Earth Metals

8/8/2019 Earth Metals

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Earth Metals - Aluminium

The elements of the third main group are named as earth metals having an s2px

valence-shell configuration. Typically they can form three covalent bonds, their

oxidation number is +3 in compounds. The earth metals in their compounds can be

considered as electrondeficient, because with the formation of three covalent bonds

they do not reach the electron octett characteristic for noble gases. Therefore they

behave as Lewis acids and are able to form complexes.

Boron (B) the earth metal of the lowest atomic number, is perhaps one of the most

special element of the periodic table, but it arouses rather the interest of the

chemists, not that of the physicians... Boron is a very hard high-melting substance

present in network crystal lattice. It occurs in the Nature in deposits of minerals, the

most important ones are borax (Na2B407) and boric acid (H3BO3). Both of them are

crystalline, water-soluble solids. The aqueous solution of boric acid is slightly acidic,

but contrary to the other oxoacids boric acid is not a proton donor, but it acts as a

Lewis acid, binding OhTions:

H3B03 + H20 ======= [B(OH)4]"+H+

The most important earth metal is aluminium (Al); aluminium is the third most

abundant element of the Earth's crust. It occurs in the nature in silicates or in

hydroxides, its most important ore is bauxite that is mostly the mixture of AIO(OH)

and Al(OH)3. Cryolite (sodium hexafluoro-aluminate(lll), Na3[AlF6]) is also an important

mineral, applied in the manufacturing of aluminium (electrolysis).

Aluminium is a bluish silvery soft metal of low density (2.7 g/cm3, light metal). It is

an excellent thermal and electrical conductor. Therefore and because of its

advantageous mechanical properties it is applied as structural material in

construction of vehicles, as well as because its good electrical conductivity as electric

cables. Generally for these purposes its alloys are used. It can be drawn to very thinfoils (aluminium foils) that can be utilized for the packing of food. Aluminium has very

advantageous properties, but its manufacturing is rather expensive (made by

electrolysis), the objects made of aluminium cost much more than that of made of

iron (Between 1970 and 1980 coins were made of aluminium, too).

Gallium (Ga) and indium (In) are applied in the semiconductor industry for the

production of transistors, LEDs, they do not have biological importance.

Thallium (Tl) and its compounds are rather toxic, their contact with skin is

dangerous, they are very toxic by inhalation and if swallowed. (It can be found in the

periodic table between the very toxic mercury and lead.) Previously its sulphate


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(TI2S04) was applied as rat poison, but this tasteless and odourless substance is too

dangerous, therefore nowadays it is not used for this purpose.


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The chemical properties of aluminium

It has a medium value of electronegativity (1.5); it can form both ionic and

covalent compounds. The pure aluminium is a very reactive metal with reducing

property.

A thin protective surface layer of aluminium oxide forms when the metal is

exposed to air, preventing further oxidation, so the objects made of aluminium are

resistant to corrosion. In order to protect aluminium from oxidation more effectively

it is oxidized by electrochemical methods (anodizing) forming a passivation layer on

the surface of the metal.

If aluminium is treated with mercury, an aluminium-amalgam (alloy) forms on its

surface inhibiting the development of the protecting layer of aluminium oxide, thus

the metal undergoes oxidation. (This happens at the treatment with

mercury(ll)chloride, too.)

At higher temperature the aluminium burns yielding aluminium oxide in an

exothermic reaction.

4 Al + 3 02 = 2 Al203

Aluminium has a high affinity toward oxygen, at high temperature it reduces

other metal oxides, this reaction is an appropriate method for the preparation of

certain metals (thermite method).

Cr203 + 2 Al = Al203 + 2 Cr

Aluminium dissolves in both acids and bases generating hydrogen gas, i.e. it is

amphoteric.

2 Al + 6 H+ = 2 Al3+ + 3 H2

2 Al + 2 OH" + 6 H20 = 2 [AI(OH)4]" + 3 H2

Therefore it is not advisable to clear the objects made of aluminium with cold

degreasing agents, because they contain mostly potassium hydroxide. The acidic

limescale remover liquids can damage the aluminium objects, too.

