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3.2: Compounds of the Alkali Metals

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    The chemistry of the alkali metals is dominated by the stability of the +1 oxidation state and the noble gas configuration of the M+ cation. The alkali metals all have low first ionization energies (Table \(\PageIndex{3}\).6) but very high second ionization energies.

    Element 1st ionization energy (kJ/mol)
    Li 526
    Na 502
    K 425
    Rb 409
    Cs 382
    Table \(\PageIndex{3}\).6: First ionization potentials for the alkali metals.

    As a consequence of the stability of M+, the Group 1 metals have the least variation in chemistry of the any Group in the periodic table. The only exceptions are the subtle trends that exist for lithium due to its small size (Table \(\PageIndex{3}\).7). All of the metals are more electropositive than hydrogen (Table \(\PageIndex{3}\).8).

    Element Atomic radius (Å) Ionic radius (Å) Covalent radius (Å) Van der Waals radius (Å)
    Li 1.52 0.68 1.52 1.82
    Na 1.86 0.97 1.53 2.27
    K 2.31 1.33 1.90 2.75
    Rb 2.44 1.47 2.47 -
    Cs 2.62 1.67 2.65 -
    Fr - 1.80 2.70 -
    Table \(\PageIndex{3}\).7: Radii of alkali metals. N.B. Some values are unknown.
    Element Electronegativity
    H 2.20
    Li 0.98
    Na 0.93
    K 0.82
    Rb 0.82
    Cs 0.79
    Fr 0.70
    Table \(\PageIndex{3}\).8: Pauling electronegativities of Group 1 elements.

    Solid State

    In the solid state, the compounds of the alkali metals generally form ionic lattices, e.g., \(\ce{Na^{+}Cl^{-}}\). These structures are essentially electrostatic in nature and the lattice energy is usually defined as the enthalpy of formation of the ionic compound from gaseous ions and as such is invariably exothermic.

    \[M^+_{(vap)} + X^-_{(vap)} \rightarrow MX_{(s)} \]

    In all cases the lattice energies is high and is found to be proportional to the ratio of the charges on the ions and the sum of the ionic radii (\(r\)).

    \[ U \propto \dfrac{z^+ z^-}{\displaystyle \sum(r)} \]

    The ionic radii for alkali metal cations are given in Table \(\PageIndex{3}\).7; those for common anions are given in Table \(\PageIndex{3}\).9.

    Anion Ionic radius (Å)
    F- 1.33
    Cl- 1.81
    Br- 1.96
    I- 2.20
    H- 1.54
    O2- 1.32
    Table \(\PageIndex{3}\).9: Ionic radii of common anions.

    The ratio of the ionic radii (r+/r-) neatly defines the structural type observed for alkali metal salts (Table \(\PageIndex{3}\).10). The unit cells for ZnS (zinc blende), NaCl (rock salt), and CsCl are shown in Figure \(\PageIndex{3}\).11, Figure \(\PageIndex{3}\).12, and Figure \(\PageIndex{3}\).13, respectively.

    r+/r- Structural type Metal coordination number
    0.225 - 0.414 ZnS (zinc blende) 4
    0.414 - 0.732 NaCl (rock salt) 6
    0.732 - CsCl 8
    Table \(\PageIndex{3}\).10: Defining MX structural types for alkali metal salts.
    Figure \(\PageIndex{3}\).11.png
    Figure \(\PageIndex{11}\): Unit cell structure of a zinc blende (ZnS) lattice.
    Figure \(\PageIndex{3}\).12.png
    Figure \(\PageIndex{12}\): Unit cell structure of a rock salt lattice.
    Figure \(\PageIndex{3}\).13.png
    Figure \(\PageIndex{13}\): Unit cell of CsCl.

    As an example, the structure of KBr can be predicted from the data in Table \(\PageIndex{3}\).7 and Table \(\PageIndex{3}\).9. The ionic radius of K+ is 1.33 Å, while that for Br- is 1.96 Å. The ratio of the ionic radii (r+/r-) is 0.67. Hence, KBr has a NaCl (rock salt) structure.

    Exercise \(\PageIndex{1}\): Sodium Hydride

    What is the structure of NaH?


    The ionic radius of Na+ is 0.97 Å, while that for H- is 1.54 Å. The ratio of the ionic radii (r+/r-) is 0.67. Hence, NaH has a NaCl (rock salt) structure

    Exercise \(\PageIndex{2}\): Rubidium Fluoride

    What is the structure of RbF?


    The ionic radius of Rb+ is 1.47 Å, while that for F- is 1.33 Å. The ratio of the ionic radii (r+/r-) is 1.10. Hence, RbF has a CsCl structure.


    The coordination complexes of the alkali metal cations (M+) involve electrostatic, or ion-dipole, interactions (Figure \(\PageIndex{3}\).14) that have no preferred direction of interaction. Thus, the ionic radius of the cation (Table \(\PageIndex{3}\).7) controls the coordination numbers of the metal in its complexes (Table \(\PageIndex{3}\).11).

    Figure \(\PageIndex{3}\).14.png
    Figure \(\PageIndex{14}\): A schematic representation of the ion-dipole or electrostatic interaction of Na+ with water.
    Aqua ion n


    [Na(H2O)n]+ 6
    [K(H2O)n]+ 6
    [Rb(H2O)n]+ 6-8
    [Cs(H2O)n.]+ 8
    Table \(\PageIndex{3}\).11: Coordination number for alkali metal ions in aqueous complexes.

