The alkali earth metals are a group of elements found in the second column (Group 2) of the periodic table. The elements of this group are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra).
The Chemistry of Alkali Earth Metals
These elements have features attributed to metals such as shiny appearance, and high reactivity, though they are relatively low density.
These group 2 elements are called alkali earth metals because they are found in the earth crust and form alkaline solution in aqueous solution
Significance of Alkali Earth Metals
The reactivity of Alkali earth metals and specific properties make them valuable in many applications. For instance, they are widely used as catalysts in chemical reactions, aiding the production of various chemicals and materials. Additionally, alkali earth metals are utilized in batteries and energy storage technologies, contributing to advancements in renewable energy.
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Occurrence and Extraction of Group 2 Elements
Beryllium (Be)
Beryllium is the lightest member of the alkali earth metals typically found in the mineral beryl (Be3Al2(SiO3)6) and bertrandite (Be4Si2O7(OH)2).
The extraction of beryllium usually involves a multi-step process. First, the beryl or bertrandite ore is crushed and heated with sulfuric acid (H2SO4)to form beryllium sulfate (BeSO4). Then, beryllium sulfate undergoes a series of chemical reactions to obtain beryllium hydroxide [Be(OH)2]. Finally, beryllium hydroxide is heated to produce pure beryllium metal.
Be3Al2(SiO3)6 + 6H2SO4 → 3BeSO4 + 2Al2(SO4)3 + 6SiO2 + 6H2O
BeSO4 + 2NaOH → Be(OH)2 + Na2SO4
Be(OH)2 → Be + 2H2O
Magnesium (Mg)
Magnesium, a more abundant element than beryllium in the Earth’s crust and the second member of the group, occurs in minerals such as magnesite (MgCO3) and dolomite (CaMg(CO3)2).
The extraction of magnesium involves mining the ore and subjecting it to various processes, including crushing, calcination, and electrolysis. In the extraction process, magnesite (MgCO3)or dolomite(CaMg(CO3)2) is heated to produce magnesium oxide (MgO), which is then converted into magnesium chloride (MgCl2) through a reaction with hydrochloric acid (HCl). Magnesium chloride is further purified and electrolyzed to obtain pure magnesium metal.
MgCO3 → MgO + CO2
CaMg(CO3)2 → CaCO3 + MgO + CO2
MgO + 2HCl → MgCl2 + H2O
MgCl2(l) → Mg(l) + Cl2(g)
Calcium (Ca)
Calcium is a relatively abundant element in the Earth’s crust. It is widely present in minerals like limestone (CaCO3) and gypsum (CaSO4·2H2O).
The extraction of calcium involves mining the limestone or gypsum deposits and subjecting them to crushing and heating processes. In the extraction of calcium from limestone, the limestone is heated to high temperatures in a kiln, leading to the formation of calcium oxide (CaO), also known as quicklime. Quicklime is then reacted with water to produce calcium hydroxide [Ca(OH)2], which can be further processed to obtain pure calcium metal. These processes such as electrolysis or the aluminothermic process involve the reduction of calcium compounds to yield calcium metal.
CaCO3 → CaO + CO2
CaO + H2O → Ca(OH)2
Strontium (Sr)
Strontium occurs in minerals such as celestite (SrSO4) and strontianite (SrCO3). The extraction process for strontium involves mining the ores and converting them into soluble forms. Celestite is typically converted to strontium sulfate through chemical processes. The strontium sulfate is then purified and undergoes electrolysis to obtain pure strontium metal.
Barium (Ba)
Barium is found primarily in the mineral barite (BaSO4). The extraction process for barium involves mining the barite ore and converting it into a soluble form, usually barium chloride (BaCl2). Barium chloride is further purified and electrolyzed to obtain pure barium metal.
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Properties of Alkali Earth Metals
Atomic Number and Symbol
Each alkali earth metal has a unique atomic number and symbol. Beryllium has an atomic number of 4 and is represented by the symbol Be. Magnesium has an atomic number of 12 and is represented by the symbol Mg. Calcium has an atomic number of 20 and is represented by the symbol Ca. Strontium has an atomic number of 38 and is represented by the symbol Sr. Barium has an atomic number of 56 and is represented by the symbol Ba. Radium has an atomic number of 88 and is represented by the symbol Ra. All these elements are located in Group 2 which is situated on the left side of the periodic table, following the alkali metals in Group 1.
