of forces. All intermolecular/van der Waals forces are anisotropic, which means that they depend on the relative orientation of the molecules.
The induction and dispersion interactions are always attractive, irrespective of orientation, but the electrostatic interaction changes sign upon rotation of the molecules. That is, the electrostatic force can be attractive or repulsive, depending on the mutual orientation of the molecules; when molecules are in thermal motion, as they are in the gas and liquid phase, the electrostatic force is averaged out to a large extent, because the molecules thermally rotate and thus probe both repulsive and attractive parts of the electrostatic force. Sometimes this effect is expressed by the statement that "random thermal motion around room temperature can overcome or disrupt them"; the thermal averaging effect is much less pronounced for the attractive induction and dispersion forces. The Lennard-Jones potential is used as an approximate model for the isotropic part of a total van der Waals force as a function of distance. Van der Waals forces are responsible for certain cases of pressure broadening of spectral lines and the formation of van der Waals molecules.
The London-van der Waals forces are related to the Casimir effect for dielectric media, the former being the microscopic description of the latter bulk property. The first detailed calculations of this were done in 1955 by E. M. Lifshitz. A more general theory of van der Waals forces has been developed; the main characteristics of van der Waals forces are: They are weaker than normal covalent and ionic bonds. Van der Waals forces can not be saturated, they have no directional characteristic. They are all short-range forces and hence only interactions between the nearest particles need to be considered. Van der Waals attraction is greater. Van der Waals forces are independent
State of matter
In physics, a state of matter is one of the distinct forms in which matter can exist. Four states of matter are observable in everyday life: solid, liquid and plasma. Many other states are known to exist, such as glass or liquid crystal, some only exist under extreme conditions, such as Bose–Einstein condensates, neutron-degenerate matter, quark-gluon plasma, which only occur in situations of extreme cold, extreme density, high-energy; some other states remain theoretical for now. For a complete list of all exotic states of matter, see the list of states of matter; the distinction is made based on qualitative differences in properties. Matter in the solid state maintains a fixed volume and shape, with component particles close together and fixed into place. Matter in the liquid state maintains a fixed volume, but has a variable shape that adapts to fit its container, its particles move freely. Matter in the gaseous state has both variable shape, adapting both to fit its container, its particles are neither close together nor fixed in place.
Matter in the plasma state has variable volume and shape, but as well as neutral atoms, it contains a significant number of ions and electrons, both of which can move around freely. The term phase is sometimes used as a synonym for state of matter, but a system can contain several immiscible phases of the same state of matter. In a solid, constituent particles are packed together; the forces between particles are so strong that the particles cannot move but can only vibrate. As a result, a solid has a stable, definite shape, a definite volume. Solids can only cut. In crystalline solids, the particles are packed in a ordered, repeating pattern. There are various different crystal structures, the same substance can have more than one structure. For example, iron has a body-centred cubic structure at temperatures below 912 °C, a face-centred cubic structure between 912 and 1,394 °C. Ice has fifteen known crystal structures, or fifteen solid phases, which exist at various temperatures and pressures.
Glasses and other non-crystalline, amorphous solids without long-range order are not thermal equilibrium ground states. Solids can be transformed into liquids by melting, liquids can be transformed into solids by freezing. Solids can change directly into gases through the process of sublimation, gases can change directly into solids through deposition. A liquid is a nearly incompressible fluid that conforms to the shape of its container but retains a constant volume independent of pressure; the volume is definite if the pressure are constant. When a solid is heated above its melting point, it becomes liquid, given that the pressure is higher than the triple point of the substance. Intermolecular forces are still important, but the molecules have enough energy to move relative to each other and the structure is mobile; this means that the shape of a liquid is determined by its container. The volume is greater than that of the corresponding solid, the best known exception being water, H2O; the highest temperature at which a given liquid can exist is its critical temperature.
A gas is a compressible fluid. Not only will a gas conform to the shape of its container but it will expand to fill the container. In a gas, the molecules have enough kinetic energy so that the effect of intermolecular forces is small, the typical distance between neighboring molecules is much greater than the molecular size. A gas occupies the entire container in which it is confined. A liquid may be converted to a gas by heating at constant pressure to the boiling point, or else by reducing the pressure at constant temperature. At temperatures below its critical temperature, a gas is called a vapor, can be liquefied by compression alone without cooling. A vapor can exist in equilibrium with a liquid, in which case the gas pressure equals the vapor pressure of the liquid. A supercritical fluid is a gas whose temperature and pressure are above the critical temperature and critical pressure respectively. In this state, the distinction between liquid and gas disappears. A supercritical fluid has the physical properties of a gas, but its high density confers solvent properties in some cases, which leads to useful applications.