    In general the alkali metal ions form complexes with hard donor such as oxygen (H2O, ROH, RCO2-, etc.) or nitrogen (e.g., NH3, NR3, etc.). The aquo complexes readily exchange the water for other ligands, (3.13); however, the equilibrium constants are small when the ligand is similar in size to water.

    \[ [M(H_2O)_n]^+ + L \leftrightharpoons [M(H_2O)_{n-1}L]^+ + H_2O\]

    As a consequence of the low equilibrium constants for monodentate ligands, the alkali metal cations, M+, favor coordination to polydentate ligands such as ethylenediaminetetraacetic acid (EDTA, Figure \(\PageIndex{3}\).15), polyethers, and even natural polyesters or polypeptides. In each case the polydentate ligand wraps itself around the cation.

    Figure \(\PageIndex{3}\).15.png
    Figure \(\PageIndex{15}\): Structure of ethylenediaminetetraacetic acid (EDTA).

    Macrocyclic ligands

    Macrocyclic ligands represent a special class of polydentate ligand. They are defined as being a cyclic compound with nine or more members including all the heteroatoms and with three or more donor atoms. and are given the special name of cryptands when they are synthetic bi- and poly-cyclic multidentate ligands. Crown ethers is the name applied to macrocyclic ligands that consist of a ring containing several ether groups. Figure \(\PageIndex{3}\).16 shows several common macrocyclic ligands.

    Figure \(\PageIndex{3}\).16.png
    Figure \(\PageIndex{16}\): The structures of some common macrocyclic ligands.

    Macrocyclic ligands are generally oxygen or nitrogen donor ligands, and they form highly stable 1:1 complexes with alkali metal ions in which all or most of the coordination sites on the metal are occupied, i.e., [M(L)]+ rather than [M(L)(H2O)n]+. Since the external surface of the macrocyclic ligands comprises of organic residue (e.g., CH2 groups) the complexes are soluble in organic solvents. Thus, crown ethers are commonly used to solubilize salts (e.g., NaCl) in organic solvents. They have also been used to create solutions of nanoparticles, such as carbon nanotubes (Figure \(\PageIndex{3}\).17), that are ordinarily highly insoluble in common organic solvents (Table \(\PageIndex{3}\).12).

    Figure \(\PageIndex{3}\).17.png
    Figure \(\PageIndex{17}\): Representation of the reaction between Na/Hg amalgam, dibenzo-18-crown-6, and puried single walled carbon nanotubes (SWNTs) in toluene and the formation of the [Na(dibenzo-18-crown6)]n[SWNT] complex. Adapted from R. E. Anderson and A. R. Barron, J. Nanosci. Nanotechnol., 2007, 7, 3436
    Solvent SWNT concentration (mg/L)
    CH2Cl2 14.05
    DMF 11.42
    hexane 4.65
    toluene 3.79
    EtOH 2.37
    MeOH 2.37
    CHCl3 0.30
    H2O 0.10
    Table \(\PageIndex{3}\).12: The concentration of Na/dibenzo-18-crown-6 solubilized reduced single walled carbon nanotubes (SWNTs) in various solvents.

    The most important factor in the coordination of various macrocyclic ligands to alkali metal ions is the relative size of the ion and the ligand cavity. For example, 4,7,13,16,21,24-hexaoxa-1,10diazabicyclo[8,8,8]hexacosane is a potentially octa-dentate ligand; however, the binding eciency is very dependant on the identity of the M+ ion (Figure \(\PageIndex{3}\).18). The low binding constant for lithium is probably as a consequence that the ligand would have to distort to coordinate to the small cation. Conversely, the lower equilibrium for caesium is due to its being to large to t completely into the ligand cavity.

    Figure \(\PageIndex{3}\).18.png
    Figure \(\PageIndex{18}\): : Relative stability constants (log K) of alkali metal ions to 4,7,13,16,21,24-hexaoxa-1,10diazabicyclo[8,8,8]hexacosane and relationship to the M+ ionic radius.

    One application of the size effect for macrocyclic ligands is the ability to selectively bind different metals. For example, the 4,7,13,16,21-pentaoxa-1,10-diazabicyclo[8,8,5]tricosane and 4,7,13,16,21,24-hexaoxa-1,10diazabicyclo[8,8,8]hexacosane ligands shown in Figure \(\PageIndex{3}\).19 have very different binding constants to Na+ and K+, as a consequence of the relative size of the cation and the ligand cavity (Table \(\PageIndex{3}\).13).

    Figure \(\PageIndex{3}\).19.png
    Figure \(\PageIndex{19}\): Structure of (a) 4,7,13,16,21-pentaoxa-1,10-diazabicyclo[8,8,5]tricosane [2,2,1] and (b) 4,7,13,16,21,24-hexaoxa-1,10-diazabicyclo[8,8,8]hexacosane [2,2,2].
    Cation [2,2,2] [2,2,1]
    Na+ 800 250,000
    K+ 250,000 700
    Table \(\PageIndex{3}\).13: Equilbrium constants for the M+ complexes for macrocyclic ligands with different cavity sizes. The structures of the ligands are shown in Figure \(\PageIndex{3}\).19.


    1. R. E. Anderson and A. R. Barron, J. Nanosci. Nanotechnol., 2007, 7, 3436.
    2. J. M. Lehn, Supramolecular Chemistry: Concepts and Perspectives. VCH (1995).

    This page titled 3.2: Compounds of the Alkali Metals is shared under a CC BY 3.0 license and was authored, remixed, and/or curated by Andrew R. Barron (CNX) via source content that was edited to the style and standards of the LibreTexts platform; a detailed edit history is available upon request.