Chemistry

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General Characteristics of the Compounds of the Alkali Earth Metals

The dipositive oxidation state (M²⁺) is the predominant valence of group 2 elements. The alkaline earth metals form compounds which are predominantly ionic. However, they are less ionic than the corresponding compounds of alkali metals. This is due to increased nuclear charge and smaller size. The general characteristics of some of the compounds of alkaline earth metals are described below.

(a) Oxides:

Generally alkaline earth metals form monoxides and peroxides.

Monoxides

Monoxides are obtained by heating the metals in oxygen. BeO and MgO are almost insoluble in water. On the other hand, oxides of other elements form hydroxides. BeO is amphoteric; MgO is weakly basic while CaO, SrO and BaO are strongly basic.

BeO oxide is covalent due to the small size of Be²⁺ ion, while other oxides are ionic in nature.

Peroxides

Except beryllium, all the remaining metals form peroxides. It is prepared by heating monoxides with oxygen at high temperature.

2 BaO + O₂ → 2 BaO₂

(b) Hydroxides:

All the oxides except BeO are basic in nature and react with water to form sparingly soluble hydroxides.

MO + H₂O → M(OH)₂

The solubility, thermal stability and the basic character of the hydroxides increase down the group. The alkaline earth metal hydroxides are, however, less basic and less stable than alkali metal hydroxides. Beryllium hydroxide is amphoteric in nature as it reacts with both acid and alkali.

Be(OH)₂ + 2 NaOH → Na₂BeO₂ + 2H₂O
Be(OH)₂ + 2HCl → BeCl₂ + 2H₂O

(c) Halides:

Alkaline earth metals form halides with general formula MX₂. They can be prepared by heating metals with halogens on heating.

M +X₂ → MX₂

Beryllium halides are covalent on account of smaller size of Be²⁺. Beryllium halides are hygroscopic, fume in moist air and soluble in organic solvents. Beryllium chloride has a chain structure in the solid state as shown in figure 5.9 (structure-a). In the vapour phase BeCl₂ tends to form a chloro-bridged dimer (structure-c) which dissociates into the linear monomer at high temperatures of the order of 1200 K. (structure-b).

Except beryllium halides, all the other halides of alkaline earth metals are ionic in nature. Chloride and fluorides of the other metals are ionic solids. These are good conductors of electricity in fused state and inaqueous solutions. The tendency to form halide hydrates gradually decreases (for example, MgCl₂.8H₂O, CaCl₂.6H₂O, SrCl₂.6H₂O and BaCl₂.2H₂O) down the group.

Salts of Oxo Acids

The alkaline earth metals form salts of oxo acids. Some of these are given below:

Carbonates:

All the carbonates decompose on heating to give carbon dioxide and the oxide.

• The solubility of carbonates in water decreases down the group.
• The thermal stability increases down the group with increasing cationic size.

Decomposition temperature of alkaline metal carbonates and sulphates

The sulphates of the alkaline earth metals are all white solids and stable to heat. BeSO₄, and MgSO₄ are
readily soluble in water; the solubility decreases from CaSO₄ to BaSO₄. The greater hydration enthalpies of
Be²⁺ and Mg²⁺ ions overcome the lattice enthalpy factor and therefore their sulphates are soluble in water.

Nitrates:

The nitrates are made by dissolution of the carbonates in dilute nitric acid. Magnesium nitrate crystallises with six molecules of water, whereas barium nitrate crystallises as the anhydrous salt. This again shows a decreasing tendency to form hydrates with increasing size. All of them decompose on heating to give the oxide.

Important Compounds of Calcium

Quick lime, CaO

Preparation

It is produced on a commercial scale by heating limestone in a lime kiln in the temperature range 1070-1270K.

CaCO₃ ⇄ CaO + CO₂

The reaction being reversible, carbon dioxide is removed as soon as it is produced to enable the reaction to proceed to completion.

Properties

Calcium oxide is a white amorphous solid. It has a melting point of 2870 K.

