Elements In Group 2 Are Called

Elements in Group 2 of the periodic table are called alkaline earth metals. These elements share similar chemical properties and exhibit a gradual change in characteristics as you move down the group. This article will comprehensively explore the alkaline earth metals, covering their properties, occurrence, reactions, uses, and more.

Introduction to Alkaline Earth Metals

Alkaline earth metals occupy Group 2 (IIA) of the periodic table. This group comprises six elements:

• Beryllium (Be)
• Magnesium (Mg)
• Calcium (Ca)
• Strontium (Sr)
• Barium (Ba)
• Radium (Ra)

Each element has two electrons in its outermost shell, which they readily lose to form doubly charged positive ions (cations) with a +2 charge. This characteristic is fundamental to their chemical behavior and the formation of various compounds.

The term "alkaline earth" comes from the fact that their oxides (also known as earths) form alkaline solutions when dissolved in water. Historically, these oxides were difficult to melt, hence the term "earth."

General Properties of Alkaline Earth Metals

Alkaline earth metals exhibit a range of physical and chemical properties that distinguish them from other groups in the periodic table.

Physical Properties

1. Appearance: Alkaline earth metals are silvery-white, lustrous metals.

2. Melting and Boiling Points: The melting and boiling points are relatively high compared to alkali metals but generally lower than those of transition metals. However, there isn't a consistent trend down the group.

3. Density: Densities are higher than those of alkali metals. Density generally increases down the group, but magnesium is an exception, being less dense than beryllium.

4. Hardness: They are harder than alkali metals but still relatively soft and can be cut with a knife.

5. Electrical Conductivity: Alkaline earth metals are good conductors of electricity due to the presence of mobile electrons.

6. Atomic and Ionic Radii: Atomic and ionic radii increase down the group due to the addition of electron shells.

7. Ionization Energy: The first ionization energy is high, but the second ionization energy is even higher, which explains why they form +2 ions. Ionization energy decreases down the group.

8. Flame Color: When heated in a flame, alkaline earth metals impart characteristic colors:

   • Calcium: Orange-red
   • Strontium: Crimson-red
   • Barium: Apple-green

   Beryllium and magnesium do not impart color to the flame because their ionization energies are too high.

Chemical Properties

1. Reactivity: Alkaline earth metals are reactive, but less so than alkali metals. Reactivity increases down the group as the ionization energy decreases.

2. Oxidation State: They primarily exhibit a +2 oxidation state in their compounds, reflecting their tendency to lose two electrons to achieve a stable electron configuration.

3. Reaction with Water: They react with water to form metal hydroxides and hydrogen gas. The reactivity increases down the group.

   • Magnesium reacts slowly with cold water and more rapidly with hot water or steam.
   • Calcium, strontium, and barium react readily with cold water.

   The general equation for the reaction is:

   M(s) + 2H₂O(l) → M(OH)₂(aq) + H₂(g)

4. Reaction with Air: They tarnish in air due to the formation of a layer of oxide or nitride on the surface.

   • Beryllium is resistant to oxidation due to the formation of a protective oxide layer.
   • Magnesium reacts slowly with oxygen at room temperature but burns vigorously when heated.
   • Calcium, strontium, and barium react readily with oxygen and nitrogen at room temperature.

5. Reaction with Acids: They react with dilute acids to form salts and hydrogen gas.

   M(s) + 2HCl(aq) → MCl₂(aq) + H₂(g)

6. Reaction with Halogens: They react directly with halogens to form metal halides.

   M(s) + X₂(g) → MX₂(s) (where X is a halogen)

7. Reducing Agents: Alkaline earth metals are strong reducing agents due to their low ionization energies.

Occurrence and Extraction

Alkaline earth metals are found in various minerals and compounds in the Earth's crust.

Occurrence

1. Beryllium: Found in minerals such as beryl (Be₃Al₂Si₆O₁₈), bertrandite (Be₄Si₂O₇(OH)₂), and chrysoberyl (BeAl₂O₄).
2. Magnesium: Abundant in the Earth's crust and seawater. Found in minerals such as magnesite (MgCO₃), dolomite (CaMg(CO₃)₂), and carnallite (KCl·MgCl₂·6H₂O).
3. Calcium: A major component of many rocks and minerals, including limestone (CaCO₃), gypsum (CaSO₄·2H₂O), and fluorite (CaF₂).
4. Strontium: Found in minerals such as celestite (SrSO₄) and strontianite (SrCO₃).
5. Barium: Occurs in minerals such as barite (BaSO₄) and witherite (BaCO₃).
6. Radium: A radioactive element found in trace amounts in uranium ores, such as pitchblende.

