The alkali metal is a group (column) in the periodic table comprising the chemical elements of lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). This group lies in the block-s of the periodic table of elements because all alkali metals have their outer electrons in orbital s: these joint electron configurations produce very similar characteristic properties. Indeed, alkali metals provide the best example of group trends in properties in the periodic table, with elements that exhibit good homologous behavior.
Alkali metals are lustrous, soft, highly reactive metals at standard temperature and pressure and easily lose their outer electrons to form cations with charge 1. They can all be cut easily with a knife because of their softness, showing shiny surfaces that tarnish quickly in the air because oxidation by atmospheric moisture and oxygen (and in the case of lithium, nitrogen). Because of their high reactivity, they must be stored under oil to prevent reactions with air, and are found naturally only in salt and never as free elements. Cesium, the fifth alkali metal, is the most reactive of all metals. In the modern IUPAC nomenclature, alkali metals consist of elements of group 1 , excluding hydrogen (H), which is nominally an element of group 1 but is not normally regarded as an alkali metal because it seldom exhibits a behavior proportional to alkali metals. All alkali metals react with water, with heavier alkaline metals reacting louder than lighter ones.
All the alkali metals found occur in nature as their compounds: in the order of abundance, sodium is the most abundant, followed by potassium, lithium, rubidium, cesium, and finally franium, which is extremely rare because of its very high radioactivity; franium occurs only in the smallest trace in nature as a further step in some obscure side branches of the natural decay chain. Experiments have been made to try the ununennium synthesis (UUE), which is likely to be the next group member, but they all have failed. However, ununennium may not be alkali metal because of its relativistic effect, which is predicted to have a major influence on the chemical properties of the superheavy element; even if it turns out to be an alkali metal, is thought to have some differences in the physical and chemical properties of the lighter homologues.
Most alkali metals have many different applications. One of the most famous applications of pure elements is the use of rubidium and cesium in atomic clocks, where cesium atomic clock is the most accurate and precise time representation. A common application of sodium compounds is the sodium-vapor lamp, which emits very efficient light. Table salt, or sodium chloride, has been used since antiquity. Sodium and potassium are also important elements, have a major biological role as electrolytes, and although other alkali metals are unimportant, they also have various effects on the body, both beneficial and harmful.
Video Alkali metal
History
Sodium compounds have been known since ancient times; salt (sodium chloride) has become an important commodity in human activities, as witnessed by the English word salary , referring to the salary , money paid to Roman soldiers for salt purchase. While potassium has been used since ancient times, it is not understood for most of its history to be fundamentally different substances from sodium mineral salts. Georg Ernst Stahl obtained experimental evidence which led him to suggest a fundamental difference in sodium and potassium salt in 1702, and Henri Louis Duhamel du Monceau was able to prove this difference in 1736. The exact chemical composition of the potassium and sodium compounds, and its status. as a chemical element of potassium and sodium, is not known then, and thus Antoine Lavoisier excludes alkali in the chemical element list in 1789.
Pure potassium was first isolated in 1807 in England by Sir Humphry Davy, derived from potash caustic (KOH, potassium hydroxide) by the use of liquid salt electrolysis with a newly discovered voltaic pile. Previous experiments on aqueous salt electrolysis did not work because of the extreme reactivity of potassium. Potassium is the first metal isolated by electrolysis. Later that same year, Davy reported the extraction of sodium from a caustic soda of similar substance (NaOH, lye) by the same technique, showing the elements, and thus the salt, being different. Later that year, the first piece of pure liquid sodium metal was also prepared by Humphry Davy through the electrolysis of liquid caustic soda (now called sodium hydroxide).
Petalite (Li Al Si 4 10 ) was discovered in 1800 by the Brazilian chemist JosÃÆ'Â © BonifÃÆ'¡cio de Andrada at a mine on the island of Uta, Sweden. However, it was not until 1817 that Johan August Arfwedson, then working in the chemistry laboratory of JÃÆ'¶ns Jacob Berzelius, detected the presence of a new element when analyzing petalit ores. This new element was recorded by him to form compounds similar to sodium and potassium, although the carbonates and hydroxides are less soluble in water and more alkaline than other alkali metals. Berzelius gave the unknown name the name "lithion / lithina ", from the Greek word ??? o? (transliterated as litos , meaning "stone"), to reflect its discovery in solid minerals, compared to potassium, which has been found in plant ash, and sodium, known in part due to its abundance high in animal blood. He named the metal in lithium "" material. Lithium, sodium, and potassium are part of the discovery of periodicity, since they are among a series of triad elements in the same group recorded by Johann Wolfgang DÃÆ'¶bereiner in 1850 as having the same properties.
Rubidium and cesium were the first elements found using spectroscopes, discovered in 1859 by Robert Bunsen and Gustav Kirchhoff. The following year, they found cesium in mineral water from Bad DÃÆ'¼rkheim, Germany. The discovery of their rubidium occurred the following year in Heidelberg, Germany, found it in lepidolite minerals. The names of rubidium and cesium come from the most prominent line in their emission spectrum: the bright red line for rubidium (from the Latin word rubidus, meaning dark red or bright red), and sky-blue lines for cesium (derived from Latin caesius , meaning blue sky).
Around 1865 John Newlands produced a series of papers in which he listed elements in order of increasing atomic weight and similar physical and chemical properties that reappeared at interval eight; he likens that periodicity to the octave of music, in which a separate octave record has the same musical function. Its version places all then known alkali metals (lithium to cesium), as well as copper, silver, and thallium (which exhibit oxidation properties 1 of alkali metals), together into groups. The table places hydrogen with halogens.