(i) It absorbs moisture and carbon dioxide on exposure to atmosphere.

CaO + H₂O → Ca(OH)₂
CaO + CO₂ → CaCO₃

(ii) The addition of limited amount of water breaks the lump of lime. This process is called slaking of lime and the product is slaked lime.

CaO + H₂O → Ca(OH)₂

(iii) The mixture of Quick lime(CaO) and sodium hydroxide is called soda lime.

(iv) It combines with acidic oxides such as SiO₂ and P₄O₁₀ to form CaSiO₃ and Ca₃(PO₄)₂, respectively.

CaO + SiO₂ → CaSiO₃
6CaO + P₄O₁₀ → 2Ca₃(PO₄)₂

Uses

Calcium oxide is used

1. To manufacture cement, mortar and glass.
2. In the manufacture of sodium carbonate and slaked lime.
3. In the purification of sugar.
4. As a drying agent.

Calcium Hydroxide

Preparation

Calcium hydroxide is prepared by adding water to quick lime, CaO.

Properties

It is a white powder. It is sparingly soluble in water. The aqueous solution is known as lime water and a suspension of slaked lime in water is known as milk of lime.

When carbon dioxide is passed through lime water, it turns milky due to the formation of calcium carbonate.

Ca(OH)₂ + CO₂ → CaCO₃ + H₂O

On passing excess of carbon dioxide, the precipitate dissolves to form calcium hydrogen carbonate.

CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂

Milk of lime reacts with chlorine to form hypochlorite, a constituent of bleaching powder.

2Ca (OH)₂ + 2Cl₂ →
CaCl₂ + Ca(OCl₂) + 2H₂O

Uses:

Calcium hydroxide is used

1. In the preparation of mortar, a building material.
2. In white wash due to its disinfectant nature.
3. In glass making, in tanning industry, in the preparation of bleaching powder and for the purification of sugar.

Gypsum (CaSO₄.2H₂O)

Gypsum beds were formed due to the evaporation of water from the massive prehistoric sea basins. When water evaporates, the minerals present in it become concentrated, and crystallise.

Properties of Gypsum

Gypsum is a soft mineral, which is moderately soluble in water. The solubility of this mineral in water is affected by temperature. Unlike other salts, gypsum becomes less soluble in water as the temperature increases. This is known as retrograde solubility, which is a distinguishing characteristic of gypsum.

Gypsum is usually white, colorless, or gray in color. But sometimes, it can also be found in the shades of pink, yellow, brown, and light green, mainly due to the presence of impurities.

Gypsum crystals are sometimes found to occur in a form that resembles the petals of a flower. This type of formation is referred to as ‘desert rose’, as they mostly occur in arid areas or desert terrains.

Gypsum is known to have low thermal conductivity, which is the reason why it is used in making drywalls or wallboards. Gypsum is also known as a natural insulator.

Alabaster is a variety of gypsum, that is highly valued as an ornamental stone. It has been used by the sculptors for centuries. Alabaster is granular and opaque.

Gypsum has hardness between 1.5 to 2 on Moh’s Hardness Scale. Its specific gravity is 2.3 to 2.4.

Uses of Gypsum

1. The alabaster variety of gypsum was used in ancient Egypt and Mesopotamia by the sculptors. The ancient Egyptians knew how to turn gypsum into plaster of Paris about 5,000 years ago. Today, gypsum has found a wide range of uses and applications in human society, some of which are enlisted below.

2. Gypsum is used in making drywalls or plaster boards. Plaster boards are used as the finish for walls and ceilings, and for partitions.

3. Another important use of gypsum is the production of plaster of Paris. Gypsum is heated to about 300 degree Fahrenheit to produce plaster of Paris, which is also known as gypsum plaster. It is mainly used as a sculpting material.

4. Gypsum is used in making surgical and orthopedic casts, such as surgical splints and casting moulds.

5. Gypsum plays an important role in agriculture as a soil additive, conditioner, and fertilizer. It helps loosen up compact or clay soil, and provides calcium and sulphur, which are essential for the healthy growth of a plant. It can also be used for removing sodium ion from soils having excess salinity.

