7.7: Group Trends For The Active Metals - Chemistry LibreTexts

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 (Na2CO3 and 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:

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

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

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

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

• React with sulfur to form solid sulfides

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

React with chlorine to form solid chlorides

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

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

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

The reaction between alkali metals and oxygen is more complex:

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

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

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

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

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

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

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