Extraction

1. Beryllium: Extracted from beryl by converting it to beryllium fluoride, which is then reduced with magnesium.
2. Magnesium: Extracted from seawater by precipitating magnesium hydroxide with lime (calcium hydroxide). The magnesium hydroxide is then converted to magnesium chloride, which is electrolyzed to produce magnesium metal. It can also be obtained by the reduction of magnesium oxide with ferrosilicon.
3. Calcium: Obtained by the electrolysis of molten calcium chloride.
4. Strontium: Prepared by the electrolysis of molten strontium chloride or by reducing strontium oxide with aluminum.
5. Barium: Produced by the electrolysis of molten barium chloride or by reducing barium oxide with aluminum.
6. Radium: Isolated from uranium ores through a complex process involving fractional crystallization and ion exchange.

Individual Alkaline Earth Metals: Properties and Uses

Beryllium (Be)

• Properties: Beryllium is a hard, lightweight, and brittle metal with high tensile strength. It has a high melting point and excellent thermal conductivity. It is also resistant to corrosion due to the formation of a protective oxide layer.
• Uses:
  • Alloys: Used in alloys with copper and nickel to increase strength, hardness, and corrosion resistance.
  • Nuclear Reactors: Used as a moderator and reflector in nuclear reactors due to its low neutron absorption cross-section.
  • X-ray Windows: Used in X-ray tubes due to its ability to transmit X-rays.
  • Aircraft and Missiles: Used in structural components of aircraft and missiles due to its lightweight and high strength.

Magnesium (Mg)

• Properties: Magnesium is a lightweight, strong, and ductile metal. It has a relatively low density and is a good conductor of heat and electricity. It burns in air with a brilliant white light.
• Uses:
  • Alloys: Used in alloys with aluminum and other metals to produce lightweight and strong materials for aerospace, automotive, and other applications.
  • Structural Material: Used in the construction of aircraft, missiles, and other vehicles.
  • Reducing Agent: Used as a reducing agent in the production of other metals, such as titanium.
  • Grignard Reagents: Used in organic synthesis as Grignard reagents.
  • Medicine: Magnesium compounds are used as antacids and laxatives.
  • Photography: Magnesium was historically used in flash photography.

Calcium (Ca)

• Properties: Calcium is a soft, silvery-white metal that tarnishes rapidly in air. It is an essential element for living organisms.
• Uses:
  • Construction: Calcium carbonate (limestone) is used in the production of cement, concrete, and other building materials.
  • Metallurgy: Used as a reducing agent in the extraction of other metals.
  • Desiccant: Used as a desiccant to remove moisture from organic liquids.
  • Nutrition: Calcium is essential for bone and teeth development and function.
  • Steelmaking: Used to remove oxygen, sulfur, and phosphorus in steelmaking.

Strontium (Sr)

• Properties: Strontium is a soft, silvery-white metal that is more reactive than calcium. It reacts readily with air and water.
• Uses:
  • Fireworks: Strontium salts, such as strontium carbonate, are used to produce red colors in fireworks and flares.
  • Television Tubes: Used in the past in the glass of cathode-ray tubes (CRTs) to block X-ray emissions.
  • Radioactive Isotope: Strontium-90 is a radioactive isotope used in thermoelectric generators.

Barium (Ba)

• Properties: Barium is a soft, silvery-white metal that is very reactive. It reacts readily with air and water.
• Uses:
  • Medical Imaging: Barium sulfate is used as a contrast agent in X-ray imaging of the digestive system.
  • Drilling Mud: Barium sulfate is used as a weighting agent in drilling mud for oil and gas wells.
  • Fireworks: Barium salts, such as barium chloride, are used to produce green colors in fireworks.
  • Pesticides: Barium carbonate is used as a rodenticide.

Radium (Ra)

• Properties: Radium is a radioactive, silvery-white metal that glows in the dark due to its radioactivity. It is highly toxic.
• Uses:
  • Historical Medical Uses: Radium was formerly used in cancer treatment, but this practice has been largely replaced by safer alternatives.
  • Luminous Paints: Radium was used in luminous paints for watch dials and instrument panels, but this practice has been discontinued due to health concerns.
  • Neutron Source: Radium is used as a neutron source in some scientific instruments.

Trends in Properties Down the Group

Several trends are observed in the properties of alkaline earth metals as you move down the group from beryllium to radium.

