Volume 3

Natural Occurrence and Physical Properties. Would Fireworks Be Possible Without Alkaline Earth Metals? • Chemical Properties of the Alkaline Earth Metals • Water Hardness • Alkaline Earth Metal Compounds Around Us • Exercises

3.2.1. Natural Occurrence and Physical Properties. The alkaline earth metals are Be (beryllium), Mg (magnesium), Ca (calcium), Sr (strontium), Ba (barium), and Ra (radium). These six elements constitute Group 2 of the periodic table. Calcium and magnesium are among the top 10 most abundant chemical elements in the Earth's crust. Barium and strontium are also abundant, but beryllium is less so. Radium is radioactive and extremely rare; the estimated abundance of Ra in the Earth's crust is on the order of one part per trillion. Selected physical properties of the alkaline earth metals are compiled in Table 2. The alkaline earth metals are harder and have higher densities and melting points than their alkali metal congeners. All of the alkaline metals are silvery-gray in color.

Table 2. Selected physical properties of the alkaline earth metals.
Three of the alkaline earth metals (calcium, strontium, and barium) are strong flame colorants. Their salts are broadly employed to make orange (Ca), red (Sr), and green (Ba) flares and fireworks. Burning magnesium metal emits dazzling white light and is also used in pyrotechnics. Without Mg, Ca, Sr, and Ba, there would have been no beautiful fireworks as we know them today (Figure 3-8).
Figure 3-8. Alkaline earth metals in fireworks. From left to right: Mg (white; source), Ca (orange; source), Sr (red; source), and Ba (green; source).

3.2.2. Chemical Properties. The alkaline earth metals have two electrons in the outermost shell, as presented below.

Be 1s2 2s2
Mg 1s2 2s2 2p6 3s2
Ca 1s2 2s2 2p6 3s2 3p6 4s2
Sr 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
Ba 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2

The two valence electrons are readily given away by an alkaline earth metal in order to attain the stable electron configuration of a noble gas. As a consequence, the oxidation state of the alkaline earth metals in their compounds is always +2.

The ability of the alkaline earth metals to donate electrons increases down the group (Be < Mg < Ca < Sr < Ba), exactly as for the alkali metals. As explained above, the weaker attraction of the more remote valence electrons to the nucleus makes the heavier (larger) atoms better electron donors. As with the alkali metals (see above), the order of reactivity of the alkaline earth metals is convincingly illustrated by their reactions with water (Figure 3-9). Watch this video to see how the vigor of the reaction with water depends on the nature of the alkaline earth metal. Beryllium reacts with H2O only at very high temperatures (>700 oC; not shown in the video). Magnesium reacts with water already at room temperature, albeit sluggishly (watch another video). Calcium reacts with H2O quite vigorously. Strontium reacts with water faster than calcium, and barium (not shown in the video) even more so. (Note that the statement that the reaction product Ca(OH)2 is insoluble in water is not entirely correct. The solubility of Ca(OH)2 in H2O at ambient temperature is moderate (1.7 g/L), but certainly not low enough to consider Ca(OH)2 an insoluble compound.)
Figure 3-9. Electron-donating ability and, consequently, reactivity of the alkaline earth metals increases down the group.

How does the electron-donating ability of the alkaline earth metals compare with that of the alkali metals? An alkaline earth metal donates its valence electrons willingly, but considerably less so compared with its alkali metal neighbor of the same period. The strongest electron donor among the conventional alkaline earth metals (Ra excluded) is barium, which is roughly as reactive as lithium, the weakest electron donor among the alkali metals.

Figure 3-10 provides an explanation for the overall lower reactivity of the alkaline earth metals relative to the alkali metals. An alkaline earth metal atom (Be in Figure 3-10) has one more electron and one more proton than the preceding alkali metal (Li in Figure 3-10). The greater electric charges and, consequently, more powerful attractive electric forces cause the stronger contraction of the electron shells, thereby making the Be atom (radius 1.12 Å; Table 2) about 33% smaller in size than the Li atom (radius 1.67 Å; Table 1). As a result, the electrons are closer to the nucleus and more tightly held in Be than in Li, which makes beryllium a weaker donor of electrons than lithium.
Figure 3-10. An alkaline earth metal is less reactive than its alkali metal neighbor of the same period.