Electronic configuration
The electronic configuration of alkali earth metals is characterized by the filling of the s orbital in their outermost energy level. For example, beryllium has an electronic configuration of 1s² 2s², magnesium has 1s² 2s² 2p⁶ 3s², calcium has 1s² 2s² 2p⁶ 3s² 3p⁶ 4s², strontium has 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s², barium has 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s², and radium has 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s s² 4f¹⁴ 5d¹⁰ 6p⁶ 7s².
Appearance and State at Room Temperature
Alkali earth metals generally have a silvery-white lustrous appearance. Beryllium and magnesium have a shiny surface, while calcium, strontium, barium, and radium exhibit a more reactive nature and may tarnish when exposed to air. At room temperature, beryllium and magnesium are solid metals, while calcium, strontium, barium, and radium are soft and malleable.
Density, Melting Point, and Boiling Point
The alkali earth metals exhibit distinctive properties such as density, melting point, and boiling point. Compared to transition metals, alkali earth metals generally have lower densities. As we progress down the group, the densities of the alkali earth metals increase sequentially.
Meanwhile, the melting and boiling points of alkali earth metals are generally high due to the strong metallic bonding between atoms. However, as we move down the group, a consistent trend is observed where the melting and boiling points decrease.
The trend of decreasing melting and boiling points as we move down the group can be attributed to the increasing atomic size and the weakening of metallic bonds. As the atoms become larger, the forces of attraction between the positively charged atomic nuclei and the delocalized electrons diminish, resulting in easier disruption of the metallic lattice structure and lower melting and boiling points.
Properties of Alkali Earth Metals
| Element | Density (g/cm³) | Melting Point (°C) | Boiling Point (°C) | Appearance | Electronic Configuration |
|---|---|---|---|---|---|
| Beryllium | 1.85 | 1287 | 2471 | Grayish white | 1s2 2s2 |
| Magnesium | 1.74 | 650 | 1090 | Silvery white | 1s2 2s2 2p6 3s2 |
| Calcium | 1.55 | 842 | 1484 | Silvery white | 1s2 2s2 2p6 3s2 3p6 4s2 |
| Strontium | 2.63 | 777 | 1387 | Silvery white | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 |
| Barium | 3.62 | 727 | 1804 | Silvery white | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p 6 5s2 4d10 5p6 6s2 |
| Radium | 5.5 | 700 | 1737 | Shiny white | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 |
Chemical Properties of Alkali Earth Metals
Reactivity of Alkaline Earth Metals with Water and Air
Alkali earth metals exhibit varying degrees of reactivity with water and air. Beryllium, being the least reactive, does not react with water or air under normal conditions. It forms a protective oxide layer on its surface that prevents further reactions.
2Be + O₂ → 2BeO
Magnesium as a metal reacts with water though slowly, producing magnesium hydroxide and hydrogen gas.
Mg + 2H₂O → Mg(OH)₂ + H₂
However, magnesium readily reacts with oxygen in the air to form a thin layer of magnesium oxide on its surface, protecting it from further oxidation.
2Mg + O₂ → 2MgO
Meanwhile, calcium, strontium, and barium are more reactive with water, vigorously producing metal hydroxides and hydrogen gas.
Calcium: Ca + 2H₂O → Ca(OH)₂ + H₂
Strontium: Sr + 2H₂O → Sr(OH)₂ + H₂
Barium: Ba + 2H₂O → Ba(OH)₂ + H₂
These metals also react with oxygen in the air, forming oxides or hydroxides depending on the conditions.
Calcium: 2Ca + O₂ → 2CaO
Strontium: 2Sr + O₂ → 2SrO
Barium: 2Ba + O₂ → 2BaO
Generally, the reactivity of alkali earth metals increases as you move down the group, with beryllium being the least reactive and barium being the most reactive. The varying reactivity of alkali earth metals with water and air can be attributed to the differences in their electronic configurations and atomic structures.
