Alkaline earth metal
The alkaline earth metals are six chemical elements in group 2 of the periodic table. They are beryllium, calcium, strontium and radium; the elements have similar properties: they are all shiny, silvery-white, somewhat reactive metals at standard temperature and pressure. Structurally, they have in common an outer s- electron shell, full. All the discovered alkaline earth metals occur in nature, although radium occurs only through the decay chain of uranium and thorium and not as a primordial element. Experiments have been conducted to attempt the synthesis of element 120, the next potential member of the group, but they have all met with failure; as with other groups, the members of this family show patterns in their electronic configuration the outermost shells, resulting in trends in chemical behavior: Most of the chemistry has been observed only for the first five members of the group. The chemistry of radium is not well-established due to its radioactivity; the alkaline earth metals are all silver-colored and soft, have low densities, melting points, boiling points.
In chemical terms, all of the alkaline earth metals react with the halogens to form the alkaline earth metal halides, all of which are ionic crystalline compounds. All the alkaline earth metals except beryllium react with water to form alkaline hydroxides and, should be handled with great care; the heavier alkaline earth metals react more vigorously than the lighter ones. The alkaline earth metals have the second-lowest first ionization energies in their respective periods of the periodic table because of their somewhat low effective nuclear charges and the ability to attain a full outer shell configuration by losing just two electrons; the second ionization energy of all of the alkaline metals is somewhat low. Beryllium is an exception: It does not react with water or steam, its halides are covalent. If beryllium did form compounds with an ionization state of +2, it would polarize electron clouds that are near it strongly and would cause extensive orbital overlap, since beryllium has a high charge density.
All compounds that include beryllium have a covalent bond. The compound beryllium fluoride, the most ionic beryllium compound, has a low melting point and a low electrical conductivity when melted. All the alkaline earth metals have two electrons in their valence shell, so the energetically preferred state of achieving a filled electron shell is to lose two electrons to form doubly charged positive ions; the alkaline earth metals all react with the halogens to form ionic halides, such as calcium chloride, as well as reacting with oxygen to form oxides such as strontium oxide. Calcium and barium react with water to produce hydrogen gas and their respective hydroxides, undergo transmetalation reactions to exchange ligands. Calcium and barium all form octacarbonyl complexes and an 18-electron valence shell; the compounds were characterized in frozen neon matrices by vibrational spectroscopy and in gas phase by mass spectrometry. Analysis of the electronic structure reveals that not only barium but strontium and calcium may use their d atomic orbitals in chemical bonding, suggesting they behave like transition metals.
The table below is a summary of the key atomic properties of the alkaline earth metals. Of the six alkaline earth metals, calcium and radium have at least one occurring radioisotope. Beryllium-7, beryllium-10, calcium-41 are trace radioisotopes. Calcium-48 is the lightest nuclide to undergo double beta decay. Calcium and barium are weakly radioactive: calcium contains about 0.1874% calcium-48, barium contains about 0.1062% barium-130. The longest lived isotope of radium is radium-226 with a half-life of 1600 years; the alkaline earth metals are named after their oxides, the alkaline earths, whose old-fashioned names were beryllia, lime and baryta. These oxides are basic. "Earth" is an old term applied by early chemists to nonmetallic substances that are insoluble in water and resistant to heating—properties shared by these oxides. The realization that these earths were not elements but compounds is attributed to the chemist Antoine Lavoisier. In his Traité Élémentaire de Chimie of 1789 he called them salt-forming earth elements.