After 1869, Dmitri Mendeleev proposed his periodic table to place lithium at the top of the group with sodium, potassium, rubidium, cesium, and thallium. Two years later, Mendeleev revised his desk, placing hydrogen in group 1 on lithium, and also transferring the thallium to the boron group. In this 1871 version, copper, silver and gold are placed twice, once as part of the IB group, and once as part of the "Group VIII" which includes the current groups of 8 to 11. After the introduction of an 18-column table, group IB elements are moved to their current position in the d-block, while the alkali metals are left in the IA group . Then the group name was changed to group 1 in 1988. The trivial name "alkali metal" comes from the fact that the hydroxides of the group 1 elements are all strong alkali when dissolved in water.
There are at least four incorrect and incomplete discoveries before Marguerite Perey of the Curie Institute in Paris, France discovered franium in 1939 by purifying the actinium-227 sample, which has been reported to have a decay energy of 220 keV. However, Perey noticed the rotting particles with energy levels below 80 keV. Perey thought this decay activity might be caused by a previously unidentified decomposition product, separated during purification, but reappeared from pure-227 aktinium. Various tests eliminate the possibility of unknown elements are thorium, radium, tin, bismuth, or thallium. The new product shows the chemical properties of alkali metals (such as coprecipitating with cesium salts), which causes Perey to believe that it is element 87, caused by the alpha-aktinium-227 decay. Perey then attempted to determine the proportion of beta decay into alpha decay in acyl-227. His first test placed alpha branching at 0.6%, a figure later revised to 1%.
- 227
223
87 Fr
? - -> 22 min 223
88 Ra
? -> 11,4 d
The next element under francium (eka-francium) in the periodic table will be ununennium (Uue), element 119. The ununennium synthesis was first attempted in 1985 by bombarding the einsteinium-254 target with calcium-48 ions at the superHILAC accelerator in Berkeley, California. No atoms are identified, leading to a result that limits 300 nb.
- 254
99 48
20 Ca
-> 302
119 Uue
* -> no atoms
It is highly unlikely that this reaction will be able to create ununennium atoms in the near future, given the extremely difficult task of making sufficient quantities of einsteinium-254, favored for the production of ultra-heavy elements because of their large, relatively long part-time mass 270 days, and significant availability of multiple micrograms, to create large enough targets to increase the sensitivity of the experiment to the required level; einsteinium has not been found in nature and is only produced in laboratories, and in smaller quantities than is required for effective synthesis of superheavy elements. However, given that the ununennium is only the first 8 period element in the extended periodic table, it may be found in the near future through another reaction, and indeed attempts to synthesise it are currently under way in Japan. Currently, no 8 period elements have been found, and it is also possible, because of the instability of the droplets, that only the element of the lower 8 period, up to about element 128, is physically possible. No synthesis attempts were made for the heavier alkali metals: because of their very high atomic numbers, they would require more powerful new methods and technologies to be created.
Maps Alkali metal
Genesis
In the Solar System
The Oddo-Harkins Rule states that elements with atomic numbers are even more common than those with odd atomic numbers, with the exception of hydrogen. This rule holds that the elements with odd atomic numbers have one unpaired proton and are more likely to catch the other, thus increasing their atomic number. In elements with even atomic numbers, protons are paired, with each pair member balancing the spin from the other, increasing stability. All alkali metals have a strange atomic number and they are unusual like elements even with atomic numbers adjacent to them (noble gases and alkaline earth metals) in the Solar System. Heavier alkali metals are also less abundant than lighter ones such as alkali metals from rubidium upwards can only be synthesized in supernovae and not in star nucleosynthesis. Lithium is also much less than sodium and potassium because it is not well synthesized on the Big Bang nucleosynthesis and in the stars: Big Bang can only produce trace amounts of lithium, beryllium and boron in the absence of a stable nucleus with 5 or 8 nucleons, and nucleosynthesis stars can only pass through these barriers with a triple-alpha process, combining three helium nuclei to form carbon, and jumping over the three elements.
On Earth
The Earth is formed from the same cloud of matter that forms the Sun, but the planets acquire different compositions during the formation and evolution of the solar system. In turn, Earth's natural history causes parts of the planet to have different concentrations of elements. The Earth's mass is about 5.98 ÃÆ' - 10 24 Ã, kg. Mostly composed of iron (32.1%), oxygen (30.1%), silicon (15.1%), magnesium (13.9%), sulfur (2.9%), nickel (1.8% ), calcium (1.5%), and aluminum (1.4%); with the remaining 1.2% consisting of trace amounts of other elements. Due to planetary differentiation, the core region is believed to consist primarily of iron (88.8%), with a smaller amount of nickel (5.8%), sulfur (4.5%), and less than 1% trace element.
Alkali metals, due to their high reactivity, do not occur naturally in their pure form in nature. They are lithophiles and therefore remain close to the Earth's surface as they join the oxygen and so closely associate with silica, forming minerals with relatively low densities that do not sink into the Earth's core. Potassium, rubidium and cesium are also incompatible elements due to the large ionic radius.