Compounds of aluminium

In its compounds aluminium has oxidation number of +3. In covalent compounds

it is strongly electron deficient, two more electrons are needed to reach the octet

configuration. Thus the aluminium compounds behave as strong Lewis acids forming

complexes with electron pair donor.

AICI3 + :NH3 = CI3AI^-NH3

AI(OH)a +OH" = [AI(OH)4]~


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Aluminium-oxide (Al203) is a colourless, crystalline substance. It is not water

soluble (neutral oxide) but it is soluble in acids or in bases (amphoteric).

Al203 + 6 H+ = 2 Al3+ + 3 H20

Al203 + 2 OH" + 3 H20 = 2 [Al(OH)4]" (complex formation)

Different network crystal structures of aluminium oxide occur in Nature as semi-

precious stones (ruby, sapphire, turquoise), their colour is connected to the metal

contaminations. Some special


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forms as corundum are utilized as grinding powder (sandpaper), because it is a very

hard material. In chemical practice aluminium oxide (alumina) is used as a polar

stationary phase in adsorption chromatography.

The aluminium hydroxide (Al(OH)3) a colourless powder, insoluble in water, butit is an amphoteric substance.

AKOHK + 3 H+ = Al3+ + 3 H20

AKQHU +OH" = [Al(OH)4]"

Aluminium hydroxide is also used in dyeing of textiles by mordant because the

aluminium hydroxide precipitated can adsorb dyes. In addition it is used in vaccines

and in toxins to absorb the active drugs.

Aluminium-chloride (AICI3) is a colourless crystalline substance, very

hygroscopic having covalent structure. If it is exposed to air it deliquesces (forms a

solution), because it has a strong affinity for moisture. Its aqueous solution is acidic

because of hydrolysis. (See equilibria and complexes).

[AI(H20)6]3+ + 3 Cf ======= [AI(H20)5OH]2+ + H+ + 3 CI" (hydrolysis)

In the preparative organic chemistry it is applied as a Lewis acid catalyst -

electrondeficient molecule - in the so-called Friedel-Crafts reactions.

Medical and biological importance of the earth metals

Boron is a trace element in our organism; it plays a role in the regulation of the

calcium metabolism. Both boric acid (H3B03) and borax (Na2B407) are used in the

medical practice. They are not harmful to the human organism, whereas they act as

bacteriostatic agents. Therefore they are constituents of disinfective baby powders

and ointments. Boric acid is also applied in eye drops and artificial tears. The dilute

solution of boric acid is a decontaminant at accidents caused by alkali burns (in Red

Cross bottle in the lab). The solutions of borax in glycerine have a role in the

treatment of the soar of infants.

Toxicity of aluminium compounds was described in the 1970s. The growth of the

plants is retarded by the contaminations caused by aluminium compounds. The

human organism does not require aluminium. In our environment the acid rains wash

out the aluminium from aluminium silicates and we ingest them with food. At the

Alzheimer patients in the grey matter of the brain there are plaques composed of

aluminium silicates but probably this is a symptom of the disease, and not its cause.

Aluminium sulphate (AI2(S04)3) and alum (potassium aluminium sulphate,

KAI(S04)2) are colourless crystalline compounds, their solutions is acidic because of

hydrolysis. Both of them are applied as astringent that is for the precipitation of

proteins and for cauterization. Aluminium sulphate has a role in waste water


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treatment. Alum is an astringent contracting the blood vessels; it is applied in after-

shave hemostatic sticks.


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Aluminium oxide (Al203) is utilized as an antacid to neutralize the excess of gastric

acid. Its advantage is that it does not produce gases and it is not absorbed, because

it precipitates in the intestines at pH=7 as aluminium hydroxide polymer.

Aluminium chloride and its hydrates are the active ingredients of antiperspirants.

Since AICI3 is a highly electron deficient molecule (Lewis acid, there are only 6

valence electrons around the aluminium), it binds the basic amine components of

sweat having intense odour (see complexes).

Self-test questions

1. Write the atomic symbol of aluminium and its valence electronic configuration.

Give the oxidation number of aluminium in its compounds. What kind of

advantageous physical properties does this metal have?

2. What do you know about the application of aluminium in the every day life? Listthree different items. Why do not aluminium objects undergo corrosion?