Reactivity of Alkaline Earth Metals with Halogens
The chemical properties of alkali earth metals include their reactivity with halogens. Alkali earth metals, which belong to Group 2 of the periodic table, react with halogens, which are elements in Group 17.
Beryllium (Be) exhibits limited reactivity with halogens due to its small size and high ionization energy. It forms primarily the covalent compound beryllium halide, such as beryllium fluoride (BeF₂):
Be + F₂ → BeF₂
Magnesium (Mg), Calcium (Ca), Strontium (Sr) & Barium (Ba), any of these metals readily reacts with halogens to form ionic compounds known as alkali earth halides.
M + X₂ → MX₂ where M = (Mg, Ca, Sr or Ba), X = (F, Cl, Br or I)
The reactivity of alkali earth metals with halogens increases as you move down the group. This is due to the decreasing ionization energy and increasing atomic size of the alkali earth metals. The resulting ionic compounds formed with halogens have a 1:2 stoichiometry, indicating the transfer of two electrons from the alkali earth metal to the halogen atom.
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The Reaction Mechanism of Alkaline Earth Metals and Halogen
The reaction between alkali earth metals and halogens involves the ionization of metal atoms and the formation of metal ions. The alkali earth metal atom loses electrons to attain a stable electronic configuration, while the halogen atom gains electrons to achieve a noble gas configuration. This electron transfer leads to the formation of ionic compounds, where the alkali earth metal ion is positively charged, and the halogen ion is negatively charged.
The Reactivity of Alkali Earth Metals with Acids
Alkali earth metals exhibit different degrees of reactivity with acids. Beryllium, due to its strong passivation layer, does not react with most acids. However, it can react with concentrated sulfuric acid under certain conditions, forming beryllium sulfate and sulfur dioxide:
Be + H₂SO₄ → BeSO₄ + SO₂ + H₂O
Other alkali earth metals, such as magnesium, calcium, strontium, and barium, react readily with dilute acids to form their respective salts and hydrogen gas. The general equation for the reaction between an alkali earth metal (M) and a dilute acid (HX) to release hydrogen gas.
M + 2HX → MX₂ + H₂↑
For example, the reaction between magnesium and hydrochloric acid is:
Mg + 2HCl → MgCl₂ + H₂↑
Reactivity of Group 2 Elements with Non-Metals
Alkali earth metals also react with some other non-metals, such as sulfur and phosphorus, to form binary compounds. The reactions involve the transfer of electrons from the alkali earth metal to the non-metal. For instance, magnesium reacts with sulfur to form magnesium sulfide:
3Mg + S → Mg₃S₂
Group 2 Elements as Reducing Agents
Alkali earth metals are strong reducing agents due to their low ionization energies. They can reduce various metal ions to their elemental form. For example, calcium can reduce copper(II) oxide to copper:
Ca + CuO → CaO + Cu
Reactivity of Group Elements with Organic Compounds
Alkali earth metals engage in reactions with organic substances, particularly in the presence of heat or catalysts. A prominent example of such reactivity is the Grignard reaction, which involves the formation of alkali earth metal organometallic compounds through the reaction of alkali earth metals with organic halides. These organometallic compounds have significant applications in organic synthesis.
The Grignard reaction typically follows the general equation:
RMgX + R’X’ → RR’ + MgX’₂
Here, RMgX represents the alkali earth metal organometallic compound (Grignard reagent) formed by reacting the alkali earth metal (M) with the organic halide (R’X’). The resulting product is the organic compound (RR’) and magnesium halide (MgX’₂).
For example, when magnesium (Mg) reacts with an alkyl halide, such as ethyl bromide (C₂H₅Br), a Grignard reagent, ethylmagnesium bromide (C₂H₅MgBr), is formed.
Mg + C₂H₅Br → C₂H₅MgBr
The ethylmagnesium bromide formed in this reaction can further react with various carbonyl compounds, such as aldehydes or ketones, to yield alcohols through nucleophilic addition reactions.