For example, supercritical carbon dioxide is used to extract caffeine in the manufacture of decaffeinated coffee. Like a gas, plasma does not have definite volume. Unlike gases, plasmas are electrically conductive, produce magnetic fields and electric currents, respond to electromagnetic forces. Positively charged nuclei swim in a "sea" of freely-moving disassociated electrons, similar to the way such charges exist in conductive metal, where this electron "sea" allows matter in the plasma state to conduct electricity. A gas is converted to a plasma in one of two ways. E.g. Either from a huge voltage difference between two points, or by exposing it to high temperatures. Heating matter to high temperatures causes electrons to leave the atoms, resulting in the presence of free electrons; this creates a so-called ionised plasma. At high temperatures, such as those present in stars, it is assumed that all electrons are "free", that a high-energy plasma is bare nuclei swimming in a
A chemical element is a species of atom having the same number of protons in their atomic nuclei. For example, the atomic number of oxygen is 8, so the element oxygen consists of all atoms which have 8 protons. 118 elements have been identified, of which the first 94 occur on Earth with the remaining 24 being synthetic elements. There are 80 elements that have at least one stable isotope and 38 that have radionuclides, which decay over time into other elements. Iron is the most abundant element making up Earth, while oxygen is the most common element in the Earth's crust. Chemical elements constitute all of the ordinary matter of the universe; however astronomical observations suggest that ordinary observable matter makes up only about 15% of the matter in the universe: the remainder is dark matter. The two lightest elements and helium, were formed in the Big Bang and are the most common elements in the universe; the next three elements were formed by cosmic ray spallation, are thus rarer than heavier elements.
Formation of elements with from 6 to 26 protons occurred and continues to occur in main sequence stars via stellar nucleosynthesis. The high abundance of oxygen and iron on Earth reflects their common production in such stars. Elements with greater than 26 protons are formed by supernova nucleosynthesis in supernovae, when they explode, blast these elements as supernova remnants far into space, where they may become incorporated into planets when they are formed; the term "element" is used for atoms with a given number of protons as well as for a pure chemical substance consisting of a single element. For the second meaning, the terms "elementary substance" and "simple substance" have been suggested, but they have not gained much acceptance in English chemical literature, whereas in some other languages their equivalent is used. A single element can form multiple substances differing in their structure; when different elements are chemically combined, with the atoms held together by chemical bonds, they form chemical compounds.
Only a minority of elements are found uncombined as pure minerals. Among the more common of such native elements are copper, gold and sulfur. All but a few of the most inert elements, such as noble gases and noble metals, are found on Earth in chemically combined form, as chemical compounds. While about 32 of the chemical elements occur on Earth in native uncombined forms, most of these occur as mixtures. For example, atmospheric air is a mixture of nitrogen and argon, native solid elements occur in alloys, such as that of iron and nickel; the history of the discovery and use of the elements began with primitive human societies that found native elements like carbon, sulfur and gold. Civilizations extracted elemental copper, tin and iron from their ores by smelting, using charcoal. Alchemists and chemists subsequently identified many more; the properties of the chemical elements are summarized in the periodic table, which organizes the elements by increasing atomic number into rows in which the columns share recurring physical and chemical properties.
Save for unstable radioactive elements with short half-lives, all of the elements are available industrially, most of them in low degrees of impurities. The lightest chemical elements are hydrogen and helium, both created by Big Bang nucleosynthesis during the first 20 minutes of the universe in a ratio of around 3:1 by mass, along with tiny traces of the next two elements and beryllium. All other elements found in nature were made by various natural methods of nucleosynthesis. On Earth, small amounts of new atoms are produced in nucleogenic reactions, or in cosmogenic processes, such as cosmic ray spallation. New atoms are naturally produced on Earth as radiogenic daughter isotopes of ongoing radioactive decay processes such as alpha decay, beta decay, spontaneous fission, cluster decay, other rarer modes of decay. Of the 94 occurring elements, those with atomic numbers 1 through 82 each have at least one stable isotope. Isotopes considered stable are those. Elements with atomic numbers 83 through 94 are unstable to the point that radioactive decay of all isotopes can be detected.
