Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 60
4.4.2 The Group 1 metals
Li
Na
K
Rb
Cs
Fr
Sodium and potassium are the sixth and seventh most abundant elements in the earth’s crust, 2.6% and 2.4%, respectively,
by mass. The abundances of Li, Rb and Cs are much lower (in the ppm range). Francium is a radioactive element, found in
minute traces in nature as a short-lived daughter isotope of the 235U decay series (it converts to the much-longer lived 223Ra
isotope). All the Group 1 elements, are highly reactive with oxygen, water, and other oxidizing agents. In all cases, compounds
of the Group 1 metals contain the element as a 1+ ion. The solubility guidelines indicate that most sodium and potassium
compounds are water-soluble, so it is not surprising that sodium and potassium salts are found on earth either in solution, in the
oceans, or in under-ground deposits that are the residue of ancient seas. To a much smaller extent, these elements also are found
in minerals, like Chilean saltpeter (NaNO3) and borax (Na 2B4O7•10H2O).
The name ‘‘sodium’’ comes from caustic soda, NaOH. The Latin name for caustic soda (NaOH), natrium, furnished the
symbol Na. The name ‘‘potassium’’ is derived from the word ‘‘potash’’ (literally, the ashes from a fire pot). Potash (K2CO3) can
be extracted from ashes with water. The pure metals were prepared first by the English scientist Sir Humphry Davy (1778–1829)
in 1807 by electrolyzing the molten carbonates Na 2CO3 and K2CO3. The Latin word for potash is kalium, and hence we get the
symbol K for this element. Lithium is named after the Greek λιθοξ for "stone", because Arfvedson obtained it from mineral
sources, in contrast to the potash from plant sources Davy used to first isolate potassium. Rubidium and cesium were first
discovered using the spectroscope, and afterwards isolated. They are named after the colors of their principal spectral lines.
Within the earth’s crust, sodium and potassium are about equally abundant; however, sea water contains about 2.8% NaCl
but only about 0.8% KCl. Why this great difference, given that compounds of these elements have similar solubilities? The
answer lies in the fact that potassium is an important factor in plant growth. Much of the potassium in ground water is taken up
by plants. Most plants contain four to six times as much potassium as sodium. You may be aware that potassium is one of the
three components in most fertilizers. Some NaCl is essential in the diet of humans and animals because many biological functions
are controlled by the concentrations of Na + and Cl– ions. Animals travel great distances to reach a salt lick, and farmers often
place large blocks of salt in fields for cattle. The fact that salt has been
important for a long time is evident in surprising ways. We are paid a ‘‘salary’’
for work done; this word is derived from the Latin word salarium, which meant
‘‘salt money’’ be-cause Roman soldiers were paid in salt. We still talk about
‘‘salting away’’ money for a rainy day, a term related to the practice of
preserving meat by salting it.
Preparation and general properties
Sodium is still produced by electrolysis, although molten NaCl is now
used. Common chemical species are not strong enough reducing agents to
convert Na + to the metal, so electrolysis is the only viable method of
preparation. Potassium can be made by electrolysis also, but there are problems
with this method, not the least of which is that molten potassium is soluble in
the molten salt, making separation difficult. The preferred preparation of
potassium uses the reaction of liquid sodium with molten KCl.
Na(l) + KCl(l) à K(g) + NaCl(l)
This reaction is done at 850°C, where potassium is a gas, while sodium is still
primarilly a liquid (bp. 766, Na and 890°C, K). This reaction is a good example of
the importance of understanding chemical equilibria. The equilibrium constant
is less than 1, indicating that reactants are favored. The green potassium gas,
however, is removed continuously from the reaction, shifting the equilibrium to
the right.
The melting points of the alkali metals decline on moving down the group:
(Li 180.6, Na 97.8, K 63.7, Rb 39.5, Cs 28.4°C. Thus, Cs, Ga (m.p. 29.8 °C), and Hg
A Downs cell for producing sodium metal by
electrolysis. The circular iron cathode is
separated from the graphite anode by an iron
screen. Because the cell operates at about
600 °C, the sodium is produced at the
cathode in the molten state. The liquid metal
has a low density, so it floats on the top of
the liquid salt. The chlorine produced at the
anode bubbles out of the cell and is
collected.