3. Write two equations to demonstrate the amphoteric character of aluminium.

4. Write two equations to demonstrate the amphoteric character of aluminium

hydroxide.

5. Write formula for aluminium sulphate and alum. For what purposes are they

applied in the medical practice? Write equation for the dissociation of alum.

6. Name the following compounds: H3B03, KAI(S04)2, AICI3, Al203, AI2(S04)3.

7. Define the term of Lewis acid. Write equation to demonstrate the Lewis acid

character of aluminium chloride

8. For what purposes is aluminium oxide applied in the medical practice? What kind

of advantages does its application have?

9. What do you know about the physiological effect of aluminium compounds?

Experiments

Dissolution of aluminium in water

Materials: a sheet of aluminium, mercury(ll) chloride solution, water

Procedure: Place a piece of aluminium sheet into a test-tube and add a small amountof 1% mercury(ll) chloride solution to it. Let the mixture stand for 1-2 minutes. Then

decant the solution and rinse the metal sometimes with distilled water. Then add

some water to the pre-treated aluminium to cover it. Place the test-tube into hot

water bath and let it there for 5 minutes. What can you observe? Why do you have

to pre-treat the metal with mercury(II) chloride solution? Comments: After a while

the solution will turn turbid because of the AI(OH)3 precipitate.

The amphoteric character of aluminium

Materials: aluminium filings, 20% hydrochloric acid, 20% NaOH


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Procedure: Put a small amount of aluminium filling (or a sheet of aluminium) into two

test-tubes and place them into the test-tube rack. Fill into one test tube 5 cm3 of 20%

hydrochloric acid, into the


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other one 5 cm3 of 20% NaOH solution. What can you observe? Write equations for

these reactions. How can you characterize the acid-base properties of aluminium?

Why is it not supposed to clean objects made of aluminium with cold degreasing

agents or limescale removers? Comments: Proteins are also amphoteric, since they

contain both acidic and basic functional groups. They can react with both acids and

bases, thus they have a role in the regulation of constant pH of the body (See chapter

'Buffer systems').

Oxidation of aluminium

Materials: sheet of aluminium, mercury(ll) chloride solution

Procedure: Clean the aluminium sheet with a piece of sandpaper. Add a drop of

mercury(ll) chloride solution onto the clean metallic surface. When the surface of the

sheet turned gray, shake off the drop of solution and disperse the remained liquid

with cotton wool or a piece of filter paper. (The drop of mercury(ll) chloride solution is

dropped with a pipette onto the aluminium sheet, and not with hand!) Let the sheet

stand on air. After some minutes observe the surface treated with mercury(ll)

chloride solution and touch the back side of the sheet (the not treated side). What do

we observe? Write equation.

Comments: At high temperature aluminium burns throwing out sparks. It has a very

high affinity to oxygen, which is the basis of the so-callet thermite reaction used for

the preparation of pure metals.

(http://www.youtubexom/watch?v=4EviwhG2nU8).

Amphoterism of aluminium hydroxide

Materials: alum solution, sodium hydroxide solution, hydrochloric acid solution

Procedure: Add dropwise 5 % of sodium hydroxide solution to 1-2 cm3 of

concentrated alum solution. Divide the reaction mixture into two test-tubes. Add

some more sodium hydroxide to one test tube, to the other one hydrochloric acid

solution. What do you observe? Write equation. Comments: In the slighly basic

medium of small intestines aluminium hydroxide precipitates and it is not absorbed.That is why aluminium oxide or hydroxide are used as antacids. However, the finely

dispersed aluminium hydroxide is able to adsorb several substances on its surface

(see colloids), therefore it is not supposed to take other medicaments together with

antacids (aluminium hydroxide can adsorb the active ingredients).
http://www.youtubexom/watch?v=4EviwhG2nU8http://www.youtubexom/watch?v=4EviwhG2nU8


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s-block elements

The elements of the s-block (alkali metals and alkaline earth metals), are soft

metals of low density and low melting point. Releasing one or two electrons they

reach the noble gas electron configuration, therefore they have very low

electronaffinity and ionization energy. They are very reactive elements; they do not

occur in nature in elementary state only in their compounds having mostly ionic

character. Their reactivity increases in the column from the top to the bottom; this

way decreases the ionization energy and the electronegativity.