Some of these elements, notably bismuth and uranium, have one or more isotopes with half-lives long enough to survive as remnants of the explosive stellar nucleosynthesis that produced the heavy metals before the formation of our Solar System. At over 1.9×1019 years, over a billion times longer than the current estimated age of the universe, bismuth-209 has the longest known alpha decay half-life of any occurring element, is always considered on par with the 80 stable elements. The heaviest elements undergo radioactive decay with half-lives so short that they are not found in nature and must be synthesized; as of 2010, there are 118 known elements (in this context, "known" means observed well enough from just a few de
Solid is one of the four fundamental states of matter. In solids particles are packed, it is characterized by structural resistance to changes of shape or volume. Unlike liquid, a solid object does not flow to take on the shape of its container, nor does it expand to fill the entire volume available to it like a gas does; the atoms in a solid are bound to each other, either in a regular geometric lattice or irregularly. Solids cannot be compressed with little pressure whereas gases can be compressed with little pressure because in gases molecules are loosely packed; the branch of physics that deals with solids is called solid-state physics, is the main branch of condensed matter physics. Materials science is concerned with the physical and chemical properties of solids. Solid-state chemistry is concerned with the synthesis of novel materials, as well as the science of identification and chemical composition; the atoms, molecules or ions that make up solids may be arranged in an orderly repeating pattern, or irregularly.
Materials whose constituents are arranged in a regular pattern are known as crystals. In some cases, the regular ordering can continue unbroken over a large scale, for example diamonds, where each diamond is a single crystal. Solid objects that are large enough to see and handle are composed of a single crystal, but instead are made of a large number of single crystals, known as crystallites, whose size can vary from a few nanometers to several meters; such materials are called polycrystalline. All common metals, many ceramics, are polycrystalline. In other materials, there is no long-range order in the position of the atoms; these solids are known as amorphous solids. Whether a solid is crystalline or amorphous depends on the material involved, the conditions in which it was formed. Solids that are formed by slow cooling will tend to be crystalline, while solids that are frozen are more to be amorphous; the specific crystal structure adopted by a crystalline solid depends on the material involved and on how it was formed.
While many common objects, such as an ice cube or a coin, are chemically identical throughout, many other common materials comprise a number of different substances packed together. For example, a typical rock is an aggregate of several different minerals and mineraloids, with no specific chemical composition. Wood is a natural organic material consisting of cellulose fibers embedded in a matrix of organic lignin. In materials science, composites of more than one constituent material can be designed to have desired properties; the forces between the atoms in a solid can take a variety of forms. For example, a crystal of sodium chloride is made up of ionic sodium and chlorine, which are held together by ionic bonds. In diamond or silicon, the atoms share form covalent bonds. In metals, electrons are shared in metallic bonding; some solids most organic compounds, are held together with van der Waals forces resulting from the polarization of the electronic charge cloud on each molecule. The dissimilarities between the types of solid result from the differences between their bonding.
Metals are strong and good conductors of both electricity and heat. The bulk of the elements in the periodic table, those to the left of a diagonal line drawn from boron to polonium, are metals. Mixtures of two or more elements in which the major component is a metal are known as alloys. People have been using metals for a variety of purposes since prehistoric times; the strength and reliability of metals has led to their widespread use in construction of buildings and other structures, as well as in most vehicles, many appliances and tools, road signs and railroad tracks. Iron and aluminium are the two most used structural metals, they are the most abundant metals in the Earth's crust. Iron is most used in the form of an alloy, which contains up to 2.1% carbon, making it much harder than pure iron. Because metals are good conductors of electricity, they are valuable in electrical appliances and for carrying an electric current over long distances with little energy loss or dissipation. Thus, electrical power grids rely on metal cables to distribute electricity.
Home electrical systems, for example, are wired with copper for its good conducting properties and easy machinability. The high thermal conductivity of most metals makes them useful for stovetop cooking utensils; the study of metallic elements and their alloys makes up a significant portion of the fields of solid-state chemistry, materials science and engineering. Metallic solids are held together by a high density of shared, delocalized electrons, known as "metallic bonding". In a metal, atoms lose their outermost electrons, forming positive ions; the free electrons are spread over the entire solid, held together by electrostatic interactions between the ions and the electron cloud. The large number of free electrons gives metals their high values of electrical and thermal conductivity; the free electrons prevent transmission of visible light, making metals opaque and lustrous. More advanced models of metal properties consider the effect of the positive ions cores on the delocalised electrons.
As most metals have crystalline structure, those ions are arranged into a periodic lattice. Mathematically, the potential of the ion cores can be treated by various models, the simplest being the nearly free electron model. Minerals are