Apart from magnesium, which elements form basic carbonates


The alkaline earth metals all have the valence electron configuration s2p0 in detail:
  • Be (beryllium): 2 s2
  • Mg (magnesium): 3 s2
  • Ca (calcium): 4 s2
  • Sr (strontium): 5 s2
  • Ba (barium): 6 s2
  • Ra (radium): 7 s2
All elements therefore form divalent ions. Exceptions to this are the subnitrides M.2N (M. = Ca, Sr, Ba) and the subhalide CaCl. Except for beryllium, there are far-reaching similarities between the elements. Radium is radioactive. The properties and some important compounds of the alkaline earth metals are shown in the following overview table 3.1.1. summarized:
Tendencies ⟹ Atomic and ionic radii, reactivity ⟹
⟹ Basicity of the oxides, solubility of the hydroxides ⟹
⟸ Radius / heat of hydration, ionization energy ⟸
⟸ Boiling and melting points, solubility of sulfates, lattice energy of salts ⟸
EN 1.5 1.2 1.0 1.0 0.9
E.0 [V] -1.85 -2.37 -2.87 -2.89 -2.91
r[6]M.2+ [pm] 45 72 100 118 135
Elements (Section 3.2.) resistant to air (passivation) gray air-sensitive metals, with H2O evolution of H2
structure h.c.p. f.c.c. b.c.c.
Mp. [OC] 1278 649 839 768 727
presentation chemically or electrochemically today only chemically
Oxides (Section 3.4.) BeO MgO (magnesia) CaO (quick lime, quicklime) SrO BaO, BaO2
Hydroxides Be (OH)2 (amphoteric)M.(OH)2 (basic)
Carbonates . MgCO3 (Magnesite) CaCO3 (Calcite, aragonite, vaterite) SrCO3 (Celestine) .
Nitrates all easily soluble
Sulfates . MgSO4.7 H.2O (Epsom Salt) CaSO4.2 H2O (plaster of paris), CaSO4.1/2 H2O (anhydrite) SrSO4 (Strontianite) BaSO4 (Barite)
Halides covalent Fluoride sparingly soluble
other verb. . MgNH4(PO4) . 6 H.2O Approx5[PO4]3(OH) (apatite) . BaCrO4
Flame color no Orange red red pale green
Line positions in [nm] . . 622.0 650-660 524.2
. . 553.3 460.7 513.9
Tab. 3.1.1. Overview of alkaline earth metals

Those in the group correspond to the ratios observed for the alkali metals, i.

  • from top to bottom increase in the periodic table:
    • the atomic and ionic radius,
    • the responsiveness,
    • the electropositive character (the normal potentials range from -1.8 V for Be to -2.9 V for Ba, see Table 3.1.2)
    • the basicity of the oxides
    • the solubility of the hydroxides (the solubility products range from LBe (OH)2=3.10-4 g / l for beryllium hydroxide up to LBa (OH)2= 40 g / l).
  • From the bottom (barium) to the top (beryllium) increase:
    • the radius of hydration, the heat of hydration and the hydration number,
    • the heat of sublimation,
    • the ionization energy,
    • the boiling and melting points (except for Mg),
    • the electronegativity,
    • the solubility of sulfates,
    • the lattice energy of the salts (which is the reverse of the ionic radii) and
    • the resistance of the elements to air and H2O.

Table 3.1.2. shows how the alkaline earth metals fit into the series of voltages.

element oxidized reduced E [V]
Fluorine (F) F.2+ 2e- 2 F- +2.87 V.
oxygen H2O2 + 2 H.3O+ + 2e- 4 H.2O +1.78
Gold (Au)Au++ e- Au +1.69 V
Au3++ 3e- Au +1.50 V
Au3++ 2e- Au++1.40 V
Chlorine (Cl) Cl2+ 2e- 2Cl-+1.36 V
Bromine (Br) Br2+ 2e- 2Br-+1.07 V.
Silver (Ag) Ag++ e- Ag +0.80 V
Iron (Fe) Fe3++ e- Fe2++0.77 V
Iodine (I) I.2+ 2e- 2I-+0.53 V
Copper (Cu) Cu++ e- Cu +0.52 V
Cu2++ 2e- Cu +0.34 V.
Cu2++ e- Cu++0.16 V.
Hydrogen (H) 2H++ 2e-H20 V
Cadmium (Cd) CD2++ 2e- CD -0.40 V
Iron (Fe) Fe2++ 2e- Fe -0.45 V.
Zinc (Zn) Zn2++ 2e- Zn -0.76 V
Hydrogen (H) 2 H2O+ 2e- H2 + 2 OH--0.83 V
Aluminum (Al) Al3++ 3e- Al -1.66 V
Beryllium (Be) Be2++ 2e- Be -1.85 V
Magnesium (Mg) Mg2++ 2e- Mg -2.37 V
Sodium (Na) N / A++ e- N / A -2.71 V
Calcium (Ca) Approx2++ 2e- Approx -2.82 V
Barium (Ba) Ba2++ 2e- Ba -2.91 V
Potassium (K) K++ e- K -2.92 V
Lithium (Li) Li++ e- Li -3.04 V
Tab. 3.1.2. Standard potentials of selected redox couples