Sodium and potassium are abundant in the earth, both of which are among the ten most common elements in the Earth's crust; sodium makes up about 2.6% of the Earth's crust measured by weight, making it the sixth most abundant element and the most abundant alkaline metal. Potassium forms about 1.5% of the Earth's crust and is the seventh most abundant element. Sodium is found in many different minerals, the most common being the common salt (sodium chloride), which occurs in large quantities dissolved in seawater. Other solid deposits include halite, amphibole, cryolite, nitratine, and zeolite. Many of these solid deposits occur as a result of the vaporized ancient oceans, still occurring in places like the Great Salt Lake in Utah and the Dead Sea. Although the abundance is almost the same in the earth's crust, sodium is much more common than potassium in the oceans, both because the larger potassium size makes the salt less soluble, and because potassium is bound by silicates in the soil and what is absorbed potassium is absorbed much more easily. by plant life rather than sodium.
Despite its chemical similarities, lithium usually does not coincide with sodium or potassium because of its smaller size. Because of its relatively low reactivity, it can be found in seawater in large quantities; It is estimated that seawater is about 0.14 to 0.25 parts per million (ppm) or 25 micromolane. Diagonal relationship with magnesium often allows it to replace magnesium in ferromagnesium minerals, where the crust concentration is about 18 ppm, proportional to gallium and niobium. Commercially, the most important lithium mineral is spodumene, which occurs in large deposits around the world.
Rubidium is approximately as abundant as zinc and more than copper. It occurs naturally in leucite, pollucite, carnallite, zinnwaldite, and lepidolite minerals, although none contain rubidium and no other alkali metals. Cesium is more abundant than some commonly known elements, such as antimony, cadmium, lead, and tungsten, but much less than rubidium.
Francium-223, the only naturally occurring francium isotope, is the product of alpha-227 alpha decay and can be found in small quantities in uranium minerals. In a given uranium sample, it is estimated that there is only one francium atom for every 10 18 uranium atoms. It has been calculated that there are at most 30 g of francium in the earth's crust at all times, due to a very short half-life of 22 minutes.
Properties
Physical and chemical
The physical and chemical properties of alkali metals can be easily explained by they having a valence electron configuration of ns 1 , which results in a weak metal bond. Therefore, all alkali metals are soft and have low density, melting and boiling points, as well as sublimation heat, vaporization, and dissociation. They all crystallize in the cubic structure of the body-centered crystal, and have a distinctive flame color because the outer electrons are very easily excited. The ns 1 configuration also produces alkali metals that have very large atomic and ionic radii, as well as very high thermal and electrical conductivity. Their chemistry is dominated by the loss of their single valence electrons in the outer s-orbitals to form oxidation state 1, because of the ease of ionizing these electrons and the very high ionisation energies. Most of the chemistry has been observed only for the first five members of the group. Chemical franium is not well established due to its extreme radioactivity; thus, the presentation of the properties here is limited. What little is known about francium shows that behavior is very close to cesium, as expected. The physical properties of franium are even more vague because the bulk element is never observed; then any data that can be found in the literature is of course a speculative extrapolation.
Alkali metals are more similar to each other than the elements in other groups to each other. Indeed, the similarity is so great that it is difficult enough to separate potassium, rubidium, and cesium, because their ionic radii are the same; lithium and sodium are more different. For example, when moving down the table, all known alkali metals exhibit increased atomic radius, decrease electronegativity, increase reactivity, and reduce melting and boiling points and heat up fusion and vaporization. In general, their density increases as they move around the table, with the exception that potassium is less dense than sodium. One of the very few properties of alkali metals that do not display a very subtle trend is their reduction potential: lithium values ​​are anomalies, becoming more negative than others. This is because Li ion has a very high hydration energy in the gas phase: although lithium ions interfere with the water structure significantly, causing a higher change in entropy, this high hydration energy is enough to make the reduction potential show it as the most electropositive alkali metal, despite the difficulty of ionizing in the gas phase.
Stable alkali metals are all silver metal except for cesium, which has a pale gold color: it is one of only three clearly colored metals (the other two are copper and gold). In addition, the alkaline earth metals that are heavy in calcium, strontium, and barium, as well as the divalent bile of europium and ytterbium, are pale yellow, although the color is much less prominent than for cesium. Their sparkle tarnishes quickly in the air due to oxidation. They all crystallize in the cubic structure of the body-centered crystal, and have a distinctive flame color because the outer electrons are very easily excited. Indeed, the color of this flame test is the most common way to identify them because all of their salts with regular ions dissolve.
All alkali metals are highly reactive and never found in elemental form in nature. Therefore, they are usually stored in mineral oil or kerosene (paraffin oil). They react aggressively with halogens to form alkali metal halides, which are white ionic crystallizing compounds which are all soluble in water except lithium fluoride (Li F). Alkali metals also react with water to form strongly alkaline hydroxides and thus must be handled with extreme caution. Heavier alkaline metals react harder than lighter ones; for example, when it falls into water, cesium produces an explosion that is larger than potassium if the same number of moles of any metal is used. Alkali metals have the first lowest ionization energies in periods of each periodic table due to low effective nuclear charges and the ability to achieve noble gas configurations by losing just one electron. Alkali metals not only react with water, but also with proton donors such as alcohol and phenol, ammonia gas, and alkalo, the latter showing a phenomenal rate of reactivity. Their great strength as a reducing agent makes them particularly useful in freeing other metals from their oxides or halides.
