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Transition Metal

Transition metal

In chemistry, the term transition metal (sometimes also called a transition element) has two possible meanings:
- It commonly refers to any element in the d-block of the periodic table, including zinc and scandium. This corresponds exactly to periodic table groups 3 to 12 inclusive.
- More strictly, it can refer to those elements which form at least one ion with a partially filled d shell of electrons. This is exactly the d-block with zinc and scandium excluded. Both definitions have their uses and supporters. The first has the attraction of apparent simplicity and is the traditional usage. However, many interesting properties of the transition elements as a group are the result of their ability to contribute valence electrons from s orbitals before d orbitals, a property which all members of the d-block except zinc and scandium share, so the more restricted definition is in many contexts the more useful. The d orbitals are contributed after the s orbitals because once the d orbital begins to fill its electrons move closer to the nucleus, leaving the s electrons as the outermost.

The 40 transition metals

The (loosely defined) transition metals are the forty chemical elements 21 to 30, 39 to 48, 71 to 80, and 103 to 112. The name transition comes from their position in the periodic table of elements. In each of the four periods in which they occur, these elements represent the successive addition of electrons to the d atomic orbitals of the atoms. In this way, the transition metals represent the transition between group 2 elements and group 13 elements.

Electronic configuration

Main group elements prior to the appearance of the transition group elements in the periodic chart (ie, elements number 1 through 20) have no electrons in d orbitals, but only in the s and p orbitals. The 3rd period p block elements have empty d orbitals. In the fourth period from scandium to zinc, d-block elements fill up their d orbitals across the period. With the exception of the copper group and the chromium group, all d-block elements in the ground state have two electrons in their outer s orbital. The electronic configuration of the d-block elements is ns2(n-1)d1-10, where n is the ground state principal quantum number. The outer s orbitals in the d-block elements are at lower energy states than the d orbitals of the n−1 levels. As atoms always strive to be in states of lowest energy, s orbitals are filled up first. The copper (4s13d10) and chromium (4s13d5) exceptions, which have one electron in their outer orbital, occur because half- and fully-filled orbitals are more stable than any other configurations (this occurs when there are 5 or 10 electrons in the d-orbitals). Scandium has one electron in its d orbital, and 2 electrons in its outer s orbital. As scandium's only ion (Sc3+) has no electrons in its d orbital it is clear that it does not have a 'partially filled d orbital', and is not a transition metal in the stricter sense. Similarly, zinc is not a transition metal in the stricter sense because its only ion, Zn2+, has a full d orbital, which does not participate in bonding.

Properties

Transition elements tend to have high tensile strength, density and melting and boiling points. As with many properties of transition metals, this is due to d orbital electrons' ability to delocalise within the metal lattice. In metallic substances, the more electrons shared between nuclei, the stronger the metal. There are several common characteristic properties of transition elements:
- They form coloured compounds
- They can have a variety of different oxidation states
- They are good catalysts
- They are silvery-blue at room temperature (except copper and gold)
- They are solids at room temperature (except mercury)
- They form complexes, which is described by crystal field theory.

Variable oxidation states

Compared to Group II elements such as calcium, transition elements form ions with a wide variety of oxidation states. The transition metals show such a range of oxidation states because their partially filled d orbitals can accept or donate electrons in chemical reactions. Calcium ions typically do not lose more than two electrons, whereas transition metals can lose up to nine. The reason for this can be obtained by studying the ionisation enthalpies of both groups. The energies required to remove electrons from calcium are low until you try to remove electrons from below its outer two s orbitals. In fact Ca3+ has an ionisation enthalpy so high that it rarely occurs naturally. However a transition element like vanadium has roughly linear increasing ionisation enthalpies throughout its s and d orbitals, due to the close energy difference between the 3d and 4s orbitals. Transition metal ions are therefore commonly found in very high states. vanadium Certain patterns can be seen to emerge across the period of transition elements:
- The number of oxidation states of each ion increases up to Mn, after which they start to drop. This drop is due to the stronger pull from the protons in the nucleus towards the electrons, making them harder to remove.
- When the elements are in lower oxidation states, they can be found as simple ions. However elements in higher oxidation states are usually bonded covalently to electronegative compounds such as O or F, often as a polyatomic ion such as chromate, vanadate, and permanganate ions. Properties with respect to the stability of oxidation states:
- Higher oxidation state ions become less stable across the period.
- Ions in higher oxidation states tend to make good oxidising agents, whereas elements in low oxidation states become reducing agents.
- The 2+ ions across the period start as strong reducing agents, and become more stable.
- The 3+ ions start stable and become more oxidizing across the period.

Catalytic activity

Transition metals form good homogeneous or heterogeneous catalysts, for example iron is the catalyst for the Haber process. Nickel or platinum is used in the hydrogenation of alkenes.

Colored compounds

We observe color as varying frequencies of electromagnetic radiation in the visible region of the electromagnetic spectrum. Different colors result from the changed composition of light after it has been reflected, transmitted or absorbed after hitting a substance. Because of their structure, transition metals form many different colored ions and complexes. Color even varies between the different ions of a single element - MnO4− (Mn in oxidation state 7+) is a purple compound, whereas Mn2+ is pale-pink. Complex formation can play a part in determining color in a transition compound. This is because of the effect that ligands have on the 3d orbital. Ligands pull on some of the 3d electrons and split them in to higher and lower (in terms of energy) groups. Electromagnetic radiation is only absorbed if its frequency is proportional to the difference in energies between two energy states present in an atom (through the formula e=hf.) When light hits an atom which has had its 3d orbitals split, some of the electrons become promoted to the higher group. Compared to an un-complexed ion, different frequencies can be absorbed, hence different colors are observed. The color of a complex depends on:
- the nature of the metal ion, specifically the number of electrons in the d orbitals
- the arrangement of the ligands around the metal ion (for example geometric isomers can display different colors)
- the nature of the ligands surrounding the metal ion. The stronger the ligands then the greater the energy difference between the split high and low 3d groups. The complex formed by the d block element zinc (though not strictly a transition element) is colorless, because the 3d orbitals are full - no electrons are able to move up to the higher group.

See also


- inner transition element — a name given to any member of the f-block Category:Periodic table
- ja:遷移元素 th:โลหะทรานซิชัน

Chemistry

Chemistry (derived from the Arabic word kimia, alchemy, where al is Arabic for the) is the science of matter that deals with the composition, structure, and properties of substances and with the transformations that they undergo. In the study of matter, chemistry also investigates its interactions with energy and itself (see physics, biology). Because of the diversity of matter, which is mostly composed of different combinations of atoms, chemists often study how atoms of different chemical elements interact to form molecules and how molecules interact with each other. molecules

Introduction

Chemistry is a large field encompassing many subdisciplines that often overlap with significant portions of other sciences. The fundamental component of chemistry is that it involves matter in some way (this explains its broad reach). It may involve the interaction of matter with non-material phenomena such as energy. More central to chemistry is the interaction of matter with other matter such as in the classic chemical reaction where chemical bonds are broken and made, forming new molecules. Matter, such as the chair you are sitting on or the air you breathe, is known today to consist of molecules. Each molecule consists of small bits of matter known as atoms that are connected together through chemical bonds. Each atom consists of smaller bits of matter known as subatomic particles. The structure of the world we commonly experience and the properties of the matter we commonly interact with are determined by the nature of this matter on the chemical level. Steel is hard because of how the atoms are bound together. Wood will burn because it can react with oxygen in a chemical reaction. Water is a liquid at room temperature because of how each molecule of water interacts with its neighbors. In fact, you are a thinking, sentient being because of an on-going series of chemical reactions and other chemical interactions. You can see this text because of how light interacts with molecules called proteins in the back of your eye. Chemistry is often called the central science because it is what connects most of the other sciences together. Chemistry is in some ways physics on a larger scale and in some ways is biology or geology on a smaller scale. Chemistry is used to understand and make better materials for engineering. It is used to understand the chemical mechanisms of disease as well as to create pharmaceuticals to treat disease. Chemistry is somehow involved in almost every science, every technology and every "thing". With such a large area of study, it is impossible to know everything about chemistry and very difficult to summarize the field concisely. Even the most knowledgable, experienced chemist only knows a very narrow area of chemistry better than others. Of course, most chemists have a broad general knowledge of many areas of chemistry as well. Chemistry is divided into many areas of study called subdisciplines in which chemists specialize. The chemistry taught at the high school or early college level is often called "general chemistry" and is intended to be an introduction to a wide variety of fundamental concepts and to give the student the tools to continue on to more advanced subjects. Many concepts presented at this level are often incomplete and technically inaccurate yet of extraordinary utility. Chemists regularly use these simple, elegant tools and explanations in their work when they suffice because the best solution possible is often so overwhelmingly difficult and the true solution is usually unobtainable. The science of chemistry is historically a recent development but has its roots in alchemy which has been practiced for millennia throughout the world. The word chemistry is directly derived from the word alchemy, however the etymology of alchemy is unclear (see alchemy).