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 61
are the only metals that are liquids at or near room temperature. The elements sodium to cesium are silvery metals that are soft
and easily cut with a knife, while lithium is harder and somewhat brittle. Their densities are lower than water. These elements are
all highly reactive. When exposed to air, the metal surface is quickly coated with an oxide film. Consequently, the metals must be
stored in a way to avoid contact with air; this is often done by placing them in kerosene or mineral oil.
They belong to "Group V" of the qualitative analysis scheme, i.e. the group that is not readily precipitated. (There are some
highly specialized tests that can distinguish between some of the alkali elements by precipitation, but these are too unusual to
be worth learning.) Instead, a general method of analysis and identification of these elements is to measure their flame colors. In
this technique, a cleaned loop of platinum wire, sealed in the tip of a soft-glass rod, is dipped in a solution containing the Group
1 element, usually as the aqueous cation, though covalent lithium compounds in non-polar solvents work very well in this
procedure. The solution is evaporated in a hot (blue colored) Bunsen-burner flame. Ionized Group 1 elements are carried away
in the flame, and undergo characteristic spectroscopic changes. The spectral properties are summarized in the following table:
Alkali metals and their flame colors
Metal
Color
Li
Crimson
Na
Yellow
K
Violet
Rb
Red-violet
Cs
Blue
wavelength of light, nm
670.8
589.2
766.5
780.0
455.5
↑
The exact spectroscopic process involved is different for each compound. What is
3p
important is that they are ionized relatively easilly, e.g. by the heat of the flame. The
process involved in the characteristic sodium yellow (also the basis of "orange" street
↑
3s
lighting) is as follows: in a flame, the sodium ion ("A" in figure at right) acquires an
electron in the 3p orbital (B); as the electron decays to the ground state 3s state, a
A
B
C
yellow light photon is emitted (C). Other elements such as heavier Group 2, Cu, etc. also
give colored flames, whereas many which have higher ionization energies do not.
The redox potential of Li is anomalous. The standard reduction potentials of Li, Na and K are –3.045, –2.714 and –2.925. In
terms of gas-phase ionization energy, we expect K to be easier to ionize than Na than Li, hence Li should have the smallest
reduction potential (in magnitude). This difference is due to solvation (hydration). As discussed in the section on acid/base
properties, the electrostatic parameter for Li is high (ξ = Z2/r) because of its very small radius. This gives it a high energy of
hydration, and it is this term in the thermochemical cycle that makes lithium a more potent reducing agent in aqueous solution.
Chemical reactivity
The Group 1 metals are highly reactive. The reaction with water generates an aqueous solution of the metal hydroxide and
hydrogen:
2 Na(s) + H2O(l) à 2 Na + (aq) + 2 OH– (aq) + H2(g)
In the laboratory you get to test the relative reactivity of some of the alkali metals with water. In fact, lithium is the least reactive,
despite the fact that it has the largest (in magnitude) reduction potential. The reason for this must be found in kinetic rather
than thermodynamic behavior. Li is a harder, higher-melting metal. At least some of the explosive reactivity of potassium (and
Rb, Cs) with water is due to the fact that the metals are quickly molten by the heat of the initial reaction, and this melted metal
blob exposes a larger surface area of the metal to the water (often multiple small balls of molten metal skim over the surface of the
water, instantly refreshing the surfaces as the metal is consumed); the explosions occur by the ignition of the produced
dihydrogen from the heat of the reaction. If sufficiently violent, the explosion can hurl burning, molten metal in all directions, so
these are definitely hazardous reactions.
Reaction with any of the halogens yields a metal halide.