6. Gypsum is used in toothpastes, shampoos, and hair products, mainly due to its binding and thickening properties.

7. Gypsum is a component of Portland cement, where it acts as a hardening retarder to control the speed at which concrete sets.

8. To sum up, gypsum is one of the most abundant minerals that have endless uses and applications. Mining of gypsum is simple and easy, as the mineral occurs in large thick beds near the Earth’s surface. However, large-scale mining of gypsum involves considerable damage to the environment. Gypsum can also be recycled, but not much importance has been given to recycle this mineral due to its abundance.

Plaster of Paris

Calcium Sulphate Hemihydrate:

CaSO₄·½ H₂O (Plaster of Paris)

It is a hemihydrate of calcium sulphate. It is obtained when gypsum, CaSO₄.2H₂O is heated to 393 K.

2CaSO₄.2H₂O(s) → 2 CaSO₄. ½ H₂O + 3H₂O

Above 393 K, no water of crystallisation is left and anhydrous calcium sulphate, CaSO₄ is formed. This is known as ‘dead burnt plaster’.

It has a remarkable property of setting with water. On mixing with an adequate quantity of water it forms a plastic mass that gets into a hard solid in 5 to 15 minutes.

Uses:

Plaster of Paris is used as/in,

1. The building industry as well as plasters.
2. For immobilising the affected part of organ where there is a bone fracture or sprain.
3. Employed in dentistry, in ornamental work and for making casts of statues and busts.

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Alkali Earth Metals

Group 2 in the modern periodic table contains the elements beryllium, magnesium, calcium, strontium, barium and radium. These elements with the exception of beryllium are commonly known as the alkaline earth metals because their oxides and hydroxides are alkaline in nature and these metal oxides are found in the earth’s crust.

General Characteristics of Alkaline Earth Metals

Physical State

Beryllium is rare and radium is the rarest of all comprising only 10 % of igneous rocks. Magnesium and calcium are very common in the earth’s crust, with calcium the fifth-most-abundant element, and magnesium the eighth. Magnesium and calcium are found in many rocks and minerals: magnesium in carnallite, magnesite, dolomite and calcium in chalk, limestone, gypsum.

Most strontium is found in the minerals celestite and strontianite. Barium is slightly less common, much of it in the mineral barite. Radium, being a decay product of uranium, is found in all uranium-bearing ores.

Electronic Configuration

These elements have two electrons in the valence shell of their atoms, preceded by the noble gas configuration. Their general electronic configuration is written as [Noble gas]ns² where ‘n’ represents the valence shell.

Atomic and Ionic Radii

The atomic and ionic radii of alkaline earth metals are smaller than the corresponding members of the alkali metals. This is due to the fact the Group 2 elements having a higher nuclear charge that allows electrons to be attracted more strongly towards the nucleus. On moving down the group, the radii increases due to gradual increase in the number of the shells and the screening effect.

Common Oxidation State

The group 2 elements have two electrons in their valence shell and by losing these electrons, they acquire the stable noble gas configuration. So these elements exhibit +2 oxidation state in their compounds.

Ionisation Enthalpy

Due to a fairly large size of the atoms, alkaline earth metals have low ionisation enthalpies when compared to ‘p’ block elements. Down the group the ionisation enthalpy decreases as atomic size increases. This is due to the addition of new shells as well as increase in the magnitude of the screening effect of inner shell electrons. Members of group 2 have higher ionization enthalpy values than group 1 because of their smaller size, with electrons being more attracted towards the nucleus of the atoms. Correspondingly they are less electropositive than alkali metals.

Although IE₁ values of alkaline earth metals are higher than that of alkali metals, the IE₂ values of alkaline
earth metals are much smaller than those of alkali metals. This occurs because in alkali metals the second electron is to be removed from a cation, which has already acquired a noble gas configuration. In the case of alkaline earth metals, the second electron is to be removed from a monovalent cation, which still has one electron in the outermost shell. Thus, the second electron can be removed more easily in the case of group 2 elements than in group 1 elements.