Subdisciplines of chemistry

Chemistry typically is divided into several major sub-disciplines. There are also several main cross-disciplinary and more specialized fields of chemistry. ; Analytical chemistry : Analytical chemistry is the analysis of material samples to gain an understanding of their chemical composition and structure. Analytical chemistry incorporates standardized experimental methods in chemistry. These methods may be used in all subdiciplines of chemistry, exluding purely theoretical chemistry. ; Biochemistry : Biochemistry is the study of the chemicals, chemical reactions and chemical interactions that take place in living organisms. Biochemistry and organic chemistry are closely related f.e. in medicinal chemistry. ; Inorganic chemistry : Inorganic chemistry is the study of the properties and reactions of inorganic compounds. The distinction between organic and inorganic disciplines is not absolute and there is much overlap, most importantly in the sub-discipline of organometallic chemistry. ; Organic chemistry : Organic chemistry is the study of the structure, properties, composition, mechanisms, and reactions of organic compounds. ; Physical chemistry : Physical chemistry or physicochemistry is the study of the physical basis of chemical systems and processes. In particular, the energetics and dynamics of such systems and processes are of interest to physical chemists. Important areas of study include chemical thermodynamics, chemical kinetics, electrochemistry, statistical mechanics, and spectroscopy. Physical chemistry has large overlap with molecular physics. ; Theoretical chemistry : Theoretical chemistry is the study of chemistry via theoretical reasoning (usually within mathematics or physics). In particular the application of quantum mechanics to chemistry is called quantum chemistry. Since the end of the second world war, the development of computers has allowed a systematic development of computational chemistry, which is the art of developing and applying computer programs for solving chemical problems. Theoretical chemistry has large overlap with molecular physics. ; Other fields : Astrochemistry, Atmospheric chemistry, Chemical Engineering, Electrochemistry, Environmental chemistry, Geochemistry, History of chemistry, Materials science, Medicinal chemistry, Molecular Biology, Molecular genetics, Nuclear chemistry, Organometallic chemistry, Petrochemistry, Pharmacology, Photochemistry, Phytochemistry, Polymer chemistry, Supramolecular chemistry, Surface chemistry, and Thermochemistry.

Fundamental concepts

Nomenclature

Nomenclature refers to the system for naming chemical compounds. There are well-defined systems in place for naming chemical species. Organic compounds are named according to the organic nomenclature system. Inorganic compounds are named according to the inorganic nomenclature system. See also: IUPAC nomenclature

Atoms

Main article: Atom. An atom is a collection of matter consisting of a positively charged core (the nucleus) which contains protons and neutrons, and which maintains a number of electrons to balance the positive charge in the nucleus.

Elements

Main article: Chemical element. An element is a class of atoms which have the same number of protons in the nucleus. This number is known as the atomic number of the element. For example, all atoms with 6 protons in their nuclei are atoms of the chemical element carbon, and all atoms with 92 protons in their nuclei are atoms of the element uranium. The most convenient presentation of the elements is in the periodic table, which groups elements with similar chemical properties together. Lists of the elements by name, by symbol, and by atomic number are also available. See also: isotope

Compounds

Main article: Chemical compound A compound is a substance with a fixed ratio of chemical elements which determines the composition, and a particular organisation which determines chemical properties. For example, water is a compound containing hydrogen and oxygen in the ratio of two to one, with the Oxygen between the hydrogens, and an angle of 104.5° between them. Compounds are formed and interconverted by chemical reactions.

Molecules

Main article: Molecule. A molecule is the smallest indivisible portion of a pure compound that retains a set of unique chemical properties. A molecule consists of two or more atoms covalently bonded together.

Ions

Main article: Ion. An ion is a charged species, or an atom or a molecule that has lost or gained an electron. Positively charged cations (e.g. sodium cation Na+) and negatively charged anions (e.g. chloride Cl-) can form neutral salts (e.g. sodium chloride NaCl). Examples of polyatomic ions that do not split up during acid-base reactions are hydroxide (OH-), or phosphate (PO43-).

Bonding

Main article: Chemical bond. A chemical bond is an interaction which holds together atoms in molecules or crystals. In many simple compounds, valence bond theory and the concept of oxidation number can be used to predict molecular structure and composition. Similarly, theories from classical physics can be used to predict many ionic structures. With more complicated compounds, such as metal complexes, valence bond theory fails and alternative approaches which are based on quantum chemistry, such as molecular orbital theory, are necessary.

States of matter

Main article: Phase (matter). A phase is a set of states of a chemical system that have similar bulk structural properties, over a range of conditions, such as pressure or temperature. Physical properties, such as density and refractive index tend to fall within values characteristic of the phase. The phase of matter is defined by the phase transition, which is when energy put into or taken out of the system goes into rearranging the structure of the system, instead of changing the bulk conditions. Sometimes the distinction between phases can be continuous instead of having a discrete boundary, in this case the matter is considered to be in a supercritical state. When three states meet based on the conditions, it is known as a triple point and since this is invariant, it is a convenient way to define a set of conditions. The most familiar examples of phases are solids, liquids, and gases. Less familiar phases include plasmas, Bose-Einstein condensates and fermionic condensates and the paramagnetic and ferromagnetic phases of magnetic materials. Even the familiar ice has many different phases, depending on the pressure and temperature of the system. While most familiar phases deal with three-dimensional systems, it is also possible to define analogs in two-dimensional systems, which is getting a lot of attention because of its relevance to biology.

Chemical reactions

Main article: Chemical reaction. Chemical reactions are transformations in the fine structure of molecules. Such reactions can result in molecules attaching to each other to form larger molecules, molecules breaking apart to form two or more smaller molecules, or rearrangement of atoms within or across molecules. Chemical reactions usually involve the making or breaking of chemical bonds.

Quantum chemistry

Main article: Quantum chemistry. Quantum chemistry describes the behavior of matter at the molecular scale. It is, in principle, possible to describe all chemical systems using this theory. In practice, only the simplest chemical systems may realistically be investigated in purely quantum mechanical terms, and approximations must be made for most practical purposes (e.g., Hartree-Fock, post Hartree-Fock or Density functional theory, see computational chemistry for more details). Hence a detailed understanding of quantum mechanics is not necessary for most chemistry, as the important implications of the theory (principally the orbital approximation) can be understood and applied in simpler terms.