Their typical feature is the flame test. The valence electrons of the atoms can be

excited easily thus the non-luminous flame of the Bunsen burner is coloured by these

substances (except magnesium and beryllium).

Alkali metals

Elements of the first main group have s1 valence-shell electron configuration.

These metals show a tendency to form monovalent cations. Hydrogen is an

exception in this column, it is a non-metal, it is discussed elsewhere.

element density (g/cm3) melting point (C) reactivitylithium Li 0.53 180.5 Ysodium Na 0.97 97.8potassium K 0.86 63.7rubidium Rb 1.53 38.9cesium Cs 1.90 28.7francium Fr - (radioactive) (27.0)

They are very reactive, exposed to the air they immediately undergo oxidation.

To prevent it they are stored under mineral oil. Their reaction with water is very

vehement yielding hydrogen gas:

K + H20->KOH + 1/2H2

This reaction is rather exothermic and the hydrogen evolved can even burst into

flame.

These metals behave as strong reducing agents having the highest negative

standard electrode potential (see electrochemistry).

In their compounds the alkali metal cations always have the oxidation number

+1. Because they have a closed electron shell and a stable octet in their compounds

they are colourless with low


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reactivity. Their ionic compounds (halogenides, sulphates, carbonates and

phosphates) have good water solubility. Only lithium, sodium and potassium

compounds are of biological importance.

The oxides of alkali metals are base anhydrides, with water they yield thecorresponding hydroxides:

Na20 + H20 -> 2 NaOH

Their hydroxides are ionic compounds of very good water solubility, they act as

strong bases, and in addition they are caustic and corrosive compounds.

NaOH-+Na+ + OH~

They sodium and potassium hydroxides are ingredients of cold degreasers and

drain cleaning unblockers, because they hydrolizes fats yielding water soluble

compounds (see organic chemistry).

The alkali hydroxides react with acids or acid anhydrides forming salts. They are

hygroscopic and readily absorb water and carbondioxide (acid anhydride) from air, so

they should be stored in an airtight container in order to prevent the formation of

carbonates:

2 KOH + C02 K2C03 + H20 (reaction of an acid anhydride with a base)

Alkali metals occur in nature as salts. Sodium exists as table salt (NaCI), washing

soda or natron (Na2C0

3), potassium as potassium chloride or muriate of potash (KG)

found in the top layer of salt mines. Sea water contains about 3% NaCI so the

evaporitic deposits formed by the evaporation of inland seas. In mineral waters there

can be an important amount of Glauber's salt (Na2S04, sodium sulphate) showing a

mild laxative effect in higher amount. An important mineral of sodium is Chile

saltpetre (NaN03, sodium nitrate); the corresponding potassium mineral is salpetre

(KN03, potassium nitrate) formerly applied for the manufacturing of gunpowder.

Nowadays it is used to prepare artificial fertilizers. The detection of the alkali metals

can be performed in an easy way using their flame test: sodium colours the non-

luminous flame of the Bunsen burner intense yellow, while potassium give a faint

blue flame test.

Medical and biological importance of the alkali metals

Some lithium salts are used in the treatment of maniac depression. Sodium and

potassium ions are responsible for the electrical properties of the cell membranes,

the Na+/K+ proportion between the living cell and its environment serves as a base of

the membrane potential, signal transductions and some active transport processes.

In case of severe diarrhoea or application of diuretics there is a significant loss of


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potassium of the organism that must be supplied with medicines containing

potassium chloride.


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Alkaline earth metals

Elements in the second main group have s2 valence-shell electron configuration.

They are metals of low electronegativity and ionization energy having reducing

properties. Their electronegativity decreases from the top to the bottom while the

atomic radius increases. In compounds their characteristic oxidation number is +2.

element density (g/cm3) reactivityberyllium Be 1.85 TmagnesiumMg

1.74

calcium Ca 1.55strontium Sr 2.6barium Ba 3.5radium Ra 5.0 (radioactive)

They form mostly ionic compounds, but beryllium and magnesium

can occur in covalent molecules, too. They react with acids, water and

alcohols, though they are less reactive than the alkali metals.

