GeoWiki.png
ChemWiki: The Dynamic Chemistry E-textbook > Inorganic Chemistry > Descriptive Chemistry > Main Group Elements > Group 2 Elements: The Alkaline Earth Metals

Group 2 Elements: The Alkaline Earth Metals

The elements in the group include Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), and Radium (Ra).

Introduction

Group 2 contains soft metals that are more metallic in character compared to Group 1 and are silver in color. Although characteristics are maintained homologous throughout the group, the heavier metals such as Ca, Sr, Ba, and Ra are almost as reactive as the Group 1 Alkali Metals. All of the elements in Group 2 have two electrons in their valence shell giving them an oxidation state of +2. This enables the metals to easily lose electrons which increases their stability and allow them to form compounds via ionic bonds. The following diagram shows the location of these metals in the Periodic Table.

periodic table.jpg

This tables below gives a detailed account of the descriptive chemistry of each of the individual elements. Notice how there is an increase down the table for atomic number, mass and atomic radius and decrease going down for ionization energy. These are common periodic trends that are consistent across the whole table. 

 

         Atomic # Mass(g) Oxidation State(s) Electron Configuration Atomic Radius(pm) Ionization Energy(kJ/mol) Melting Point Boiling Point Flame Color Magnetic Order Crystal Structure
Be 4 9.012 +2 1s22s2 105 899.5

1560K

2742K None Diamagnetic Hexagonal
Mg 12 24.31 +1, +2 [Ne]3s2 150 737.7 923K 1363K Bright White Paramagnetic Hexagonal
Ca 20 40.08 +2 [Ar]4s2 180 589.8 1115K 1757K Orange/ Red Diamagnetic Face Centered Cubic
Sr 38 87.62 +2 [Kr]5s2 200 549.5 1655K 1655K Scarlet Paramagnetic Face Centered Cubic
Ba 56 137.3 +2 [Xe]6s2 215 502.9 1000K 2170K Green Paramagnetic Body Centered Cubic
Ra 88 226.0 +2 [Rn]7s2 215 509.3

973K

2010K ~ Non-magnetic Body Centered Cubic

Alkaline Earth Metal Reactions

The Alkaline Earth Metals are different in their reactions compared to Group I metals. Ra is radioactive and is not considered in these reactions.  

  • Hydrogen: All of the Alkaline Earth Metals react with hydrogen to create metallic hydrides. Here is an example of a reaction:

\[ Ca(s) + H_2(g)  \rightarrow CaH_2(s) \]

  • Oxygen: The alkaline earth metals react with oxygen to produce metal oxides. An oxide is just a compound involving oxygen. This is another example reaction of an alkaline earth metal with oxygen. Be does not react with oxygen but otherwise all reactions are similar.

\[ Sr(s) + O_2 (g) \rightarrow SrO_2(s) \]

  • Nitrogen: These reactions can not occur without extreme circumstances. For example, a compound may be created via really high temperatures. If they do react however, the equation would look similar to this:

 \[ 3Mg(s) + N_2(g) \rightarrow Mg_3N_2 (s) \]

  • Halogens: When reacting with halogens, these metals create metal halides. Halide refers to a compound that is one part halogen. These reactions would look similar to:

\[  Mg(s) + Cl_2(g) \rightarrow MgCl_2(s) \]

  • Water: Be is shown to not react with water, however Mg, Ca, Sr, and Ba do react to form hydroxides, or bases. These reactions can be generalized with an example:

\[ Ba(s) + 2H_2O (l) \rightarrow Ba(OH)_2(aq) + H_2(g) \]

Properties of Individual Alkaline Earth Metals

Beryllium (Be)

  • Atomic Number = 4         Mass = 9.012g/mol     Electron Configuration =1s22s2            Density= 1.85gcm-3

Be.jpg

Beryllium was first identified in 1798 by Louis-Nicolas Vauquelin, after performing a chemical analysis on aluminium silicates. The element was originally named Glucinum and it was first isolated in 1828 by Antoine Bussy and Friedrich Wohler. In 1898, Paul Lebeau was able to produce the first pure samples of Beryllium by electrolyzing molten beryllium fluoride and sodium fluoride. It was later renamed beryllium (beryl means "to become pale" in Greek) after the fact that its gemstone was a pale color.

Beryllium is the very first element and has the highest melting point (1560 K) in Group 2. It is very rare on Earth as well as in the universe and is not considered important for plant or animal life. In nature, it can only be found in combination with other elements and it tends to only exist in solutions with a pH below 5.5. Beryllium is extremely light with high ionization energy and is used mostly to strengthen alloys.