All alkaline earth metals react with oxygen. In the previous subsection, we touched on the tricks of the reactions of the alkali metals with O2. The alkaline earth metals are oxidized by O2 in a more straightforward manner, furnishing "normal" oxides.

2 M + O2 = 2 MO
(M = Be, Mg, Ca, Sr, Ba)

There is one notable exception though. Depending on what temperature is used for the reaction of barium with oxygen, the oxidation can furnish barium oxide (BaO) or peroxide (BaO2) or a mixture thereof. When heated at 500-600 oC in air, BaO is transformed into barium peroxide, BaO2, a stable white solid (Figure 3-11). Above 800 oC, BaO2 cleanly decomposes back to BaO and O2. This uncommon property of BaO laid the foundation for the development and commercialization of the first industrial production of pure oxygen. Known as the Brin process, this method was broadly used to make O2 from air in the 19th century.
Figure 3-11. Oxidation of Ba metal to BaO and the process for making pure O2 from air using the BaO/BaO2 cycle.

Do not be misled by the presence of the Ba metal oxidation reaction at the top of Figure 3-11. Barium does react with oxygen to give BaO at as low as room temperature and even below. However, making BaO from Ba metal and oxygen is not economical. The industrial production of BaO is based on the barium mineral barite as the raw material (Figure 3-12). Barite, the naturally occurring barium sulfate (BaSO4), is reduced to barium sulfide (BaS) using coke (carbon). The treatment of the BaS thus produced with sodium carbonate, Na2CO3 (soda ash) gives poorly soluble barium carbonate (BaCO3), which precipitates out and is then separated and thermally decomposed to BaO and CO2.
Figure 3-12. Industrial method to make BaO.

You are expected to remember that CaO (quicklime) is made from limestone rock, CaCO3, a reaction we have discussed earlier. The most common and broadly used alkaline metal oxides are CaO (quicklime) and MgO (magnesia). Both react with water to give the corresponding hydroxides (Figure 3-13). While MgO reacts with water sluggishly, the reaction between CaO and H2O is highly exothermic and can be even violent (watch this demonstration). The amount of heat produced in the hydration of CaO is even enough to cook an egg!
Figure 3-13. Reactions of MgO and CaO with water.

In contrast with Ca(OH)2 that is modestly soluble in water, Mg(OH)2 is virtually insoluble. The common name of calcium hydroxide is slaked lime. Its solution in water is called limewater, and its suspension in limewater is known as milk of lime. Limewater is a classic reagent for CO2 detection. If limewater grows turbid on bubbling a gas through it, the gas contains carbon dioxide because CO2 reacts with Ca(OH)2 to give insoluble CaCO3 (Figure 3-14). You can watch a demonstration of the limewater test here.
Figure 3-14. In the presence of CO2, a solution of Ca(OH)2 (limewater) turns cloudy due to the formation of insoluble CaCO3.

The strength of alkaline earth metal hydroxides increases in the order Be(OH)2 < Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba(OH)2. The trend is the same as for alkali metal hydroxides (see above), yet not all of the alkaline earth hydroxides are strong bases. While Ba(OH)2 is a very strong base, Ca(OH)2 is a medium strength base, Mg(OH)2 is a weak base, and Be(OH)2 is even weaker. Moreover, Be(OH)2 is amphotheric, meaning that it can react both as a weak base when treated with a strong acid and as a weak acid when treated with a strong base.

3.2.3. Water Hardness. Many of us have experienced the annoying problem of hard water. Hard water can clog plumbing pipes and results in scale buildup in water heaters and on appliances. Soap does not work well in hard water. Instead of producing lather, soap in hard water makes soap scam, a grayish-white precipitate.

