7.7: Group Trends for the Active Metals

Group 1A: The Alkali Metals

1A

3 Li

11 Na

19 K

37 Rb

55 Cs

87 Fr

The word "alkali" is derived from an Arabic word meaning "ashes". Many sodium and potassium compounds were isolated from wood ashes (Na₂CO₃ and K₂CO₃ are still occasionally referred to as "soda ash" and "potash").

As we move down the group (from Li to Fr) we find the following trends:

• 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)

Note

The alkali metals have the lowest I₁ 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 -> M⁺ + e^-

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) + H₂(g) -> 2MH(s)

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

• React with sulfur to form solid sulfides

2M(s) + S(s) -> M₂S(s)

• React with chlorine to form solid chlorides

2M(s) + Cl₂(g) -> 2MCl(s)

• Alkali metals react with water to produce hydrogen gas and alkali metal hydroxides (very exothermic)

2M(s) + 2H₂O(l) -> 2MOH(aq) + H₂(g)

The reaction between alkali metals and oxygen is more complex:

• A common reaction is to form metal oxides which contain the O^2- ion

4Li(s) + O₂ (g) -> 2Li₂O(s) (lithium oxide)

• Other alkali metals can form metal peroxides (contains O₂^2- ion)

2Na(s) + O₂ (g) -> Na₂O₂(s) (sodium peroxide)

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

K(s) + O₂ (g) -> KO₂(s) (potassium superoxide)

Note:

• 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.
• 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 2A: The Alkaline Earth Metals

2A

4 Be

12 Mg

20 Ca

38 Sr

56 Ba

88 Ra

Compared with the alkali metals, the alkaline earth metals are typically:

• harder
• more dense
• melt at a higher temperature

The first ionization values (I₁) of the alkaline earth metals are not as low as the alkali metals:

Note

The alkaline earth metals are therefore less reactive than the alkali metals (Be and Mg are the least reactive of the alkaline earth metals)

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

Ca(s) + 2H₂O(l) -> Ca(OH)₂(aq) + H₂(g)

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) + Cl₂(g) -> MgCl₂(s)

2Mg(s) + O₂(g) -> 2MgO(s)

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

Mike Blaber (Florida State University)