Alkaline Earth Metals for IIT JEE

Properties and Applications - Alkaline Earth Metals

Alkaline earth metals (Group 2) majorly includes six elements namely beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). This group usually contain smooth, silver metals with a less metallic character than the elements of Group 1. Although many features are common throughout the group, heavier metals like Ca, Sr, Ba, and Ra are just as reactive as Alkali Metals Group 1. All elements in alkaline earth metals have in their valence shells two electrons, giving them a + 2 oxidation state. This allows metals to easily lose electrons, which enables them to form compounds via ionic bonds making them more stable. 

Physical Properties of Alkali Earth Metals


Atomic radius


In alkaline earth metals, atomic radius increases as we move down the group. Beryllium has small radius as compared with other alkaline earth metals.

The radius of an atom is governed by firstly the number of layers of electrons around the nucleus and secondly by attraction the outermost electrons felt by the nucleus.

First ionization energy


The first ionization energy is the energy needed to remove the most loosely held electron from every single mole of gaseous atoms to make a single mole of gaseous ions. Note that the first energy of ionization decreases down the group. The ionization energy is governed by three factors:

  • • The charge on nucleus,

  • • The amount of screening by inner electrons,

  • • The distance between outer electrons and nucleus.

  • By increasing the number of internal electrons, the increase in nuclear charge down the group is exactly offset. The outer electrons get a net charge of +2 from the center, as mentioned earlier. However, the distance between the nucleus and the outer electrons increases down the group, making it easier to remove them; hence the ionization energy decreases.

    Electronegativity


    Electronegativity is a measurement of an atom's tendency to attract an electrons bonding pair. It is typically measured on the Pauling scale, where a 4.0 electronegativity is assigned to the most electronegative element (fluorine). Just as the metal atoms grow larger, any bonding pair is gradually pulled away from the metal nucleus, thus becoming somewhat less attracted to the metal. In simple words, the elements are less electronegative down the group. The bonds between these elements and other elements such as chlorine are becoming more ionic when moving down the group. The bonding pair is ever more attracted to the more electronegative element from the Group 2 element.

    Melting Points, Boiling Points, and Atomization Energies


  • Melting points: Apart from magnesium irregularities, each element's melting point decreases down the group.

  • • Boiling points: No trend

  • • Atomization energy: The atomization energy is the energy required by the element in its standard state to produce 1 mole of atoms in the gas phase. As with boiling points, the alkaline earth elements do not have a simple pattern for atomizing energies. 

  • Alkaline Earth Metal Reactions

  • The reactions of the alkaline earth metals are as follow:

  • Reactions with Hydrogen: Alkaline earth metals react with hydrogen to form metallic hydrides.

  • Example: Ca(s)+ H2(g)→ CaH2(s)

  • Reactions with Oxygen: Alkaline earth metals react with oxygen to form metal oxides. An oxide is a compound in a state of -2 oxidation containing oxygen.

  • Example: Sr(s)+ O2(g)→ SrO2(s)
  • Reactions with Nitrogen: Reactions with nitrogen cannot occur in ordinary condition; very high temperatures are required.

  • Example:3Mg(s)+ N(g)→ Mg3N2(s)
  • Reactions with Halogens: Reaction between alkaline earth metal with halogens leads to metal halides. A halide is a compound that contains an ionic halogen.

  • Example: Mg(s)+ Cl2(g)→ MgCl2(s)
  • Reactions with Water: Magnesium, calcium, strontium, and barium do react to form metal hydroxides and hydrogen gas.

  • Example: Ba(s)+ 2H2O(l)→ Ba(OH)2(aq)+ H2(g)

    Properties of Individual Alkaline Earth Metals


    Beryllium (Be)


    Beryllium may be the first alkaline earth metal element and has the highest melting point of any element in the group. On Earth as well as in the universe, it is very rare and is not considered important for plant or animal life. It can be found only in compounds with other elements in nature. In solutions, only pH values below 5.5 remain in elemental form. Beryllium is extremely light with high ionizing energy and is mainly used to reinforce alloys.

    Applications:

  • • It has many mechanical uses because beryllium is relatively light and has a wide temperature range.

  • • It can be used in the manufacture of aircraft in liquid-fuelled spacecrafts nozzles and meteorological satellite mirrors.

  • • The production of radiation windows is one of the most important applications of beryllium.