Hydration Enthalpies

Compounds of alkaline earth metals are more extensively hydrated than those of alkali metals, because the hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions.

Like alkali metal ions, the hydration enthalpies of alkaline earth metal ions also decrease with increase in ionic size down the group.

Be > Mg > Ca > Sr > Ba

e.g., Magnesium chloride and calcium chloride exist as their hydrated crystals MgCl₂.6H₂O and CaCl₂.6H₂O
respectively. whereas NaCl and KCl do not form such hydrates.

Electronegativity

In alkaline earth metals the electronegativity values decrease as we go down the group as seen in the alkali metals.

Flame Colour and the Spectra:

When the alkaline earth metal salts moistened with concentrated hydrochloric acid are heated on a platinum wire in a flame, they show characteristic coloured flame as shown below.

Table 5.10 Flame Color and Wavelength

The heat in the flame excites the valence electron to a higher energy level. when it drops back to its actual energy level, the excess energy is emitted as light, whose wavelength is in the visible region as shown in the above table.

Distinctive Behaviour of Beryllium

The anomalous properties of beryllium is mainly due to its small size, high electronegativity, high ionisation energy and high polarising power compared to the other elements in the block. The anomalous properties of beryllium compared to other elements of the group are mentioned in Table 5.11

Diagonal Relationship:

As observed in alkali metals, beryllium (the first member of group 2) shows a diagonal relationship with aluminium. In this case, the size of these ions (rBe²⁺ = 0.45 Å and r_Al³⁺ = 0.54 Å) is not as close. However, their charge per unit area is closer (Be²⁺ = 2.36 and Al³⁺ = 2.50). They also have same electronegativity values (Be = 1.5; Al = 1.5).

Table 5.12 Similarities Between Beryllium and Aluminium

Chemical Properties of Alkaline Earth Metals

The alkaline earth metals are less reactive than the alkali metals. The reactivity of these elements increases on going down the group.

Reactivity Towards the Halogens:

All the alkaline earth metals combine with halogen at elevated temperatures to form their halides.

M + X₂ → MX₂
(M = Be, Mg, Ca, Sr, Ba, Ra, X = F, Cl, Br, l)

Thermal decomposition of (NH₄)₂BeF₄ is the best route for the preparation of BeF₂. BeCl₂ is conveniently made from the oxide.

Reactivity Towards Hydrogen:

All the elements except beryllium, combine with hydrogen on heating to form their hydrides with general formula MH₂. BeH₂ can be prepared by the reaction of BeCl₂ with LiAlH₄.

2BeCl₂ + LiAlH₄ → 2BeH₂ + LiCl + AlCl₃

Uses of Alkaline Earth Metals

Uses of Beryllium

1. Because of its low atomic number and very low absorption for X-rays, it is used as radiation windows for X-ray tubes and X-ray detectors.
2. The sample holder in X-ray emission studies usually made of beryllium
3. Since beryllium is transparent to energetic particles, it is used to build the ‘beam pipe’ in accelerators.
4. Because of its low density and diamagnetic nature, it is used in various detectors.

Uses of Magnesium

1. Removal of sulphur from iron and steel.
2. Used as photoengrave plates in printing industry.
3. Magnesium alloys are used in aeroplane and missile construction.
4. Mg ribbon is used in synthesis of Grignard reagent in organic synthesis.
5. It alloys with aluminium to improve its mechanical, fabrication and welding property.
6. As a desiccant.
7. As sacrificial anode in controlling galvanic corrosion.

Uses of Calcium

1. As a reducing agent in the metallurgy of uranium, zirconium and thorium.
2. As a deoxidiser, desulphuriser or decarboniser for various ferrous and non-ferrous alloys.
3. In making cement and mortar to be used in construction.
4. As a getter in vacuum tubes.
5. In dehydrating oils
6. In fertilisers, concrete and plaster of paris.