Laws

The most fundamental concept in chemistry is the law of conservation of mass, which states that there is no detectable change in the quantity of matter during an ordinary chemical reaction. Modern physics shows that it is actually energy that is conserved, and that energy and mass are related; a concept which becomes important in nuclear chemistry. Conservation of energy leads to the important concepts of equilibrium, thermodynamics, and kinetics. Further laws of chemistry elaborate on the law of conservation of mass. Joseph Proust's law of definite composition says that pure chemicals are composed of elements in a definite formulation; we now know that the structural arrangement of these elements is also important. Dalton's law of multiple proportions says that these chemicals will present themselves in proportions that are small whole numbers (i.e. 1:2 O:H in water); although in many systems (notably biomacromolecules and minerals) the ratios tend to require large numbers, and are frequently represented as a fraction. Such compounds are known as Non-Stoichiometric Compounds More modern laws of chemistry define the relationship between energy and transformations.
- In equilibrium, molecules exist in mixture defined by the transformations possible on the timescale of the equilibrium, and are in a ratio defined by the intrinsic energy of the molecules—the lower the intrinsic energy, the more abundant the molecule.
- Transforming one structure to another requires the input of energy to cross an energy barrier; this can come from the intrinsic energy of the molecules themselves, or from an external source which will generally accelerate transformations. The higher the energy barrier, the slower the transformation occurs.
- There is a hypothetical intermediate, or transition structure, that corresponds to the structure at the top of the energy barrier. The Hammond-Leffler Postulate states that this structure looks most similar to the product or starting material which has intrinsic energy closest to that of the energy barrier. Stabilizing this hypothetical intermediate through chemical interaction is one way to achieve catalysis.
- All chemical processes are reversible (law of microscopic reversibility) although some processes have such an energy bias, they are essentially irreversible.

History of chemistry


- Alchemy
- Discovery of the chemical elements
- History of chemistry
- Nobel Prize in chemistry
- Timeline of chemical element discovery

Etymology

Old French: alkemie; Arab al-kimia: the art of transformation. See also: alchemy

See also


- American Chemical Society
- Chemical engineering
- Chemist and list of chemists
- International Union of Pure and Applied Chemistry
- List of chemistry topics
- List of compounds
- List of important publications in chemistry
- Periodic table
- Chemistry resources
- Valency number

External links


- [http://www.allchemicals.info/ Chemical Glossary]
- [http://chem.sis.nlm.nih.gov/chemidplus/ Chemistry Information Database includes basic information and some toxicity]
- [http://www.chem.qmw.ac.uk/iupac/ IUPAC Nomenclature Home Page], see especially the "Gold Book" containing definitions of standard chemical terms
- [http://www.cci.ethz.ch/index.html Experiments] videos and photos of the techniques and results
- [http://physchem.ox.ac.uk/MSDS/ Material safety data sheets for a variety of chemicals]
- [http://www.flinnsci.com/search_MSDS.asp Material Safety Data Sheets]

Further reading


- Chang, Raymond. Chemistry 6th ed. Boston: James M. Smith, 1998. ISBN 0071152210. Category:School subjects als:Chemie ko:화학 ms:Kimia ja:化学 simple:Chemistry th:เคมี



Scandium

Scandium is a chemical element in the periodic table that has the symbol Sc and atomic number 21. A soft, silvery, white transition element, scandium occurs in rare minerals from Scandinavia and it is sometimes classified along with yttrium and the lanthanides as a rare earth.

Notable characteristics

Scandium is a rare, soft, silvery, trivalent, very light metallic element that develops a slightly yellowish or pinkish cast when exposed to air. This element resembles yttrium and rare earth metals more than it resembles aluminium or titanium (which are closer on the periodic table). The most common oxidation state of scandium is +3 and this metal is not attacked by a 1:1 mixture of HNO3 and 48% HF.

Applications

Approximately 20 kg (as Sc2O3) of scandium are used annually in the United States to make high-intensity lights. Scandium iodide added to mercury-vapor lamps produces a highly efficient artificial light source that resembles sunlight and allows good color reproduction with TV cameras. About 80 kg of scandium is used in lightbulbs globally per year. The radioactive isotope Sc-46 is used in oil refinery crackers as a tracing agent. The main application by volume is in aluminium-scandium alloys for the aerospace industry and for sports equipment (bikes, baseball bats, etc.) which rely on high performance materials. When added to aluminium, scandium can produce improvements in strength (at ambient and elevated temperature), ductility, aging response and grain refinement through the formation of the Al3Sc phase. Furthermore, it has been shown to reduce solidification cracking during the welding of high strength aluminium alloys.

History

Dmitri Mendeleev used his periodic law, in 1869, to predict the existence and some properties of three unknown elements including one he called ekaboron . Lars Fredrick Nilson and his team, apparently unaware of that prediction in the spring of 1879, were looking for rare earth metals; using spectrum analysis he found a new element within the minerals euxenite and gadolinite. He named it Scandium, from the Latin Scandia meaning "Scandinavia", and by way of isolating the element he processed 10 kilograms of euxenite with other rare-earth residues, obtaining about 2 grams of very pure scandium oxide (Sc2O3). Per Teodor Cleve concluded that scandium corresponded well to the hoped-for ekaboron, and notified Mendeleev of this in August. Metallic scandium was prepared for the first time in 1937, by electrolysis of a eutectic melt of potassium, lithium, and scandium chlorides at 700 to 800° C. Tungsten wire in a pool of liquid zinc were the electrodes in a graphite crucible. The first pound of 99% pure scandium metal wasn't produced until 1960.

Occurrence

Rare minerals from Scandinavia and Malagasy such as thortveitite, euxenite and gadolinite are the only known concentrated sources of this element (which is never found as a free metal). Element 21 is the 23rd most abundant element in the sun and similar stars but on earth it is only the 50th most abundant element. Scandium is distributed widely on earth, occurring in trace quantities in over 800 minerals. The blue color of the aquamarine variety of beryl is thought to be caused by scandium. It is an important part of the rare mineral thortveitite and is found in residues that remain after tungsten is extracted from Zinnwald wolframite. Thortveitite is the primary source of scandium with uranium mill tailings by-products also being an important source. Pure scandium is commercially produced by reducing scandium fluoride with calcium metal. The main source of scandium is from military stockpiles from the former Soviet Union, which were themselves extracted from uranium tailings. There is no primary production in the Americas or Europe.

Isotopes

Naturally occurring scandium is composed of 1 stable isotope Sc-45. 13 radioisotopes have been characterized with the most stable being Sc-46 with a half-life of 83.79 days, Sc-47 with a half-life of 3.3492 days, and Sc-48 with a half-life of 43.67 hours. All of the remaining radioactive isotopes have half-lifes that are less than 4 hours and the majority of these have half lifes that are less than 2 minutes. This element also has 5 meta states with the most stable being Scm-44 (t½ 58.6 h). The isotopes of scandium range in atomic weight from 39.978 amu (Sc-40) to 53.963 amu (Sc-54). The primary decay mode before the only stable isotope, Sc-45, is electron capture and the primary mode after is beta emission. The primary decay products before Sc-45 are element 20 (calcium) isotopes and the primary products after are element 22 (titanium) isotopes.

Precautions

Scandium metal powder is combustible and presents a fire hazard.

References


- [http://periodic.lanl.gov/elements/21.html Los Alamos National Laboratory – Scandium]

External links


- [http://www.webelements.com/webelements/elements/text/Sc/index.html WebElements.com – Scandium] Category:Chemical elements Category:Transition metals ja:スカンジウム th:สแคนเดียม

Periodic table group

A periodic table group is a vertical column in the periodic table of the chemical elements. There are 18 groups in the standard periodic table. It is no accident that several of these correspond directly to chemical series: the periodic table was originally created to organize the known chemical series into a single coherent scheme. The modern explanation of the pattern of the periodic table is that the elements in a group have similar configurations of the outermost electron shells of their atoms: as most chemical properties are dominated by outer electron interactions, this tends to give elements in the same group similar physical and chemical properties.