2 Na(s) + Cl2(g) à 2 NaCl(s)
2 K(s) + Br2(l) à 2 KBr(s)
The reactions with elemental oxygen are surprising. Group 1 metal oxides, M 2O, are known, but it is the principal product of
reaction between the elements and oxygen only for lithium. The primary product of the reaction of sodium and oxygen is sodium
peroxide, Na 2O2, and not sodium oxide, Na 2O. The principal product from the reaction
You were wondering…
of potassium, rubidium and cesium with oxygen is the superoxide.
How can you make sense of a
2 Na(s) + O2(g) à Na 2O2(s)
formula such as KO2? Try to write a
K(s) + O2(g) à KO2(s)
Lewis structure for a molecular
Both Na 2O2 and KO2 are ionic, with Group 1 cations paired with either the peroxide ion
oxygen ion.
(O22–) or the superoxide ion (O2–). These are not just laboratory curiosities. They are
used in oxygen generation devices in places where people are confined, such as
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 62
submarines, aircraft, and space craft. When a person breathes, for every liter of O2 inhaled, 0.82 L of CO2 is exhaled. A
requirement of an O2 generation system, therefore, is that it should produce a larger volume of O2 than the volume of CO2 taken
in. This requirement is met with peroxides and superoxides.
4 KO2(s) + 2 CO2(g) à 2 K2CO3(s) + 3 O2(g)
K2CO3(s) + CO2(g) H2O (g) à 2 KHCO3(s)
Lithium (along with magnesium from Group 2) is a rare example of an element that reacts with the normally inert elemental
nitrogen to form a reactive, sometimes explosive, metallic nitride:
6 Li + N2 (g) à 2 Li3N (s)
Sodium compounds of commercial importance
Electrolysis of aqueous sodium chloride is the basis of the chlor-alkali industry, one of the largest chemical industries in the
world. The major commercial products from this process are chlorine and sodium hydroxide. We will consider this process in
detail as part of our study of the Group 17 elements, since it forms part of the story of the complex redox states of chlorine.
2 NaCl(aq) + 2 H2O(l) à 2 NaOH(aq) + H2(g) + Cl2(g)
In 1993, 11.7 billion kg of NaOH and 11.1 billion kg of Cl2 were produced in the United States. Sodium carbonate is another
commercially important compound. In 1993, 9.0 billion kg of Na 2CO3 was produced in the United States, making it the 12th ranked
industrial chemical. This compound has the common name soda ash, or washing soda, and it has been obtained, since
prehistoric times, from naturally occurring deposits of Na 2CO3•10H2O. Trona, Na 2CO3•NaHCO3•2H2O, however, estimated at 6
× 1010 tons, was recently discovered in Wyoming, and virtually all Na 2CO3 in the United States now comes from that source.
About 40% of the soda ash is used in the manufacture of glass, but large amounts are also used in water treatment, in pulp and
paper manufacture, and in cleaning materials.
The use of chlorine has come under attack from environmental groups; therefore, considerable interest has arisen in
manufacturing sodium hydroxide without the co-product chlorine, as in the case of salt electrolysis. This has led to a revival of
the old ‘‘lime-soda process,’’ which produces NaOH from inexpensive lime (CaO) and soda (Na 2CO3).
Na 2CO3(aq) + CaO(s) + H2O(l) à 2 NaOH(aq) + CaCO3(s)
The calcium carbonate byproduct is filtered off and is recycled into the process by heating it (calcining) to recover lime.
CaCO3(s) à CaO(s) + CO2(g)
Sodium bicarbonate, NaHCO3, is produced in small amounts from soda ash. You are probably more aware of the bicarbonate
under the common name bak ing soda Not only is NaHCO3 used in cooking, but it is also added in small amounts to table salt.
This is because NaCl is often contaminated with small amounts of MgCl2. The magnesium salt is hygroscopic; that is, it picks
water up from the air and, in doing so, causes the NaCl to clump on damp days. Adding NaHCO3 converts MgCl2 to magnesium
carbonate, a nonhygroscopic salt.