Uses of Strontium

1. ⁹⁰Sr is used in cancer therapy.
2. ⁸⁷Sr/⁸⁶Sr ratios are commonly used in marine investigations as well as in teeth, tracking animal migrations or in criminal forensics.
3. Dating of rocks.
4. As a radioactive tracer in determining the source of ancient archaeological materials such as timbers and coins.

Uses of Barium

1. Used in metallurgy, its compounds are used in pyrotechnics, petroleum mining and radiology.
2. Deoxidiser in copper refining.
3. Its alloys with nickel readily emits electrons hence used in electron tubes and in spark plug electrodes.
4. As a scavenger to remove last traces of oxygen and other gases in television and other electronic tubes.
5. An isotope of barium 133Ba, used as a source in the calibration of gamma ray detectors in nuclear chemistry.

Uses of Radium

Used in self-luminous paints for watches, nuclear panels, aircraft switches, clocks and instrument dials.

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Biological Importance of Sodium and Potassium

Monovalent sodium and potassium ions are found in large proportions in biological fluids. These ions perform important biological functions such as maintenance of ion balance and nerve impulse conduction. A typical 70 kg man contains about 90 g of sodium and 170 g of potassium compared with only 5 g of iron and 0.06 g of copper.

Sodium ions are found primarily on the outside of cells, being located in blood plasma and in the interstitial fluid which surrounds the cells. These ions participate in the transmission of nerve signals, in regulating the flow of water across cell membranes and in the transport of sugars and amino acids into cells.

Sodium and potassium, although so similar chemically, differ quantitatively in their ability to penetrate cell membranes, in their transport mechanisms and in their efficiency to activate enzymes.

Thus, potassium ions are the most abundant cations within cell fluids, where they activate many enzymes, participate in the oxidation of glucose to produce ATP and, with sodium, are responsible for the transmission of nerve signals. Sodium-potassium pump play an important role in transmitting nerve signals.

Sodium maintains the electrolyte balance in the body. Potassium ions are primarily found inside the cell. Potassium ions maintain the osmolarity (the concentration of a solution expressed as the total number of solute particles per litre) of the cell. They also regulate the opening and the closing of the stomata.

Sodium is found mainly in body fluids. It plays a major role in maintaining blood volume and blood pressure by attracting and holding water. Sodium is also important in cellular osmotic pressure (the passage of fluids in and out of the cells) and in transmitting nerve impulses.

Potassium is the main intracellular ion for all types of cells, while having a major role in maintenance of fluid and electrolyte balance. Potassium is necessary for the function of all living cells, and is thus present in all plant and animal tissues.

Potassium and sodium are electrolytes that help your body maintain fluid and blood volume so it can function normally. However, consuming too little potassium and too much sodium can raise your blood pressure.

Calcium and Magnesium are alkaline earth metals, which have become a necessity in our everyday lives and elements which sustain the human body and help us function properly. These two elements are extremely effective alloy mediums and are used in different industries to serve various purposes.

Function:

The body uses sodium to control blood pressure and blood volume. Your body also needs sodium for your muscles and nerves to work properly.

Along with sodium, potassium regulates the water balance and the acid-base balance in the blood and tissues, and plays a critical role in the transmission of electrical impulses in the heart. The active transport of potassium into and out of the cells is crucial to cardiovascular and nerve function.

The more potassium you eat, the more sodium you lose through urine. Potassium also helps to ease tension in your blood vessel walls, which helps further lower blood pressure. Increasing potassium through diet is recommended in adults with blood pressure above 120/80 who are otherwise healthy.

Getting enough potassium from your diet can help you maintain healthy nerve function. Summary: This mineral plays an essential role in activating nerve impulses throughout your nervous system. Nerve impulses help regulate muscle contractions, the heartbeat, reflexes and many other processes.

Potassium is necessary for the normal functioning of all cells. It regulates the heartbeat, ensures proper function of the muscles and nerves, and is vital for synthesizing protein and metabolizing carbohydrates.