Group numbers

There are three ways of numbering the groups of the periodic table, one using Arabic numerals and the other two using Roman numerals. The Roman numeral names are the original traditional names of the groups; the Arabic numeral names are those recommended by the International Union of Pure and Applied Chemistry (IUPAC) to replace the old names in an attempt to reduce the confusion generated by the two older, but mutually confusing, schemes. There is considerable confusion surrounding the two old systems in use (old IUPAC and CAS) that combined the use of Roman numerals with letters. In the old IUPAC system the letters A and B were designated to the left (A) and right (B) part of the table, while in the CAS system the letters A and B were designated to main group elements (A) and transition elements (B). The former system was frequently used in Europe while the latter was most common in America. The new IUPAC scheme was developed to replace both systems as they confusingly used the same names to mean different things. The periodic table groups are as follows (in the brackets are shown the old systems: European and American):
- Group 1 (IA,IA): the alkali metals
- Group 2 (IIA,IIA): the alkaline earth metals
- Group 3 (IIIA,IIIB)
- Group 4 (IVA,IVAB)
- Group 5 (VA,VB)
- Group 6 (VIA,VIB)
- Group 7 (VIIA,VIIB)
- Group 8 (VIII)
- Group 9 (VIII)
- Group 10 (VIII)
- Group 11 (IB,IB): the coinage metals (not a IUPAC-recommended name)
- Group 12 (IIB,IIB)
- Group 13 (IIIB,IIIA): the boron group
- Group 14 (IVB,IVA): the carbon group
- Group 15 (VB,VA): the pnictogens (not a IUPAC-recommended name) or nitrogen group
- Group 16 (VIB,VIA): the chalcogens
- Group 17 (VIIB,VIIA): the halogens
- Group 18 (Group 0): the noble gases ---- Note: Wikipedia style should be to replace the old names of the groups with the new IUPAC names throughout, with a historical mention of the old name where appropriate.
- Category:Periodic table ko:주기율표 족 ja:元素の族 th:หมู่ในตารางธาตุ

Ion

: This article is about the electrically charged particle. For other uses of this word, see ion (disambiguation). An ion is an atom or group of atoms with a net electric charge. A negatively charged ion, which has more electrons in its electron shell than it has protons in its nucleus, is known as an anion, for it is attracted to anodes, and a positively charged ion, which has fewer electrons than protons, is known as a cation (pronounced cat-eye-on), for it is attracted to cathodes. The process of converting into ions and the state of being ionized is called ionization. The recombining of ions and electrons to form neutral atoms is called recombination. Polyatomic anions which contain oxygen are sometimes known as oxyanion. Atomic and polyatomic ions are denoted by a superscript with the sign of the net electric charge and the number of electrons lost or gained, if more than one. For example: H+, SO32−. A collection of non-aqueous ions, or even a gas containing a proportion of charged particles, is called a plasma, which is called the fourth state of matter because its properties are quite different from solids, liquids, and gases.

Ionization potential

The energy required to detach an electron in its lowest energy state from an atom or molecule of a gas with less net electric charge is called the ionization potential, or ionization energy. The nth ionization energy of an atom is the energy required to detach its nth electron after the first n − 1 electrons have already been detached. Each successive ionization energy is markedly greater than the last. Particularly great increases occur after any given block of atomic orbitals is exhausted of electrons. For this reason, ions tend to form in ways that leave them with full orbital blocks. For example, sodium has one valence electron, in its outermost shell, so in ionized form it is commonly found with one lost electron, as Na+. On the other side of the periodic table, chlorine has seven valence electrons, so in ionized form it is commonly found with one gained electron, as Cl−. Francium has the lowest ionization energy of all the elements and fluorine has the greatest.

Other ions

A dianion is a species which has two negative charges on it. For example, the dianion of pentalene is aromatic. A zwitterion is an ion with a net charge of zero, but has both a positive and negative charge on it.

History

Ions were first theorized by Michael Faraday around 1830, to describe the portions of molecules that travel either to an anode or to a cathode. However, the mechanism by which this was achieved was not described until 1884 by Svante August Arrhenius in his doctoral dissertation to the University of Uppsala. His theory was initially not accepted but his dissertation won the Nobel Prize in Chemistry in 1903.

Etymology

The word ion is a name given by Michael Faraday, from Greek , neutral present participle of , "to go", thus "a goer". So, anion, , and cation, κ, mean "(a thing) going up" and "(a thing) going down", respectively, and anode, , and cathode, κ, mean "a going up" and "a going down", respectively, from , "way".

Applications

Ions are essential to life. Sodium, potassium, calcium and other ions play an important role in the cells of living organisms, particularly in cell membranes. They have many practical, everyday applications in items such as smoke detectors and are also finding use in unconventional technologies such as ion engines and ion cannons. Category:Physical chemistry ko:이온 ms:Ion ja:イオン simple:Ion th:ไอออน

Electron configuration

In atomic physics, the electron configuration is the arrangement of electrons in an atom, molecule or other body. Specifically, it is the placement of electrons into atomic, molecular, or other forms of electron orbitals. electron orbital

Electron configuration in atoms

The discussion below presumes knowledge of material contained at Atomic orbital.

Summary of the quantum numbers

The state of an electron in an atom is given by four quantum numbers. Three of these are properties of the atomic orbital in which it sits (a more thorough explanation is given in that article).
- The principal quantum number is denoted n, and can take any integer value greater than or equal to 1. It represents in part the overall energy of the orbital, and by extension its general distance from the nucleus.
- The azimuthal quantum number is denoted l, and can take any integer value in the range 0 \le l \le n-1. It determines the orbital's angular momentum.
- The magnetic quantum number is denoted m, and can take any integer value in the range -l \le m \le l. This number determines the energy shift of an atomic orbital due to an external magnetic field (Zeeman effect).
- The spin quantum number is denoted s and can take the values +1/2 or -1/2 (sometimes referred to as "up" and "down"). The spin is an intrinsic property of the electron and independent of the other numbers. The spin (in conjunction with the angular momentum) in part determines the magnetic dipole moment of the electron.

Shells and subshells

Shells and subshells are defined by the aforementioned quantum numbers, NOT by the distance of its electrons from the nucleus. As a matter of fact, in large atoms shells above the second shell overlap (see Aufbau principle). States with the same value of n are related, and said to lie within the same electron shell. States with the same value of n and also l are said to lie within the same electron subshell. If the states also share the same value of m, they are said to lie in the same atomic orbital. Because electrons have only two possible spin states, an atomic orbital cannot contain more than two electrons (Pauli exclusion principle). From this, we can deduce that the total electron capacity of a subshell is 4l+2, and that of a shell is 2n2.

Worked example

Here is the electron configuration for a filled fifth shell: This information can be written as 5s2 5p6 5d10 5f14 5g18 (see below for more details on notation). The subshell labels s, p, d, and f originate from a now-discredited system of categorizing spectral lines as "sharp", "principal", "diffuse", or "fundamental", based on their observed fine structure. When the first four types of orbitals were described, they were associated with these spectral line types, but there were no other names. The designation g was derived by following alphabetical order. Shells with more than five subshells are theoretically permissible, but this covers all discovered elements.

Notation

Physicists and chemists use a standard notation to describe atomic electron configurations. In this notation, a subshell is written in the form nxe, where n is the shell number, x is the subshell label and e is the number of electrons in the subshell. An atom's subshells are written in order of increasing energy - in other words, the sequence in which they are filled (see Aufbau principle below). For instance, ground-state hydrogen has one electron in the s subshell of the first shell, so its configuration is written 1s1. Lithium has two electrons in 1s subshell and one in the (higher-energy) 2s subshell, so its ground-state configuration is written 1s2 2s1. Phosphorus (atomic number 15), is as follows: 1s2 2s2 2p6 3s2 3p3. For atoms with many electrons, this notation can become lengthy. It is often abbreviated by noting that the first few subshells are identical to those of one or another noble gas. Phosphorus, for instance, differs from neon (1s2 2s2 2p6) only by the presence of a third shell. Thus, the electron configuration of neon is pulled out, and phosphorus is written as follows: [Ne]3s2 3p3. An even simpler version is simply to quote the number of electrons in each shell, e.g. (again for phosphorus): 2-8-5.

