Group 1 Alkali Metals

10 Key excerpts on "Group 1 Alkali Metals"

• 

  eBook - PDF

  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.

  Sign up to read

  Learn more about book
• 

  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.

  Sign up to read

  Learn more about book
• 

  eBook - PDF

  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.

  Sign up to read

  Learn more about book
• 

  eBook - PDF

  Chemistry

  An Industry-Based Introduction with CD-ROM

  • John Kenkel, Paul B. Kelter, David S. Hage(Authors)
  • 2000(Publication Date)
  • CRC Press

    (Publisher)

  We now provide a summary of the descriptive chemistries of selected families and other elements and also present at least one important fact about each of these elements. 1 4.4.1 Alkali Metals Lithium, sodium, potassium, rubidium, and cesium are all highly reactive metals. They all react violently with water, forming hydrogen gas and the metal hydroxide with the formula MOH (“M” symbolizing the metal). They all react readily with oxygen in the air forming the metal oxide with the formula M 2 O, but can also form superoxides of the formula MO 2 . These latter compounds are used in self-contained breathing devices, such as those that deep-sea divers use to generate oxygen for breathing. Each alkali metal must be stored under mineral oil or kerosene in order to preserve them in the elemental state. Since they are so reactive, they are never found in nature in the free elemental state, only in compounds. Each is a silvery, soft, pliable metal that can be cut with a knife. Lithium is used in lithium batteries because it is a light element and produces a higher battery voltage than other metals. Sodium forms a plethora of compounds used in con-sumer products important to our livelihood, including NaCl (table salt), NaHCO 3 (baking soda), and NaOH (drain cleaner, oven cleaner, soap making). Potassium is important for plant growth, and in the form of potassium nitrate (KNO 3 ) and other compounds, is an component of fertilizers. Rubidium is of little value. It is found as a trace impurity in ores 1 See also Stwertka, Albert, A Guide to the Elements, revised edition, Oxford University Press, 1998. Li Na K Rb Cs Fr The Periodic Table 85 of other alkali metals. The only naturally occurring isotope of cesium, cesium-133, is used as the world’s official measure of time. Francium is radioactive (all isotopes) and unstable. 4.4.2 Alkaline Earth Metals Beryllium, magnesium, calcium, strontium, barium, and radium are the alkaline earth metals.

  Sign up to read

  Learn more about book
• 

  eBook - ePub

  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+

  Sign up to read

  Learn more about book
• 

  eBook - PDF

  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

  Sign up to read

  Learn more about book
• 

  eBook - PDF

  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

  Sign up to read

  Learn more about book
• 

  eBook - PDF

  Inorganic Chemistry

  Butterworths Intermediate Chemistry

  • C. Chambers, A. K. Holliday(Authors)
  • 2016(Publication Date)
  • Butterworth-Heinemann

    (Publisher)

  110 6.1 The elements 6.1.1 General characteristics These elements form two groups, often called the alkali (Group I) and alkaline earth (Group II) metals. Some of the physical properties usually associated with metals — hardness, high m.p. and b.p. — are noticeably lacking in these metals, but they all have a metallic appearance and are good electrical conductors. Table 6.1 gives some of the physical properties. From Table 6.1, it is easy to see that Group II metals are more dense, are harder and have higher m.p. and b.p. than the corresponding Group I metals. In Chapter 2, a discussion of the theory of metallic bonding indicated that the strength of such bonding generally depends on the ratio (number of electrons available for bonding)/(atomic radius). The greater this ratio is, the stronger are the bonds between the metal atoms. In the pre-transition metals, this ratio is small and at a minimum in Group I with only one bonding electron. Metallic bond strength is greater in Group II but there are still only two bonding electrons available, hence the metals are still relatively soft and have low melting and boiling points. Hardness, m.p. and b.p. all decrease steadily down Group I, the metallic bond strength decreasing with increasing atomic radius. These changes are not so well marked in Group II but note that beryllium and, to a lesser extent, magnesium are hard metals, as a result of their small atomic size; this property, when coupled with their low density, makes them of some technological importance (p. 113). A full discussion of the changes in ionization energy with group and period position has been given in Chapter 2. These data are given again in Table 6.2. 6.1.2 Formation of ions We note first that the elements are all electropositive, having relatively low ionization energies, and are, in consequence, very reactive.

  Sign up to read

  Learn more about book
• 

  eBook - PDF

  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.

  Sign up to read

  Learn more about book
• 

  eBook - ePub

  Chemistry For Dummies

  • John T. Moore(Author)
  • 2016(Publication Date)
  • For Dummies

    (Publisher)

  alkali metals . In reactions, these elements all tend to lose a single electron. This family contains some important elements, such as sodium (Na) and potassium (K). Both of these elements play an important role in the chemistry of the body and are commonly found in salts.
• The IIA family is made up of the alkaline earth metals . All these elements tend to lose two electrons. Calcium (Ca) is an important member of the IIA family (you need calcium for healthy teeth and bones).
• The VIIA family is made up of the halogens . They all tend to gain a single electron in reactions. Important members in the family include chlorine (Cl), used in making table salt and bleach, and iodine (I). Tincture of iodine is sometimes used as a disinfectant.
• The VIIIA family is made up of the noble gases. These elements are very unreactive. For a long time, the noble gases were called the inert gases, because people thought that these elements wouldn’t react at all. Later, a scientist named Neil Bartlett showed that at least some of the inert gases could be reacted, but they required very special conditions. After Bartlett’s discovery, the gases were then referred to as noble gases.
FIGURE 5-5: Some important chemical families.

What valence electrons have to do with families

Chapter 4 explains that an electron configuration shows the number of electrons in each orbital in a particular atom. The electron configuration forms the basis of the concept of bonding and molecular geometry and other important stuff that I cover in the various chapters of this book.

Tables 5-1 through 5-4 show the electron configurations for the first three members of the families IA, IIA, VIIA, and VIIIA.

TABLE 5-1 Electron Configurations for Members of IA (Alkali Metals)

Element	Electron Configuration
Li	1s2 2s1
Na	1s2 2s2 2p6 3s1
K	1s2 2s2 2p6 3s2 3p6 4s1

TABLE 5-2 Electron Configurations for Members of IIA (Alkaline Earth Metals)

Element	Electron Configuration
Be	1s2 2s2
Mg	1s2 2s2 2p6 3s2
Ca	1s2 2s2 2p6 3s2 3p6 4s2

TABLE 5-3