S-Block Elements: Groups 1 & 2

• The s-block elements are those in which the last electron enters the s-orbital
• The s-block elements are located in Groups 1 and 2 of the periodic table
• Group 1 elements are known as alkali metals and Group 2 elements are known as alkaline earth metals

Electron Configuration

• Alkali metals have the general electronic configuration of ns¹
• Alkaline earth metals have the general electronic configuration of ns²
• The s-block elements are characterized by having their outermost electrons in the s-orbital
• The filling of the s-orbitals determines many of their chemical and physical properties

Alkali Metals

• Alkali metals include Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), and Francium (Fr)
• Alkali metals are silver-white, soft, and metallic
• Alkali metals are highly reactive due to their low ionization enthalpies
• They readily lose their one valence electron to form univalent cations (M⁺)
• All alkali metals react vigorously with water to form hydrogen gas and hydroxides

Reactivity Trends of Alkali Metals

• Reactivity increases down the group
• This is due to the decrease in ionization enthalpy
• Lower ionization enthalpy means it's easier to lose the outermost electron
• Cesium is the most reactive alkali metal, while Lithium is the least reactive

Physical Properties of Alkali Metals

• Low densities compared to other metals
• Low melting and boiling points
• These properties are attributed to their relatively weak metallic bonding

Alkaline Earth Metals

• Alkaline earth metals include Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), and Radium (Ra)
• They are also reactive, but less so than alkali metals
• Alkaline earth metals have higher ionization enthalpies compared to alkali metals
• They tend to lose their two valence electrons to form divalent cations (M²⁺)
• Alkaline earth metals are harder and denser than alkali metals

Properties Down the Group

• Atomic and ionic radii increase down the group for both alkali and alkaline earth metals
• Ionization enthalpy decreases down the group
• Electronegativity decreases down the group
• Metallic character increases down the group

Reaction with Water

• Alkaline earth metals react with water to form hydrogen gas and hydroxides, though less vigorously than alkali metals
• Beryllium does not react with water unless at very high temperatures
• Magnesium reacts very slowly with cold water
• Calcium, strontium, and barium react more readily

Flame Coloration

• Alkali and alkaline earth metals impart characteristic colors to a flame
• This is due to the excitation of their valence electrons to higher energy levels
• Lithium (Li) gives a crimson red flame
• Sodium (Na) gives a yellow flame
• Potassium (K) gives a lilac flame
• Calcium (Ca) gives a brick red flame
• Strontium (Sr) gives a crimson flame
• Barium (Ba) gives an apple green flame

Oxides and Hydroxides

• Alkali metals form various oxides upon reaction with oxygen
• They include oxides (O²⁻), peroxides (O₂²⁻), and superoxides (O₂⁻)
• Lithium mainly forms the monoxide (Li₂O)
• Sodium forms the peroxide (Na₂O₂)
• Potassium, rubidium, and cesium form superoxides (KO₂, RbO₂, CsO₂)
• Alkaline earth metals form monoxides (MO) upon reaction with oxygen
• These monoxides react with water to form hydroxides M(OH)₂
• The hydroxides of alkaline earth metals are alkaline in nature

Uses

• Sodium is used in street lighting and as a heat exchange fluid in nuclear reactors
• Potassium is an essential nutrient for plants and animals
• Magnesium is used in lightweight alloys and in the production of Grignard reagents
• Calcium is important for bones and teeth and used in cement

Important Compounds

• Sodium chloride (NaCl) is common table salt
• Sodium hydroxide (NaOH) is used in the manufacture of soap and paper
• Calcium carbonate (CaCO₃) is found in limestone and marble
• Magnesium sulfate (MgSO₄) is known as Epsom salt and used in medicine