Mg + 2 HCI MgCl2 + H2 Ca + 2 CH3OH Ca(CH30)2 + H2

Their most important compounds are the hydroxides acting as strong bases. Thesolubility of the hydroxides increases as the column is descended: Mg(OH)2 and

Ca(OH)2 are precipitates of low solubility; the other hydroxides are well soluble in

water. The saturated solution of calcium hydroxide is the lime water.

Solubility of their sulphates decreases as the column is descended, while their

carbonates are not soluble in water.

The in earth's crust their compounds are very abundant: CaC03 (calcium

carbonate: limestone, marble, calcite or chalk), CaC03 MgC03 (dolomite), CaS04

2 H20 (gypsum, alabaster), MgC03 (magnesite). The carbonates and sulphates of

strontium and barium also occur in nature. All of the isotopes of radium are

radioactive, in small amounts they occur in the uranium ores.

Magnesium and calcium

Both of them are silvery light metals. Magnesium is protected by its tough oxide

layer, therefore it can be stored in air. Calcium is covered by a thin calcium nitride

(Ca3N2) layer avoiding oxidation.

On ignition magnesium burns with a brilliant white light. Contrary to the other

alkali and alkaline earth metals its flame colouration is not in the visible region of the


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electromagnetic spectrum. Calcium's flame test is brick-red, that of strontium is

crimson and that of barium is green.


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Metallic magnesium is applied as a reducing agent, in the organic chemical

practice it is used to prepare organometallic compounds. Metallic calcium has an

importance in the preparation of anhydrous organic solvents. Calcium reacts with

water vigorously, while magnesium reacts only slowly at ambient temperature, at anelevated temperature the reaction is faster.

Ca + 2 H20 = Ca2+ + 2 OH" + H2

Mg + 2 H20 == MgiOH^ + H2

Calcium compounds

Calcium oxide (CaO), quick lime, is a base anhydride. It is a colourless caustic

solid. With water it yields the corresponding base, calcium hydroxide in a very

vehement exothermic reaction.

CaO+ H20 = Ca(OH)2 (lime slaking)

It is a starting material in building industry, ceramic industry and in

manufacturing of cement.

Calcium hydroxide, Ca(OH)2, (slaked lime, its aqueous solution is the lime

water) is a strong base with low water solubility. Lime water is a corrosive, caustic

agent, it destroys the tissues, especially avoid the contact of eyes with it.

In the laboratory practice the saturated solution of Ca(OH)2 is used to detect

carbon dioxide:

Ca(OH)2 + C02 = CaCOq+ H20 (a base reacts with an acid anhydride)

The same reaction happens at the hardening of lime and mortar (mixture of lime

and sand), applied in building industry. In addition calcium hydroxide is used for the

softening of water, manufacturing of paper and steel.

By dissolving copper(ll)sulphate (CuS04 5 H20) in lime water Bordeaux mixture

forms having fungicide effect used as a fungicide spray in agriculture.

Calcium carbonate (CaC03) occurs in nature as a mineral, it is a colourless water

insoluble crystalline substance. In addition it is widespread in the teeth, in bones, in

tartar, in egg-shell, in pearl and in scale. Calcium carbonate is the main component of

solid skeleton of shells, snails and corals. Thermal decomposition of limestone

(calcium carbonate) yields slaked lime, which is an indispensable material for the

building industry.

CaC03 ====== CaO + C02 (~8-900C, lime burning)

Acids dissolve it releasing carbon dioxide.

CaCQ3 + 2 H+ = Ca2+ + H20 + C02 (strong acid liberates the weaker one from its

salt)


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This is the theoretical base of the scaling, in the household the most widely used

scale removing agents contain phosphoric or citric acid. Calcium carbonate is also

applied to neutralize the excess of gastric acid.

Calcium carbonate also dissolves in carbonic acid forming calcium

hydrogen carbonate. CaCQ3 + H2C03 ===== Ca2+ + 2 HC03"


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In this way carbon dioxide containing ground water can excavate holes and caves

in the limestone mountains.