Water hardness is caused by the ions of two alkaline earth metals, Ca2+ and Mg2+. Soap is a sodium or potassium salt of a fatty acid such as stearic acid, C17H35CO2H. The main component of bar soap is sodium stearate, C17H35CO2Na, and that of liquid soap is potassium stearate, C17H35CO2K. Stearic acid and the way soap works will be discussed in detail in Volume 4. At this point, we do not need to know the structure of stearic acid. All we need to know for now is that while both K and Na stearates (which is soap) are soluble in water, Ca and Mg stearates are not. When soap is used with hard water, the insoluble Mg and/or Ca stearates (soap scam) are formed, precipitating out right away (watch this video). An ionic equation for the sequestering of the stearate anion by Mg2+ and/or Ca2+ is given in Figure 3-15. Again, please do not be intimidated by the stearate formula. Just view it as an anion that forms insoluble salts with Mg2+ and Ca2+.
Figure 3-15. Stearate, the active ingredient of soap, is precipitated from solution by Mg2+ and Ca2+ present in hard water.

There are two main types of water hardness, temporary and permanent. Temporary hardness of water is due to the presence of calcium and magnesium bicarbonates, M(HCO3)2 (M = Ca, Mg). These soluble bicarbonates are formed from the naturally occurring insoluble carbonates, moisture, and atmospheric CO2. Figure 3-16 displays an equation for the formation of calcium bicarbonate from limestone (CaCO3). In exactly the same manner, magnesium bicarbonate is formed from Mg minerals such as magnesite, MgCO3.
Figure 3-16. Soluble Ca(HCO3)2 is formed in nature from limestone, water, and atmospheric CO2.

The reaction in Figure 3-16 can be reversed by simply boiling the water. As the water boils, the equilibrium is shifted to the right because the CO2 bubbles off and the insoluble MgCO3 and CaCO3 precipitate out (Figure 3-17). The water turns soft as it no longer contains the Ca2+ and Mg2+ cations.
Figure 3-17. Temporary water hardness is removed by boiling the water.

Permanent water hardness is due to the presence of Mg and Ca salts of sulfuric and hydrochloric acids, MSO4 and MCl2 (M = Mg, Ca). This type of hardness cannot be eliminated by simply boiling the water. The main solutions for the permanent water hardness problem include using a water softening method or just replacing soap with a different detergent that is not affected by magnesium and calcium ions.

3.2.4. Alkaline Earth Metal Compounds Around Us. We have learned what role some of the alkaline earth metals play in fireworks and in creating the problem of hard water. Here are some other interesting facts about the alkaline earth metals.

Beryllium. If you have seen a genuine emerald, you have seen a beryllium compound. Emeralds are high-quality natural crystals of the mineral beryl. The formula of beryl is Be3Al2(SiO3)6. Pure beryl is colorless. The beautiful green color of emeralds is due to trace amounts of chromium (Cr) and sometimes vanadium (V) that some natural forms of beryl contain. Trace quantities of other metals give beryl other colors (Figure 3-18). Chances also are that you use beryllium each time you play music through your headphones or speakers. Many manufacturers use beryllium to build tweeter drivers that provide the highs that are superbly crisp and clear.
Figure 3-18. Left to right: morganite, aquamarine, and emerald are all forms of beryl (source).

Beryllium is highly toxic. Yet no one has ever experienced beryllium poisoning from wearing emerald jewelry or using speakers or headphones equipped with beryllium drivers. Beryllium poisoning can result only from exposure to soluble beryllium salts or long-term inhalation of fine particles of a beryllium compound such as BeO dust. So, rock your emeralds and enjoy music from your headphones and speakers with beryllium in them, as using these beryllium-containing products is safe. Beryllium has many other important applications that you can read about here.

Magnesium is essential for plants. Being at the core of the complex molecule of chlorophyll, magnesium is vital for photosynthesis, the process used by plants to make carbohydrates (sugars) from CO2 and water in the presence of light. The carbohydrates produced by photosynthesis are then used by plants as a source of energy and as the main construction material for building their tissues. Magnesium is critical for animals, too. Bones and teeth cannot be healthy without magnesium. Hundreds of enzymes in the human body cannot function properly without magnesium. Magnesium and its compounds are irreplaceable in the production of many important alloys, cements, flame retardants, and high-temperature resistant (refractory) materials for furnaces to make iron, steel, and glass. Certain magnesium compounds are used in the textile and pulp and paper industries. Organomagnesium compounds, broadly known as Grignard reagents, play a very important role in organic synthesis.