He suggested that the alkaline earths might be metal oxides, but admitted that this was mere conjecture. In 1808, acting on Lavoisier's idea, Humphry Davy became the first to obtain samples of the metals by electrolysis of their molten earths, thus supporting Lavoisier's hypothesis and causing the group to be named the alkaline earth metals; the calcium compounds calcite and lime have been used since prehistoric times. The same is true for the beryllium compounds emerald; the other compounds of the alkaline earth metals were discovered starting in the early 15th century. The magnesium compound magnesium sulfate was first discovered in 1618 by a farmer at Epsom in England. Strontium carbonate was discovered in minerals in the Scottish village of Strontian in 1790; the last element is the
A molecule is an electrically neutral group of two or more atoms held together by chemical bonds. Molecules are distinguished from ions by their lack of electrical charge. However, in quantum physics, organic chemistry, biochemistry, the term molecule is used less also being applied to polyatomic ions. In the kinetic theory of gases, the term molecule is used for any gaseous particle regardless of its composition. According to this definition, noble gas atoms are considered molecules as they are monatomic molecules. A molecule may be homonuclear, that is, it consists of atoms of one chemical element, as with oxygen. Atoms and complexes connected by non-covalent interactions, such as hydrogen bonds or ionic bonds, are not considered single molecules. Molecules as components of matter are common in organic substances, they make up most of the oceans and atmosphere. However, the majority of familiar solid substances on Earth, including most of the minerals that make up the crust and core of the Earth, contain many chemical bonds, but are not made of identifiable molecules.
No typical molecule can be defined for ionic crystals and covalent crystals, although these are composed of repeating unit cells that extend either in a plane or three-dimensionally. The theme of repeated unit-cellular-structure holds for most condensed phases with metallic bonding, which means that solid metals are not made of molecules. In glasses, atoms may be held together by chemical bonds with no presence of any definable molecule, nor any of the regularity of repeating units that characterizes crystals; the science of molecules is called molecular chemistry or molecular physics, depending on whether the focus is on chemistry or physics. Molecular chemistry deals with the laws governing the interaction between molecules that results in the formation and breakage of chemical bonds, while molecular physics deals with the laws governing their structure and properties. In practice, this distinction is vague. In molecular sciences, a molecule consists of a stable system composed of two or more atoms.
Polyatomic ions may sometimes be usefully thought of as electrically charged molecules. The term unstable molecule is used for reactive species, i.e. short-lived assemblies of electrons and nuclei, such as radicals, molecular ions, Rydberg molecules, transition states, van der Waals complexes, or systems of colliding atoms as in Bose–Einstein condensate. According to Merriam-Webster and the Online Etymology Dictionary, the word "molecule" derives from the Latin "moles" or small unit of mass. Molecule – "extremely minute particle", from French molécule, from New Latin molecula, diminutive of Latin moles "mass, barrier". A vague meaning at first; the definition of the molecule has evolved. Earlier definitions were less precise, defining molecules as the smallest particles of pure chemical substances that still retain their composition and chemical properties; this definition breaks down since many substances in ordinary experience, such as rocks and metals, are composed of large crystalline networks of chemically bonded atoms or ions, but are not made of discrete molecules.
Molecules are held together by ionic bonding. Several types of non-metal elements exist only as molecules in the environment. For example, hydrogen only exists as hydrogen molecule. A molecule of a compound is made out of two or more elements. A covalent bond is a chemical bond; these electron pairs are termed shared pairs or bonding pairs, the stable balance of attractive and repulsive forces between atoms, when they share electrons, is termed covalent bonding. Ionic bonding is a type of chemical bond that involves the electrostatic attraction between oppositely charged ions, is the primary interaction occurring in ionic compounds; the ions are atoms that have lost one or more electrons and atoms that have gained one or more electrons. This transfer of electrons is termed electrovalence in contrast to covalence. In the simplest case, the cation is a metal atom and the anion is a nonmetal atom, but these ions can be of a more complicated nature, e.g. molecular ions like NH4+ or SO42−. An ionic bond is the transfer of electrons from a metal to a non-metal for both atoms to obtain a full valence shell.