The second ionisation energy of all alkali metals is very high because it is in full shell which is also closer to the nucleus; thus, they almost always lose a single electron, forming a cation. Alkalides are exceptions: they are unstable compounds containing alkali metals in the highly unusual oxidation state -1 as before the invention of alkalides, alkali metals are not expected to form anions and are thought to appear in salts only as cations. Anion alkalides have filled s-subshells, which gives them sufficient stability to exist. All alkali metals are stable unless lithium is known to form alkalides, and alkalides have many theoretical interests due to unusual stoichiometry and low ionization potential. Alkaloids are chemically similar to electrons, ie salts with trapped electrons act as anions. The most striking example of an alkalide is "inverse sodium hydride", H Na - (the two ions are complexed), compared with ordinary sodium hydride, Na H - : unstable in isolation, because of its high energy resulting from the transfer of two electrons from hydrogen to sodium, although some of the derivatives are predicted to be metastable or stable.
In aqueous solutions, alkali metal ions form aqueous ions of the formula [M (H 2 O) n , where n is the solvation number. Their coordination numbers and shape are as expected of their ionic radii. In an aqueous solution water molecules directly attached to a metal ion are said to belong to the first coordinating sphere, also known as the first, or primary, shell solvation. The bond between the water molecule and the metal ion is the dativ covalent bond, with the oxygen atom donating the two electrons to the bond. Each coordinated water molecule can be bonded with hydrogen bonds to other water molecules. The latter is said to be in the second coordination room. However, for alkali metal cations, the second coordination sphere is not well defined because the 1 c charge on the cation is not high enough to polarize water molecules in the primary solvent skin sufficient for them to form strong hydrogen bonds with those within the second coordination sphere, a more stable entity. The solvation number for Li has been experimentally determined to be 4, forming a tetrahedral [Li (H 2 O) 4 ] : while solvation figures 3 to 6 have been found for aqua lithium ions, solvation numbers less than 4 may result from the formation of ion contact pairs, and higher solvation numbers can be interpreted in water form. molecules that close to the face of the tetrahedron, although molecular dynamic simulations may indicate the presence of octahedral hexaaqua ions (Li = H 2 O) 4 ] There are also possibly six water molecules in the main solvation sphere of the sodium ions, forming an octahedral [ions] (subcool) ion. Although it was previously suspected that heavier alkali metals also formed octahedral hexaaqua ions, it has since been found that potassium and rubidium may form [K (H 2 O) 8 > and [Rb (H 2 O) 8 ] ion, which has a square antiprismatic structure, and that cesium forms ion 12 -coordinate [Cs (H 2 O) 12 ] .
Lithium
The chemistry of lithium shows some difference from the rest of the clusters as polymic anions Li small cations and gives the compound a more covalent character. Lithium and magnesium have diagonal relationships because of the radius of the same atom, so they show some similarities. For example, lithium forms a stable nitride, a property common among all alkaline earth metals (magnesium groups) but unique among alkali metals. In addition, among each group, only lithium and magnesium form organometallic compounds with significant covalent characters (eg LiMe and MgMe 2 ).
Lithium fluoride is the only alkali metal halide that is poorly soluble in water, and lithium hydroxide is the only non-light alkali metal hydroxide. In contrast, lithium perchlorate and other lithium salts with large un-polarized anions are much more stable than analogue compounds of other alkali metals, perhaps because Li has high solvent energy. This effect also means that the simplest lithium salt is commonly found in hydrated form, since the anhydrous shape is highly hygroscopic: it allows salts such as lithium chloride and lithium bromide to be used in a reducer and AC.
Francium
Francium is also predicted to show some differences due to its high atomic weight, causing its electrons to travel at a considerable fraction of the speed of light and thus create a more prominent relativistic effect. Unlike the declining tendency of electronegativity and ionization energy of alkali metals, francium ionization electronegativity and energy is predicted to be higher than cesium because of the relativistic stabilization of 7s electrons; also, the radius of the atom is thought to be very low. Thus, contrary to expectations, cesium is the most reactive of alkali metals, not franium. All known physical properties of franium also deviate from the apparent tendency of lithium to cesium, such as the first ionisation energy, electron affinity, and anion polarization, although due to the lack of known data about francium, many sources provide an extrapolation value, ignoring that relativistic effect of making trends from lithium to cesium becomes unusable in franium. Some of the predicted properties of franium taking into account relativity are electron affinity (47.2 kJ/mol) and fractional separation entries Fr 2 (42.1 kJ/mol). The CsFr molecule is polarized as Cs Fr - , indicating that the 7s subshell of franium is much stronger influenced by the relativistic effect than the 6s subshell of cesium. In addition, francium superoxide (FrO 2 ) is expected to have significant covalent character, unlike other alkali metal superoxides, due to the bonding contribution of 6p francium electrons.
Nuclear
All alkali metals have an odd number of atoms; therefore, the isotope must be strange (odd number of protons and neutrons) or odd (odd proton number, but even number of neutrons). The odd-numbered nuclei even have mass numbers, whereas even-odd nuclei have a strange mass number. Odd primordial nuclides are rare because most odd-bizarre nuclei are highly unstable with respect to beta decay, due to even decomposition products, and are therefore more strongly bound, due to the effects of nuclear pairs.
Because of the large scarcity of odd nuclei, almost all of the primordial isotopes of odd alkali metals (exceptions are stable isotopes of lithium-6 light and long-lived potassium potassium-40). For the amount of odd mass given, there is only one beta-stable nuclide, since there is no difference in the binding energy between even and odd evenly proportional to even and odd, leaving the other nobles of the same free (isobar) mass to the beta decay toward the lowest mass nuclides. The odd-number instability effect of both types of nucleons is that odd-numbered elements, such as alkali metals, tend to have fewer stable isotopes than even-numbered elements. Of the 26 monoisotopic elements which have only one stable isotope, all but one have an odd number of atoms and all but one also have the same number of neutrons. Beryllium is the sole exception for both rules, because of its low atomic number.