MgCl2(s) + 2 NaHCO3(s) à MgCO3(s) + 2 NaCl(s) + H2O(l) + CO2(g)
Ammonium, the pseudo-alkali metal ion
The ammonium ion, NH4+ , can substitute for alkali elements in most compounds and reactions. Like the Group 1 elements,
ammonium has a 1+ charge in the ionized form, and most of its salts are very soluble. Moreover the effective ionic radius of this
molecular ion (1.51 Å) is almost identical to the potassium ion (1.50 Å). However, whereas the Group 1 cations in virtually all
their compounds are inert (have the noble gas ion configuration), ammonia can undergo additional reactions involving the four
hydrogen atoms (e.g. acid-base reactions) or the nitrogen atom (redox reactions). For example, contrast the thermal
decomposition of the nitrates of sodium and ammonium:
2 NaNO3 (s) à 2 NaNO2 (s) + O2 (g)
NH4NO3 (s) à N2O (s) + H2O (g)
On the resemblance of lithium to magnesium
Because so many periodic trends move from the bottom left to the top right of the periodic table, chemical resemblance
within a particular group is sometimes less striking than expected, and similarities often exist with the elements located on group
left and one period down in the periodic table. (Consider for example the familiar zig-zag sequence of the metalloids across the p
block of the periodic table.) Often these elements have similar electronegativity or electrostatic parameters. This is the case for
lithium and magnesium in groups 1 and 2. Thus, to list some examples, lithium is quite hard and brittle, much more like
magnesium than the soft remaining Group 1 elements; it forms a normal oxide on reaction with O2, rather than a peroxide or a
superoxide. Also lithium is the only one in its group to form a nitride, while this is quire common for Group 2. Lithium carbonate,
phosphate and fluoride have anomalously low solubility, but the analogous Group 2 salts are also of low solubility.
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 63
Biological aspects
Sodium and potassium are essential elements. Plants in particular need high levels of potassium, and hence potassium salts
are a big part of the fertilizer industry (in Canada, the potash mining industry of Saskatchewan is worthy of mention, since it is
responsible for about 24% of world production. Note that it is not KOH that is mined! The term potash is incorrectly applied to
this industry and in fact it is KCl (sylvite) and mixed NaCl,KCl (sylvanite) that are actually mined.) Most of the fertilizer is
applied to the soil as KCl, usually in formulations that include potassium and nitrogen (e.g. mixtures of KCl and NH4NO3, etc).
Fertilizers are usually designated with three numbers, indicating respectively their nitrogen, phosphorus and potassium
content. The balance is an inert filler, including stabilizers that minimize clumping of the active ingredients. Note that the
chemical identity of the components is not specified, and is to some extent irrelevant. For specialized applications, however,
certain chemicals must be avoided. For example, in soils in contact with ocean water (e.g. in the Netherlands), chloride salts
should be avoided. Thus potassium sulfate must be used instead of the more common KCl.
Fertilizer labelling system
10
10
% N by mass
5
% P as mass of P2O 5 % K as mass of K2O
Human nutrition requires both sodium and potassium. Although our bodies contain much more sodium than potassium, but
western diets tend to be so sodium-rich (over-salting) that most people take in much more than the ≈ 1 g sodium ion required to
maintain optimum levels of the element. However, potassium deficiency is quite common; coffee and banana are good dietary
sources of potassium. The main use of the alkali metal ions is to balance the negative charges associated with many of the
protein units in the body. They also help maintain the osmotic pressure within cells, preventing them from collapse. Of
particular importance, the total concentrations of potassium and sodium ions are strikingly different inside and outside of cells,
as shown in the following table:
Concentrations of alkali ions in the body (mmol/L)
Ion
[Na+]
[K+]
Red blood cells
11
92
Blood plasma
160
10
Since the volume of bulk body fluids is much greater than the intracellular fluids, there is less potassium present in mamals than
there is sodium. The difference in ion concentration inside and outside the cells produces an electrical potential across cell
membranes, which is essential for the functioning of nerves and muscles. The cell walls contain biological ion pumps that move
K+ into cells and remove excess Na + from cells. Going into shock after a traumatic accident is caused by the massive leakage of
alkali metal ions through the cell walls.