Aufbau principle

In the ground state of an atom (the condition in which it is ordinarily found), the electron configuration generally follows Aufbau principle. According to this principle, electrons enter into states in order of the states' increasing energy; i.e., the first electron goes into the lowest-energy state, the second into the next lowest, and so on. The order in which the states are filled is as follows: A pair of electrons with identical spins has slightly more energy than a pair of electrons with opposite spins. Since two electrons in the same orbital must have opposite spins, this causes electrons to prefer to occupy different orbitals. This preference manifests itself if a subshell with l>0 (one that contains more than one orbital) is less than full. For instance, if a p subshell contains four electrons, two electrons will be forced to occupy one orbital, but the other two electrons will occupy both of the other orbitals, and their spins will be equal. This phenomenon is called Hund's rule. The Aufbau principle can be applied, in a modified form, to the protons and neutrons in the atomic nucleus (see the shell model of nuclear physics).

Exceptions

A d subshell that is half-filled or full (ie 5 or 10 electrons) is more stable than the s subshell of the next shell. For instance, copper (atomic number 29) has a configuration of [Ar]4s1 3d10, not [Ar]4s2 3d9 as one would expect by the Aufbau principle. Likewise, chromium (atomic number 24) has a configuration of [Ar]4s1 3d5, not [Ar]4s2 3d4. This can be most easily understood by stepping through the electron configuration shown at [http://www.webelements.com/webelements/elements/text/Zn/econ.html]

Relation to the structure of the periodic table

Electron configuration is intimately related to the structure of the periodic table. The chemical properties of an atom are largely determined by the arrangement of the electrons in its outermost ("valence") shell (although other factors, such as atomic radius, atomic mass, and increased accessibility of additional electronic states also contribute to the chemistry of the elements as atomic size increases).

Electron configuration in molecules

In molecules, the situation becomes more complex, as each molecule has a different orbital structure. See the molecular orbital article and the linear combination of atomic orbitals method for an introduction and the computational chemistry article for more advanced discussions.

Electron configuration in solids

In a solid, the electron states become very numerous. They cease to be discrete, and effectively blend together into continuous ranges of possible states (an electron band). The notion of electron configuration ceases to be relevant, and yields to band theory.

See also


- Atomic electron configuration table
- Periodic table (electron configurations)
- Atomic orbital
- Energy level
- Molecular term symbol
- HOMO/LUMO Category:Chemical properties Category:Atomic physics Category:Molecular physics Category:Quantum chemistry ja:電子配置

Valence electron

In chemistry, valence electrons are the electrons located within the outermost energy level of an atom. These electrons participate in chemical reactions. Elements with a full outer shell are very unreactive. Elements with an almost full or almost empty outer shell, such as the alkali metals and halogens, tend to be very reactive.

Determining the number of valence electrons

To determine the quantity of valence electrons an element has, one must look at the family (vertical column) in which the element is categorized. With the exception of families 3–12 (transition metals), the number within the unit's place identifies how many valence electrons are contained within the elements listed under that particular column. The asterisk (
- ) is used to indicate that all of the noble gases contain eight valence electrons with the exception of helium (He), which only contains two.

Valence electrons in chemical reactions

The chemical behavior of atoms is largely due to interactions between electrons. Electrons of an atom remain within certain, predictable electron configurations. Electrons fall into shells based on their relative energy level which is usually visualized as their mean distance from the nucleus. The valence electrons have the greatest influence on chemical behavior. Core electrons (those not in the outer shell) play a role, but it is usually in terms of a secondary effect due to screening of the positive charge in the atomic nucleus. shells of hydrogen. The principal quantum number is at the right of each row and the azimuthal quantum number is denoted by letter at top of each column.]] Each shell, numbered from the one closest to the nucleus (lowest in energy), can hold up to a specific number of electrons due to its differing sublevel and orbital capacity: To determine the electron capacity of a shell, the formula 2n² is used, where n is the shell number or principal quantum number. Electrons fill orbitals and shells from the inside out, beginning with shell one. Whichever occupied shell is currently most outward is the valence shell, even if it only has one electron. The reason why shells fill up in order is that the energy levels of electrons in the innermost shells are significantly lower than the energy levels of electrons in outer shells. So if the inner shells were not completely full, the electron in an outer shell would quickly "fall" into the inner shell (with the emission of a photon that would carry away the difference in the energy). The number of electrons in an atom's outermost valence shell governs its bonding behavior. Therefore, elements with the same number of valence electrons are grouped together in the periodic table of the elements. As a general rule, the fewer electrons in an atom's valence shell, the more reactive it is. Group 1 metals are therefore very reactive, with caesium, rubidium, and francium being the most reactive of all metals. Every atom is much more stable, or less reactive, with a full valence shell. This can be achieved one of two ways: an atom can either share electrons with neighboring atoms, a covalent bond, or it can remove electrons from other atoms, an ionic bond. Another form of ionic bonding involves an atom giving some of its electrons to another atom; this also works because it can end up with a full valence by giving up its entire outer shell. By moving electrons, the two atoms become linked. This is known as chemical bonding and serves to build atoms into molecules or ionic compounds. Five major types of bonds exist:
- ionic bonds
- covalent bonds
- coordinate covalent bonds
- hydrogen bonds
- metallic bonds

External links


- [http://encarta.msn.com/dictionary_/valence%2520electrons.html MSN Encarta Definition]
- [http://dl.clackamas.cc.or.us/ch104-06/valence_electrons.html Additional Information] Category:Chemical bonding Category:Electron

Atomic orbital

orbitals]] A less formal description of the electrons in atoms can be found at Electron configuration. In quantum mechanics, the state of an atom, i.e. the eigenstates of the atomic Hamiltonian, are expanded (see configuration interaction expansion and basis (linear algebra)) into linear combinations of anti-symmetrized products (Slater determinants) of one-electron functions. The spatial components of these one-electron functions are called atomic orbitals. (When one considers also their spin component, one speaks of atomic spin orbitals.) In atomic physics, the atomic spectral lines correspond to transitions (quantum leaps) between quantum states of an atom. These states are labelled by a set of quantum numbers summarized in the term symbol and usually associated to particular electron configurations, i.e. by occupations schemes of atomic orbitals (e.g. 1s^2 2s^2 2p^6 for the ground state of Neon -- term symbol: ^1S_0). This notation means that the corresponding Slater determinants have a clear higher weight in the configuration interaction expansion. The atomic orbital concept is therefore a key concept for visualizing the excitation process associated to a given transition. One can say for example for a given transition that it corresponds to the excitation of an electron from an occupied orbital to a given unoccupied orbital. Nevertheless one has to keep in mind that electrons are fermions ruled by Pauli exclusion principle and cannot be distinguished from the other electrons in the atom. Moreover, it sometimes happens that the configuration interaction expansion converges very slowly and that one cannot speak about simple one-determinantal wave function at all. This is the case when electron correlation is large.

Hydrogen-like atoms

The simplest atomic orbitals are those that occur in an atom with a single electron, such as the hydrogen atom. In this case the atomic orbitals are the eigenstates of the hydrogen Hamiltonian. They can be obtained analytically (see Hydrogen atom). An atom of any other element ionized down to a single electron is very similar to hydrogen, and the orbitals take the same form. For atoms with two or more electrons, the governing equations can only be solved with the use of methods of iterative approximation. Orbitals of multi-electron atoms are qualitatively similar to those of hydrogen, and in the simplest models, they are taken to have the same form. For more rigorous and precise analysis, the numerical approximations must be used. A given (hydrogen-like) atomic orbital is identified by unique values of three quantum numbers: n, l, and ml. The rules restricting the values of the quantum numbers, and their energies (see below), explain the electron configuration of the atoms and the periodic table. The stationary states (quantum states) of the hydrogen-like atoms are its atomic orbital. However, in general, an electron's behavior is not fully described by a single orbital. Electron states are best represented by time-depending "mixtures" (linear combinations) of multiple orbitals. See Linear combination of atomic orbitals molecular orbital method. The quantum number n first appeared in the Bohr model. It determines, among other things, the distance of the electron from the nucleus; all electrons with the same value of n lay at the same distance. Modern quantum mechanics confirms that these orbitals are closely related. For this reason, orbitals with the same value of n are said to comprise an "shell". Orbitals with the same value of n and also the same value of l are even more closely related, and are said to comprise a "subshell".