Calcium is indispensable for bone and teeth, making up approximately 2% of the human body by mass. There are many other vital roles that calcium plays in living organisms. We use calcium compounds on a daily basis. Each time we brush our teeth, we use a calcium compound, CaCO3, the most important abrasive ingredient of toothpaste. Each time we write with chalk on a chalkboard, we use CaCO3, too, as chalk is made of the form of limestone named calcite. Calcium oxide is manufactured from limestone (CaCO3) on a tremendous scale for use in cement, a key component of concrete. The production of iron and steel is inconceivable without CaCO3 flux (see section 3.8 below). Calcium sulfate (CaSO4) is used to make gypsum casting plaster. Calcium minerals fluorite (CaF2), and apatite (Ca5F(PO4)3) are used to make hydrogen fluoride (HF) and phosphoric acid (H3PO4). Calcium metal has a number of applications in metallurgy.

Strontium. Up until the year of 2000, strontium could be found in almost every household that had a color TV. Color TV cathode ray tubes were made of the special glass that strontium was a key component of. This once very important application has lost its significance with the massive replacement of the previous generation heavy and bulky TV sets with liquid crystal, plasma, and OLED flat-panel displays. Besides pyrotechnics, strontium is used to make toys that glow in the dark (SrAl2O4) and as an additive for toothpastes for sensitive teeth (SrCl2). Strontium compounds are nontoxic.

Barium. Apart from pyrotechnics, perhaps the most widely known application of barium is in X-ray diagnostics. As mentioned above, BaSO4 is an excellent radiocontrast agent for X-ray imaging of esophagus, stomach, and intestines. The so-called barium meal taken by the patient prior to the test is a suspension of BaSO4 in water. Although soluble barium salts are toxic, barium meals are perfectly safe to take because of the extremely low solubility of BaSO4.

Radium is a very rare element whose isotopes are all radioactive. Due to its radioactivity and high toxicity, radium has almost no modern application with just a few exceptions. One is the use of the isotope 223Ra to treat certain types of cancer. Before the 1970s, radium-based self-luminous paints were used to make compasses, gauges, watches, and alarm clocks with dials, hands, and numerals that glowed in the dark. Read the tragic story of the Radium Girls and watch this video to understand why the use of radium paints has long been discontinued. However, the often claimed danger of using a clock or a wristwatch containing radium-painted parts is greatly exaggerated, "A radium watch becomes hazardous only when someone opens one and tinkers with the dials, inhaling radioactive dust particles".

3.2.5. Exercises.

1. Magnesium metal reacts with water more vigorously than barium metal. True or false? Answer

2. Ca(OH)2 is a stronger base than Mg(OH)2. True or false? Answer

3. A flame test on a sample of an individual alkaline earth metal compound produced red color. The metal was (a) magnesium; (b) strontium; (c) barium. Answer

4. Provide a plausible explanation for the fact that MgO occurs in nature (mineral periclase), whereas CaO does not. Answer

5. Write a balanced chemical equation for the CO2 test with limewater. [Answer: Figure 3-14]

6. Of the alkaline earth metals, the toxic ones are (a) Be and Sr; (b) Be, Sr, and Ba; (c) Be, Ba, and Ra. Answer

7. Bubbling a small quantity of CO2 through limewater results in precipitation of CaCO3. However, if the bubbling is continued, the precipitate eventually disappears. Why? Answer

8. Permanent hardness of water is caused by Mg(HCO3)2 and Ca(HCO3)2 and can be reduced by simply boiling the water. True or false? Answer

9. A 100 g sample of a pure alkaline earth carbonate was treated with hydrochloric acid in excess. The volume of CO2 produced was 22.4 L. Carbonate of what metal was used for the reaction? Try to solve the problem mentally using only the periodic table. Answer

10. Write balanced chemical equations for the following reactions.

(a) MgCl2 + KOH

(b) Ca(NO3)2 + Na2CO3

(c) CaO + H2O

(d) thermal decomposition of MgCO3

(e) BaCl2 + H2SO4

(f) Sr + H2O