The temporary hardness of natural waters is bought about by their calcium

hydrogencarbonate and magnesium hydrogencarbonate content. On heating hardwater the equilibrium above is shifted to the left, i.e. insoluble calcium carbonate

(limescale) forms. This process has a role in the formation of stalagmites and

stalactites.

Ca(HC03)2 ===== CaCO, + H20 + C02

The permanent hardness of waters arises from other dissolved calcium and

magnesium salts (chlorides, sulphates) that can not be broken down by heating.

Though drinking of hard water (till 30 German degrees*) physiologically is preferable,

but it is less appropriate for household and industrial purposes: on heatingprecipitate (scale, fur) forms, the traditional soaps also form a scum, precipitate with

calcium ions without cleaning activity.

Calcium sulphate (CaS04 2 H20), gypsum or alabaster is a colourless

crystalline water insoluble mineral. On heating it yields a derivative, calcium

sulphate hemihydrate (CaS04 1/4H20) containing half mole of crystalline water. This

substance when mixed with water liberates heat and then hardens.

2 CaS04 J4 H20 + 3 H20 = 2 CaS04 2 H20

It is applied in surgery for plaster casts, in dentistry for mounting casts and inaddition for making statues, mouldings. Special mortars also contain plaster (CaS04

34 H20).

Calcium chloride (CaCl2) is a colourless crystalline water-soluble compound. It is

used as a desiccant for drying liquids and gases (see Drying p.13.), as well as mixed

with ice it forms freezing mixtures (the CaCl2-ice mixture has as low temperature as

-50C). In winter streets it is applied as an defrosting agent harmless to the

environment.

The calcium phosphates are not soluble in water. They play important role in

the organism as components of the bone.

Magnesium compounds

The magnesium oxide or magnesia (MgO) is a colourless powder, insoluble in

water. It reacts with the carbon dioxide content of the air yielding magnesium

carbonate.

MgO + C02 = MgC03 (reaction of a base anhydride with an acid anhydride)

MgO is an ingredient of tooth-pastes, baby powders and antacids, in addition

athletes apply it as a grip improving agent, while in the chromatographic practice it

is used as a basic stationary phase.


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The water hardness is expressed in German degrees (GD). 1 GD hardness of water

means the amount of calcium or magnesium salts in 100 g water equivalent with 1 mg of CaO

(i.e. it contains 1.78 10"2 mmol Ca2+ and/or Mg2+-ion).


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Magnesium hydroxide (Mg(OH)2) is a water insoluble colourless crystalline

material acting as a strong base. It serves as a starting material for the preparation

of the other magnesium compounds.

Magnesium sulphate (bitter salt, MgS04 7 H20) is a colourless crystallinewater soluble salt of bitter taste. Anhydrous magnesium sulphate is applied as a

drying agent for liquids.

Medical, biological importance of the alkaline earth metals

The water soluble compounds of beryllium and barium are toxic, while magnesium

and calcium ions are essential for the living organisms. The toxic effect of

beryllium(ll) probably connected to the fact that it replaces magnesium in the

magnesium containing enzymes.

Magnesium ion is the constituent of the prosthetic group of chlorophyll playing

essential role for the green plants. In the humans it functions as an activator of

enzymes and takes part in the electrolytic balance of the body fluids. It catalyzes the

hydrolysis of the phosphates (e.g. ATP, nucleotides and phosphoproteines, see.

Bioinorganic chemistry). It can reduce the cramps, muscle contractions; therefore it

is recommended a higher intake for sportsman and pregnant women.

The magnesium sulphate (MgS04, bitter salt), an important component of mineral

waters, is applied as a laxative in medicine.

The majority of calcium of the body is deposited in the bones or in the teeth as

water insoluble hydroxyapatite (Ca5(P04)3OH) or fluorapatite (Ca5(P04)3F). Calcium ion

plays a vital role in several physiological processes, e.g. muscle contraction, blood

clotting, because it is a secondary messenger and activator of several enzymes.

Calcium carbonate (CaC03) has a medical role as calcium supplement and antacid.

The basic magnesium carbonate (3MgC03 Mg(OH)2 3 H20) is a water insoluble

powder acting as a base, it has a similar application. Acids dissolve it by releasing

carbon dioxide; it has a mild laxative effect, too.