Lithium ions, administered as lithium carbonate, remains the standard treatment for manic depressive illness (now known as
bipolar illness). Care must be taken to keep the total [Li+ ] low enough to avoid serious side-effects, including cardiac arrest. A
blood level of 1 mmol/L seems to be optimal. Although the effect has been known since 1938, the origin of the lithium ion
therapy is not understood. The disease itself is also poorly understood. It is possible that if affects either the Na + /K+ or the
Mg 2+/Ca 2+ balances, or both. Note that the mobility of aquated lithium ions are much lower than that of sodium or potassium, so
that significant differences in the solution behavior of the elements is expected. Obviously the affinity of various biological
receptor sites designed for any of the above-mentioned metals towards lithium ion will also be very different.
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 64
4.4.3 The Group 2 metals
Be
Mg
Ca
Sr
Ba
Ra
The Group 2 elements are called alkaline earths. The ‘‘earth’’ part of the group name is left over from the days of medieval
alchemy. To alchemists, any solid substance that did not melt and was not changed by fire into another substance was called an
‘‘earth.’’ Various compounds of Group 1 and 2 elements that were known in those times, such as NaOH and CaO, were alkaline
according to the experimental tests of the alchemists: they had a bitter taste and could be shown to neutralize acids. Group 1
compounds, however, melted in a fire or combined with the clay containers in which they were heated. Melting points of most
Group 2 compounds are very high, due to the strong attractive forces between the M2+ cation with the anions present (page
119).
For example, CaO melts at 2572°C, a temperature well beyond the range of an ordinary fire. Calcium compounds such as lime
(CaO) were known and used in ancient times. Calcium metal, however, was first prepared in 1808 by Sir Humphry Davy, who also
prepared magnesium, strontium, and barium in the same year. As with sodium and potassium, Davy’s preparation of these
elements was accomplished by electrolysis of a molten salt. The great abundance of calcium and magnesium on earth leads to
their occurrence in plants and animals, and both elements form many commercially important compounds. It is on this chemistry
that we want to focus our attention.
Similar to the Group 1 elements, the Group 2 elements are very reactive, so they are only found in nature combined with
other elements. Unlike Group 1 metals, however, many compounds of the Group 2 elements have low water-solubility. This
explains their common occurrence in various minerals. Limestone (CaCO3), gypsum (CaSO4•2 H2O), and fluorspar (CaF2) are
examples of common calcium-containing minerals. Magnesite (MgCO3), talc or soapstone (3MgO•SiO2•H2O), and asbestos (3
MgO•4SiO2•H2O) are magnesium-containing minerals, whereas dolomite is a compound with both elements (MgCO3•CaCO3).
Limestone, a sedimentary rock, is found widely on the surface of the earth. These deposits are the fossilized remains of
marine life and are the most common of the calcite forms of the compound. You also know other forms of calcite, however.
Marble is fairly pure calcite formed by crystallization of CaCO3 under high pressure. High-quality deposits of marble are found in
Italy and in the United States in Vermont, Georgia, and Colorado. Another form of calcite is Icelandic spar, which forms large,
clear crystals.
Preparation and general properties
Beryllium is a rare element with metalloid properties. It has a steel-gray colour, and is very hard. It is obtained from the
mineral beryl, Be 3Al2Si6O18, the main constituent of the gem-stones aquamarine and emerald. Needless to say, less-valuable
brown or colorless beryl is used to produce the metal. One of its main uses is in the production of vacuum-stable, X-ray
transparent windows for X-ray tubes. X-ray absorption depends on the number of electrons (i.e. the atomic number), and Be is
the lightest element strong enough to make such windows. Beryllium oxide is amphoteric, dissolving in strong acid to give the
aqua ion, but in strong alkali to produce the soluble [Be(OH)4]2– ion which has Td symmetry.