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General Characteristics of the Compounds of Alkali Metals

All the common compounds of the alkali metals are generally ionic in nature. General characteristics of some of their compounds are discussed here.

Oxides and Hydroxides

On combustion in excess of air, alkali metals forms normal oxides with formula M₂O. They react with water to form corresponding hydroxides which are basic in nature.

M₂O + H₂O → 2 MOH

Alkali metals apart from lithium form peroxides in addition to normal oxides upon combustion with excess air. These peroxides produce hydroxides and H₂O₂ upon reacting with water.

M₂O₂ + 2 H₂O → 2MOH + H₂O₂
(M = Na, K, Rb, Cs)

Except lithium and sodium, all the other alkali metals form superoxides also. These superoxides also gives basic hydroxides upon treatment with water.

2 MO₂ + 2OH₂ → 2 MOH + H₂O₂ + O₂
(M = K, Rb, Cs)

Under appropriate conditions pure compounds M₂O, M₂O₂ or MO₂ may be prepared.

Properties of Oxides and Hydroxides:

The oxides and the peroxides are colourless when pure, but the superoxides are yellow or orange in colour. The peroxides are diamagnetic while the superoxides are paramagnetic. Sodium peroxide is widely used as an oxidising agent.

The hydroxides which are obtained by the reaction of the oxides with water are all white crystalline solids. The alkali metal hydroxides are strong bases. They dissolve in water with evolution of heat on account of intense hydration.

Halides:

The alkali metal halides, MX, (X=F, Cl, Br, I) are colourless crystalline solids with high melting points. They can be prepared by the reaction of the appropriate oxide, hydroxide or carbonate with aqueous hydrohalic acid (HX).

As the electropositive character of alkali metal increases from Li to Cs, the ease with which the metals form halides increases from Li to Cs. All halides are ionic in nature except LiBr and LiI. Except LiF, all other halides are soluble in water. The low solubility of LiF in water is due to its high lattice enthalpy (small size of Li⁺ and F^–). Due to the presence of covalent nature both LiBr and LiI are soluble in organic solvents.

Salts of Oxo-Acids

Alkali metals form salts with all the oxo-acids. Most of these salts are soluble in water and are thermally stable. As the electropositive character increases down the group, the stability of the carbonates and bicarbonates increases.

This is due to the decrease in polarising power of alkali metal cations. The carbonates (M₂CO₃) of alkali metals are remarkably stable up to 1273 K, above which they first melt and then eventually decompose to form oxides. However, Li₂CO₃ is considerably less stable and decomposes readily.

This is presumably due to large size difference between Li⁺ and CO₃^2- which makes the crystal lattice unstable. Being strongly basic, alkali metals except lithium form solid bicarbonates. No other metal forms solid bicarbonates.

M₂CO₃ + CO₂ + H₂O → 2 MHCO₃
(M = Na, K, Rb, Cs)

Important Compounds of Alkali Metals:

Sodium Carbonate Na₂CO₃.10H₂O (Washing Soda):

Sodium carbonate is one of the important inorganic compounds used in industries. It is prepared by Solvay process. In this process, ammonia is converted into ammonium carbonate which then converted to ammonium bicarbonate by passing excess carbon dioxide in a sodium chloride solution saturated with ammonia.

The ammonium bicarbonate thus formed reacts with the sodium chloride to give sodium bicarbonate and ammonium chloride. As sodium bicarbonate has poor solubility, it gets precipitated. The sodium bicarbonate is isolated and is heated to give sodium carbonate. The equations involved in this process are,

2NH₃ + H₂O + CO₂ → (NH₄)₂CO₃
(NH₄)₂CO₃ + H₂O + CO₂ → 2 NH₄HCO₃
NH₄HCO₃ + NaCl → NH₄Cl + NaHCO₃
2NaHCO₃ → Na₂CO₃ + CO₂ + H₂O

The ammonia used in this process can be recovered by treating the resultant ammonium chloride solution with calcium hydroxide. Calcium chloride is formed as a by-product.