Most molecules are far too small to be seen with the naked eye. DNA, a macromolecule, can reach macroscopic sizes, as can molecules of many polymers. Molecules used as building blocks for organic synthesis have a dimension of a few angstroms to several dozen Å, or around one billionth of a meter. Single molecules cannot be observed by light, but small molecules and the outlines of individual atoms may be traced in some circumstances by use of an atomic force microscope; some of the largest molecules are supermolecules. The smallest molecule is the diatomic hydrogen, with a bond length of 0.74 Å. Effective molecular radius is the size; the table of permselectivity for different substances contains examples. The chemical formula for a molecule uses one line of chemical element symbols and sometimes al
The calorie is a unit of energy. The Calorie is 1,000 calories; that capital C, distinguishing Calorie from calorie, is a long-established scientific convention but is not always understood more widely. Where the context is about food and exercise, the term appears without the capital C; the Calorie is termed the large calorie or kilocalorie — symbols: Cal, kcal — or food calorie, defined as the heat energy involved in warming one kilogram of water by one degree Celsius. The small calorie was defined as the heat energy to raise the temperature of one gram of water — rather than a kilogram — by the same amount. Although both units relate to the metric system, they have been considered obsolete, or deprecated, in scientific usage, since the adoption of the SI system, but the small calorie is still used in laboratory measurements and calculations, with the values thus established being reported in kilocalories. The calorie was first defined by Nicolas Clément in 1824 as a unit of heat energy, it entered French and English dictionaries between 1841 and 1867.
The word comes from Latin calor, meaning'heat'. The small calorie was introduced by Pierre Antoine Favre and Johann T. Silbermann in 1852. In 1879, Marcellin Berthelot introduced the convention of capitalizing the large Calorie to distinguish the senses; the use of the calorie for nutrition was introduced to the American public by Wilbur Olin Atwater, a professor at Wesleyan University, in 1887. The alternate spelling calory is archaic; the energy needed to increase the temperature of a given mass of water by 1 °C depends on the atmospheric pressure and the starting temperature. Accordingly, several different precise definitions of the calorie have been used; the pressure is taken to be the standard atmospheric pressure. The temperature increase can be expressed as one kelvin, which means the same as an increment of one degree Celsius; the two definitions most common in older literature appear to be the 15 °C calorie and the thermochemical calorie. Until 1948, the latter was defined as 4.1833 international joules.
The calorie was first defined to measure energy in the form of heat in experimental calorimetry. In a nutritional context, the kilojoule is the SI unit of food energy, although the kilocalorie is still in common use; the word calorie is popularly used with the number of kilocalories of nutritional energy measured. To avoid confusion, it is sometimes written Calorie to make the distinction, although this is not understood. To facilitate comparison, specific energy or energy density figures are quoted as "calories per serving" or "kilocalories per 100 g". A nutritional requirement or consumption is expressed in calories per day. One gram of fat in food contains nine calories, while a gram of either a carbohydrate or a protein contains four calories. Alcohol in a food contains seven calories per gram. In other scientific contexts, the term calorie always refers to the small calorie. Though it is not an SI unit, it is still used in chemistry. For example, the energy released in a chemical reaction per mole of reagent is expressed in kilocalories per mole.
This use was due to the ease with which it could be calculated in laboratory reactions in aqueous solution: a volume of reagent dissolved in water forming a solution, with concentration expressed in moles per liter, will induce a temperature change in degrees Celsius in the total volume of water solvent, these quantities can be used to calculate energy per mole. It is occasionally used to specify energy quantities that relate to reaction energy, such as enthalpy of formation and the size of activation barriers. However, its use is being superseded by the SI unit, the joule, multiples thereof such as the kilojoule. In the past a bomb calorimeter was utilised to determine the energy content of food by burning a sample and measuring a temperature change in the surrounding water. Today this method is not used in the USA and has been succeeded by calculating the energy content indirectly from adding up the energy provided by energy-containing nutrients of food; the fibre content is subtracted to account for the fact fibre is not digested by the body
The Kelvin scale is an absolute thermodynamic temperature scale using as its null point absolute zero, the temperature at which all thermal motion ceases in the classical description of thermodynamics. The kelvin is the base unit of temperature in the International System of Units; until 2018, the kelvin was defined as the fraction 1/273.16 of the thermodynamic temperature of the triple point of water. In other words, it was defined such that the triple point of water is 273.16 K. On 16 November 2018, a new definition was adopted, in terms of a fixed value of the Boltzmann constant. For legal metrology purposes, the new definition will come into force on 20 May 2019; the Kelvin scale is named after the Belfast-born, Glasgow University engineer and physicist William Thomson, 1st Baron Kelvin, who wrote of the need for an "absolute thermometric scale". Unlike the degree Fahrenheit and degree Celsius, the kelvin is not referred to or written as a degree; the kelvin is the primary unit of temperature measurement in the physical sciences, but is used in conjunction with the degree Celsius, which has the same magnitude.