In organic synthesis, Grignard reagents are widely utilized to introduce alkyl, aryl, or vinyl groups into organic molecules. The reactivity of alkali earth metals in these reactions arises from their low ionization energies, which facilitate the formation of the organometallic compounds. The resulting compounds can undergo further reactions to create more complex organic structures.
| Alkali Earth Metal | Non-Metals | Acids | Reducing Agent | Organometallics |
|---|---|---|---|---|
| Beryllium (Be) | Forms covalent compounds | Limited reactivity | Not a strong reducing agent | Can form beryllium organometallic compounds |
| Magnesium (Mg) | Forms ionic compounds | Reacts with dilute acids (e.g., 2Mg + 2HCl → 2MgCl2 + H2) | Moderate reducing agent | Participates in Grignard reactions (e.g., 2Mg + 2RX → R2Mg + 2LiX) |
| Calcium (Ca) | Forms ionic compounds | Reacts with dilute acids (e.g., Ca + 2HCl → CaCl2 + H2) | Moderate reducing agent | Can engage in various organic reactions |
| Strontium (Sr) | Forms ionic compounds | Reacts with dilute acids (e.g., Sr + 2HCl → SrCl2 + H2) | Moderate reducing agent | Participates in organic synthesis |
| Barium (Ba) | Forms ionic compounds | Reacts with dilute acids (e.g., Ba + 2HCl → BaCl2 + H2) | Moderate reducing agent | Utilized in various organic transformations |
Anomalous Behaviour of Alkali Earth Metals
The anomalous behavior of beryllium, in comparison to other alkali earth metals, is a noteworthy characteristic. While alkali earth metals generally display similar trends in their properties and reactivity, beryllium stands out due to its unique features.
One significant anomaly is beryllium’s relatively high ionization energy compared to other alkali earth metals. This attribute is because of its small atomic size and high effective nuclear charge. This higher energy barrier makes it more difficult to remove electrons from beryllium atoms, resulting in a greater stability of its valence electrons.
Moreover, beryllium exhibits covalent bonding characteristics, unlike the predominantly ionic bonding observed in other alkali earth metals. Beryllium’s small atomic size and high electronegativity contribute to its ability to form covalent compounds. These compounds often exhibit directional bonding, allowing for the formation of stable covalent networks.
Beryllium shares similarities in its chemistry with aluminum, a group 13 element. Both elements have comparable electronegativities and tend to form covalent compounds. Beryllium oxide (BeO) and aluminum oxide (Al₂O₃), for example, exhibit similar structures and bonding characteristics. This is called a diagonal relationship between the two elements.
Additionally, beryllium’s toxicity and its ability to form strong complexes with ligands are notable aspects of its chemistry, which have practical implications in various fields.
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Compounds of Alkali Earth Metals
The compounds of alkali earth metals are categorized and their properties and reactions are mentioned below:
Oxides of Alkaline Earth Metals
Alkaline Earth metals combine with oxygen to form this basic oxide
These compounds such as Magnesium oxide (MgO), Calcium oxide (CaO) are White powders or crystals with high melting points, insoluble in water
These elements react with water to form metal hydroxides:
MgO + H2O → Mg(OH)2
Alkali earth metals react with acids to produce salts and water:
CaO + 2HCl → CaCl2 + H2O
These compounds have basic properties, capable of neutralizing acids.
Meanwhile, these oxide is used as refractory materials and in the production of glass.
Hydroxides of Alkaline Earth Metals
These compounds such as Barium hydroxide (Ba(OH)2), Strontium hydroxide (Sr(OH)2) are white crystalline solids
These hydroxides are sparingly soluble in water.
They are strong bases, capable of neutralizing acids,
Ba(OH)2 + 2HCl → BaCl2 + 2H2O
These hydroxides react with carbon dioxide to form carbonates
Mg(OH)2 + CO2 à MgCO3 + H2O
These hydroxide which are also called alkali is used in agriculture and water treatment. Specifically, calcium hydroxide, also known as slaked lime, is commonly used to adjust soil pH and in construction materials, such as mortars and plasters.