Properties:

Sodium carbonate, commonly known as washing soda, crystallises as decahydrate which is white in colour. It is soluble in water and forms an alkaline solution. Upon heating, it looses the water of crystallisation to form monohydrate. Above 373 K, the monohydrate becomes completely anhydrous and changes to a white powder called soda ash.

Na₂CO₃.10H₂O → Na₂CO₃.H₂O + 9H₂O
Na₂CO₃.H₂O → Na₂CO₃ + H₂O

Uses:

1. Sodium carbonate known as washing soda and is used mainly for laundering.
2. It is also used in water treatment to convert the hard water to soft water.
3. It is used in the manufacturing of glass, paper, paint etc…

Sodium Chloride NaCl (Cooking Salt or Table salt):

Sodium chloride is isolated by evaporation from sea water which contains 2.7 to 2.9% by mass. Approximately 50 lakh tons of salt are produced annually in India by solar evaporation. Crude sodium chloride can be obtained by crystallisation of brine solution which contains sodium sulphate, calcium sulphate, calcium chloride and magnesium chloride as impurities. Pure sodium chloride can be obtained from crude salt as follows.

Firstly removal of insoluble impurities by filtration from the crude salt solution with minimum amount of water. Sodium chloride can be crystallised by passing HCl gas into this solution. Calcium and magnesium chloride, being more soluble than sodium chloride, remain in solution.

Sodium chloride melts at 1081K. It has a solubility of 36.0 g in 100 g of water at 273 K. The solubility does not increase appreciably with increase in temperature.

Uses:

1. It is used as a common salt or table salt for domestic purpose.
2. It is used for the preparation of many inorganic compounds such as NaOH and Na₂CO₃

Sodium Hydroxide:

Sodium hydroxide is prepared commercially by the electrolysis of brine solution in Castner-Kellner cell using a mercury cathode and a carbon anode. Sodium metal is discharged at the cathode and combines with mercury to form sodium amalgam. Chlorine gas is evolved at the anode. The sodium amalgam thus obtained is treated with water to give sodium hydroxide.

At Cathode:

Na⁺ + e^– → Na(amalgam)

At Anode:

Cl^– → ½ Cl₂↑+ e^–
2Na(amalgam) + 2H₂O → 2NaOH + 2Hg + H₂↑

Sodium hydroxide is a white, translucent and deliquescent solid, that dissolves in water to give a strong alkaline solution. It melts at 591 K. The sodium hydroxide solution at the surface reacts with the CO₂ in the atmosphere to form Na₂CO₃

Uses:

1. Sodium hydroxide is used in the purification of bauxite (ore of Aluminium) and petroleum refining
2. It is used in the textile industries for mercerising cotton fabrics
3. It is used in the manufacture of soap, paper and artificial silks.

Sodium Bicarbonate NaHCO₃ (Baking Soda):

Sodium hydrogen carbonate or sodium bicarbonate is used in baking cakes pastries etc. It is called so because it decomposes on heating to generate bubbles of carbon dioxide, leaving holes in cakes or pastries and making them light and fluffy.

This compound is prepared by saturating a solution of sodium carbonate with carbon dioxide. The white crystalline powder of sodium bicarbonate, being less soluble, precipitated out.

Uses:

1. Primarily used as an ingredient in baking.
2. Sodium hydrogen carbonate is a mild antiseptic for skin infections.
3. It is also used in fire extinguishers.

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Alkali Metals

The word “alkali” is derived from the word al-qalīy meaning the plant ashes, referring to the original source of alkaline substances. A water-extract of burnt plant ashes, called potash contain mainly potassium carbonate. Alkali metal group consists of the elements: lithium, sodium, potassium, rubidium, caesium and francium. They are all metals, generally soft and highly reactive. They form oxides and hydroxides and these compounds are basic in nature.