The definition implies that absolute zero is equivalent to −273.15 °C. In 1848, William Thomson, made Lord Kelvin, wrote in his paper, On an Absolute Thermometric Scale, of the need for a scale whereby "infinite cold" was the scale's null point, which used the degree Celsius for its unit increment. Kelvin calculated; this absolute scale is known today as the Kelvin thermodynamic temperature scale. Kelvin's value of "−273" was the negative reciprocal of 0.00366—the accepted expansion coefficient of gas per degree Celsius relative to the ice point, giving a remarkable consistency to the accepted value. In 1954, Resolution 3 of the 10th General Conference on Weights and Measures gave the Kelvin scale its modern definition by designating the triple point of water as its second defining point and assigned its temperature to 273.16 kelvins. In 1967/1968, Resolution 3 of the 13th CGPM renamed the unit increment of thermodynamic temperature "kelvin", symbol K, replacing "degree Kelvin", symbol °K. Furthermore, feeling it useful to more explicitly define the magnitude of the unit increment, the 13th CGPM held in Resolution 4 that "The kelvin, unit of thermodynamic temperature, is equal to the fraction 1/273.16 of the thermodynamic temperature of the triple point of water."In 2005, the Comité International des Poids et Mesures, a committee of the CGPM, affirmed that for the purposes of delineating the temperature of the triple point of water, the definition of the Kelvin thermodynamic temperature scale would refer to water having an isotopic composition specified as Vienna Standard Mean Ocean Water.
In 2018, Resolution A of the 26th CGPM adopted a significant redefinition of SI base units which included redefining the Kelvin in terms of a fixed value for the Boltzmann constant of 1.380649×10−23 J/K. When spelled out or spoken, the unit is pluralised using the same grammatical rules as for other SI units such as the volt or ohm; when reference is made to the "Kelvin scale", the word "kelvin"—which is a noun—functions adjectivally to modify the noun "scale" and is capitalized. As with most other SI unit symbols there is a space between the kelvin symbol. Before the 13th CGPM in 1967–1968, the unit kelvin was called a "degree", the same as with the other temperature scales at the time, it was distinguished from the other scales with either the adjective suffix "Kelvin" or with "absolute" and its symbol was °K. The latter term, the unit's official name from 1948 until 1954, was ambiguous since it could be interpreted as referring to the Rankine scale. Before the 13th CGPM, the plural form was "degrees absolute".
The 13th CGPM changed the unit name to "kelvin". The omission of "degree" indicates that it is not relative to an arbitrary reference point like the Celsius and Fahrenheit scales, but rather an absolute unit of measure which can be manipulated algebraically. In science and engineering, degrees Celsius and kelvins are used in the same article, where absolute temperatures are given in degrees Celsius, but temperature intervals are given in kelvins. E.g. "its measured value was 0.01028 °C with an uncertainty of 60 µK." This practice is permissible because the degree Celsius is a special name for the kelvin for use in expressing relative temperatures, the magnitude of the degree Celsius is equal to that of the kelvin. Notwithstanding that the official endorsement provided by Resolution 3 of the 13th CGPM states "a temperature interval may be expressed in degrees Celsius", the practice of using both °C and K is widespread throughout the scientific world; the use of SI prefixed forms of the degree Celsius to express a temperature interval has not been adopted.