Strontium and its compounds have no chemical and biological importance.

Strontium has radioactive isotopes, e.g. the90

Sr is a by-product of nuclear fission ofuranium (explosion of uranium bomb). It can substitute for calcium in bone; the

isotope is a (B-emitter causing bone sarcoma.

The water soluble barium compounds are toxic. The water insoluble barium

sulphate (BaS04) suspension is used as an X-ray contrast medium (barium-cocktail) of

the gastrointestinal system, it absorbs X-rays but it is not absorbed from the stomach

at all.

The alkali and alkaline earth metals ions are excreted from the body through the

urine and sweat. At an important loss of liquid or at surgical event the so called

Ringer solution -in infusion- is applied as a supplement of these ions. This solution


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has a physiological osmotic pressure and contains Na+, K+, Ca2+ and Mg2+-ions with

chloride and hydrogencarbonate counterions.


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Self-test questions

1. Write names and chemical symbols for alkali metals. Give the valence shell

configuration of the alkali metals. How many valence electrons do they have?

2. Write names and chemical symbols for alkaline earth metals. Give the valenceshell configuration of the alkaline earth metals. How many valence electrons do

they have?

3. Compare the electronegativity, ionization energy and atomic radii of the s-block

elements. Characterize the reactivity of the s-block elements. What is their

characteristic oxidation number?

4. Characterize the water solubility of the hydroxides, sulphates and carbonates of

the s-block metals.

water solubility alkali metals alkaline earth metals

hydroxiodes

sulphates

carbonates

5. Describe of the chemical properties of the s-field elements. Write equation for the

reaction of sodium and water.

6. Describe the solubility and the chemical properties of the alkali metal hydroxides.

What happens to sodium hydroxide stored in the air? Write the equation.

7. Write name and empirical formula for three minerals of calcium and magnesium.

8.Write equation for the reaction of calcium with water. How is it possible to detect

the products?

9. Write equation for the reaction of calcium oxide with water. For what purposes

are calcium oxide and calcium hydroxide applied in the everyday life?

10. How is it possible to detect carbon dioxide in the laboratory practice? Write an

equation.

11. What kind of reaction occurs, if the mortar hardens? Write an equation.

12. 12. How does calcium carbonate react with the gastric juice (hydrochloric

acid)? Write an equation. What kind of forms of calcium carbonate can be found

in the nature or in our organism?13. 13. Which alkaline earth metal compounds are applied as antacids (names

and formulas)?

14. How does limestone react with carbonic acid? Write an equation. Which ions

are responsible for the temporary hardness of water?

15. Write equations for the reactions describing the formation and dissolution of

limescale.

16. Write equation for the thermal decomposition of calcium carbonate. What is

the practical importance of the product?


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17. What is calcium sulphate hemihydrate applied in the medical practice for?

Give an equation showing its reaction with water.

18. Give the chemical composition of the following substances: lime water,

gypsum, chalk, Bordeaux mixture, quick lime.


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Experiments

Reaction of sodium with water (demonstration)

Materials: metallic sodium, Phenolphthalein solution, distilled water

Procedure: Fill half a crystallisation dish of ~25 cm diameter with water and add

some drops of Phenolphthalein solution to it. Take a small piece (pea size) of sodium

and remove sodium oxide coat from it. Drop the shiny piece of sodium into the dish.

Observe the behaviour of the sodium, its situation, its change in shape. Write

equation. What kind of conclusions can you draw connected to the density and

melting point of sodium?

Comments: In case of fire caused by alkali metals water is forbidden to use for fire-

fighting, since it just adds fuel to the fire. Only dry-powder estinguisher can be used

(burning magnesium reacts with C02/ thus estinguisher with carbon dioxide are not

convenient!)

Reaction of calcium with water

Materials: metallic calcium, distilled water, Phenolphthalein solution

Procedure: Fill a test tubes to about half with distilled water and add some drops of

Phenolphthalein solution and a small piece of calcium (pea size). What can you

observe? Explains your observations, write equation.

Comments: The reaction is less violent than with sodium or potassium. Calcium has

to be stored not under petroleum only in a well sealed bottle.