In the alkaline earth metals have greater ionization energies. Most salts are less soluble than the corresponding alkali metal salts because of the higher cation charge and the resulting increased lattice energy.

The detection of the heavy alkaline earth metals is usually still good spectroscopically. Atomic absorption spectrometry is possible without restrictions. Except for magnesium, the alkaline earth metals show characteristic flame colors: Ca: orange, Sr: red and Ba: green. These flame colors can no longer be explained by simple electron transitions in the atom.

Flame colors of Ca - Sr - Ba

To discover the elements:

  • Beryllium was first produced by Wöhler in 1828 by reducing BeCl2 made with potassium.
  • Magnesium, calcium, strontium and barium were first obtained as elements by the amalgam method (cf. Na production) by H. Davy (1810) and J. J. Berzelius. When reducing the salts at the Hg cathode, 500OC the elemental mercury in an H2-Atmosphere are withdrawn and the pure metals are left behind.
  • Radium was obtained by P. and M. Curie from the pitchblende in the processing of UO2 First received in 1898.
(‣SVG on element frequencies in the earth's crust)
  • Beryllium is quite rare, but it is concentrated in the mineral beryl (see Fig. 3.1.1.), A hexacyclosilicate of the formula Be3Al2[Si6O18] in front. Colored varieties of this are the emerald (green by doping with Cr3+) and the aquamarine (light blue due to Fe doping).
    Fig. 3.1.1. Beryl (left) and aquamarine (middle) with their structure (right; green octahedron: AlO6; yellow tetrahedron: BeO4) ‣VRML
  • Magnesium and calcium are very common and occur in the form of carbonates, sulfates, silicates, and magnesium, also dissolved in the sea. The most important minerals are:
    • Mg (CO3) (Magnesite)
    • MgCa (CO3)2 (Dolomite, Mg / Ca always in an exactly stoichiometric 1: 1 ratio)
    • MgAl2O4 (Spinel)
      Fig. 3.1.2. Spinel (the little red octahedron!)
    • (Mg / Fe)2SiO4 (Olivine)
    • CaCO3, whereby there are three modifications to be distinguished:
      1. Calcite (limestone, calcite, marble, chalk, pearls) (Fig. 3.1.3; see also building materials in Chapter 3.8 of this lecture). Calcite is the most stable and by far the most common form of CaCO3.
      2. Aragonite (in warmer areas such as Spain) (see Fig. 3.1.4.)
      3. Vaterite (in biological systems, e.g. eggshells; otherwise only synthetic; metastable to calcite at any temperature).
    Fig. 3.1.3. Kalkspar (Calcite, 2x)Fig. 3.1.4. Aragonite
    • CaSO4 . 2 H2O (gypsum) (see Fig. 3.1.5.) (VRML of the structure, see also for building materials in Chapter 3.8. Of this lecture)
    • CaF2 (Fluorspar) (see Fig. 3.1.6.)
    • Approx5(PO4)3F (apatite) (see Fig. 3.1.7.)
    • Anorthite, the Ca feldspar Ca [Al2Si2O4] (see chapter 8.6. of silicate chemistry)
    • CaSiO3 (Wollastonite, a pyroxene; see chapter 6.1. Of silicate chemistry)
    ... and HERE from the LMU (Prof. Klüfers) another nice page on the importance of calcium minerals in a biological context (biomineralization).
  • Strontium and barium are relatively rare, the most important minerals are:
    • for strontium: SrCO3 (Strontianite, Fig. 3.1.8 left, with elemental sulfur) and SrSO4 (Celestine, Fig. 3.1.8 right)
      Fig. 3.1.8. Strontianite (left) and Celestine (right)
    • for barium: BaSO4 (Barite, barite, Fig. 3.1.9.) (Mining: 5 million t / year, e.g. also in Oberwolfach / Kinzigtal)
    Fig. 3.1.9. Barite, on the right in chisel spar form from Oberwolfach
  • Radium only has radioactive isotopes, 226Ra arises as the decay product of 238U (see nuclide map (from Wikipedia)

External links to minerals (partly with explanation)

Links to web elements

with very comprehensive information on the elements: