GROUP 2A ELEMENTS [ALKALINE EARTH METALS]

Contents

• Introduction
• Group 2A Elements
• Occurrence of Alkaline Earth metals
• Physical Properties of Alkaline Earth Metals
• Chemical Properties of Alkaline earth metals
• Beryllium
• Magnesium
• Calcium
• Strontium
• Barium
• Radium
• Differences between Alkali metals and Alkaline earth metals
• Summary
• FAQs
• Recommended Videos
• Test Questions
• Discuss And Explain

Learning Objectives

By the end of this section, you should be able to:

• Explain Physical and Chemical properties of group 2 elements.
• Describe each alkaline earth metals, their properties and their reactions.
• Explains the differences between alkali metals and alkaline earth metals.

Introduction

I believe that this is not your first time to hear the word 'Periodic Table'. The Periodic Table contains all known elements; and these elements are arranged in columns and rows - technically known as groups and periods. 

Elements in the same group share a common characteristic - presence of the same number of electrons in their outermost or valence shell. On the other hand, elements in the same period have only the same number of electron shells. In our last discussion, we talked on elements in Group 1A, right now we're to discuss the elements in Group 2 of the periodic table.

Periodic table of elements

Group 2A Elements

Group 2 / 2A elements are elements that have 2 electrons in their valence shells. The 2 valence electrons are usually found in the s subshell, so we say group 2 elements always have a filled up s subshell in their atoms - (recall, s shell can take maximum of 2 electrons).

Group 2 elements are reactive elements, though their reactivity is usually less than the alkali metals. They usually lose their valence electrons to form compounds with +2 oxidation state, thereby achieving a stable configuration. 

Elements in Group 2 of the periodic table include: Beryllium, Magnesium, Calcium, Strontium, Barium and Radium. 

These elements are referred to as Alkaline earth metals.

Q: Why are group 2 elements referred to as Alkaline earth metals?

A: The plain meaning of Alkaline earth metal is a metal whose oxide is resistant to heat, and also appears to be insoluble or slightly soluble in water to form a basic solution.

This name was proposed by early chemists, and if we were to strictly follow this principle, beryllium wouldn't be called an alkaline earth metal, because its oxide is completely insoluble in water and can behave as either alkaline or acidic (this means it is amphoteric). Also, Strontium and Barium would be referred to as Alkaline metals, not alkaline earth metals, because their oxides dissolve in water without any hassle. However, all elements in group 2 remain to be called alkaline earth metals because IUPAC has chosen to let them be called so. Whenever they change their minds, I'll update this page😉

Occurrence of Alkaline earth metals

Just like the alkali metals, the alkaline earth metals cannot be found in native state (or pure form in nature). They are found in combined forms, and this is because of their high reactivity.

Pure forms of alkaline earth metals are extracted from their minerals or ores. The oxides of alkaline earth metals are called alkaline earths- 

Radium on the other hand is gotten by radioactive decay- usually from uranium minerals.

Alkaline earth metals (except Beryllium and magnesium) are usually stored in paraffin oil/wax or in inert gases to prevent oxidation in air.

Physical Properties of Alkaline Earth Metals

• Alkaline earth metals are silvery white solids at room temperature.

• Alkaline earth metals are harder, denser and stronger than alkali metals.

• The atomic and ionic radii of the alkaline earth metals are smaller than the alkali metals present in their respective periods. This is because: one more proton has been added to the nucleus and one electron is added to the atomic orbitals, therefore, the force of attraction between the protons (in the nucleus) and the electrons (in electron orbitals) become stronger, thus, atomic radius become smaller.

• The melting and boiling points of alkaline earth metals are higher than the alkali metals present in their respective periods. This is due to larger attractive forces between electrons and the protons in the nucleus of their atoms, making the atoms smaller in size. The trend of the melting and boiling points are however not constantly maintained as we move down the group.

• Some Alkaline earth metals have face centered cubic crystalline structure while some others have body centered cubic structure.

• Alkaline earth metals, except Beryllium and Magnesium, give off characteristic flame colours when heated up in flames. I believe I explained why this happens in the previous post- In flame, the electrons are excited to higher energy levels and when they drop back to the ground state, energy is emitted in the form of visible light.

Chemical Properties of Alkaline earth metals

•Alkaline earth metals are strongly electropositive; electropositivity increases down the group.

• Their general electronic configuration may be represented as [noble gas] ns² where n = period number.

• Beryllium and Magnesium form thin layer of oxides on contact with air or water. The oxides are usually non-porous so that there will be no further oxidation. This is known as Passivation. Other alkaline earth metals do not do this.

• Alkaline earth metals combine with non-metals to form ionic compounds. Beryllium acts weird and it forms compounds that are more covalent in nature. You will find out more later in this post.

• Ionic compounds of alkaline earth metals containing monovalent anions (e.g NO₃^-, Cl^-, HCO₃^-, anions with charge of -1) e.g CaCl₂, Mg(NO₃)₂ etc tend to be soluble in water. On the other hand, their hydroxides (OH^-) and fluorides (F^-) tend to be insoluble, but their solubility increases as we move down the group.

• Ionic compounds of alkaline earth metals containing divalent anions (e.g CO₃^2-, SO₄^2-, anions with charge of -2) tend to be insoluble in water, or at best, slightly soluble in water. Ionic salts of SO₄^2- become increasingly insoluble as we move down the group.

• Alkaline earth metals (except Beryllium) react with water (with increasing violence down the group) to form hydroxides and hydrogen gas. The hydroxides of these alkaline earth metals behave as bases.

$${\text{M}}_{(s)}+2{\text{H}}_2{\text{O}}_{(l)}\to {\text{M(OH)}}_{2(aq/s)}+{\text{H}}_{2(g)}$$

PRO TIP: basic behaviour of the hydroxides of alkaline earth metals increases down the group. This is because:

- the solubility of alkaline earths increase down the group

- there is a lower attraction between the OH¯ ions and larger positive ions

- the ions will split away from each other more easily.

- hence, there will be a greater concentration of OH¯ ions in water.

• Alkaline earth metals burn in oxygen to form their corresponding oxides= alkaline earths.

$$2{\text{M}}_{(s)}+{\text{O}}_{2(g)}\to 2{\text{MO}}_{(s)}$$

• Alkaline earth metals react with hydrogen to form ionic hydrides. Beryllium forms covalent polymeric Beryllium hydride and is not gotten by direct reaction with hydrogen.

$${\text{M}}_{(s)}+{\text{H}}_{2(g)}\to {\text{MH}}_{2(s)}$$

These always dissociate at high temperatures.

• Unlike alkali metals (excluding Li), alkaline earth metals react with Nitrogen to form nitrides.

$$3{\text{M}}_{(s)}+{\text{N}}_{2(g)}\to {\text{M}}_3{\text{N}}_{2(s)}$$

The nitrides react with water to form ammonia.

$${\text{M}}_3{\text{N}}_{2(s)}+6{\text{H}}_2{\text{O}}_{(l)}\to 3{\text{M(OH)}}_{2(aq)}+3{\text{NH}}_{3(g)}$$

These Nitrides can also dissociate into the respective metals and nitrogen when heated.

• Alkaline earth metals will react when heated with sulphur or selenium to form sulphides or selenides.

$${\text{M}}_{(s)}+{\text{S}}_{(s)}\to {\text{MS}}_{(s)}$$

$${\text{M}}_{(s)}+{\text{Se}}_{(s)}\to {\text{MSe}}_{(s)}$$

• Alkaline earth metals react with halogens to form halides.

$${\text{M}}_{(s)}+{\text{X}}_{2(g/l)}\to {\text{MX}}_{2(s)}$$

Where X= F, Cl, Br, I

• The nitrates of alkaline earth metals decompose on heating to form their corresponding oxides, nitrogen(IV) oxide and oxygen gas.

$$2\text{M}({\text{NO}}_3{)}_{2(s)}+\text{heat}\to 2{\text{MO}}_{(s)}+4{\text{NO}}_{2(g/l)}+{\text{O}}_{2(g)}$$

• The sulphates of alkaline earth metals decompose on heating to form their corresponding oxides, sulphur(IV)oxide and oxygen gas.

$${\text{MSO}}_{4(s)}+\text{heat}\to 2{\text{MO}}_{(s)}+2{\text{SO}}_{2(s)}+{\text{O}}_{2(g)}$$

• When the carbonates of alkaline earth metals are heated, they decompose into their corresponding oxides and liberate carbon(IV) oxide gas.

$${\text{MCO}}_{3(s)}+\text{heat}\to {\text{MO}}_{(s)}+{\text{CO}}_{2(g)}$$

PRO TIP: as we go down the group, tendency of thermal decomposition of carbonates, nitrates and sulphates of group 2 elements decreases.

• Hard water (=water that fail to form lather with soap) contain divalent cations, alkaline earth metals (Ca²⁺ and Mg²⁺) are usually the ones here. They interfere with action of soaps and detergents and form precipitates (insoluble solids, instead of lather).

Hard water could either be temporary or permanent.

Temporary hard water contain the carbonates of divalent cations, and can be removed by boiling.

$$2{\text{HCO}}_{3(aq)}^{-}+\text{heat}\to {\text{CO}}_{3(aq)}^{2-}+{\text{CO}}_{2(g)}+{\text{H}}_2{\text{O}}_{(l)}$$

If this is left untreated, it usually leads to furring of water heaters, kettles, pipes, etc...

$${\text{M}}_{(aq)}^{2+}+{\text{CO}}_{3(aq)}^{2-}\to {\text{MCO}}_{3(s)}$$

On the other hand, permanent hard water contain anions (or counterions) other than carbonate, e.g sulphates etc... The alkaline earth metal compounds here cannot be removed by boiling, they must be removed by chemical exchangers or ion exchangers; usually permutit (or zeolite) or synthetic resins are used. Synthetic resins are usually preferred. What these resins do is to replace the divalent cations with Na⁺ ions which will make it easier to form lather with soap.

• Alkaline earth metals (except Beryllium) directly react with carbon to form ionic carbides.

$${\text{M}}_{(s)}+2{\text{C}}_{(s)}\to {\text{MC}}_{2(s)}$$

The carbides react with water to give ethyne gas.

$${\text{MC}}_{2(s)}+2{\text{H}}_2{\text{O}}_{(l)}\to {\text{M(OH)}}_{2(s/aq)}+{\text{C}}_2{\text{H}}_{2(g)}$$

• Alkaline earth metal ions have high hydration enthalpies (or hydration energy= energy released when one mole of non-metallic ions dissolve in water to form a dilute solution), thus, compounds of alkaline earth metals usually form hydrates.

PRO TIP: hydration energy decreases down the group. Hydration energy, together with lattice energy affect the solubility of group 2 metals' salts; learn more on this in the recommended Videos section.

• Like alkali metals, alkaline earth metals also dissolve in liquid ammonia to form coloured solutions. Dilute solutions have bright blue colour due to solvated electrons. These solutions decompose very slowly forming amides and evolving hydrogen.

$${\text{M}}_{(s)}+2{\text{NH}}_{3(l)}\to \text{M}({\text{NH}}_2{)}_{2(aq)}+{\text{H}}_{2(g)}$$

When the solution is evaporated, hexammoniate, M(NH₃)₆ is formed. This usually slowly decompose to give amides.

$$\text{M}({\text{NH}}_3{)}_6\to \text{M}({\text{NH}}_2{)}_2+4{\text{NH}}_3+{\text{H}}_2$$

Concentrated solutions of the metals in ammonia are bronze coloured.

• Alkaline earth metals and their oxides react completely with hydrochloric acid, HCl and nitric acid, HNO₃ but their reaction with sulphuric acid, H₂SO₄ becomes increasingly slower from Ca to Ba; this is due to the insolubility of their sulphate salts, thereby preventing the remaining parts of the metal or oxide from taking part in the reaction.

$$\text{MO}+2\text{HCl}\to \text{MCl}+{\text{H}}_2\text{O}$$

$$\text{MO}+{\text{H}}_2{\text{SO}}_4\to {\text{MSO}}_4+{\text{H}}_2\text{O}$$

Reaction with sulphuric acid decreases from Mg - Ba

PRO TIP: Solubility of the sulphate salts of group 2 elements decreases as we move down the group. MgSO₄ is the most soluble sulphate while BaSO₄ is the most insoluble sulphate.

• Alkaline earth metals combine with mercury to form amalgams.

• Alkaline earth metals do not react with noble gases.

M= Be, Mg, Ca, Sr, Ba

Beryllium

Discovery

In 1798, a French chemist named Louis Vauquelin reported that he had found a new 'earth' and he named it glucinium (or glucine, because of its sweet taste). The name of the element was later changed to Berylina, and finally Beryllium in 1828. That year (1828), small samples of Beryllium were gotten by displacement reaction between pure Potassium and Beryllium chloride.

$$2{\text{K}}_{(s)}+{\text{BeCl}}_{2(s)}+\text{heat}\to 2{\text{KCl}}_{(s/aq)}+{\text{Be}}_{(s)}$$

This isolation was carried out by Friedrich Wöhler and Antoine Bussy, the Beryllium samples were however not 100% pure.

Pure samples of Beryllium was not gotten until about 70 years later when Paul LeBeau carried out electrolysis of molten mixture of Beryllium fluoride, BeF₂, and Sodium fluoride, NaF. 

Occurrence

Beryllium occurs in natural minerals like:

• Beryl, emerald, aquamarine, Be₃Al₂[Si₆O₁₈]
• Chrysoberyl, Al₂[BeO₄]
• Phenakite, Be₂SiO₄

Beryllium ores account for 0.00028% of the earth's crust.

Extraction

The extraction of Beryllium from its compounds is a pretty difficult task. Beryllium is usually extracted from Beryl, which will undergo series of reactions until Beryllium chloride is gotten, molten Beryllium chloride then undergoes electrolysis to yield Beryllium.

Beryllium can also be gotten from Beryllium fluoride through reduction of BeF₂ with Mg in graphite crucibles at 1300°C.

$${\text{BeF}}_2+\text{Mg}\to \text{Be}+{\text{MgF}}_2$$

Physical Properties

• Beryllium is a whitish-gray solid at room temperature.

Beryllium metal with passivation layer coat

• Beryllium is a lightweight metal, its density is 1.85 g/cm³.

• Beryllium is a hard and thermally stable metal.

Because of the high thermal stability of Beryllium, its melting and boiling points are very high! The melting point temperature of Beryllium is 1287°C and its boiling point temperature is 2471°C; the highest in group 2A elements.

• Beryllium does not emit any coloured flame when heated in a hot Bunsen burner.

Chemical Properties

• Beryllium is the fourth element in the periodic table; one atom contains 4 electrons and 4 protons in ground state.

• The atomic mass of Beryllium is 9. One atom of Beryllium will therefore have; 9 - 4 = 5 neutrons.

• The electronic configuration of Beryllium is 1s² 2s² or [He] 2s²

• Beryllium displays characteristics that are anomalous to other group 2 elements; this is because of Beryllium's size; it exhibits diagonal relationship with Aluminium. I believe I explained diagonal relationship in the previous blogpost so it is not new to you.

• Beryllium forms covalent compounds with non-metals unlike other elements in group 2; hence, beryllium compounds are insoluble in water but in organic solvents.

• Beryllium compounds are very toxic and can damage the lungs when inhaled.

• Beryllium is a poor reducing agent. This is due to its reluctance to lose its valence electrons.

• Salts of 'Beryllium + large anions' are unstable and those that are stable are hydrated. Examples are BeCO₃·4H₂O, BeSO₄·4H₂O, both of these salts will decompose on heating to give Beryllium oxide BeO.

• There's nothing like beryllium peroxide or Beryllium superoxide, only beryllium oxide exists.

• Beryllium is amphoteric. I believe you know what amphoteric means by now.

• Beryllium compounds tend to have coordination numbers (=number of atoms, molecules or ions bonded to central atom) of 4. For example, [Be(H₂O)₄]²⁺ occurs in hydrated beryllium salts; if you noticed all hydrated beryllium salts, the number of water molecules of hydration is usually 4 or less than, Beryllium is too small to form higher coordination numbers.

• Beryllium chloride, BeCl₂, exists as a polymer. It is a covalent chain polymer, existing as dimers, Be₂Cl₄ or (BeCl₂)₂ when it is in the vapour phase, and it will break-up into monomers, BeCl₂, at higher temperatures (AlCl₃ is similar, but the difference is its shorter chains held together by weak intermolecular van der Waal's forces).

• Beryllium oxide, BeO, is very hard, has a wurtzite structure (hexagonal crystal structure) and a coordination number of 4. Be 4:4 O. On the other hand, oxides of other group 2 elements have coordination numbers of 6 and NaCl structure (face centered cubic).

Reactions of Beryllium

• On exposure to air, Beryllium quickly forms a tough and chemically resistant passivation layer of BeO on its surface which will prevent further oxidation.

• When Beryllium is put in water, it quickly forms the same hardened BeO layer which will prevent the insides from reacting with water, hence beryllium does not react with water even when heated up to very high temperatures. This is because BeO is thermally stable.

• Beryllium reacts freely with acids, displacing hydrogen.

$$\text{Be}+{\text{H}}_2{\text{SO}}_4\to {\text{BeSO}}_4+{\text{H}}_2$$

$$\text{Be}+2\text{HCl}\to {\text{BeCl}}_2+{\text{H}}_2$$

Hydrated Beryllium sulphate

• Unlike other elements in its group, Beryllium can react with alkalies to liberate hydrogen.

$$\text{Be}+\text{NaOH}\to {\text{Na}}_2{\text{BeO}}_2+{\text{H}}_2$$

• Ability of Beryllium (and its oxides) to react with both acids and bases makes them to be referred to as amphoteric. Beryllium hydroxide is also amphoteric.

• While other alkaline earth metals can react directly with hydrogen, Beryllium cannot. Beryllium hydride, BeH₂ is gotten by reacting Beryllium chloride and Lithium aluminium hydride.

$$2{\text{BeCl}}_2+{\text{LiAlH}}_4\to 2{\text{BeH}}_2+\text{LiCl}+{\text{AlCl}}_3$$

Beryllium hydride is a covalent polymer.

• Beryllium burns completely (beyond its passivation layer) in oxygen when heated to temperatures over 600°C. This forms BeO.

Beryllium oxide powder

• Beryllium forms covalent compounds with halogens.

$$\text{Be}+{\text{I}}_2\to {\text{BeI}}_2$$

• Beryllium halides can also be gotten by action of halogen acids on Beryllium metal, its oxide, hydroxide and carbonate.

$$\text{Be}+2\text{HCl}\to {\text{BeCl}}_2+{\text{H}}_2$$

$$\text{BeO}+2\text{HCl}\to {\text{BeCl}}_2+{\text{H}}_2\text{O}$$

$${\text{Be(OH)}}_2+2\text{HI}\to {\text{BeI}}_2+2{\text{H}}_2\text{O}$$

$${\text{BeCO}}_3+2\text{HCl}\to {\text{BeCl}}_2+{\text{H}}_2\text{O}+{\text{CO}}_2$$

• Beryllium powder burns in Nitrogen to yield volatile (covalent character) beryllium nitride while other nitrides are not volatile as they are ionic crystalline solids.

$$3\text{Be}+{\text{N}}_2\to {\text{Be}}_3{\text{N}}_2$$

Beryllium nitride can be hydrolysed with water to liberate ammonia.

$${\text{Be}}_3{\text{N}}_2+6{\text{H}}_2\text{O}\to 3{\text{Be(OH)}}_2+2{\text{NH}}_3$$

• When BeO is heated with carbon at about 2000°C, a brick red coloured carbide, Be₂C, is formed. This on hydrolysis evolves methane and is, thus, called methanide.

$${\text{Be}}_2\text{C}+4{\text{H}}_2\text{O}\to 2{\text{Be(OH)}}_2+{\text{CH}}_4$$

Be(OH)₂ acts as an amphoteric substance.

Insoluble Beryllium hydroxide

Uses of Beryllium

• Beryllium has low and carefully monitored commercial uses because of its toxicity when inhaled. It can cause berylliosis, pneumonia and other associated respiratory illnesses.

• Beryllium is used in the production of alloys. Be-Cu alloys are hard and strong and do not create sparks when they strike other metals.

• Beryllium is used in mechanics, production of aerospace components, fast and guided missiles, and engines of motor cars (but not for long, was later banned I think).

• Beryllium is transparent or translucent to several wavelengths of X-rays and gamma rays, hence it is used as a window material for X-ray equipment and components of particle detectors.

• Beryllium and its oxide are used in electronics, as an insulator and good heat conductor.

Magnesium

Discovery

Magnesium has long been discovered in its salts many years ago, but it was first isolated by Sir Humphry Davy in 1808 by electrolysis of a mixture of magnesia and mercuric oxide.

Occurrence

• Dolomite, CaCO₃·MgCO₃
• Magnesite
• Brucite
• Carnallite
• Talc, Mg₃(OH)₂[Si₄O₁₀]
• Olivine, (Mg,Fe)[SiO₄]

Magnesium ores account for 2.33% of the earth's crust.

Extraction

Several ways can be used to get magnesium from its ores.

Magnesium ores are treated until magnesium chloride is gotten, which will undergo electrolysis to yield magnesium. This is known as dow process.

Magnesium is also gotten by thermal reduction of magnesium oxide with carbon at about 2000°C.

$$\text{MgO}+\text{C}\rightleftharpoons \text{Mg}+\text{CO}$$

The same way, magnesium can also be gotten by action of Ferrosilicate on reduced dolomite. These are known as pidgeon process.

$${\text{CaCO}}_3+{\text{MgCO}}_3\to \text{CaO}+\text{MgO}+2{\text{CO}}_2$$

$$2\text{MgO}+2\text{CaO}+\text{FeSi}\to 2\text{Mg}+{\text{Ca}}_2{\text{SiO}}_4+\text{Fe}$$

Physical Properties

• Magnesium is a shiny greyish-white solid at room temperature.

Metallic Magnesium

• Magnesium has a density of about 1.74 g/cm³.

• The boiling and melting points of magnesium are the lowest in group 2. Its melting point temperature is 650°C and its boiling point temperature is 1090°C.

• Magnesium does not emit any coloured flame when heated in a hot Bunsen burner. It burns with a bright white flame during reaction.

Chemical Properties

• Magnesium has an atomic number of 12, hence its atoms each have 12 electrons and 12 protons in ground state.

• The atomic mass number of magnesium is about 24.307. An atom of Magnesium therefore has; 24.307 - 12 ≈ 12 neutrons.

• The electronic configuration of magnesium is 1s² 2s² 2p⁶ 3s² or [Ne] 3s²

• Magnesium forms ionic compounds that have partial covalent character.

• Magnesium oxide, magnesium hydroxide are slightly soluble in water.

• Magnesium compounds do have coordination number of 6.

• Magnesium hydride is partially covalent, also some halides of magnesium are intermediate between ionic and covalent.

• Mg is found on the same diagonal as Li; therefore, Mg has similar properties to Li.

Both Li and Mg are used as reagents in organic synthesis (Grignard’s reagent).

• Mg²⁺ is an important ion in bio-inorganic chemistry:

- Metal centre in chlorophyll (used for photosynthesis) of plants.

- In the active centres of ATPases and other enzymes.

- PCR (Polymerase Chain Reaction).

- Intracellular fluids.

Reactions of Magnesium

• Magnesium form passivation layer when exposed to air which prevents further oxidation.

• Magnesium reacts very slowly with liquid water, however the reaction can be much faster with steam. When warm water is used, magnesium fizzes and reacts slowly to form magnesium hydroxide and hydrogen gas. There's usually no flame.

$${\text{Mg}}_{(s)}+2{\text{H}}_2{\text{O}}_{(l)}\to {\text{Mg(OH)}}_{2(s/aq)}+{\text{H}}_{2(g)}$$

When steam is used, magnesium burns with a bright white flame to form hydrogen and magnesium oxide, a white powder.

$${\text{Mg}}_{(s)}+{\text{H}}_2{\text{O}}_{(g)}\to {\text{MgO}}_{(s)}+{\text{H}}_{2(g)}$$

Magnesium hydroxide

• Ignited Mg metal will also react with carbon dioxide in the absence of air.

$$2{\text{Mg}}_{(s)}+{\text{CO}}_{2(g)}\to 2{\text{MgO}}_{(s)}+{\text{C}}_{(s)}$$

• Formation of peroxides is possible with magnesium.

Magnesium oxide powder

• When treated with ethyne (or acetylene), Mg forms the acetylide, MgC₂.

$$\text{Mg}+{\text{C}}_2{\text{H}}_2\to {\text{MgC}}_2+{\text{H}}_2$$

• To Prevent Reiteration and an unnecessarily long post, read other reactions that magnesium undergo in the chemical properties of Alkaline earth metals

Uses of magnesium

• Magnesium is used as a reducing agent during extraction of titanium in Kroll process.

• Magnesium compounds are used in medicine as common laxatives and antacids. Magnesium compounds are usually preferred in antacids due to the low solubility of its hydroxide and carbonate, which makes it a weak base, making it less dangerous in cells.

• Magnesium is used in the production of aluminium alloys.

• Magnesium alloys are used in the body frames and engine of automotives.

• Magnesium is also used in the production of phones, laptops, cameras, and other electronic components because of its light weight.

• Magnesium is used in die casting where it is alloyed with zinc.

Hydrated magnesium sulphate/Epsom salt

• Magnesium is used to remove sulphur in production of iron and steel.

• Magnesium is used to prepare Grignard reagents, which are useful in organic synthesis.

Calcium

Discovery

Calcium was first isolated by Sir Humphry Davy in 1808.

Occurrence

Calcium occurs in nature as:

• Gypsum, CaSO₄
• Calcite, CaCO₃
• Dolomite, CaCO₃·MgCO₃
• Fluorite, CaF₂
• Apatite, Ca₅(PO₄)₃(OH,F)

Calcium ores account for about 4.15% of the earth's crust.

Extraction

Calcium is produced only in small amounts by electrolysis. Calcium is also produced from its chloride salts by reduction with sodium.

Physical Properties

• Calcium is a silvery solid (with tints of pale yellow) at room temperature.

Calcium

• Calcium is a very ductile metal.

• Calcium has a density of about 1.55 g/cm³.

• The melting and boiling points of calcium are 842°C and 1484°C respectively.

• Calcium emits brick-red coloured flame when heated in a hot Bunsen burner.

Chemical Properties

• Calcium has an atomic number of 20, hence, it has 20 electrons and 20 protons in an atom.

• The atomic mass number of calcium is 40. A Calcium atom will therefore have; 40 - 20 = 20 neutrons.

• The electronic configuration of calcium is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² or [Ar] 4s²

• Calcium forms only ionic compounds.

• Oxides and hydroxides of calcium are more soluble than that of magnesium and beryllium, and they act as stronger bases.

• Calcium compounds are usually more extensively hydrated than Magnesium compounds.

Calcium oxide powder, common name is quicklime

• Calcium oxide does not react completely in sulphuric acid. It starts normally but stops after sometime with traces of leftover calcium. This owes to the insolubility of calcium sulphate in water.

Reactions of Calcium

See chemical properties of Alkaline earth metals

Find out about stalactites and stalagmites in the calcium section of the Recommended Videos.

Uses of Calcium

• The largest use of metallic calcium is in production of steel.

• Calcium is an important component in bone minerals of animals.

Calcium hydroxide, slacked lime

• Limestone (CaCO₃) is used as a flux in the extraction of iron from its ore.

• Calcium compounds are important components of cement, building and construction materials.

• Calcium is used to strengthen aluminium alloys used for bearing.

• Calcium chloride is used as a dehydrating agent in laboratory.

• Plaster of Paris contains calcium compounds and is used in dentistry, bone damage repair, building and many other commercial uses.

• Bleaching powder, a calcium salt, is used for several industrial purposes as an oxidising agent.

• Calcium is used to remove traces of oxygen and nitrogen, thus it acts as a getter.

• Tricalcium phosphate is used as a polishing agent in toothpaste and antacids.

• Calcium phosphate is used as a leavening agent in baking.

Calcium sulphate

• Calcium is an important component of liming rosin and it is used to make soaps and synthetic resins.

• A type of Calcium compound is used as a bleach in paper making and also as a disinfectant.

Strontium

Discovery

Strontium is named after a small Scottish village, Strontian, where its ore was discovered. It was not isolated until 1808 when Sir Humphry Davy carried out electrolysis of a mixture of Strontium chloride and mercuric oxide. It was then, he gave the element its permanent name, Strontium.

Occurrence

Strontium occurs in nature as:

• Cölestine or Celestine, SrSO₄
• Strontianite, SrCO₃

Strontium ores account for about 0.038% of the earth's crust.

Extraction

Strontium is produced commercially by reduction of Strontium oxide with aluminium.

Strontium is also produced on a small scale by electrolysis of a mixture of Strontium chloride and potassium chloride.

Physical Properties

• Strontium is a silvery metal with a pale yellow tint.

Strontium, stored in Argon

• It is softer than calcium and harder than Barium.

• Strontium has a density of 2.54 g/cm³.

• Its melting point temperature is 777°C and its boiling point temperature is 1382°C.

• Strontium burns with a crimson red flame when heated in a hot Bunsen burner.

Chemical properties

• Strontium has an atomic number of 38.

• Strontium has an atomic mass number of 87.6 Thus, Strontium will have; 87.6 - 38 ≈ 50 neutrons in one atom.

• Electronic configuration of Strontium is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² or [Kr] 5s²

• Like calcium, Strontium forms only ionic compounds.

• Oxides and hydroxides of Strontium are very much soluble in water than those of the metals above it in the group.

Strontium oxide

Reaction of Strontium

• Strontium tarnishes on exposure to air and it is therefore stored in paraffin. It forms only oxides at room temperature because it cannot react with nitrogen below 380°C.

• Strontium can form oxides, and peroxides at higher temperatures.

Read chemical properties of Alkaline earth metals to see other reactions of Strontium

Strontium hydroxide

Uses of Strontium

• Strontium, together with barium, was mostly used in glass for colour television cathode ray tubes where it prevents X-ray emission/leakage. Because CRTs are now replaced by different kinds of displays, the use of Strontium has declined.

• Strontium salts are used in fireworks to give deep red colours.

Insoluble Strontium sulphate/Celestine

• Strontium chloride is used in some toothpastes for sensitive teeth.

• Strontium is also used during refining of zinc to remove small amounts of lead impurities.

Barium

Discovery

Barium was discovered by Carl Wilhelm Scheele in 1774 but was first isolated by Sir Humphry Davy in 1808 by electrolysis. Before then, scientists had tried to isolate barium, but they were only able to get its oxide and nothing beyond that, until the breakthrough of electrolysis.

Occurrence

Barium occurs in nature as:

• Baryte, BaSO₄
• Witherite, BaCO₃

Barium accounts for about 0.042% of the earth's crust.

Extraction

Electrolysis is not feasible for the extraction of Barium. Reduction of treated ore is used instead. Barium metal is produced by reduction of Barium oxide with aluminium at 1,100°C.

$$3{\text{BaO}}_{(s)}+2{\text{Al}}_{(s)}\to 3{\text{Ba}}_{(l)}+{\text{Al}}_2{\text{O}}_{3(s)}$$

Physical Properties

• Barium appears to be a silvery-white metal with tints of pale yellow.

Metallic Barium

• Barium has a density of 3.51 g/cm³ and has a body centered cubic crystalline structure.

• The melting point temperature of Barium is 727°C and its boiling point is 1897°C.

• Barium burns with an apple green coloured flame when heated in a hot Bunsen burner.

Chemical Properties

• Barium is number 56th element on the periodic table. One atom in ground state has 56 electrons and 56 protons.

• Barium has an atomic mass number of 137.3. Hence, Barium has 137.3 - 56 ≈ 81 neutrons in one atom.

• Electronic configuration of Barium is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² or [Xe] 6s²

• Like calcium and Strontium, Barium forms only ionic compounds.

• Oxides and hydroxides of Barium are readily soluble in water, more than the other group 2 elements above it.

• Barium hydroxide has more basic character than others above it in the group.

• Barium can form peroxides and maybe superoxides when burnt in excess oxygen at very high temperatures.

• Solubility of Group 2 sulphates reduces down the group. So systematically, Barium sulphate, BaSO₄ is the least soluble in water.

• If barium metal is reacted with sulphuric acid it will only react slowly as the insoluble barium sulphate produced will cover the surface of the metal (almost same like a passivation layer) and act as a barrier to further attack.

$$\text{Ba}+{\text{H}}_2{\text{SO}}_4\to {\text{BaSO}}_4+{\text{H}}_2$$

Insoluble Barium sulphate powder

The same effect will happen to a lesser extent with metals going up the group.

The same effect does not happen with other acids like hydrochloric or nitric as they form soluble group 2 salts.

• Barium carbonate, Barium nitrate and Barium sulphate decompose the hardest on heating; remember, ease of decomposition of carbonates, nitrates and sulphates reduces down the group.

Reaction Of Barium

• On exposure to air, barium in powder form burst into flames.

• Barium reacts fastest with water, forming Barium hydroxide - the strongest base formed by group 2 elements.

• When barium or its compounds are dissolved in sulphuric acid, H₂SO₄, the reaction is usually slow because of the formation of coats of insoluble BaSO₄ which hinders the reaction from fully occuring.

Read other reactions of Barium in the chemical properties of Alkaline earth metals

Uses of Barium

• Water soluble barium compounds are poisonous and have been used as rat poisons and rodenticides. This is because they dissolve in the body and form wicked, strong bases.

• Metallic Barium is used as a getter (to remove unwanted traces of gases from vacuum tubes) e.g box TV cathode ray tubes. This usage has declined since the release of tubeless TVs and Plasma sets.

• Baryte is used in petroleum industry in drilling fluids.

• Barium sulphate is the mostly used compound of barium. It is used as a radiocontrast agent in X-ray imaging of the digestive system. Its usability is due to its insolubility, thus, it cannot form poisonous strong base.

• Barium nitrate or Barium monochloride give fireworks an apple green or green colour.

• Barium peroxide is used as a catalyst for aluminothermic reactions for welding rail tracks.

• Barium fluoride has a wide transparency range and is used in optics.

Radium

Discovery

Radium was discovered in 1898 by Pierre and Marie Curie but was not isolated until 1911 by Marie Curie and Louis Debierne.

Occurrence

Radium is a radioactive element and can only be gotten from uranium minerals, mostly uraninite (commonly known as pitchblende).

Extraction

Radium has been extracted using three known methods.

First is by electrolysis of Radium chloride, RaCl₂, using mercury as cathode. The Radium-mercury amalgam that was gotten is then heated in hydrogen, leaving pure Radium.

Second is thermal decomposition of Radium azide, Ra(N₃)₂.

Third is the thermal reduction of radium oxide by aluminium at 1200°C.

Physical Properties

• Pure Radium is a silvery metal but it quickly turns black on exposure to air.

Metallic Radium

• Radium has a density of about 5.5 g/cm³ and has a body centered cubic crystalline structure.

• Radium is radioactive. Its most stable isotope, Radium 226 has an half life of about 1602 years. Other isotopes have half lives that span between few days to very few years.

• Known Melting and boiling points of radium are 700 – 970°C and 1140 – 1730°C respectively; reports claim that these values are not yet well established.

• All isotopes of Radium decay into Radon (by alpha decay) or Actinium (by beta decay).

• The people who have at one time handled radium report that it burns with a red flame.

Chemical Properties

• Radium is the 88th element on the periodic table.

• Radium has an atomic mass number of 226. Its most stable isotope is Radium 226; other isotopes are Radium 223, 224, 225, 228.

• Electronic configuration of radium is [Rn] 7s²

• Radium forms only ionic compounds like Barium and Strontium.

• Radium, on exposure to air, reacts with nitrogen to form black nitride, not oxygen.

• Radium usually does not impart any colour to its compounds. Barium impurities usually cause rose colours some other times.

• Radium hydroxide is the most readily soluble amongst the alkaline earth metals hydroxides and it also acts as the strongest base.

• Compounds of radium are harmful and radioactive and may auto break up after sometime.

• Radium halides are colourless, luminous and soluble in water. The Ionizing radiation emitted (from them) excites nitrogen molecules in the air, making it glow (radioluminiscence).

PRO TIP: Ionizing radiation are responsible for causing cancer, hence exposure to radium or its compounds leads to high risks of cancer. This is unlike the non-ionizing radiation released by Bluetooth and WiFi radio waves, which do not cause cancer-

• Radium sulphate (RaSO₄, the most insoluble known sulphate), radium nitrate (Ra(NO₃)₂), radium carbonate (RaCO₃), radium chromate (RaCrO₄), radium iodate (Ra(IO₃)₂), radium tetrafluoroberyllate (RaBeF₄), radium phosphate, etc are all insoluble salts in water.

Uses of Radium

Radium was used for several purposes many years before now, but all uses have been banned because of deaths, caused by radiation. It was unknown (during those years) that radiation was harmful, thus, Radiation from Radium and its compounds has caused quite a number of deaths; you should run when you see one. Some people however, report that the radioactivity of barium is still exploited in the treatment of cancers; I do not know how true that is, so, I cannot tell if it is truly, being used.

Differences between Alkali metals and Alkaline earth metals

Properties	Alkali Metals	Alkaline Earth Metals
Electronic Configuration	One electron is present in valence shell. The configuration is ns1.	Two electrons are present in valence shell. The configuration is ns2.
Valency	Monovalent.	Divalent.
Electropositive Nature	More Electropositive.	Less Electropositive.
Hydroxides	Strong bases, highly soluble and stable towards heat.	Weak bases, less soluble and decompose on heating.
Bicarbonates	These are known in solid states. LiHCO3 is an exception.	These are not known in free state, exist only in solutions.
Carbonates	Soluble in water, do not decompose on heating. Li2CO3 is an exception.	Insoluble in water, decompose on heating.
Action of Nitrogen	Do not directly combine with Nitrogen, Li is an exception.	Directly combine with Nitrogen and form nitrides.
Action of Carbon	Do not directly combine with carbon, Li is an exception.	Directly combine with carbon and form carbides.
Nitrates	Decompose on heating, evolving only Oxygen, Li is an exception.	Decompose on heating, evolving mixture of NO2 and Oxygen.
Solubility of Salts	Sulphates, phosphates, fluorides, chromates, oxalates, etc., are soluble in water. Li is an exception.	Sulphates, phosphates, fluorides, chromates, oxalates, etc., are insoluble in water.
Physical properties	Soft, low melting points. Paramagnetic (=valence electron is unpaired).	Are less reactive and comparatively harder metals. High melting points. Diamagnetic (=valence electrons are paired).
Hydration of compounds	The compounds are less hydrated. NaCl, KCl and RbCl form non-hydrated chlorides.	The compounds are extensively hydrated. MgCl2·6H2O, CaCl2·6H20 and BaCl2·2H2O are hydrated chlorides.
Reducing power	Stronger, as ionisation potential values are low and oxidation potential values are high.	Weaker, as ionisation potential values are high and oxidation potential values are low.

Key Points

📌 Group 2 elements are elements with two valence electrons. Their electronic configuration is usually [Noble gas] ns² where n is the period number.

📌 Group 2 elements are called alkaline earth metals.

📌 Group 2 elements include Beryllium, Magnesium, Calcium, Strontium, Barium and Radium.

📌 Group 2 elements readily lose their valence electrons to form compounds with +2 oxidation states.

📌 Alkaline earth metals are silvery metals and some of them form passivation layer when exposed to air.

📌 Beryllium and Magnesium do not burn with flames when exposed to heat, calcium burns with a brick red flame, Strontium burns with a crimson red flame while Barium burns with a yellow green flame.

📌 The atomic radius and ionic radius of the alkaline earth metals increase as we go down the group while ionization energy, hydration energy and electronegativity decrease as we move down the group.

📌 Reactivity and solubility of the alkaline earth metals increase as we go down the group while hardness and ability to withstand heat decrease as we go down the group.

📌 Alkaline earth metals, except Beryllium, react with water to form their respective hydroxides. Magnesium, unlike the others, would react with boiling water or steam.

📌 The basicity of their hydroxides increases as we go down the group. Beryllium hydroxide on the other hand is amphoteric.

📌 Solubility of the hydroxides of group 2 elements increases as we move down the group.

📌 The halides (sometimes excluding fluorides) and bicarbonates of alkaline earth metals are usually soluble in water. The carbonates and sulphates on the other hand gradually become insoluble in water as we move down the group; these decompose on heating, but the ease of decomposition decreases as we go down the group.

📌 Solubility of the sulphates of alkaline earth metals reduces as we move down the group. Barium sulphate is an insoluble sulphate and this makes it suitable for medicine, unlike its other compounds that dissolves and end up being poisonous because they form strong bases.

📌 Alkaline earth metals form hexammoniate ions when reacted with liquid ammonia; blue when dilute and bronze when concentrated.

📌 Beryllium displays some anomalous characters owing to its small atomic size. It exhibits diagonal relationship with Aluminium in group 3.

📌 While other alkaline earth metals form ionic solids with non-metals, Beryllium forms covalent compounds and magnesium forms some compounds that are partly ionic and partly covalent.

FAQs

Q: Why is (BeCl₂)_n covalent and BaCl₂ is ionic?

A: (BeCl₂)_n is covalent because of a smaller radius, large polarizing power and high electron density of Be²⁺, which makes it harder to lose its electrons, it rather shares them. BaCl₂ is ionic because of the larger radius of Ba²⁺, a heavier and more electropositive group 2 metal ion.

Q: Give a balanced equation for the extraction of magnesium from dolomite.

A: CaCO_3(s) + MgCO_3(s) →CaO_(s) + MgO_(s) + 2CO_2(g)

   

    2CaO_(s) + 2MgO_(s) + FeSi_(s) →2Mg_(s) + Ca₂SiO_4(s) + Fe_(s)

Q: Why is magnesium hydroxide a much more effective antacid than calcium or barium hydroxide?

A: Mg(OH)₂ is sparingly soluble and mildly basic; Ca(OH)₂ is more soluble and so moderately basic; Ba(OH)₂ is very soluble and is also strongly basic, making it a poison.

Q: Explain Why group 1 hydroxides are much more corrosive than group 2 hydroxides? 

A: Group 1 hydroxides are more soluble than group 2 hydroxides, and therefore have higher concentrations of OH^-. The increase in OH^- concentration increases the corrosiveness of group 1 hydroxides. 

Recommended Videos

Alkaline Earth Metals

Solubility of Alkaline earth metals

Beryllium

Magnesium

Calcium

Strontium

Barium

Radium

Test Questions

Discuss And Explain

1. (a) State the trends in solubility of the hydroxides and of the sulphates of the Group II elements from Mg–Ba.

(b) Give the formula of the least soluble hydroxide of the Group 2 metals from Mg to Ba.

2. Magnesium hydroxide can be used to neutralise excess stomach acid .

(a) Explain why it would be preferable to use magnesium hydroxide rather than barium hydroxide for this use.

(b) What advantage does magnesium hydroxide have over calcium carbonate in neutralising stomach acid?

(c) Explain the reason why sugars are often added to antacid tablets.

3. State how barium sulphate is used in medicine. Explain why this use is possible, given that solutions containing barium ions are poisonous.

4. Acidified barium chloride solution is used as a reagent to test for sulphate ions.

(a) Generally, hydrochloric acid is used to acidify the solution. Why is it used?

(b) State why sulphuric acid should not be used to acidify the barium chloride.

(c) Write the simplest balanced equation for the reaction that occurs when acidified barium chloride solution is added to a solution containing sulphate ions.

5. The following pair of compounds BaCl_2(aq) and MgCl_2(aq) can be distinguished by observing what happens in test-tube reactions.

(a) Give a suitable aqueous reagent that could be added separately to each compound.

(b) Describe what you would observe in each case.

6. An aqueous solution of sodium chloride may be distinguished from an aqueous solution of sodium sulphate using a simple chemical test.

(a) Identify a reagent for this test. (The reagent is a group 2 element).

(b) State the observations you would expect to make if the reagent identified in part (a) is added to a separate sample of each solution. Write an equation for any reaction which occurs.

7. This question concerns the chemistry of the Group II metals Mg to Ba.

An aqueous solution of a Group II metal chloride, XCl₂, forms a white precipitate when dilute aqueous sodium hydroxide is added. A separate sample of the solution of XCl₂ does not form a precipitate when dilute aqueous sodium sulphate is added.

An aqueous solution of a different Group II metal chloride, YCl₂, does not form a precipitate when dilute aqueous sodium hydroxide is added. A separate sample of the solution of YCl₂ forms a white precipitate when dilute aqueous sodium sulphate is added.

Suggest identities for the Group II metals X and Y. Write balanced equations, including state symbols, for the reactions.

*Hint: Solubility of hydroxides increases down the group, while solubility of sulphates decreases down the group.

8. Barium metal reacts vigorously with hydrochloric acid but with sulphuric acid the reaction is slow.

(a) Write equations for the reactions of Barium with both acids.

(b) Explain why Barium reacts more slowly with sulphuric acid.

9. Pure magnesium reacts completely with an excess of dilute sulphuric acid.

The reaction of pure calcium with an excess of dilute sulphuric acid is very rapid to start with and then the reaction slows down and stops before all of the calcium has been used up.

Explain why these reactions of magnesium and calcium with dilute sulphuric acid are different. 

10. Both strontium carbonate and strontium sulphate are white solids which are insoluble in water. Strontium carbonate reacts with hydrochloric acid to produce a solution of strontium chloride. Strontium sulphate does not react with hydrochloric acid.

Describe how you would obtain strontium sulphate from a mixture of strontium carbonate and strontium sulphate.

11. What happens when

(a) magnesium is burnt in air.

(b) quick lime, CaO is heated with silica.

(c) chlorine reacts with slaked lime, Ca(OH)₂

(d) calcium nitrate is heated.

12. Predict the products of the following reactions and balance them:

(a) Be(s) + Br₂(l) → ?

(b) Sr(s) + H₂O(l) → ?

(c) Mg(s) + O₂(g) → ?

13. What product do you think is formed by reaction of magnesium with sulphur, a group 6A element? What is the oxidation number of sulphur in the product?

14. Briefly explain why solutions of Ca²⁺ and Mg²⁺ in the presence of carbonate leave deposits, but Na⁺ does not.