All alkali metals except lithium and cesium have at least one natural radioisotope: sodium-22 and sodium-24 are cosmogenic, potassium-40 and rubidium-87 radioisotope traces have a very long half-life and thus occur naturally, and all francium isotope is radioactive. Cesium was also considered radioactive at the beginning of the 20th century, although it did not have a natural radioisotope. (Francium has not been found at the time.) The long-lived, potassium-40 long-potassium radioisotopes make up about 0.012% of the natural potassium, and thus natural potassium is a weak radioactive. This natural radioactivity became the basis for the incorrect claim of the invention for element 87 (subsequent alkali metal after cesium) in 1925. Natural rubidium was also slightly radioactive, with 27.83% being a long-lived radioisotope of rubidium-87.
Cesium-137, with a half-life of 30.17 years, is one of the two main medium-life fission products, along with strontium-90, which is responsible for most of the spent nuclear fuel radioactivity after several years of cooling, up to several hundred years after used. This is largely the remaining radioactivity of the Chernobyl accident. Cesium-137 has high-energy beta decay and eventually becomes stable barium-137. It is a powerful emitter of gamma radiation. Cesium-137 has a very low level of neutron capture and can not be disposed of properly in this way, but it should be left to rot. Cesium-137 has been used as a tracer in hydrological studies, analogous to the use of tritium. A small amount of cesium-134 and cesium-137 were released into the environment for almost all nuclear weapons tests and several nuclear accidents, notably the GoiÃÆ'nia nia accident and the Chernobyl disaster. In 2005, cesium-137 was the main source of radiation in the alien zones around the Chernobyl nuclear power plant. Its chemical properties as one of the alkali metals make it one of the most problematic of the fission products that are short to medium-sized because they move easily and spread in nature due to the high water solubility of their salts, and are taken up by the body, the fault for sodium and potassium essential congeners his.
Periodic trends
Alkali metals are more similar to each other than the elements in other groups to each other. For example, when moving down the table, all known alkali metals exhibit increased atomic radius, decrease electronegativity, increase reactivity, and reduce melting and boiling points and heat up fusion and vaporization. In general, their density increases as they move around the table, with the exception that potassium is less dense than sodium.
Atomic and ionic radius
The atomic radius of the alkali metal rises below the group. Because of the shielding effect, when an atom has more than one electron shell, each electron feels electrocution from other electrons as well as the electrical attraction of the nucleus. In alkali metals, the outer electrons only feel a net charge of 1, since some nuclear charges (equal to the atomic number) are canceled by the inner electrons; the number of inner electrons of the alkali metal is always less than the nuclear charge. Therefore, the only factor affecting the atomic radius of an alkali metal is the number of electron shells. Since this number rises to the bottom of the group, the radius of the atom must also rise to the bottom of the group.
The ionic radii of alkali metals are much smaller than their atomic radius. This is because the outer electrons of the alkali metal reside in the skin of electrons different from the inner electrons, and thus when removed the resulting atom has a fewer and smaller electron shell. In addition, the effective nuclear charge has increased, and thus the electrons are attracted stronger toward the nucleus and the ionic radius decreases.
The first ionisation energy
The first ionisation energy of an element or molecule is the energy required to move the most loose electrons from one mole of gas atom from an element or molecule to form one mole of gaseous ions with an electric charge 1. The factors affecting the first ionisation energy are nuclear charge, the number of shields by the inner electrons and the distance from the loosest electrons held from the nucleus, which is always the outer electron in the main group elements. The first two factors change the effective nuclear charge that most feels the electron of an electron. Since the outer electrons of the alkali metal always feel the same effective nuclear charge (1), the only factor affecting the first ionisation energy is the distance from the outer electron to the nucleus. Because this distance rises down the group, the outer electrons feel less attractive than the nucleus and thus the first ionisation energy decreases. (This trend breaks out in franium due to relativistic stabilization and contraction of the 7s orbital, carrying the francial valence electrons closer to the nucleus than expected from non-relativistic calculations.This makes the francium outer electrons feel more attractive than the nuclei, increasing their first ionization energy slightly more than cesium.)
The second ionisation energy of the alkali metal is much higher than the first as the second most loose electron is part of the fully charged electron shell and is thus difficult to remove.
Reactivity
The alkali metal reactivity is increasing down the group. This is the result of a combination of two factors: the first ionisation energy and the atomisation energy of the alkali metal. Since the first ionisation energy of the alkali metal drops down the group, it is easier for the outer electrons to be removed from the atoms and participate in chemical reactions, thereby increasing the reactivity down the group. The atomization energy measures the strength of the metal bond of an element, which falls into the group when the atom rises within a radius and thus the metal bond must grow long, making the electrons delocalized further away from the gravitational attraction of the heavier nuclei. alkali metals. Adding atomisation and the first ionisation energies provide a quantity closely related to (but not equal to) the activation energy of the alkali metal reaction with other substances. This quantity decreases in the group, and so does the activation energy; thus, chemical reactions can occur faster and the reactivity decreases to the bottom of the group.
electronegativity
Electronegativity is a chemical trait that describes the tendency of an atom or functional group to attract electrons (or electron densities) to itself. If the bond between sodium and chlorine in sodium chloride is covalent, the pair of electrons together will be attracted to chlorine because the effective nuclear charge on the outer electron is 7 in chlorine but only 1 in sodium. The electron pair is attracted so close to the chlorine atom that it is practically transferred to a chlorine atom (ionic bond). However, if the sodium atoms are replaced by lithium atoms, the electron will not be as close to the chlorine atom as before because the lithium atom is smaller, making the electron pair more attracted to the effective nuclear charge closer to lithium. Therefore, larger alkali metal atoms (further down the group) will be less electronegative because the bonding pair is less attracted towards them. As mentioned earlier, franium is expected to be an exception.
