Alkali Metals

10 Key excerpts on "Alkali Metals"

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  eBook - PDF

  Essentials of Inorganic Chemistry

  For Students of Pharmacy, Pharmaceutical Sciences and Medicinal Chemistry

  • Katja A. Strohfeldt(Author)
  • 2014(Publication Date)
  • Wiley

    (Publisher)

  2 Alkali Metals Members of group 1 of the periodic table (first vertical column) with exception of hydrogen are called Alkali Metals. Under the term Alkali Metals, the following elements are included: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). Generally, francium is not included in fur- ther discussions, as only artificial isotopes are known with 223 Fr having the longest half-life T 1/2 = 21.8 min (Figure 2.1) [1]. In terms of a clinical use, sodium and potassium are essential ions for the human body and any imbalance in them has to be corrected. Lithium is medically used to treat bipolar disorder (BD), and the application of lithium salts is further discussed within this chapter. 2.1 Alkali metal ions This group of elements belongs to the so-called s-block metals as they only have one electron in their outer shell, which is of s type. The chemistry of the metals is characterised by the loss of this s electron to form a monocationic ion M + , which results from the relatively low ionisation energy of this electron (Table 2.1). The term ionisation energy (IE) is defined as the energy that is required to remove the outer electron of an atom or molecule. The tendency to lose the outer electron is directly correlated to the ionisation energy – the lower the ionisation energy, the easier the removal of the electron. Within the group of Alkali Metals, the ionisation energy for the removal of the outer electron decreases as a result of the increasing distance of this electron from the nucleus. The loss of the outer s electron within the group of Alkali Metals results in the formation of the M + ion as mentioned. Consequently, most of the compounds of group 1 elements tend to be ionic in nature and form salts. In all pharmaceutical applications, only the salts of Alkali Metals are used, as most of the pure metals react violently with water.

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  Comparative Inorganic Chemistry

  • Bernard Moody(Author)
  • 2013(Publication Date)
  • Arnold

    (Publisher)

  17 Group I: the Alkali Metals Lithium, sodium, potassium, rubidium and caesium 3 The Alkali Metals are a dis-Li tinctive family of the chemi-2,1 cally most reactive metals, 11 showing a progressive in-Na crease in electropositive 2,8, 1 character with increasing 19 atomic number. K The overall picture shows 2,8,8,1 the steady gradation of 37 properties of very similar Rb elements and their com-2,8,18,8,1 pounds. Sodium and potas-55 sium compounds in general Cs use are described in detail. 2,8 18,18,8,1 The radioactive element, 87 francium, will not be men-Fr tioned further but rubidium 2,8,18 32, 18,8, 1 and caesium are included in a general way to emphasize the close family similarity. The elements of the second period of the Periodic Classification com-prise the first members, except for the noble (inert) gases, of the periodic groups and they show, in varying degrees, anomalies in behaviour although clearly belonging to their assigned family groups. Compounds of lithium are included to illustrate this point. The Alkali Metals are all soft, white and lustrous but rapidly tarnish in air. They are usually stored under oil or solvent naphtha. With increasing atomic number, the metals become softer, easier to fuse and more volatile. Rubidium and caesium are denser than water while the others are less dense. Lithium has the lowest density and highest specific heat capacity of any solid element at room tem-perature. With increase in atomic number, the general reactivity increases markedly, lithium being somewhat apart from the others in this respect. This is illustrated by the action on water which ranges in intensity from the quiet evolution of hydrogen with lithium to the violence of caesium. In the original Mendeléeff Periodic Table, lithium and sodium were the typical elements and the family branched into subgroups: la comprised potassium, rubidium and caesium, and lb, copper, silver and gold.

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  The Chemistry of the Metallic Elements

  The Commonwealth and International Library: Intermediate Chemistry Division

  • David J. Steele, J. E. Spice(Authors)
  • 2017(Publication Date)
  • Pergamon

    (Publisher)