1. Atomic and Ionic Radii: Increase due to the addition of electron shells.
2. Ionization Energy: Decreases because the outermost electrons are farther from the nucleus and are easier to remove.
3. Electronegativity: Decreases as the attraction between the nucleus and the valence electrons decreases.
4. Reactivity: Increases due to the decrease in ionization energy, making it easier for the metals to lose electrons and form compounds.
5. Melting and Boiling Points: Generally decrease, but there are exceptions.
6. Density: Generally increases, but magnesium is an exception.
7. Solubility of Hydroxides: Increases due to the increasing ionic character of the metal-hydroxide bond.
8. Thermal Stability of Carbonates: Decreases due to the increasing size of the cation, which weakens the lattice energy of the carbonate.

Compounds of Alkaline Earth Metals

Alkaline earth metals form a variety of compounds with different properties and uses.

1. Oxides: Alkaline earth metals react with oxygen to form oxides (MO). These oxides are basic and react with water to form hydroxides.
2. Hydroxides: Alkaline earth metal hydroxides (M(OH)₂) are strong bases. The solubility of hydroxides increases down the group.
3. Halides: Alkaline earth metals react with halogens to form halides (MX₂). These halides are generally soluble in water.
4. Carbonates: Alkaline earth metal carbonates (MCO₃) are insoluble in water. The thermal stability of carbonates decreases down the group.
5. Sulfates: Alkaline earth metal sulfates (MSO₄) have varying solubilities in water. Barium sulfate is highly insoluble and is used as a contrast agent in medical imaging.
6. Nitrates: Alkaline earth metal nitrates (M(NO₃)₂) are soluble in water.

Biological Importance

Several alkaline earth metals play crucial roles in biological systems.

1. Magnesium: Essential for plant and animal life. It is a component of chlorophyll in plants and is involved in many enzymatic reactions in animals. Magnesium ions are also important for muscle and nerve function.
2. Calcium: Essential for bone and teeth development and function. Calcium ions are also involved in blood clotting, muscle contraction, and nerve transmission.
3. Strontium: While not considered an essential element, strontium is found in bones and can be used to treat osteoporosis.
4. Barium: Not considered to have any biological role and is toxic.
5. Beryllium and Radium: Toxic and have no known biological role.

Safety Considerations

Some alkaline earth metals and their compounds pose health and safety risks.

1. Beryllium: Beryllium and its compounds are toxic and can cause berylliosis, a chronic lung disease.
2. Magnesium: Magnesium dust is flammable and can cause explosions.
3. Calcium: Calcium compounds are generally safe, but excessive intake can lead to hypercalcemia.
4. Strontium: Radioactive isotopes of strontium are hazardous and can cause bone cancer.
5. Barium: Barium compounds, except for barium sulfate, are toxic.
6. Radium: Radium is highly radioactive and can cause cancer and other health problems.

Frequently Asked Questions (FAQ)

1. Why are Group 2 elements called alkaline earth metals?

   • They are called alkaline earth metals because their oxides form alkaline (basic) solutions when dissolved in water, and these oxides were historically difficult to melt (hence "earth").

2. What are the main uses of alkaline earth metals?

   • Their uses vary: Beryllium is used in alloys, magnesium in lightweight structures, calcium in construction, strontium in fireworks, barium in medical imaging, and radium historically in cancer treatment.

3. How does reactivity change down Group 2?

   • Reactivity increases down the group because the ionization energy decreases, making it easier for the metals to lose electrons.

4. Are alkaline earth metals harmful?

   • Some, like beryllium and radium, are highly toxic. Others, like magnesium and calcium, are essential for life in appropriate amounts.

5. What is the oxidation state of alkaline earth metals in their compounds?

   • They typically exhibit a +2 oxidation state, reflecting their tendency to lose two electrons to achieve a stable electron configuration.

6. Why are calcium and magnesium important for health?

   • Calcium is vital for bone and teeth health, nerve transmission, and muscle function. Magnesium is essential for enzymatic reactions, muscle function, and nerve function.

Conclusion

Alkaline earth metals are a fascinating group of elements with distinct properties and diverse applications. From the lightweight alloys of beryllium and magnesium to the essential biological roles of calcium and magnesium, these elements play a significant role in our world. Their consistent trend of increasing reactivity down the group, coupled with their unique chemical behaviors, makes them a fundamental topic in chemistry. Understanding their properties, occurrence, and uses provides valuable insights into the broader landscape of the periodic table and the chemical elements that shape our environment.