Qualitative characterization

Limitations on the quantum numbers

An atomic orbital is uniquely identified by the values of the three quantum numbers, and each set of the three quantum numbers corresponds to exactly one orbital, but the quantum numbers only occur in certain combinations of values. The rules governing the possible values of the quantum numbers are as follows: The principal quantum number n is always a positive integer. In fact, it can be any positive integer, but for reasons discussed below, large numbers are seldom encountered. Each atom has, in general, many orbitals associated with each value of n; these orbitals together are sometimes called a shell. The orbital angular momentum quantum number \ell is a non-negative integer. Within a shell where n is some integer n0, \ell ranges across all (integer) values satisfying the relation 0 \le \ell \le n_0-1. For instance, the n = 1 shell has only orbitals with \ell=0, and the n = 2 shell has only orbitals with \ell=0, and \ell=1. The set of orbitals associated with a particular value of \ell are sometimes collectively called a subshell. The magnetic quantum number m_\ell is also always an integer. Within a subshell where \ell is some integer \ell_0, m_\ell ranges thus: -\ell_0 \le m_\ell \le \ell_0. The above results may be summarized in the following table. Each cell represents a subshell, and lists the values of m_\ell available in that subshell. Empty cells represent subshells that do not exist. Subshells are usually identified by their n- and \ell-values. n is represented by its numerical value, but \ell is represented by a letter as follows: 0 is represented by 's', 1 by 'p', 2 by 'd', 3 by 'f', and 4 by 'g'. For instance, one may speak of the subshell with n=2 and \ell=0 as a '2s subshell'.

The shapes of orbitals

Any discussion of the shapes of electron orbitals is necessarily imprecise, because a given electron, regardless of which orbital it occupies, can at any moment be found at any distance from the nucleus and in any direction due to the Uncertainty Principle. However, the electron is much more likely to be found in certain regions of the atom than in others. Given this, a boundary surface can be drawn so that the electron has a high probability to be found anywhere within the surface, and all regions outside the surface have low values. The precise placement of the surface is arbitrary, but any reasonably compact determination must follow a pattern specified by the behavior of \psi^2, the square of the wavefunction. This boundary surface is what is meant when the "shape" of an orbital is mentioned. Generally speaking, the number n determines the size and energy of the orbital: as n increases, the size of the orbital increases. Also in general terms, \ell determines an orbital's shape, and m_\ell its orientation. However, since some orbitals are described by equations in complex numbers, the shape sometimes depends on m_\ell also. s-orbitals (\ell=0) are shaped like spheres. p-orbitals have the form of two ellipsoids with a point of tangency at the nucleus (sometimes referred to as a dumbbell). The three p-orbitals in each shell are oriented at right angles to each other, as determined by their respective values of m_\ell. Four of the five d-orbitals look similar, each with four pear-shaped balls, each ball tangent to two others, and the centers of all four lying in one plane, between a pair of axes. Three of these planes are the xy-, xz-, and yz-planes, and the fourth has the centres on the x and y axes. The fifth and final d-orbital consists of three regions of high probability density: a torus with two pear-shaped regions placed symmetrically on its z axis.

Orbital energy

In atoms with a single electron (essentially the hydrogen atom), the energy of an orbital (and, consequently, of any electrons in the orbital) is determined exclusively by n. The n=1 orbital has the lowest possible energy in the atom. Each successively higher value of n has a higher level of energy, but the difference decreases as n increases. For high n, the level of energy becomes so high that the electron can easily escape from the atom. In atoms with multiple electrons, the energy of an electron depends not only on the intrinsic properties of its orbital, but also on its interactions with the other electrons. These interactions depend on the detail of its spatial probability distribution, and so the energy levels of orbitals depend not only on n but also on \ell. Higher values of \ell are associated with higher values of energy; for instance, the 2p state is higher than the 2s state. When \ell = 3, the increase in energy of the orbital becomes so large as to push the energy of orbital above the energy of the s-orbital in the next higher shell; when \ell = 4 the energy is pushed into the shell two steps higher. The energy order of the first 24 subshells is given in the following table. Each cell represents a subshell with n and \ell given by its row and column indices, respectively. The number in the cell is the subshell's position in the sequence. Empty cells represent subshells that either do not exist or stand farther down in the sequence.

Electron placement and the periodic table

Several rules govern the placement of electrons in orbitals (electron configuration). The first dictates that no two electrons in an atom may have the same set of values of quantum numbers (this is the Pauli exclusion principle). These quantum numbers include the three that define orbitals , as well as (the hitherto unmentioned) s. Thus, two electrons may occupy a single orbital, so long as they have different values of s. However, only two electrons, because of their spin, can be associated with each orbital. Additionally, an electron always tries to occupy the lowest possible energy state. It is possible for it to occupy any orbital so long as it does not violate the Pauli exclusion principle, but if lower-energy orbitals are available, this condition is unstable. The electron will eventually lose energy (by releasing a photon) and drop into the lower orbital. Thus, electrons fill orbitals in the order specified by the energy sequence given above. This behavior is responsible for the structure of the periodic table. The table may be divided into several rows (called 'periods'), numbered starting with 1 at the top. The presently known elements occupy seven periods. If a certain period has number i, it consists of elements whose outermost electrons fall in the ith shell. The periodic table may also be divided into several numbered rectangular 'blocks'. The elements belonging to a given block have this common feature: their highest-energy electrons all belong to the same \ell-state (but the n associated with that \ell-state depends upon the period). For instance, the leftmost two columns constitute the 's-block'. The outermost electrons of Li and Be respectively belong to the 2s subshell, and those of Na and Mg to the 3s subshell. The number of electrons in a neutral atom increases with the atomic number. The electrons in the outermost shell, or valence electrons, tend to be responsible for an element's chemical behavior. Elements that contain the same number of valence electrons can be grouped together and display similar chemical properties.

See also


- List of Hund's rules

Related topics


- Electron configuration
- Atomic electron configuration table
- Molecular orbital
- Energy level

External links


- [http://www.shef.ac.uk/chemistry/orbitron/ The Orbitron], a visualization of all common and uncommon atomic orbitals, from 1s to 7g
- David Manthey's [http://www.orbitals.com/orb/index.html Orbital Viewer] renders orbitals with n ≤ 30
- [http://www.falstad.com/qmatom/ Java orbital viewer applet]
- [http://wps.prenhall.com/wps/media/objects/602/616516/Chapter_07.html Covalent Bonds and Molecular Structure]

References


- Tipler, Paul & Ralph Llewellyn (2003). Modern Physics (4th ed.). New York: W. H. Freeman and Company. ISBN 0-7167-4345-0 Category:Chemical bonding Category:Atomic physics