  C H A P T E R 4 Group la: the Alkali Metals Li, Na, K, Rb, Cs, Fr THE elements in this group are more closely related than those in any other and the variation in physical and chemical properties is the most regular. The electronic configuration of the atoms are given in Table 4.1. TABLE 4.1. THE Alkali Metals: THEIR ELECTRONIC CONFIGURATION. Lithium Sodium Potassium Rubidium Caesium Francium Is 2 2 2 2 2 2 2s 1 2 2 2 2 2 2p 6 6 6 6 6 3s 1 2 2 2 2 3p 6 6 6 6 4s 1 2 2 2 3d 10 10 10 4p 6 6 6 5s 1 2 2 4d 10 10 5p 6s 6 1 6 2 4f 5d 6p 7s 5f 6d 14 10 6 1 The elements each have a single s-electron in the outer shell. The attraction of the nucleus for this outer electron is slight owing to the shielding effect of completed electron shells (see Chapter 2, p. 13). The low ionisation energy (see Table 4.2) shows the readiness with which the atoms lose the ^-electron and thereby form the stable unipositively charged ion, isoelectronic with the previous rare gas. There is no tendency to lose more than the outer ^-electron hence the metals are all univalent only. The stability of the ion, shown in ionisation energy and (in aqueous solution) electrode potential data, leads to the predominance of ionic bonding in the com-pounds of this group. The mobility of the outer electron in the ion lattice of the metallic state accounts for the good conductance of heat and electricity by the metals and also their malleability and softness (see Chapter 1). The variations in these data are regular; the few exceptions, such as the density of potassium, are due to changes in the stability and arrangement in the metal lattice. 31 32 The Chemistry of the Metallic Elements In general the variations in physical properties are in accord with those predicted from the electronic configurations and nuclear charges of the atoms. For example, the addition of electrons to higher energy levels causes an increase in atomic size with atomic number.

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  Chemistry

  With Inorganic Qualitative Analysis

  • Therald Moeller(Author)
  • 2012(Publication Date)
  • Academic Press

    (Publisher)

  ion is the principal ion in the fluids that surround cells, and K+

  is the principal ion in the fluids within the cells. A deficiency or overabundance of any of the cations listed above has serious consequences for health. Calcium is also a major component of bone, and magnesium is required for the activity of several enzymes. Zinc is an essential trace element in both plant and animal life, and some recent studies suggest that inadequate zinc in the diet may increase susceptibility to heart disease.

  Representative Groups I and II: The alkali and alkaline earth metals

  26.1 Properties of Group I and II metals

  The Alkali Metals—the members of Representative Group I (Table 26.1 )—have the highest reactivity of all metals. Atoms of these elements easily give up their single valence electrons to form monopositive ions such as Na+ and K+ , which have noble gas configurations. We refer to such metals as highly electropositive. The Alkali Metals form ionic compounds in virtually all of their reactions.

  TABLE 26.1 Properties of the Representative Group I and II metals

  The trends of physical properties expected with increasing atomic weight and atomic size are clearly displayed by Group I metals. The atomic and ionic radii (Table 26.1 ) increase with increasing atomic weight. Paralleling this trend, the melting points and boiling points decrease as the atoms become larger and heavier, and the forces holding the atoms in the metal lattice decrease. The Alkali Metals have lower melting and boiling points and also lower densities, than most other metals. In addition, they are soft; all except lithium easily can be cut with a knife. When freshly cut they have a characteristic metallic luster, but on exposure to air they soon tarnish by rapidly reacting with atmospheric gases.

  The alkaline earth metals—Representative Group II , also known as the beryllium family elements—also have a high degree of reactivity, second only to that of their Group I neighbors. Beryllium is covalent in most of its compounds because of the large amount of energy needed to form Be2+ . Magnesium sometimes forms at least partially covalent bonds, but the remaining alkaline earth elements react almost exclusively to form dipositive ions, for example, Ca2+ and Ba2+

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  The History and Use of Our Earth's Chemical Elements

  A Reference Guide

  • Robert E. Krebs(Author)
  • 2006(Publication Date)
  • Greenwood

    (Publisher)

  Its atom is the smallest of the alkali earth metals and thus is the least reactive because its valence electron is in the K shell, which is held closest to its nuclei. Characteristics While classified as an alkali metal, lithium also exhibits some properties of the alkali earth metals found in group 2 (IIA). Lithium is the lightest in weight and softest of all the metals and is the third lightest of all substances listed on the periodic table, with an average atomic weight of about 7. (The other two are hydrogen and helium.) Although it will float on water, it reacts with water, liberating explosive hydrogen gas and lithium hydroxide (2Li + 2H 2 O → 2LiOH + H 2 ∆). It will also ignite when exposed to oxygen in moist air (4Li + O 2 → 2Li 2 O). It is electropositive and thus an excellent reducing agent because it readily gives up electrons in chemical reactions. Lithium is the only metal that reacts with nitrogen at room temperature. When a small piece of the metal, which is usually stored in oil or kerosene, is cut, the new surface has a bright, shiny, silvery surface that soon turns gray from oxidation. 48 | The History and Use of Our Earth’s Chemical Elements Abundance and Source Lithium ranks 33rd among the most abundant elements found on Earth. It does not exist in pure metallic form in nature because it reacts with water and air. It is always combined with other elements in compound forms. These lithium mineral ores make up only about 0.0007%, or about 65 ppm, of the Earth’s crust. Lithium is contained in minute amounts in the mineral ores of spodumene, lepidolite, and amblygonite, which are found in the United States and several countries in Europe, Africa, and South America. High temperatures are required to extract lithium from its compounds and by electrolysis of lithium chloride. It is also concentrated by solar evaporation of salt brine in lakes.