Because of the higher lithium electronegativity, some of its compounds have more covalent characters. For example, lithium iodide (Li I) will dissolve in an organic solvent, a property of most covalent compounds. Lithium fluoride (LiF) is the only water-insoluble alkaline halide, and lithium hydroxide (LiOH) is the only non-deliquescent alkali metal hydroxide.
Melting and boiling points
The melting point of a substance is the point at which it changes the state from solid to liquid while the boiling point of a substance (in liquid state) is the point at which the vapor pressure of the fluid is equal to the environmental pressure around the liquid and all the liquids. change the state to gas. As the metal is heated to its melting point, the metal bonds keep the atoms weakened so that the atoms can move, and the metal bond eventually breaks completely at the boiling point of the metal. Therefore, melting and boiling points falling from the alkali metal indicate that the metal bond strength of the alkali metal decreases down the group. This is because the metal atoms are held together by the electromagnetic attraction of the positive ion to the delocalized electron. As the atoms increase in size down to the group (because their atomic radius increases), the ion nuclei travels farther than the delocalized electrons and hence the metal bond becomes weaker so that the metal can more easily melt and boil, thus lowering the melting point and boiling point. (Increased nuclear charge is not a relevant factor because of the shielding effect.)
Density
Alkali metals all have the same crystal structure (body-centered cubes) and thus the only relevant factor is the number of atoms that can enter into a particular volume and the mass of one atom, since density is defined as mass per unit. volume. The first factor depends on the volume of the atom and thus the radius of the atom, which increases down the group; thus, the volume of alkali metal atoms rises to the bottom of the group. The mass of alkali metal atoms also rises down the group. Thus, the tendency of an alkali metal density depends on the atomic weight and radius of the atom; if the numbers for these two factors are known, the ratio between the densities of the alkali metals can be calculated. The resulting trend is that the density of alkali metals increases below the table, with the exception of potassium. Since it has the lowest atomic weight and the largest atomic radius of all elements in their period, alkali metals are the most dense metals in the periodic table. Lithium, sodium, and potassium are the only three metals in the less dense periodic table than water: in fact, lithium is the most dense solid known at room temperature.
​​Compound
Alkali metals form a series of compounds complete with all the usual anions found, which also describe the group trend. These compounds can be described as involving alkali metals that lose electrons to acceptor species and form monopositive ions. This description is most accurate for alkaline halides and becomes less and less accurate as an increase in cationic and anionic charge, and as anions become larger and more polariable. For example, ionic bonds provide a way for metal bonding along the NaCl, Na 2 series, Na 3 P, Na 3 As, Na 3 Sb, Na 3 Bi, Na.
Hydroxide
All alkali metals react strongly or explosively with cold water, producing an aqueous solution of alkaline alkali metal hydroxides and releasing hydrogen gas. This reaction becomes stronger in the group: lithium reacts with froth, but sodium and potassium can ignite and rubidium and cesium sink in water and produce hydrogen gas so quickly that shock waves are formed in water that can destroy glass containers. When the alkali metal falls into the water, it produces an explosion, in which there are two separate stages. The metal reacts with water first, breaking the hydrogen bond in the water and producing hydrogen gas; this progresses faster for more reactive alkaline metals. Second, the heat generated by the first part of the reaction often ignites hydrogen gas, causing it to explode explosively into the surrounding air. This secondary hydrogen gas explosion generates a flame over a water bowl, lake or other water body, not the initial reaction of metal with water (which tends to occur mostly under water). Alkali metal hydroxides are the most basic hydroxides known.
Recent research shows that the explosive behavior of alkali metals in water is driven by Coulomb explosions not solely by the rapid generation of hydrogen itself. All alkali metals melt as part of reaction with water. The water molecules ionize the smooth metal surfaces of the molten metal, leaving a positively charged metal surface and negatively charged water ions. The attraction between charged metal and water ions will rapidly increase the surface area, causing an increase in exponential ionization. When the dull force within the liquid metal surface exceeds the strength of the surface tension, it explodes strongly.
The hydroxide itself is the most basic hydroxide known, reacting with acids to provide salts and with alcohols to produce oligomeric alkoxides. They readily react with carbon dioxide to form carbonates or bicarbonates, or with hydrogen sulfide to form sulphides or bisulfides, and can be used to separate thiols from petroleum. They react with amphoteric oxide: for example, oxides of aluminum, zinc, lead, and lead react with alkali metal hydroxides to produce aluminate, sting, stannate, and plumbate. Silicon dioxide is acidic, and thus alkali metal hydroxide can also attack silicate glass.
Intermetallic compound
Alkali metals form many intermetallic compounds with each other and elements from groups 2 through 13 in the periodic table of various stoichiometry, such as sodium amalgam with mercury, including Na 5 Hg 8 and Na 3 Hg. Some of them have ionic characteristics: taking alloys with gold, the most electronegative of metal, for example, NaAu and KAu ​​are metallic, but RbAu and CsAu are semiconductors. NaK is a sodium and potassium alloy which is very useful because it is liquid at room temperature, although precautions must be taken because of extreme reactivity to water and air. The eutectic mixture melts at -12.6 ° C. An alloy of 41% cesium, 47% sodium, and 12% potassium has the lowest known melting point of any metal or alloy, -78 ° C.