Group 2 element

The alkaline earth metals are the series of elements in Group 2 (IUPAC style) of the periodic table: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra) (not always considered due to its radioactivity and very short half-life). The alkaline earth metals are named after their oxides, the alkaline earths, whose old-fashioned names were beryllia, magnesia, lime, strontia and baryta. These were named alkaline earths because of their intermediate nature between the alkalis (oxides of the alkali metals) and the rare earths (oxides of rare earth metals). The classification of some apparently inert substances as 'earths' is millennia old. The earliest known system used by the ancient Greeks consisted of four elements, including earth. This system was later refined by philosophers and alchemists such as Aristotle (4th century BC), Paracelsus (first half of 16th century), John Becher (mid 17th century) and Georg Stahl (late 17th century), with later thinkers subdividing 'earth' into three or more types. The realization that 'earths' were not elements but compounds is attributed to the chemist Antoine Lavoisier. In his Traité Élémentaire de Chimie (Elements of Chemistry) of 1789 he called them Substances simples salifiables terreuses, or salt-forming earth elements. Later, 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. The alkaline earth metals are silvery colored, soft, low-density metals, which react readily with halogens to form ionic salts, and with water, though not as rapidly as the alkali metals, to form strongly alkaline (basic) hydroxides. Beryllium is an exception: It does not react with water or steam, and its halides are covalent. For example, where sodium and potassium react with water at room temperature, magnesium reacts only with steam and calcium with hot water. These elements all have two electrons in their outermost shell, so the energetically preferred state of achieving a filled electron shell is to lose two electrons to form doubly charged positive ions. Category:Periodic table Category:Alkaline earth metals ko:알칼리 토금속 ms:Alkali Bumi ja:第2族元素 th:โลหะแอลคาไลน์เอิร์ธ

Main group element

In chemistry and atomic physics, main group elements are groups whose lightest members are represented by hydrogen, beryllium, boron, carbon, nitrogen, oxygen, fluorine, and helium as arranged in the periodic table of the elements. The main group elements (with some of the lighter transition metals) are the most abundant elements on the earth, in the solar system, and in the universe. Category:Chemical element groups

Zinc

Zinc (from German Zink, perhaps ultimately from Old Persian) is a chemical element in the periodic table that has the symbol Zn and atomic number 30.

Notable characteristics

Zinc is a moderately reactive metal that will combine with oxygen and other non-metals, and will react with dilute acids to release hydrogen. The one common oxidation state of zinc is +2.

Applications

Zinc is the fourth most common metal in use, trailing only iron, aluminium, and copper in annual production.
- Zinc is used to galvanise metals such as steel to prevent corrosion.
- Zinc is used in alloys such as brass, nickelled silver, typewriter metal, various soldering formulas and German silver.
- Zinc is the primary metal used in making American pennies since 1982.
- Zinc is used in die casting noteably in the automobile industry.
- Zinc is used as part of the containers of batteries.
- Zinc oxide is used as a white pigment in watercolours or paints, and as an activator in the rubber industry. As an over-the-counter ointment, it is applied as a thin coating on the exposed skin of the face or nose to prevent dehydration of the area of skin. It can protect against sunburn in the summer and windburn in the winter. Applied thinly to a baby's diaper area (perineum) with each diaper change, it can protect against rash. As determined in the Age-Related Eye Disease Study, it's part of an effective treatment for age-related macular degeneration in some cases.
- Zinc chloride is used as a deodorant and can be used as a wood preservative.
- Zinc sulfide is used in luminescent pigments such as on the hands of clocks and other items that glow in the dark.
- Zinc methyl (Zn(CH3)2) is used in a number of organic syntheses.
- Zinc stearate is a lubricative plastic additive.
- Lotions made of calamine, a mix of Zn-(hydroxy-)carbonates and silicates, are used to treat skin rash.
- Zinc metal is included in most single tablet over-the-counter daily vitamin and mineral suplements. It is believed to possess anti-oxidant properties, which protect against premature aging of the skin and muscles of the body. In larger amounts, taken as zinc alone in other proprietaries, it is believed by some to speed up the healing process after an injury. Preparations include zinc acetate and zinc gluconate.
- Zinc gluconate glycine is used as a lozenge in an attempt to remedy the common cold.

Popular misconceptions

The characteristic metal counters of traditional French bars are often referred to as zinc bars or simply zinc, but in fact zinc has never been used for this purpose and the counters are actually made of an alloy of lead and tin. In Argentina some people wrongly believe that zinc is a poison, and some of them are avoiding food which is known to include zinc. In 1997 a municipality north of the centre of Buenos Aires posted advertisements in popular magazines explaining the usefulness of zinc in the human body.

History

human body Zinc alloys have been used for centuries, as brass goods dating to 1000-1400 BC have been found in Palestine and zinc objects with 87% zinc have been found in prehistoric Transylvania. Because of the low boiling point and high chemical reactivity of this metal (isolated zinc would tend to go up the chimney rather than be captured), the true nature of this metal was not understood in ancient times. The manufacture of brass was known to the Romans by about 30 BC, using a technique where calamine and copper were heated together in a crucible. The zinc oxides in calamine were reduced, and the free zinc metal was trapped by the copper, forming an alloy. The resulting calamine brass was either cast or hammered into shape. Smelting and extraction of impure forms of zinc was being accomplished as early as 1000 AD in India and China. By the end of the 14th century, the Hindus were aware of the existence of zinc as a metal separate from the seven known to the ancients. In the West, impure zinc as a remnant in melting ovens was known since Antiquity, but usually thrown away as worthless. Strabo mentions it as pseudo-arguros "mock silver". The Berne Zinc tablet is a votive plaque dating to Roman Gaul, probably made from such zinc remnants. The discovery of pure metallic zinc is most often credited to the German Andreas Marggraf, in the year 1746, though the whole story is considerably more involved. Descriptions of brass manufacture are found in Western Europe in the writings of Albertus Magnus, c. 1248, and by the 16th century, the understanding and awareness of the new metal broadened considerably. Georg Agricola observed, in 1546, that a white metal could be condensed and scraped off the walls of a furnace when zinc ores were smelted. He added in his notes that a similar metal called "zincum" was being produced in Silesia. Paracelsus (died 1541) was the first in the West to say that "zincum" was a new metal and that it had a separate set of chemical properties from other known metals. The upshot is that zinc was known by the time Marggraf made his discoveries and in fact zinc had been isolated two years earlier by another chemist, Anton von Swab. However, Marggraf's reports were exhaustive and methodical and the quality of his research cemented his reputation as the discoverer of zinc. Before the discovery of the zinc sulfide flotation technique, calamine was the mineral source of zinc metal. calamine

Biological role

Zinc is an essential element, necessary for sustaining all life. It is estimated that 3000 of the hundreds of thousands of proteins in the human body contain zinc.

Food Sources

The best and most abundant natural food source of zinc is oysters, although these bottom scavengers also accumulate toxic metals. Zinc is found in most animal proteins such as beef, pork and poultry. Other food sources of zinc include beans, nuts, whole grains, pumpkin seeds and sunflower seeds. Phytates, which are found in whole grain breads, cereals, legumes and other products, have been known to decrease zinc absorption. This, coupled with the fact that the human body absorbs zinc more easily from animal protein than from plant protein means that vegetarians are required to eat many more food sources containing zinc than non-vegetarians.

Zinc Deficiency

Zinc deficiency results from inadequate intake of zinc, or inadequate absorption of zinc into the body. Signs of zinc deficiency includes hair loss, skin lesions, diarrhea, wasting of body tissues, and, eventually, death. Eyesight, taste, smell and memory are also connected with zinc and a deficiency in zinc can cause malfunctions of these organs and functions. Obtaining a sufficient zinc intake during pregnancy and in young children is a very real problem, especially among those who cannot afford a good supply of meat and a varied diet. Brain development is stunted by zinc insufficiency in utero and in youth. There is zinc in semen. As much as 0.25 milligram of zinc will be found in 1 mL of seminal fluid.

Zinc Toxicity

Even though zinc is almost an essential requirement for a healthy body, too much zinc can be harmful. Excessive absorption of zinc can also suppress copper and iron absorption.

Psoriasis

Ionic zinc is a potent antimicrobial, used since 2500 BC in topical creams. Calomine lotion, diaper creams, and dandruff treatments are just some of the common antimicrobial applications. At low concentrations, zinc ions promote wound healing. Zinc ions also directly stimulate zinc receptors on skin cells, promoting wound healing.