Reaction of magnesium with water

Materials: magnesium strip, distilled water, Phenolphthalein solution

Procedure: Fill a beaker to about half with distilled water and add 1-2 drops of

Phenolphthalein solution. Burnish a piece of magnesium strip shiny (by sandpaper)

and drop it into the beaker. Observe the reaction mixture for some minutes. Then

start heating it on hot plate. When the water is hot enough finish heating and

extensively study again the reaction mixture. What do you observe? What is the

turbidity caused by? Explain your observations, write equations.

Flame test with alkali metals and alkaline earth metals

Szksges vegyszerek: ntrium-klorid oldat, klium-klorid oldat, kalcium-klorid oldat,

stroncium-Materials: sodium chloride, potassium chloride, calcium chloride, strontium

chloride and barium chloride solutions, 20% hydrochloric acid, zinc granules

Procedure: Pour 1-1 ml of the corresponding sample into a porcelain crucible and

fulfill it almost with 20 % hydrochloric acid solution. Place 1-2 granules of zinc into

the crucible. Place the crucible onto a clay triangle supported by an iron tripod (see


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figure 5.19.) The nonluminous flame of the Bunsen burner held horizontally above the

intensively bubbling solution. What do you observe? Observations:

Na: K: Ca: Sr: Ba: Mg:


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Figure 5.21. Flame test (The figure was drawn by

Sachie Saito 1st

year medical student)

Comments: Light emitted by excited alkali metal atoms has a characteristic line

spectrum, which can be made monochromatic by appropriate filters. This is used in

sodium lamps (polarimetry). In everyday life sodium lamps of very intense yellow

light are used for the night illumination of busy road-crossings. By investigating the

spectrum in atomic absorption spectrophotometry the elemental composition of a

sample can be established.

Examination of the composition of egg shell

Materials: marble, Phenolphthalein solution, distilled water, eggshell, dilute

hydrochloric acid Procedure: a. Place a piece of marble into a porcelain dish and pour

onto it some water and 2 drops of indicator. What do you experience? Why?

b. Into a test-tube pour 2-3 cm3 of dilute hydrochloric acid solution and place a

piece of marble into it. What do you notice?

c. Hold a piece of marble with tweezers into the darting flame of Bunsen burner

for 2-3 minutes. Cool it down in the air and drop it into a test-tube containing distilled

water and add some drops of Phenolphthalein solution to the test-tube. Explain your

experiences, write equations.d. Into three test-tube pour small amounts of dilute hydrochloric acid solution,

distilled water and distilled water with one drop of Phenolphthalein solution,

respectively. To each test-tube add one-one piece of eggshell. What do you notice?

e. Heat a piece of eggshell thoroughly with Bunsen burner. After cooling place it

into a test-tube containing distilled water with one drop of Phenolphthalein. What do

you observe?

Based on the experiments performed what is your conclusion connected to the

composition of the eggshell? State your reason.


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Comments: The thermal decomposition of calcium carbonate occurs during lime-

burning, too. Its product - quick lime - reacts with water affording slaked lime, theprocedure is referred to as lime slaking.


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Reaction of calcium hydroxide with carbon dioxide

Materials: lime water, carbon dioxide (from Kipp's apparatus)

Procedure: a. Pour 5-6 ml distilled water into a beaker about and introduce into

the water carbon dioxide from Kipp's apparatus by a gas inlet tube. What do you

observe? Give the reaction equation. b. Continue the gas introduction until there

is a no more change (~ 2 min). Then heat the reaction mixture on a hot plate for

5 minutes and notice the change. What can you observe? Give the reacton

equation. How could you clean the beaker? Suppose a reagent with reaction

equation. Explanation: Calcium hydroxide as a strong base reacts with carbon

dioxide (acid anhydride). The first product, calcium carbonate will be converted to

a water soluble compound by the further treatment with carbon dioxide (carbonic

acid). Heating of the formed calcium hydrogencarbonate solution yields again

calcium carbonate (scale). Write equation to each step.

Comments: The dissolution of limestone and the formation of cave decorations

(stalactites and stalagmites) are based on the same principle.

Documents

Earth Metals