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  eBook - ePub

  Science and Technology of Liquid Metal Coolants in Nuclear Engineering

  • Thiagarajan Gnanasekaran(Author)
  • 2024(Publication Date)
  • Woodhead Publishing

    (Publisher)

  Since the atomic size increases with increase in atomic number, there is a continuous decrease in the strength of the metallic bond among the Alkali Metals themselves. Therefore, the softness in their solid state increases from lithium to cesium. Except lithium, they can be cut with the help of a knife. Also, Alkali Metals are highly malleable and ductile. Melting and boiling points of Alkali Metals also decrease with increase in atomic number, showing again a steady decrease in their cohesive forces in solid as well as liquid states (refer Chapter 1). Alkali Metals are chemically very reactive and this reactivity depends on their ionization enthalpy. Lower the ionization enthalpy, higher would be the rate of these reactions. Lithium has the highest first ionization enthalpy (520 kJ mol − 1) while that of cesium is the lowest (376 kJ mol − 1). Hence, the reactivity of Alkali Metals increases down the group and is in the following order: Cs> Rb>K>Na>Li. While cesium and rubidium would ignite instantaneously on exposure to air, sodium at room temperature can be handled in air with low humidity for a short time with no great difficulty. Lithium can be handled in air with low humidity for somewhat prolonged periods of time. 2.2.3.1. Reactions with air and oxygen All Alkali Metals react with oxygen in dry air and produce their oxides. In case of solid lithium and sodium, a dense oxide layer is formed on their surfaces which retard further reaction of the underlying metal. With increase in humidity level in air, rate of reaction of all Alkali Metals increases. In moist air, formation of hydroxides and carbonates are favored

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  Encyclopedia of the Alkaline Earth Compounds

  • Richard C. Ropp(Author)
  • 2012(Publication Date)
  • Elsevier

    (Publisher)

  2+ . The alkaline earth metals are silver-colored, soft metals that react readily with halogens to form ionic salts. They also react with water, though not as rapidly as the Alkali Metals, to form strongly alkaline (basic) hydroxides. For example, whereas Na and K react with water at room temperature, Mg reacts only with steam and Ca with hot water:

  Be is an exception. It does not react with water or steam, and its halides are covalent.

  The alkaline earth metals are named after their oxides, the alkaline earths , whose old-fashioned names were Beryllia, Magnesia, Lime, Strontia and Baryta. “Earth” is the old term applied by early chemists to nonmetallic substances that were insoluble in water and resistant to heating, properties shared by these oxides. The realization that these earths were not elements but compounds is attributed to the chemist Antoine Lavoisier. In his “Traité Élementaire de Chemie” (Elements of Chemistry ) of 1789, he called them “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, Humphrey Davy became the first to obtain samples of the metals by electrolysis of their molten “earths”.

  If the alkaline earths are compared to the alkalis, many similarities are apparent. The main difference is the electron configuration, which is ns2 for alkaline earth metals and ns1 for Alkali Metals. But for the alkaline earth metals, the nucleus also contains an additional positive charge. Also, the elements of Group 2 (alkaline earths) have much higher melting points and boiling points compared to those of Group 1 (Alkali Metals). The alkalis are softer and more lightweight than the alkaline earth metals that are much harder and denser.

  The second valence electron is very important when it comes to comparing chemical properties of the alkaline earth and the Alkali Metals. The second valence electron is in the same “sublevel” as the first valence electron. Therefore, the Zeff is much greater. This means that the elements of Group 2 have a smaller atomic radius and much higher ionization energy than those of Group 1. Even though the Group 2 contains a much higher ionization energy, they still form ionic compounds containing 2+ cations. Beryllium, however, behaves differently. This is due to the fact that in order to remove two electrons from this particular atom, significantly more energy is required. It never forms the Be2+

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  Foundations of Chemistry

  An Introductory Course for Science Students

  • Philippa B. Cranwell, Elizabeth M. Page(Authors)
  • 2021(Publication Date)
  • Wiley

    (Publisher)