Compound with group element 13
The intermetallic compounds of alkali metals with heavier 13 group elements (aluminum, gallium, indium, and thallium), such as NaTl, are poor conductors or semiconductors, unlike normal alloys with the preceding element, implying that the involved alkali metal has a loss of electrons to anion Zintl involved. However, while elements in groups 14 and so tend to form discrete anionic groups, the group 13 elements tend to form polymeric ions with alkali metal cations located between the giant ionic lattices. For example, NaTl consists of polymeric anions (--Tl - -) n with covalent cubic structures with Na ions located between the anionic lattice. Larger alkaline metals can not enter in the same way into the anionic lattice and tend to force the heavier group of 13 elements to form anionic groups.
Boron is a special case, being the only nonmetal in group 13. Borides of alkali metals tend to be boron-rich, involving boron-boron bonds involving deltahedral structures, and unstable thermal because alkali metals have extremely high vapor pressures. at high temperatures. This makes synthesis instantly problematic because the alkali metal does not react with boron below 700 ° C, and thus this must be completed in a sealed container with an excess of alkali metal. Furthermore, remarkably in this group, reactivity with boron decreases in the group: lithium reacts perfectly at 700 ° C but sodium at 900 ° C and potassium is not up to 1200 ° C, and its reaction is immediate to lithium but requires long hours. for potassium. Rubidium and cesium borides have not even been characterized. Various phases are known, such as LiB 10 , NaB 15 , and KB 6 . Under high pressure, the boron-boron bond in the lithium boride changes from following Wade's rule to form a Zintl anion like the rest of the group 13.
Compound with group element 14
Lithium and sodium react with carbon to form acetylides, Li 2 C 2 and Na 2 C 2 , which can also obtained by reaction of the metal with acetylene. Potassium, rubidium, and cesium react with graphite; their atoms are intercalated between layers of hexagonal graphite, forming a graphite intercalation compound of formula MC 60 (dark gray, almost black), MC 48 (dark gray, almost black ), MC 36 (blue), MC 24 (blue steel), and MC 8 (bronze) (M = K, Rb, or Cs). This compound is more than 200 times more conductive than pure graphite, showing that the valence electrons of the alkali metal are transferred to the graphite layer (eg M
C -
8 ). After heating KC 8 , the removal of the potassium atoms results in successive conversion into KC 24 , KC 36 , KC 48 and finally KC 60 . KC 8 is a very powerful and pyrophoric reducing agent and explodes when in contact with water. While larger alkali metals (K, Rb, and Cs) initially form MC 8 , smaller ones initially form MC 6 , and indeed they require a reaction of metal with graphite at high temperatures of about 500 ° C to form. Apart from this, alkali metals are strong reducing agents so they can even reduce buckminsterfullerene to produce solid fullerides M n C 60 ; sodium, potassium, rubidium, and cesium can form fullerides where n = 2, 3, 4, or 6, and rubidium and cesium can also reach n = 1.
When the alkali metal reacts with heavier elements in the carbon group (silicon, germanium, lead, and lead), ionic substances with cage-like structures are formed, such as silicide M 4 Si 4 (M = K, Rb, or Cs), containing M and tetrahedral Si 4 -
4 ion. Chemical alkali metal germanides, involving germanide ions Ge 4 - and other cluster ions (Zintl) such as Ge 2 -
4 , 4 - 9 , 2 -
9 , and [(Ge 9 ) 2 ] 6 - , very similar to that of the corresponding silicide. Stannides of alkali metals are mostly ionic, sometimes with stannide ions (Sn 4 - ), and sometimes with more complex Zintl ions such as Sn 4 -
9 , which appears in tetrapotassium nonastannide (K 4 Sn 9 ). The monetomic plumbide (Pb 4 - ) ion is unknown, and indeed its formation is not expected to be energetically energetic; alkali metal plumbides have complex Zintl ions, such as
Nitride and pnictides
Lithium, the lightest of alkali metals, is the only alkali metal that reacts with nitrogen under standard conditions, and its nitrid is the only stable alkali metal nitride. Nitrogen is a non-reactive gas because breaking a strong triple bond in a nitrogen molecule (N 2 ) takes a lot of energy. The formation of an alkali metal nitride will consume the ionization energy of the alkali metal (forming M ion), the energy required to break the triplicate in N 2 and the formation of N 3 ions - , and all energy released from the formation of an alkali metal nitride is derived from the lattice energy of the alkali metal nitride. The lattice energy is maximized with small, high-charged ions; alkali metals do not form high-charged ions, only form ions with a charge of 1, so only lithium, the smallest alkali metal, can release enough lattice energy to make a reaction with exothermic nitrogen, forming lithium nitride. The reaction of other alkali metals with nitrogen will not release enough lattice energy and thus will become endothermic, so they do not form nitrides under standard conditions. Sodium nitride (Na 3 N) and potassium nitride (K 3 N), while present, are highly unstable, susceptible to decomposition to their constituent elements, and can not be produced by reacting elements with each other under standard conditions. Steric hindrance prohibits the presence of rubidium or cesium nitride. However, sodium and potassium form a colored azide salt involving linear N - > 3 anion; because of the large size of alkali metal cations, they are sufficiently thermally stable to be able to thaw before decaying.