Carbonates of Group 2 Elements
The carbonates of group 2 elements are white powders or crystals, Insoluble in water and react with acids to produce carbon dioxide gas, water, and a salt.
For example: CaCO3 + 2HCl → CaCl2 + H2O + CO2
From the above reaction, effervescence is observed to explain the release of CO2
These compounds exist in nature as minerals like limestone and marble
Sulfates of Group 2 Elements
CaSO4, MgSO4 of permanent hard water with Barium sulfate (BaSO4), Strontium sulfate (SrSO4) which are White crystalline solids are examples of sulfates of alkaline earth metals.
These compounds react with carbonates to form insoluble precipitates,
e.g., BaSO4 + Na2CO3 → BaCO3 + Na2SO4
They are used in qualitative analysis for the detection of alkali earth metals. For example, calcium sulfate is commonly used in the production of gypsum products, including plaster of Paris and wallboard. Strontium sulfate, due to its low solubility and ability to absorb X-rays, is utilized in medical imaging as a contrast agent.
Halides (Chlorides, Bromides, and Iodides) of Group 2 Elements
These halides such as Calcium chloride (CaCl2), Barium iodide (BaI2) which are colorless or white crystalline solids are soluble in water.
The halides of group 2 elements react with silver nitrate to form insoluble precipitates:
CaCl2 + 2AgNO3 → Ca(NO3)2 + 2AgCl
These compounds are used in various applications, including industrial processes and laboratory reagents
Solubility Rules and Precipitation Reactions
When it comes to alkali earth metal compounds, solubility rules provide valuable insights into their solubility behavior. Compared to alkali metal compounds, compounds of alkali earth metals are generally less soluble in water.
One significant example is the solubility of alkali earth metal carbonates. Carbonates of alkali earth metals, such as calcium carbonate (CaCO3) and magnesium carbonate (MgCO3), exhibit lower solubilities in water compared to alkali metal carbonates. For instance, sodium carbonate (Na2CO3) and potassium carbonate (K2CO3) are highly soluble in water and readily dissolve to form aqueous solutions. However, calcium carbonate and magnesium carbonate have relatively lower solubilities and are known to precipitate out of solution under certain conditions.
Also, the solubility of alkali earth metal phosphates such as calcium phosphate (Ca3(PO4)2) and magnesium phosphate (Mg3(PO4)2) islower solubilities compared to alkali metal phosphates. This property is often utilized in analytical chemistry for the selective precipitation of alkali earth metal ions from a solution containing various metal ions. By adding a suitable reagent, such as ammonium phosphate ((NH4)3PO4), to the solution, the less soluble alkali earth metal phosphates can be selectively precipitated, allowing for their separation and identification.
Similarly, sulfates of alkali earth metals also exhibit lower solubilities. For example, barium sulfate (BaSO4) and strontium sulfate (SrSO4) have limited solubilities in aqueous solution.
Soluble and Insoluble Salts of Alkali Earth Elements
| Element | Compound | Solubility in Water |
|---|---|---|
| Beryllium (Be) | Carbonates (e.g., beryllium carbonate) | Insoluble |
| Phosphates (e.g., beryllium phosphate) | Insoluble | |
| Sulfates (e.g., beryllium sulfate) | Insoluble | |
| Hydroxides (e.g., beryllium hydroxide) | Insoluble | |
| Magnesium (Mg) | Carbonates (e.g., magnesium carbonate) | Insoluble |
| Phosphates (e.g., magnesium phosphate) | Insoluble | |
| Sulfates (e.g., magnesium sulfate) | Soluble | |
| Hydroxides (e.g., magnesium hydroxide) | Insoluble | |
| Calcium (Ca) | Carbonates (e.g., calcium carbonate) | Insoluble |
| Phosphates (e.g., calcium phosphate) | Insoluble | |
| Sulfates (e.g., calcium sulfate) | Insoluble | |
| Hydroxides (e.g., calcium hydroxide) | Insoluble | |
| Strontium (Sr) | Chlorides (e.g., strontium chloride) | Soluble |
| Barium (Ba) | Sulfates (e.g., barium sulfate) | Insoluble |
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