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7.7: Group Trends for the Active Metals

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    The elements within the same group of the periodic table tend to exhibit similar physical and chemical properties. Four major factors affect reactivity of metals: nuclear charge, atomic radius, shielding effect and sublevel arrangement (of electrons). Metal reactivity relates to ability to lose electrons (oxidize), form basic hydroxides, form ionic compounds with non-metals. In general, the bigger the atom, the greater the ability to lose electrons. The greater the shielding, the greater the ability to lose electrons. Therefore, metallic character increases going down the table, and decreases going across -- so the most active metal is towards the left and down.

    Group 1: The Alkali Metals

    The word "alkali" is derived from an Arabic word meaning "ashes". Many sodium and potassium compounds were isolated from wood ashes (\(\ce{Na2CO3}\) and \(\ce{K2CO3}\) are still occasionally referred to as "soda ash" and "potash"). In the alkali group, as we go down the group we have elements Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs) and Francium (Fr). Several physical properties of these elements are compared in Table \(\PageIndex{1}\). These elements have all only one electron in their outermost shells. All the elements show metallic properties and have valence +1, hence they give up electron easily.

    Table \(\PageIndex{1}\): General Properties of Group I Metals
    Element Electronic Configuration Melting Point (°C) Density (g/cm3) Atomic Radius Ionization Energy (kJ/mol)
    Lithium \([He]2s^1\) 181 0.53 1.52 520
    Sodium \([Ne]3s^1\) 98 0.97 1.86 496
    Potassium \([Ar]4s^1\) 63 0.86 2.27 419
    Rubidium \([Kr]5s^1\) 39 1.53 2.47 403
    Cesium \([Xe]6s^1\) 28 1.88 2.65 376

    As we move down the group (from Li to Fr), the following trends are observed (Table \(\PageIndex{1}\)):

    • All have a single electron in an 's' valence orbital
    • The melting point decreases
    • The density increases
    • The atomic radius increases
    • The ionization energy decreases (first ionization energy)

    The alkali metals have the lowest \(I_1\) values of the elements

    This represents the relative ease with which the lone electron in the outer 's' orbital can be removed.

    The alkali metals are very reactive, readily losing 1 electron to form an ion with a 1+ charge:

    \[M \rightarrow M^+ + e- \nonumber \]

    Due to this reactivity, the alkali metals are found in nature only as compounds. The alkali metals combine directly with most nonmetals:

    • React with hydrogen to form solid hydrides

    \[2M_{(s)} + H_{2(g)} \rightarrow 2MH(s) \nonumber \]

    (Note: hydrogen is present in the metal hydride as the hydride H- ion)

    • React with sulfur to form solid sulfides

    \[2M_{(s)} + S_{(s)} \rightarrow M_2S_{(s)} \nonumber \]

    React with chlorine to form solid chlorides

    \[2M_{(s)} + Cl_{2(g)} \rightarrow 2MCl_{(s)} \nonumber \]

    Alkali metals react with water to produce hydrogen gas and alkali metal hydroxides; this is a very exothermic reaction (Figure \(\PageIndex{1}\)).

    \[2M_{(s)} + 2H_2O_{(l)} \rightarrow 2MOH_{(aq)} + H_{2(g)} \nonumber \]

    Figure \(\PageIndex{1}\): A small piece of potassium metal explodes as it reacts with water. (CC SA-BY 3.0; Tavoromann)

    The reaction between alkali metals and oxygen is more complex:

    • A common reaction is to form metal oxides which contain the O2- ion

    \[4Li_{(s)} + O_{2 (g)} \rightarrow \underbrace{2Li_2O_{(s)}}_{\text{lithium oxide}} \nonumber \]

    Other alkali metals can form metal peroxides (contains O22- ion)

    \[2Na(s) + O_{2 (g)} \rightarrow \underbrace{Na_2O_{2(s)}}_{\text{sodium peroxide}} \nonumber \]

    K, Rb and Cs can also form superoxides (O2- ion)

    \[K(s) + O_{2 (g)} \rightarrow \underbrace{KO_{2(s)}}_{\text{potassium superoxide}} \nonumber \]

    Colors via Absorption

    The color of a chemical is produced when a valence electron in an atom is excited from one energy level to another by visible radiation. In this case, the particular frequency of light that excites the electron is absorbed. Thus, the remaining light that you see is white light devoid of one or more wavelengths (thus appearing colored). Alkali metals, having lost their outermost electrons, have no electrons that can be excited by visible radiation. Alkali metal salts and their aqueous solution are colorless unless they contain a colored anion.

    Colors via Emission

    When alkali metals are placed in a flame the ions are reduced (gain an electron) in the lower part of the flame. The electron is excited (jumps to a higher orbital) by the high temperature of the flame. When the excited electron falls back down to a lower orbital a photon is released. The transition of the valence electron of sodium from the 3p down to the 3s subshell results in release of a photon with a wavelength of 589 nm (yellow)

    Flame colors:

    • Lithium: crimson red
    • Sodium: yellow
    • Potassium: lilac

    Group 2: The Alkaline Earth Metals

    Compared with the alkali metals, the alkaline earth metals are typically harder, more dense, melt at a higher temperature. The first ionization energies (\(I_1\)) of the alkaline earth metals are not as low as the alkali metals. The alkaline earth metals are therefore less reactive than the alkali metals (Be and Mg are the least reactive of the alkaline earth metals). Several physical properties of these elements are compared in Table \(\PageIndex{2}\).

    Table \(\PageIndex{2}\): General Properties of Group 2 Metals
    Element Electronic Configuration Melting Point (°C) Density (g/cm3) Atomic Radius Ionization Energy (kJ/mol)
    Beryllium \([He]2s^2\) 1278 1.85 1.52 899
    Magnesium \([Ne]3s^2\) 649 1.74 1.60 738
    Calcium \([Ar]4s^2\) 839 1.54 1.97 590
    Strontium \([Kr]5s^2\) 769 2.54 2.15 549
    Barium \([Xe]6s^2\) 725 3.51 2.17 503

    Calcium, and elements below it, react readily with water at room temperature:

    \[Ca_{(s)} + 2H_2O_{(l)} \rightarrow Ca(OH)_{2(aq)} + H_{2(g)} \nonumber \]

    The tendency of the alkaline earths to lose their two valence electrons is demonstrated in the reactivity of Mg towards chlorine gas and oxygen:

    \[Mg_{(s)} + Cl_{2(g)} \rightarrow MgCl_{2(s)} \nonumber \]

    \[2Mg_{(s)} + O_{2(g)} \rightarrow 2MgO_{(s)} \nonumber \]

    The 2+ ions of the alkaline earth metals have a noble gas like electron configuration and are thus form colorless or white compounds (unless the anion is itself colored). Flame colors:

    • Calcium: brick red
    • Strontium: crimson red
    • Barium: green

    Contributors and Attributions

    7.7: Group Trends for the Active Metals is shared under a CC BY-NC-SA 3.0 license and was authored, remixed, and/or curated by LibreTexts.