All alkali metals react readily with phosphorus and arsenic to form phosphides and arsenides of the formula M 3 Pn (where M represents alkali metal and Pn represents pnictogen - phosphorus, arsenic, antimony or bismuth). This is due to the larger P 3 - and As 3 - sizes of the ion, so that less lattice energy needs to be released to allow salt to form. This is not the only phosphide and arsenide of alkali metal: for example, potassium has nine known phosphides, with the formula K 3 P, K 4 P 3
4 , KP, K 4 P 6 , K 3 sub , KP 10.3 , and KP 15 >. While most metals form arsenides, only alkali and alkali metals make up most of the ionic arsenic. Structure of Na 3 Because of the complex with short short Na-Na distance of 328-330 pm which is shorter than sodium metal, and this shows that even with electropositive metal this bond can not be directly ionic.. Other alkali metal arsenides are inconsistent with the formula M 3 As is well known, such as LiAs, which have metallic luster and electrical conductivity indicate some metal bonding. Antimony is unstable and reactive because the Sb 3 - ion is a strong reducing agent; their reaction with acids forms a toxic and unstable gas stibin (SbH 3 ). Indeed, they have some metallic properties, and the alkali metal antimony of the MSB stoichiometry involves an antimony atom bound in a spiral Zintl structure. Bismuthides are not even fully ionic; they are intermetallic compounds containing partially bonded metals and some ions.Oxides and chalcogenides
All alkali metals react vigorously with oxygen under standard conditions. They form various types of oxides, such as simple oxides (containing O 2 - ) ions, peroxides (containing O 2 -
2 ion, where there is a single bond between two oxygen atoms), superoxide (contains O - 2 ion); and many others. Lithium burns in the air to form lithium oxide, but sodium reacts with oxygen to form a mixture of sodium oxide and sodium peroxide. Potassium forms a mixture of potassium peroxide and potassium superoxide, while rubidium and cesium form superoxide exclusively. Their reactivity increases with decreasing groups: while lithium, sodium and potassium burn only in the air, rubidium and cesium are pyrophoric (spontaneously burning in the air).
Smaller alkali metals tend to polarize larger anions (peroxides and superoxides) because of their small size. This attracts electrons in a more complex anion toward one of its constituent oxygen atoms, forming oxide ions and oxygen atoms. This causes lithium to form oxides exclusively in the reaction with oxygen at room temperature. These effects become very weak for greater sodium and potassium, allowing them to form more stable peroxides. Rubidium and cesium, at the bottom of the group, are so large that even the most stable superoxide can be formed. Because superoxide releases most of the energy when it is formed, superoxide is preferably formed for larger alkaline metals where more complex anions are not polarized. (Oxides and peroxides for these alkali metals do exist, but do not form in direct reaction of metals with oxygen under standard conditions.) In addition, small size of Li and O 2 - ions contribute to form stable ionic lattice structures. However, under controlled conditions, all alkali metals, with the exception of franium, are known to form oxides, peroxides, and superoxides. Alkaline and superoxide metal peroxides are strong oxidizing agents. Sodium peroxide and superoxide potassium react with carbon dioxide to form alkali metal carbonate and oxygen gas, allowing them to be used in subsea air cleaners; The presence of water vapor, naturally present in the breath, makes the removal of carbon dioxide by potassium superoxide even more efficient. All stable alkali metals except lithium can form red ozonide (MO 3 ) through low temperature reaction of powdered anhydride hydroxide with ozone: ozonide can be extracted using liquid ammonia. They slowly decompose under standard conditions into superoxide and oxygen, and hydrolysis immediately into the hydroxide when in contact with water. Potassium, rubidium and cesium also form sesquioxides M 2 O 3 , which might be better considered peroxide disuperoxides, [(M )
4 (O 2 -
2 ) (O -
2 )
2 ] .
Rubidium and cesium can form various suboxides with metals in formal oxidation form below 1. Rubidium can form Rb 6 O and Rb 9 O 2 (copper-colored) on the oxidation in the air, while cesium forms a very large variety of oxides. , such as CsO ozonide 3 and some brightly colored suboxides, such as Cs 7 O (bronze), Cs 4 O (red-violet)), CsO, Cs 3 O 3 (violet), Cs 3 O (dark green) sub> 2 , and Cs 7 O 2 . The latter can be heated under vacuum to produce Cs 2 O.
Alkali metals may also react analogously to heavier chalcogens (sulfur, selenium, tellurium, and polonium), and all alkali metal chalcogenides are known (with the exception of francium). Reactions with excess chalcogen may also produce lower chalcogenides, with chalcogenic ions containing the corresponding chalcogen atom chains. For example, sodium may react with sulfur to form sulphides (Na 2 S) and various polysulfides with the formula Na 2 S x ( x from 2 to 6), contains 2 -
x ion. Because of the alkalinity of the Se 2 - and Te 2 - ions, the alkali metal selena and the telluride are basic in solution; when it reacts directly with selenium and tellurium, poliselenides alkali metal and polytellurides formed together with selenides and tellurides with Se x and Te 2 - x ion. They can be obtained directly from the elements in liquid ammonia or when the air does not exist, and colorless, water-soluble compounds that oxidize the air to rapidly return to the selenium or tellurium. Polonide alkali metals are all ionic compounds containing Po 2 - ions; they are very stable chemically and can be produced by the direct reaction of the elements of about 300-400 ° C.
halides, hydrides, and pseudohalides
Alkali metals are one of the most electropositive elements in the periodic table and thus tend to ionically bond to the most electronegative elements in the periodic table,
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