In 2005 the CIPM embarked on a programme to redefine the kelvin using a more experimentally rigorous methodology. In particular, the committee proposed redefining the kelvin such that Boltzmann's constant takes the exact value 1.3806505×10−23 J/K. The committee had hoped tha
Pressure is the force applied perpendicular to the surface of an object per unit area over which that force is distributed. Gauge pressure is the pressure relative to the ambient pressure. Various units are used to express pressure; some of these derive from a unit of force divided by a unit of area. Pressure may be expressed in terms of standard atmospheric pressure. Manometric units such as the centimetre of water, millimetre of mercury, inch of mercury are used to express pressures in terms of the height of column of a particular fluid in a manometer. Pressure is the amount of force applied at right angles to the surface of an object per unit area; the symbol for it is p or P. The IUPAC recommendation for pressure is a lower-case p. However, upper-case P is used; the usage of P vs p depends upon the field in which one is working, on the nearby presence of other symbols for quantities such as power and momentum, on writing style. Mathematically: p = F A, where: p is the pressure, F is the magnitude of the normal force, A is the area of the surface on contact.
Pressure is a scalar quantity. It relates the vector surface element with the normal force acting on it; the pressure is the scalar proportionality constant that relates the two normal vectors: d F n = − p d A = − p n d A. The minus sign comes from the fact that the force is considered towards the surface element, while the normal vector points outward; the equation has meaning in that, for any surface S in contact with the fluid, the total force exerted by the fluid on that surface is the surface integral over S of the right-hand side of the above equation. It is incorrect to say "the pressure is directed in such or such direction"; the pressure, as a scalar, has no direction. The force given by the previous relationship to the quantity has a direction, but the pressure does not. If we change the orientation of the surface element, the direction of the normal force changes accordingly, but the pressure remains the same. Pressure is distributed to solid boundaries or across arbitrary sections of fluid normal to these boundaries or sections at every point.
It is a fundamental parameter in thermodynamics, it is conjugate to volume. The SI unit for pressure is the pascal, equal to one newton per square metre; this name for the unit was added in 1971. Other units of pressure, such as pounds per square inch and bar, are in common use; the CGS unit of pressure is 0.1 Pa.. Pressure is sometimes expressed in grams-force or kilograms-force per square centimetre and the like without properly identifying the force units, but using the names kilogram, kilogram-force, or gram-force as units of force is expressly forbidden in SI. The technical atmosphere is 1 kgf/cm2. Since a system under pressure has the potential to perform work on its surroundings, pressure is a measure of potential energy stored per unit volume, it is therefore related to energy density and may be expressed in units such as joules per cubic metre. Mathematically: p =; some meteorologists prefer the hectopascal for atmospheric air pressure, equivalent to the older unit millibar. Similar pressures are given in kilopascals in most other fields, where the hecto- prefix is used.
The inch of mercury is still used in the United States. Oceanographers measure underwater pressure in decibars because pressure in the ocean increases by one decibar per metre depth; the standard atmosphere is an established constant. It is equal to typical air pressure at Earth mean sea level and is defined as 101325 Pa; because pressure is measured by its ability to displace a column of liquid in a manometer, pressures are expressed as a depth of a particular fluid. The most common choices are water; the pressure exerted by a column of liquid of height h and density ρ is given by the hydrostatic pressure equation p = ρgh, where g is the gravitational acceleration. Fluid density and local gravity can vary from one reading to another depending on local factors, so the height of a fluid column
A supercritical fluid is any substance at a temperature and pressure above its critical point, where distinct liquid and gas phases do not exist. It can effuse through solids like a gas, dissolve materials like a liquid. In addition, close to the critical point, small changes in pressure or temperature result in large changes in density, allowing many properties of a supercritical fluid to be "fine-tuned". Supercritical fluids occur in the atmospheres of the gas giants Jupiter and Saturn, in those of the ice giants Uranus and Neptune. In a range of industrial and laboratory processes, they are used as a substitute for organic solvents. Carbon dioxide and water are the most used supercritical fluids, being used for decaffeination and power generation, respectively. In general terms, supercritical fluids have properties between those of a liquid. In Table 1, the critical properties are shown for some substances that are used as supercritical fluids. Table 2 shows density and viscosity for typical liquids and supercritical fluids.