Several hundred thousand tons of magnesium are produced annually because the low density of the metal (1.74 g/cm3)
makes it useful in lightweight alloys. For example, most aluminum used today contains about 5% magnesium to improve its
mechanical properties and to make it more resistant to corrosion. Other alloys having more magnesium than aluminum are used
when a high strength-to-weight ratio is needed and when corrosion resistance is important, such as in aircraft and automotive
parts ("Mag wheels") and in lightweight tools. Although there are many magnesium-containing minerals, most magnesium is
obtained from sea water, in which it is present in a concentration of about 0.05 M. To obtain the element, magnesium is first
precipitated from sea water as the relatively insoluble hydroxide (Ksp for Mg(OH)2 = 1.5 × 10–11 ). The base used in this reaction,
CaO, is prepared from sea shells, CaCO3. Heating CaCO3 gives CO2 and CaO, and addition of water to CaO gives calcium
hydroxide. When Ca(OH)2 is added to sea water, Mg(OH)2 precipitates.
Mg 2+ (aq) + Ca(OH)2(s) à Mg(OH)2(s) + Ca2+ (aq)
Magnesium hydroxide is isolated by filtration and then neutralized with hydrochloric acid.
Mg(OH)2 (s) + 2 HCl (aq) à MgCl2(aq) + 2 H2O(l)
After evaporating the water, solid magnesium chloride is left. The anhydrous MgCl2 melts at 708°C, and the molten salt is
electrolyzed to give the metal and chlorine.
MgCl2 (l) à Mg(s)
+ Cl2(g)
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 65
Calcium and magnesium are fairly high-melting, silvery metals. The chemical properties of Group 2 elements present few
surprises. All are oxidized by a wide range of oxidizing agents to form ionic compounds that contain the M2+ ion. For example,
these elements combine with halogens to form MX2, and with oxygen or sulfur to form MO or MS. All except beryllium react
with water to form hydrogen and the metal hydroxide, M(OH)2. With acids, hydrogen is evolved and a salt of the metal and the
anion of the acid results.
Chemical reactivity
The elements are very reactive, and all can be burned in air or oxygen, although high temperature (e.g. a flame) is needed to
initiate the reaction. Recall the burning of magnesium ribbon in the lab experiments. This reaction is so vigorous that it is not
even extinguished by carbon dioxide, the constituent of common fire extinguishers. Thus magnesium continues to burn:
2 Mg (s) + CO2 (g) à 2 MgO (s) + C(s)
The oxides are precursors to the hydroxides. Note that the hydroxides of Sr and Ba (as well as the radioactive radium) are
quite soluble, so that these act as strong bases, and are sometimes used in analytical titration in cases where the Group 1
hydroxides are not suitable. Barium reacts with excess oxygen to form a peroxide (as does sodium). This can be explained by the
large size, and hence lower charge density, of the Ba 2+ ion, which forms stable crystal lattices with the larger peroxide ion and will
not decompose the weakly-bonded molecular ion as would the smaller Group 2 cations which have a higher charge density due
to smaller radius.
Most of the Group 2 elements react with nitrogen at elevated temperatures to form nitrides, for example:
3 Ca (s) + 2 N2 (g) à Ca3N2 (s)
Calcium oxide (lime) reacts with carbon (in the form of coke) to produce a saline carbide, CaC2. As a charge analysis shows, this
contains the C22—diatomic ion, which is better known as a diacetylide. It reacts with water to form acetylene, HC≡CH, and indeed
the carbide ion is best considered to be –C≡C–.
Calcium compounds of commercial importance
The most important fluoride of the alkaline earth metals is fluorspar, CaF2, but fluorapatite (CaF2•3Ca3(PO4)2) is becoming
increasingly important as a commercial source of fluorine. Almost half of the CaF2 mined is used in the steel industry where it is
added to the mixture of materials that are melted to make crude iron. The CaF2 serves to remove some impurities and improves the
separation of molten metal from slag, the layer of silicate impurities and byproducts that comes from reducing iron ore to the
metal. The other use of fluorspar is in the manufacture of hydrofluoric acid by reaction of the mineral with concentrated sulfuric
acid.
CaF2(s) + H2SO4 (l) à 2 HF(g) + CaSO4 (s)
Hydrogen fluoride is very reactive and is an extremely important chemical. It is used to make cryolite, Na 3AlF6, a material needed
in aluminum production, and fluorocarbons such as polytetrafluoroethylene (Teflon). Apatites are collectively referred to as
phosphate rock. Over 100 million tons are mined annually; Florida alone accounts for about one-third of the world’s output.
Much of this rock is converted to phosphoric acid by reaction with sulfuric acid. Phosphoric acid is used to manufacture a
multitude of products (fertilizers and detergents), and its reaction products are found in baking powder, in frozen fish, and in
many other food products.
CaF2•3 Ca3(PO4)2 (s) + 10 H2SO4 (aq) à CaSO4 (s) + 6 H3PO4 (aq) + 2 HF (g)
Because of their economic importance, calcium carbonate and calcium oxide are of special interest. The thermal
decomposition of CaCO3 to lime, CaO, is one of the oldest chemical transformations known. Among industrial chemicals
produced today, lime is fifth with 16.7 billion kg produced in 1993. Most limestone and lime are used in the chemicals industry,
and one third of the lime is used to make steel by the basic oxygen process. Limestone has been used in agriculture for centuries.
It is spread on fields to neutralize acidic compounds in the soil and supply the essential nutrient Ca 2+ . Because magnesium
carbonate also is often present in limestone, ‘‘liming’’ a field also supplies Mg2+ , another important plant nutrient. For several
thousand years, lime has been used as mortar (a lime, sand, and water paste) to secure stones to one another in building houses,
walls, and roads. The Chinese used it in setting stones in the Great Wall. The Romans perfected its use, and the fact that many
of their constructions still stand is testament both to their skill and the usefulness of lime. In 312 BC, the famous Appian Way, a
Roman highway stretching from Rome to Brindisi (a distance of about 350 miles) was begun, and lime mortar was used between
several layers of its stones. The utility of mortar depends on some simple chemistry. Mortar consists of one part lime to three
parts sand, with water added to make a thick paste. The first reaction that occurs is the formation of Ca(OH)2, the process
referred to as ‘‘slaking,’’ the product being called slaked lime. When the mortar is placed between bricks or stone blocks, it
slowly absorbs CO2 from the air, and the slaked lime reverts to calcium carbonate.
Ca(OH)2 (s) + CO2 (g) à CaCO3 (s) + H2O(l)
Although sand mixed into mortar is chemically inert, the grains are bound together by the particles of calcium carbonate, and a
hard material results.
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 66
In modern construction, mortar is much less important than Portland cement. This essential construction material is the
basis of all modern mortars and all concrete. Modern industrial society could not exist without the incredibly useful material that
concrete is, providing a material which in most respects acts as "synthetic rock", and moreover which can be formed in situ to
mold virtually any shape imaginable. The building we sit in is constructed largely out of this marvelous material.
The production of cement involves the grinding together of limestone and shales (mixtures of aluminosilicates), and heating the
finely-divided solids in a furnace to 1500°C. Carbon dioxide is released, and the partly molten silicates on cooling form solid
lumps called clinker. The clinker is ground to an extremely fine powder, and stabilized by the addition of small amounts of
calcium sulfate. This material, sold as Portland cement contains 26% dicalcium silicate, Ca 2SiO4; 51% tricalcium silicate, Ca 3SiO5;
and 11% tricalcium aluminate, Ca 3Al2O6. When water is added, a number of complex hydration reactions set in. One idealized
equation which gives an idea of what is going on is:
2 CaSiO4 (s) + 4 H2O (l) à Ca3Si2O7.3H2O (aq) + Ca(OH)2 (aq)
The chief product, a hydrated calcium silicate, is called tobermorite gel, forms strong crystals that bond strongly to the silicates
in the sand and gravel that are added to give the concrete volume and strength.