Immune System

Zinc salts are effective against pathogens in direct application. Gastrointestinal infections are also strongly attenuated by ingestion of zinc, and this effect could be due to direct antimicrobial action of the zinc ions in the GI tract, or to absorption of the zinc and re-release from immune cells (all granulocytes secrete zinc) or both. The direct effect of zinc (as in lozenges) on bacteria and viruses is also well-established, and has been used since at least 2000 BC, from when zinc salts in palliative salves are documented. However, exactly how to deliver zinc salts against pathogens without injuring one's own tissues is still being investigated.

Abundance

Zinc is the 23rd most abundant element in the Earth's crust. The most heavily mined ores (sphalerite) tend to contain roughly 10% iron as well as 40-50% zinc. Minerals from which zinc is extracted include sphalerite (zinc sulfide), smithsonite (zinc carbonate), hemimorphite (zinc silicate), and franklinite (a zinc spinel).

Zinc production

There are zinc mines throughout the world, with the largest producers being Australia, Canada, China, Peru and the U.S.A. Mines in Europe include Vieille Montagne in Belgium, Tara in Ireland, and Zinkgruvan in Sweden. Zinc metal is produced using extractive metallurgy. Zinc sulfide (sphalerite) minerals are concentrated using the froth flotation method and then usually roasted using pyrometallurgy to oxidise the zinc sulfide to zinc oxide. The zinc oxide is leached in sulfuric acid and the resulting solution is purified using zinc dust. The metal is then extracted by electrowinning as cathodic deposits. Zinc cathodes can be directly cast or alloyed with aluminium. Another process to produce zinc is flash smelting, a pyrometallurgical process. Then zinc oxide is obtained, usually producing zinc of lesser quality than the hydrometallurgical process. Zinc oxide treatment has much fewer applications, but high grade deposits have been successful in producing zinc from zinc oxides and zinc carbonates using hydrometallurgy.

Compounds

Zinc oxide is perhaps the best known and most widely used zinc compound, as it makes a good base for white pigments in paint. It also finds industrial use in the rubber industry, and is sold as opaque sunscreen. A variety of other zinc compounds find use industrially, such as zinc chloride (in deodorants), zinc sulfide (in luminescent paints), and zinc methyl in the organic laboratory. Roughly one quarter of all zinc output is consumed in the form of zinc compounds.

Isotopes

Naturally occurring zinc is composed of the 5 stable isotopes Zn-64, Zn-66, Zn-67, Zn-68, and Zn-70 with 64 being the most abundant (48.6% natural abundance). 21 radioisotopes have been characterised with the most being Zn-65 with a half-life of 244.26 days, and Zn-72 with a half-life of 46.5 hours. All of the remaining radioactive isotopes have half-lives that are less than 14 hours and the majority of these have half lives that are less than 1 second. This element also has 4 meta states.

Precautions

Metallic zinc is not considered to be toxic, but free zinc ions in solution (like copper or iron ions) are highly toxic. There is also a condition called zinc shakes or zinc chills that can be induced by the inhalation of freshly formed zinc oxide. Excessive intake of zinc can promote deficiency in other dietary minerals.

References


- [http://periodic.lanl.gov/elements/30.html Los Alamos National Laboratory - Zinc]

External links


- [http://www.vanderkrogt.net/elements/elem/zn.html History & Etymology of Zinc]
- [http://www.best-home-remedies.com/minerals/zinc.htm Zinc Information - Benefits, Deficiency Symptoms And Food Sources]
- [http://www.webelements.com/webelements/elements/text/Zn/index.html WebElements.com – Zinc]
- [http://chinese-school.netfirms.com/Zinc-information.html Zinc – History, sources, production, uses, health, and Zinc deficiency]
- [http://www.iza.com/zwo_org/Publications/Discovering/0202.htm Discovering the 8th metal]
- [http://minerals.er.usgs.gov/minerals/pubs/commodity/zinc/ Statistics and Information from the U.S. Geological Survey] Category:Chemical elements Category:Transition metals Category:Pyrotechnic chemicals ja:亜鉛 simple:Zinc th:สังกะสี

Chromium

Chromium is a chemical element in the periodic table that has the symbol Cr and atomic number 24.

Notable characteristics

Chromium is a steel-gray, lustrous, hard metal that takes a high polish, melts with difficulty, and tarnishes. The most common oxidation states of chromium are +2, +3, and +6, with +3 being the most stable. +4 and +5 are rare. Chromium compounds of oxidation state 6 are powerful oxidants. Chromium(0) is unstable in oxygen, immediately producing a thin oxide layer that is impermeable to oxygen and protects the metal below.

Applications

Uses of chromium:
- In metallurgy, to impart corrosion resistance and a shiny finish:
  - as an alloy constituent, e.g. in stainless steel used in cutlery, etc.,
  - in chrome plating,
  - in anodized aluminium (literally turning the surface of aluminium into ruby).
- As dyes and paints.
  - Chromium(III) Oxide is a metal polish known as Green rouge.
  - Chromium salts color glass an emerald green.
  - Chromium is what makes a ruby red, and therefore is used in producing synthetic rubies.
- As a catalyst.
- Chromite is used to make molds for the firing of bricks.
- Chromium salts are used in the tanning of leather.
- Potassium dichromate is a chemical reagent, used in cleaning laboratory glassware and as a titrating agent. It is also used as a mordant (i.e. a fixing agent) for dyes in fabric.
- Chromium(IV) oxide (CrO2) is used to manufacture magnetic tape, where its higher coercivity than iron oxide tapes gives better performance.
- In well drilling muds as an anti-corrosive.

History

In 1761, Johann Gottlob Lehmann found an orange-red mineral in the Ural Mountains which he named Siberian red lead. Though misidentified as a lead compound with selenium and iron components, the material was in fact lead chromate (PbCrO4), now known as the mineral crocoite. In 1770, Peter Simon Pallas visited the same site as Lehmann and found a red "lead" mineral that had very useful properties as a pigment in paints. The use of Siberian red lead as a paint pigment developed rapidly. A bright yellow made from crocoite became a very fashionable color. In 1797, Nicolas-Louis Vauquelin received samples of crocoite ore. He was able to produce chromium oxide (CrO3) by mixing crocoite with hydrochloric acid. In 1798, Vauquelin discovered that he could isolate metallic chromium by heating the oxide in a charcoal oven. He was also able to detect traces of chromium in precious gems, such as ruby, or emerald. During the 1800s chromium was primarily used as a component of paints but now the primary use (85%) is for metal alloys, with the remainder used in the chemical industry and refractory and foundry industries. Chromium was named based on the Greek word "chroma" meaning color, because of the many colorful compounds made from it.

Biological role

See Chromium deficiency.

Occurrence

Chromium is mined as chromite (FeCr2O4) ore. Chromium is obtained commercially by heating the ore in the presence of aluminium or silicon. Roughly half the chromite ore in the world is produced in South Africa. Kazakhstan, India and Turkey are also substantial producers. Untapped chromite deposits are plentiful, but geographically concentrated in Kazakhstan and southern Africa. Approximately 15 million tons of marketable chromite ore were produced in 2000, and converted into approximately 4 million tons of ferro-chrome with an approximate market value of 2.5 billion US dollars. Though native chromium deposits are rare, some native chromium metal has been discovered. The Udachnaya Mine in Russia produces samples of the native metal. This mine is a kimberlite pipe rich in diamonds, and the reducing environment so provided helped produce both elemental chromium and diamond.

Compounds

Potassium dichromate is a powerful oxidizing agent and is the preferred compound for cleaning laboratory glassware of any possible organics. Chrome green is the green oxide of chromium, Cr2O3, used in enamel painting, and glass staining. Chrome yellow is a brilliant yellow pigment, PbCrO4, used by painters. Chromic acid has the hypothetical structure H2CrO4