  Because of this, Group 1 elements do not easily form M 2+ ions. Melting points and boiling points The melting points of Group 1 elements are very low for metallic materials. In fact caesium melts just above room temperature (29 C) (Figure 11.5). The low melting points are due to the weak metallic bonding in the structures. Because Group 1 elements contribute only one electron per atom to the electron sea, the bonding is fairly weak. As the metal atoms get larger on descending the group, the distance between the atomic nuclei and the delocalised electrons increases; thus the attrac-tive forces decrease, making the melting points lower on going down the group. A similar argument can be used to explain the boiling points of the metals, which generally also decrease on going down the group. These weak metallic bonds explain why Alkali Metals are so soft and easily cut, unlike most other metals. 11.2.2 Chemical properties of Group 1 elements The chemical properties of Group 1 elements are dominated by their tendency to form M + cations by losing the outer electron. As the first ionisation energy decreases down the group the tendency to form positive ions increases down the group, so the reactivity of the elements increases from Li to Cs. Reaction with water This is a classic reaction that can be demonstrated with care in the lab for lithium, sodium, and potassium but should not be attempted with either rubidium or caesium. 200 150 100 Melting point/°C 50 0 Li Na K Rb Cs Element Figure 11.5 Melting points of the Group 1 metals. For a reminder of metallic bonding, see Chapter 2. The reactions of Li, Na, and K with water should only be carried out by experienced chemists wearing appro-priate personal protective equipment (PPE). The reaction should be well-screened by conducting it in a fume hood or behind a safety screen, and a full risk assessment should be writ-ten and approved. The reactions of Rb and Cs can be enjoyed on YouTube. 356 The periodic table

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  Concise Chemistry of the Elements

  • S C Siekierski, J Burgess(Authors)
  • 2002(Publication Date)
  • Woodhead Publishing

    (Publisher)

  This is because electrons supplied by the hydrogen atoms are accommodated in the conduction band until this is filled. The Group 10 elements are used as hydrogenation catalysts, which implies formation of hydrides on the surface of the catalyst. However, at moderate pressures only palladium forms a stable PdH_,. (x < 1) bulk phase. The explanation is that of the Group 10 elements 70 Group 1. H~·drogen and the Alkali Metals [Ch. 7 palladium has the lowest atomization enthalpy, which favours rupture of Pd-Pd bonding in the metal. 7.2 THE Alkali Metals 7.2.1 General properties The Alkali Metals h;l'e very low first ionization potentials, 1 1 . This results in high chemical reactivity, particularly with respect to electronegative elements. Reactivity increases down the Group i.e. with decreasing first ionization potential and electron affinity. For instance lithium reacts only slowly with water, whereas sodium reacts vigorously, potassium inflames, and rubidium and caesium react explosively. Reactivity towards liquid bromine also increases very markedly down the Group. Because of very high second ionization potentials, 1 2 , which range from 76 eV for Li to 23 eV for Cs, the Alkali Metals show only the oxidation number+ I. As they are electron-deficient the Alkali Metals can form only metallic solid phases. Their large metallic radii and small number of bonding electrons (two electrons per eight bonds in the case of a body-centered cubic lattice) result in small lattice energies and. therefore. low boiling and melting points. Yhich decrease down the Group. If francium could occur in weighable amounts it would probably be a liquid under normal conditions. The highly elcctropositive character of the Alkali Metals means they react mainly with electronegati,·e elements and form typical ionic compounds. Covalent bonding is found in the M 2 molecules and in the organometallic compounds of lithium. The stability of the M 2 molecules decreases from lithium to caesium. i.e. with increasing ionic radius. 7.2.2 Changes of properties down the Group Fundamental properties of atoms such as (r ns>· ri, r met, E::ns and 1 1 (Table 7. I) change down the Group in a remarkable way. Fig. 7.2. In particular we notice: -Relatively large changes in 1 1 • l:.ns· r 111

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  Chemistry

  The Molecular Nature of Matter

  • Neil D. Jespersen, Alison Hyslop(Authors)
  • 2018(Publication Date)
  • Wiley

    (Publisher)

  The elements in the longer columns (the A groups) are known as the representative elements or main group elements. Those that fall into the B groups in the center of the table are called transition elements. The elements in the two long rows below the main body of the table are the inner transition elements, and each row is named after the element that it follows in the main body of the table. Thus, elements 58–71 are called the lanthanide elements because they follow lanthanum, and elements 90–103 are called the actinide elements because they follow actinium. Some of the groups have acquired common names. For example, except for hydrogen, the Group 1A elements are metals. They form compounds with oxygen that dissolve in water to give solutions that are strongly alkaline, or caustic. As a result, they are called the Alkali Metals or simply the alkalis. The Group 2A elements are also metals. Their oxygen compounds are alkaline, too, but many compounds of the Group 2A elements are unable to dissolve in water and are found in deposits in the ground. Because of their properties and where they occur in nature, the Group 2A elements became known as the alkaline earth metals. On the right side of the table, in Group 8A, are the noble gases. They used to be called the inert gases until it was discovered that the heavier members of the group show a small degree of chemical reactivity. The term noble is used when we wish to suggest a very limited degree of chemical reactivity. Gold, for instance, is often referred to as a noble metal because so few chemicals are capable of reacting with it. Finally, the elements of Group 7A are called the halogens, derived from the Greek word meaning “sea” or “salt.” Chlorine (Cl), for example, is found in familiar table salt, a compound that accounts in large measure for the salty taste of seawater

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