In Alberta there is a large Lafarge Canada cement factory located at Exshaw, just east of Banff national park. The limestone
is harvested directly behind the plant, and the shales come from Seebe, just 7 km downstream along the Bow river. The major
advantage of cement over lime mortar, is that it sets without external carbon dioxide. The hardening of concrete is a natural
polymerization reaction, and one that is quite exothermic. Thus a few hours after pouring, cement or concrete are normally warm
to the touch. Note that both the calcium hydroxide by-product and the calcium silicates and aluminosilicates are quite alkaline,
so that cement is quite corrosive until it is fully set.
‘‘Hard water’’ contains dissolved ions, chiefly Ca 2+ and Mg 2+ . Water containing dissolved CO2 reacts with limestone.
CaCO3 (s) + H2O (l) + CO2 (g) à Ca2+ (aq) + 2 HCO3– (aq)
This reaction is reversible, however. When hard water is heated, the solubility of CO2 drops, and the equilibrium shifts to the left.
If this happens in a heating system or steam-generating plant, the walls of the hot water pipes can become coated or even
blocked with solid CaCO3. In your house, you may notice a coating of calcium carbonate on the inside of cooking pots. The last
equation describes the chemistry of caves as well. The acidic oxide CO2 reacts with Ca(OH)2 to produce white, solid CaCO3.
When further CO2 is available, however, the CaCO3 can re-dissolve because of the formation of aqueous Ca 2+ and HCO3– ions.
Photograph of the Lafarge Cement Plant at Exshaw, Alberta
(Photo by Prof. T.W. Swaddle, Department of Chemistry, University of Calgary)
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 67
Biological aspects
Magnesium is present in chlorophyll, which is the active site in green plant leaves for the conversion of carbon dioxide and
water into sugars and oxygen. The energy to drive this non-spontaneous reaction is obtained from sunlight, so that it is a
photochemical process. Photons of visible light are absorbed and their energy content is harnessed by the plant through
complex protein systems called photsystem 1 and photosystem 2. The reaction catalyzed by this process can be written simply
for the synthesis of glucose:
6 CO2 (g) + 6 H2O (l) à C6H12O6 (aq) + 6 O2
Magnesium sits in the middle of the chlorophyll, and maintains the molecule in a specific configuration. Its inertness to
further reaction (i.e. Mg2+ is the only oxidation state for magnesium) means that the high-energy electron transfer reactions that
take place in photosynthesis will not affect the magnesium ions.
Both magnesium and calcium are present in mammalian system. Magnesium is more concentrated within the cells, and less
without, while calcium has a higher outside concentration. Calcium ions are important in blood clotting, and they are required to
trigger the contraction of muscles. Eating certain foods, e.g. spinach or rhubarb, which are rich in oxalic acid, (COOH)2 cause the
precipitation of calcium oxalate, Ca(C2O4), a highly insoluble calcium salt. This precipitate is responsible for the "dry taste" one
gets in the mo uth from eating either of these foods.
Calcium is a chief component of bones and skeletons, which in mammalians systems are composed of apatite,
Ca5(OH)(PO4)3. Apatite is also the material of tooth enamel. Fluoride treatment of teeth involves the replacement of OH– by F–
to form fluorapatite Ca5(F)(PO4)3. This material is much harder, and resists cavity-forming bacteria more effectively than the
parent apatite. Excessive fluoride consumption, as occurs when small children swallow their nicely-flavored tooth-paste leads to
the same substitution reaction taking place in skeletal bones, and here it is destructive because it weakens the hydrogen-bonded
structural strength of bones, making them more brittle. This condition is known as fluorosis. Major public relations battles rage
from time to time over the issue of fluoridation of drinking water supplies. Most of the arguments that I have heard both from
the medical fraternity promoting the practice, and the environmentalists opposed to it are woefully ignorant of the basic science
involved. Fluoride is extremely useful as a topical treatment, e.g. at the dentist and in toothpaste, but it should not be ingested.
This simple fact immediately suggests what the scientifically correct attitude to water fluoridation should be!