Beryllium has a strong attraction for oxygen at high temperatures, and thus it is extremely difficult to extract from ores. This man-made beryllium is not easily available as it could not be economically mass produced. However, since 1957, the majority of beryllium is produced by reducing BeF2 with magnesium, and so it is becoming more available.    

Image of beryllium copper tips: http://www.semiprobe.com/store/images/tips.jpg

Image of a beryllium solid: http://periodictable.com/Samples/004.1/s9s.JPG

Applications:

Because beryllium is relatively light and has a wide temperature range, it has many mechanical uses. For example it can be used in aircraft production to make the nozzle of liquid-fueled spacecrafts, and make mirrors in meteorological satellites. The famous Spitzer Space Telescope's optics are composed entirely of beryllium.

One of the most important applications of Beryllium is the production of radiation windows. As Beryllium is almost transparent to x-rays, it can be used to make windows for x-ray tubes. The minimal absorption by Beryllium greatly reduces the heating effects due to intense radiation. 

 

Beryllium Isotopes:

Beryllium is a monoisotopic element because it only has one stable isotope, 9Be. Another notable isotope is Cosmogenic 10Be, which is produced by cosmic ray spallation of oxygen and nitrogen. This isotope has a relatively long half-life of 1.51 million years, and is useful in examining soil erosion and formation, as well as the age of ice cores. 

 

Beryllium Compounds:

Beryllium forms compounds with most non-metals. The most common compound is beryllium oxide (BeO) which does not react with water and dissolves in strongly basic solutions. Because of its high melting point, BeO is a good heat conductor in electrical insulators. It is also an amphoteric oxide, meaning it can react with both strong acids and bases.

Cation:   H2O(l) + BeO(s) + 2H3O+(aq) ------> [Be(H2O)4]2+(aq) 

Anion:    H2O(l) + BeO(s) + 2OH-(aq) -------> [Be(OH)4]2-(aq)

Magnesium (Mg)

  • Atomic Number = 12         Mass = 24.31g/mol         Electrion Configuration = [Ne]3s2            Density=1.738gcm-3

Mg.jpg

Magnesium was first discovered in 1808 by Sir Humphry Davy in England by the electrolysis of magnesia and mercury oxide and later, Antoine Bussy was the first to produce it in its consistent form in 1831.

It is the 8th most abundant element in the Earth's crust, constituting 2% by mass. It is also the 11th most common element in the human body.  50% of Magnesium ions are found in bones, and it is a required catalyst for over 300 different enzymes. Magnesium has a melting point of 923K and reacts with water at room temperature, although extremely slowly. It is also highly flammable and extremely difficult to put out once ignited. As a precaution, when burning or lighting pure magnesium, U.V protected goggles should be worn, as the bright white light produced can have permanent effects on the retinas of the eyes. 

It can be found in over 100 different minerals, but most commercial magnesium is extracted from dolomite and olivine. The Mg2+ ion is extremely common in seawater and can be filtered and then electrolyzed to produce pure magnesium. 

Applications:

As a metal, magnesium is used for structural purposes to make car engines, pencil sharpeners, and many electronic devices such as laptops, and cell phones. Due to it's bright white flame color, magnesium is also often used in fireworks. 

In a biological sense, magnesium is vital to the body's health because Mg2+ ion is a component in every cell type. Magnesium can be obtained by eating food which is rich in Magnesium, such as nuts and certain vegetables or by eating supplementary diet pills. Chlorophyll, the pigment that absorbs light in plants, is based around magnesium and is necessary for photosynthesis. 

Image of bright magnesium flame: http://www.daviddarling.info/images/magnesium_burning.jpg

Magnesium Isotopes:

The three stable isotopes of magnesium are: 24Mg, 25Mg, and 26Mg. 24Mg makes up about 74% of Mg and 26Mg is associated to meteorites in the solar system.

28Mg is the only known radioactive isotope of Magnesium and it has a half-life of around 21 hours.

Magnesium Compounds:

Magnesium ions are essential for all life on Earth, and can be found mainly in seawater and the mineral carnallite. Some examples of magnesium compounds include: magnesium carbonate (MgCO3), a white powder used by athletes and gymnasts to dry their hands for a firm grip, and magnesium hydroxide (Mg(OH)2), or milk of magnesia, used as a common component of laxatives.

Image of magnesium carbonate: http://cache.daylife.com/imageserve/04N2ang2ZxgU3/610x.jpg

Image of Milk of Magnesia: http://skinverse.com/wp-content/milk-of-magnesia.JPG

Calcium (Ca)

  • Atomic Number = 20     Mass = 40.08g/mol         Electron Configuration = [Ar]4s2        Density= 1.55gcm-3

Ca.jpg

Calcium was isolated in 1808 by Sir Humphry Davy by the electrolysis of lime and mercuric oxide. In nature, it is only found in combination with other elements. It is the 5th most abundant element in the Earth's crust, and is essential for living organisms. Calcium, with the presence of Vitamin D, is well known for it's role in building stronger, denser bones early in the lives of humans and other animals. Calcium can be found in products such as milk, cheese, and other dairy products. Calcium has a melting point of 1115K and gives off a red flame when ignited. Calcium was not readily available until the early 20th Century. 

Image of a calcium flame: http://www.dkimages.com/discover/previews/890/25053807.JPG

 

Applications:

Calcium is an important component in cement and mortars, and thus is necessary for construction. It is also used to aid cheese production. 

Calcium Isotopes:

The four stable isotopes of calcium are: 40Ca, 42Ca, 43Ca, and 44Ca. The most abundant isotope, 40Ca, makes up about 97% of naturally occurring calcium. 41Ca is the only radioactive isotope of Calcium and it has a half life of 103,000 years. 

Calcium Compounds:

The most common calcium compound is calcium carbonate (CaCO3). Calcium carbonate is a component of shells in living organisms, and also as an antacid.  It is also the main component of limestone. Three steps are required to to obtain pure CaCO3 from limestone: calcination, slaking, and carbonation.

CaCO3(s) → CaO(s) + CO2(g)              (calcination)

CaO(s) + H2O(l) → Ca(OH)2(s)             (slaking)

  Ca(OH)2 + CO2(g) → CaCO3(s) + H2O(l)      (carbonation)

   Image of limestone caves: http://www.molecreekgh.com.au/kingsolomon.jpg

Another important Calcium compound is Calcium Hydroxide (Ca(OH)2). Often referred to as 'slack lime', it can be refined to form cement. 

CaO(s)+H2O(l)→Ca(OH)2

Strontium (Sr)

  • Atomic Number = 38         Mass = 87.62g/mol         Electron Configuration =  [Kr]5s2          Density=2.64gcm-3

Sr.jpg

Strontium was first discovered in 1790 by Adair Crawford in Scotland and is named after the village it was discovered in, Strontian. In nature, it is only found in combination with other elements as it is extremely reactive.  It is the 15th most abundant element on Earth and is commonly found in the form of mineral celestite. Strontium metal is a slightly softer than calcium and has a melting point of 1655K.

     Image of a strontium solid: http://www.interhomeopathy.org/images/gallery/213-strontium.jpg 

Applications:

In it's pure form, Strontium can be used to make alloys. It can also be used in fireworks as it produces a scarlet flame color. strontium ranelate (C12H6N2O8SSr2) can be used to treat sufferers of osteoporosis and strontium chloride (SrCl2) is used to make toothpaste for sensitive teeth. 

Strontium Isotopes:

Strontium has four stable isotopes: 84Sr, 86Sr, 87Sr, and 88Sr. About 82% of naturally occurring strontium comes in the form of 88Sr. 

Strontium Compounds:

Some applications of strontium compounds include strontium carbonate (SrCO3), strontium sulfate (SrSO4), and strontium nitrate (Sr(NO3)2), which can be used as a red flame in fireworks.

Image of red strontium flame:http://image48.webshots.com/48/2/43/11/376124311VrrAKu_ph.jpg

 

Barium (Ba)

  • Atomic Number = 56         Mass = 137.3g/mol         Electrion Configuration = [Xe]6s2            Density=3.51gcm-3

Ba.jpg

Barium was first discovered in 1774 by Carl Scheele, but was not isolated as a pure metal until 1808 when Sir Humphry Davy electrolyzed molten barium salts. The name Barium comes from the Greek word 'barys', which means heavy. Barium is a soft, silvery white metal, and has a melting point of 1000K. Because of its reaction to air, barium cannot be found in nature in its pure form but can be extracted from the mineral barite.

Image of barium solid: http://radio.weblogs.com/0101365/images/elements/barium.jpg

Barium Isotopes:

There are seven stable isotopes of naturally occurring barium: 130Ba, 132Ba, 134Ba, 135Ba, 136Ba, 137Ba, and 138Ba. In total, twenty-two isotopes are known to exist, but a majority of them are highly radioactive and have relatively short half-lives. 

Barium Compounds:

Barium sulfate (BaSO4), or barite, is the most common mineral that is abundant with barium. This mineral has a density of 4.5g/cm3 and is extremely insoluble in water. Uses of barium sulfate include being a radiocontrast agent for X-ray imaging of the digestive system. Barium carbonate (BaCO3) is also commonly used as a rat poison. 

Image of barium sulfate outlining an esophagus x-ray:http://www.ispub.com/xml/journals/ija/vol4n2/x_ray_7b.jpg

 

Radium (Ra)

  •  Atomic Number = 88         Mass = 226.0g/mol         Electron Configuration = [Rn]7s2       Density=5.5gcm-3

     

Ra.jpg

Radium was first discovered in 1898 by Marie Sklodowska-Curie, and her husband, Pierre Curie,  in a pitchblende uranium ore in North Bohemia in the Czech Republic, however, it was not isolated as a pure metal until 1902.

It is the heaviest and most radioactive of the alkaline earth metals and it reacts explosively with water. Radium appears pure white but when exposed to air it immediately oxidizes and turns black.

Because radium is a decay product of uranium, it can be found in trace amounts in all uranium ores. The exposure or inhalation of radium can cause great harm, such as cancer and other disorders.

Radium Isotopes

There are 25 isotopes of radium are that known to exist, but only 4 are found in nature. However, none are stable. The isotope which has the longest half life is 226Ra, which is produced by the decay of Uranium.

The four most stable isotopes are: 223Ra, 224Ra, 226Ra, and 228Ra. 223Ra, 224Ra and 226Ra decay emitting alpha particles, while 228Ra decays emitting a beta particle. Most radium isotopes have relatively short half-lives.

Radium Compounds:

Radium compounds are extremely rare in nature because of its short half-lives and intense radioactivity.  As such, radium compounds are found almost entirely in uranium and thorium ores. All known radium compounds have a crimson colored flame. 

The most important compound of radium is radium chloride (RaCl2). Previously, it had only been found in a mixture with barium chloride, but as RaCl2 appeared to be less soluble than barium chloride, the mixture could continually be treated to form a precipitate. This procedure was repeated several times until the radioactivity of the precipitate no longer increased, as radium chloride could be electrolyzed using a mercury cathode to produce pure radium. Currently, RaCl2 is still used to separate radium from barium and it is also used to produce Radon gas, which can be used to treat cancer. 

References

  • Petrucci, Ralph H, William Harwood, and F. Herring. General Chemistry: Principles and Modern Applications. 8th Ed. New Jersey: Pearson Education Inc, 2001.
  • Maguire, Michael E. "Alkaline Earth Metals." Chemistry: Foundations and Applications. The Gale Group, Inc. 2004.
  • Zumdahl, Steven S. World of Chemistry. Thomson Gale, 2006.
  • Grew, Edward S. Beryllium: mineralogy, petrology, and geochemistry. Mineralogical Society of America. 2002
  • Cowen, James A. The Biological Chemistry of Magnesium. VCH, 1995

Problems

  1. Which of the following is NOT an Alkaline Earth Metal?
    1. (a) Ba    (b) K     (c) Mg    (d) Be    (e) Ra
  2. True or False: Alkaline Earth Metals do not react vigorously with water.
  3. What alkaline metal is a main component in our bones?
  4. Which group 2 metal has the largest ionization energy?
  5. Which element has the lowest melting point?
    1. (a) Ba     (b) Ra    (c)Ca    (d) Mg

       

  6. Why is the term alkaline used to describe the Group 2 elements?
  7. Which group 2 element has a scarlet flame color?
  8. What is the oxidation state of all group 2 elements?
  9. Which group 2 element is the most radioactive?
  10. Which substance is acting as the reducing agent?

  Be(s)+F2→BeF

Answers

  1. B (Potassium)
  2. True. 
  3. Calcium (Ca)
  4. Ba
  5. Ra
  6. The name "Alkaline" is from its slight solubility to water, while "Earth" is derived from its inability to decompose when exposed to heat.
  7. Strontium
  8. +2
  9. Radium
  10. Be

 

You must to post a comment.
Last Modified
11:23, 30 Dec 2013

Page Rating

Was this article helpful?

Tags

This page has no custom tags set.
Module Vet Level:
Module Target Level:

Creative Commons License UC Davis GeoWiki by University of California, Davis is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 3.0 United States License. Permissions beyond the scope of this license may be available at copyright@ucdavis.edu. Terms of Use