In addition, there is no surface tension in a supercritical fluid, as there is no liquid/gas phase boundary. By changing the pressure and temperature of the fluid, the properties can be "tuned" to be more liquid-like or more gas-like. One of the most important properties is the solubility of material in the fluid. Solubility in a supercritical fluid tends to increase with density of the fluid. Since density increases with pressure, solubility tends to increase with pressure; the relationship with temperature is a little more complicated. At constant density, solubility will increase with temperature. However, close to the critical point, the density can drop with a slight increase in temperature. Therefore, close to the critical temperature, solubility drops with increasing temperature rises again. All supercritical fluids are miscible with each other so for a mixture a single phase can be guaranteed if the critical point of the mixture is exceeded; the critical point of a binary mixture can be estimated as the arithmetic mean of the critical temperatures and pressures of the two components, Tc = × TcA + × TcB.
For greater accuracy, the critical point can be calculated using equations of state, such as the Peng Robinson, or group contribution methods. Other properties, such as density, can be calculated using equations of state. Figures 1 and 2 show two-dimensional projections of a phase diagram. In the pressure-temperature phase diagram the boiling separates the gas and liquid region and ends in the critical point, where the liquid and gas phases disappear to become a single supercritical phase; the appearance of a single phase can be observed in the density-pressure phase diagram for carbon dioxide. At well below the critical temperature, e.g. 280 K, as the pressure increases, the gas compresses and condenses into a much denser liquid, resulting in the discontinuity in the line. The system consists of 2 phases in a dense liquid and a low density gas; as the critical temperature is approached, the density of the gas at equilibrium becomes higher, that of the liquid lower. At the critical point, there is no difference in density, the 2 phases become one fluid phase.
Thus, above the critical temperature a gas cannot be liquefied by pressure. At above the critical temperature, in the vicinity of the critical pressure, the line is vertical. A small increase in pressure causes a large increase in the density of the supercritical phase. Many other physical properties show large gradients with pressure near the critical point, e.g. viscosity, the relative permittivity and the solvent strength, which are all related to the density. At higher temperatures, the fluid starts to behave like a gas, as can be seen in Figure 2. For carbon dioxide at 400 K, the density increases linearly with pressure. Many pressurized gases are supercritical fluids. For example, nitrogen has a critical point of 3.4 MPa. Therefore, nitrogen in a gas cylinder above this pressure is a supercritical fluid; these are more known as permanent gases. At room temperature, they are well above their critical temperature, therefore behave as a gas, similar to CO2 at 400 K above. However, they can not be liquified by pressure.
In recent years, a significant effort has been devoted to investigation of various properties of supercritical fluids. This has been an exciting field with a long history since 1822 when Baron Charles Cagniard de la Tour discovered supercritical fluids while conducting experiments involving the discontinuities of the sound in a sealed cannon barrel filled with various fluids at high temperature. More supercritical fluids have found application in a variety of fields, ranging from the extraction of floral fragrance from flowers to applications in food science such as creating decaffeinated coffee, functional food ingredients, cosmetics, powders, bio- and functional materials, nano-systems, natural products, biotechnology and bio-fuels, microelectronics and environment. Much of the excitement and interest of the past decade is due to the enormous progress made in increasing the power of relevant experimental tools; the development of new experimental methods and improvement of existing ones continues to play an important role in this field, with recent research focusing on dynamic properties of fluids.
The Fisher-Widom line, the Widom line, or the Frenk