  • • Since beryllium is almost x - ray transparent, it can be used for x - ray tubes. Because of intense radiation, minimal absorption by Beryllium significantly reduces heating effects.

  • Beryllium Isotopes:

    Beryllium is a monoisotope— it only has one stable isotope, 9Be. Another notable isotope is cosmogenic 10Be, produced by oxygen and nitrogen spalling from cosmic rays. This isotope has a relatively long half-life of 1, 51 million years and is useful for the examination of soil erosion and formation and the age of ice cores. 

    Magnesium (Mg)


    It is the Earth's 8th most abundant element, 2 percent by mass. It is also the 11th most common element in the human body: fifty percent of magnesium ions are found in bones and for more than three hundred different enzymes, it is a necessary catalyst. Magnesium has a 923 K melting point and reacts very slowly with water at room temperature. It is also extremely flammable and once ignited it is extremely hard to extinguish. UV - protected goggles should be worn as a precaution when burning or lighting pure magnesium, as the bright white light produced can seriously damage the retina.

    Applications:

  • • Magnesium is used in its elementary form in automotive engines, pencil sharpeners, and many electronic devices such as laptops and cell phones for structural purposes. Magnesium is also frequently used in fireworks because of its bright white flame colour. 

  • • Magnesium is vital to the health of the body in a biological sense: the Mg2 + ion is a component of any type of cell. Magnesium can be obtained by eating magnesium - rich foods such as nuts and certain vegetables or by eating additional dietary pills. 

  • • Chlorophyll, the pigment which absorbs light in plants, strongly interferes with magnesium and is required for photosynthesis. 

  • Calcium (Ca)
  • • Calcium, in the presence of vitamin D, is well known for its role in early human and other animal-life building stronger, denser bones.

  • • In leafy green vegetables as well as milk, cheese and other dairy products, calcium can be found.

  • • Calcium has 1115 K melting point and when ignited, it gives a red flame. Until the early 20th century, calcium was not readily available. 

  • Applications:

    Calcium in cement and mortars is an important component and is therefore essential for construction. It is also used to support the production of cheese. 

    Calcium Isotopes:

    Calcium have four stable isotopes of calcium are 40Ca, 42Ca, 43Ca, 44Ca. 
    40Cais the most abundant isotope and composes about 97% of naturally occurring calcium. 
    41Ca is the radioactive isotope of calcium with a half-life of 103,000 years. 

    Strontium (Sr)


    It is the fifteenth most important element on Earth and is usually found in the mineral celestite form. Strontium metal is slightly weaker than calcium and has a 1042 K melting point.

    Applications:

  • • Strontium is used in alloys in its pure form. As it produces a scarlet flame colour, it can also be used in fireworks.

  • • To treat patients with osteoporosis, strontium ranelate is used to make toothpaste for sensitive teeth. 

  • Strontium Isotopes:

    It has four isotopes those are stable: 84Sr, 86Sr, 87Sr, and 88Sr.
    Majority of naturally abundant strontium comes in the form of 88Sr. 

    Radium (Ra)


    Radium is the alkaline earth metals' heaviest and most radioactive metal and it reacts explosively with water. Radium appears to be pure white, but it immediately oxidizes and becomes black when exposed to air. Since radium is a declining uranium product, it can be found in all uranium ores in trace amounts.

    Radium Isotopes

    Radium have four most stable isotopes are 223Ra, 224Ra, 226Ra, and 228Ra.
    The three most abundant: 223Ra, 224Ra, and 226Ra wane by emitting alpha particles. 228Ra declines by emitting beta. Most isotopes of radium have relatively short half - lives.


            Electron ConfigurationBoiling PointFlame ColorAtomic #Mass (g)Oxidation State(s)Atomic Radius (pm)Ionization Energy (kJ/mol)Crystal StructureMagnetic Order
    Be1s22s22742 KNone49.0122105899.5HexagonalDiamagnetic
    Mg[Ne]3s21363 KBright White1224.31+1, +2150737.7HexagonalParamagnetic
    Ca[Ar]4s21757 KOrange/ Red2040.082180589.8Face Centered CubicDiamagnetic
    Sr[Kr]5s21655 KScarlet3887.622200549.5Face Centered CubicParamagnetic
    Ba[Xe]6s22170 KGreen56137.32215502.9Body Centered CubicParamagnetic
    Ra[Rn]7s22010 K~882262215509.3Body Centered CubicNon-magnetic