General Characteristics of Alkali Metals:

Alkali metals are highly reactive and are found in nature only as compounds. Rubidium and caesium are found associated in minute quantities with minerals of other alkali metals. Francium is radioactive and does not occur appreciably in nature. Francium is highly radioactive; its longest-lived isotope has a half-life of only 21 minutes.

Table 5.1 Abundance of important alkali metals and their sources

Electronic Configuration

The general valence shell electronic configuration of alkali metals is ns¹, where ‘n’ represents the period number.

Common Oxidation State

All these elements are highly electropositive in nature. They readily lose their valence electron to give monovalent cations (M⁺). Alkali metals have only one oxidation state which is +1.

Atomic and Ionic Radii

Being the first element of each period, alkali metals have the largest atomic and ionic radii in their respective periods. On moving down the group, there is an increase in the number of shells and, therefore, atomic and ionic radii increase. The monovalent ions (M⁺) are smaller than the respective parent atoms as expected.

Ionisation Enthalpy

Alkali metals have the lowest ionisation enthalpy compared to other elements present in the respective period. As we go down the group, the ionisation enthalpy decreases due to the increase in atomic size. In addition, the number of inner shells also increases, which in turn increases the magnitude of screening effect and consequently, the ionisation enthalpy decreases down the group.

The second ionisation enthalpies of alkali metals are very high. The removal of an electron from the alkali metals gives monovalent cations having stable electronic configurations similar to the noble gas. Therefore, it becomes very difficult to remove the second electron from the stable configurations already attained.

Hydration Enthalpy

Lithium salts are more soluble than the salts of other metals of group 1. eg. LiClO₄ is upto 12 times more soluble than NaClO₄. Other salts KClO₄, RbClO₄ and CsClO₄ have solubilities only 10^-3 times of that of LiClO₄. The high solubility of Li salts is due to strong solvation of small size Li⁺ ion.

Electronegativity:

Alkali metals have comparatively smaller value of electronegativity than the other elements in the respective period. When they react with other elements, they usually produce ionic compounds. For example, they react with halogens to form ionic halides.

Flame Colour and the Spectra:

When the alkali metal salts moistened with concentrated hydrochloric acid are heated on a platinum wire in a flame, they show characteristic coloured flame as shown below.

Table 5.4 Flame Colour and Wavelength

The heat in the flame excites the valence electron to a higher energy level. When it drops back to its actual energy level, the excess energy is emitted as light, whose wavelength is in the visible region as shown in the above table.

The distinctive behaviour of Li⁺ ion is due to its exceptionally small size, high polarising power, high hydration energy and non availability of d-orbitals.

Diagonal Relationship:

Similarity between the first member of group 1 (Li) and the diagonally placed second element of group 2 (Mg) is called diagonal relationship. It is due to similar size (r_Li⁺ = 0.766 Å and rMg²⁺ = 0.72 Å) and comparable electronegativity values (Li = 1.0; Mg = 1.2).

Chemical Properties of Alkali Metals

Alkali metals exhibit high chemical reactivity. The reactivity of alkali metals increases from Li to Cs, since the ionisation energy decreases down the group. All alkali metals are highly reactive towards the more electronegative elements such as oxygen and halogens. Some characteristic chemical properties of alkali metals are described below.

Reaction With Oxygen

All the alkali metals on exposure to air or oxygen burn vigorously, forming oxides on their surface. Lithium forms only monoxide, sodium forms the monoxide and peroxide and the other elements form monoxide, peroxide, and superoxides. These oxides are basic in nature.

4Li + O₂ → 2Li₂O (Simple Oxide)
2Na + O₂ → Na₂O₂ (Peroxide)
M + O₂ → MO₂ (Superoxide) (M=K, Rb, Cs)

Reaction with Hydrogen

All alkali metals react with hydrogen at about 673 K (lithium at 1073 K) to form the corresponding ionic hydrides. Reactivity of alkali metals with hydrogen decreases from Li to Cs.

2M + H₂ → 2M⁺H^–
(M = Li, Na, K, Rb, Cs)

The ionic character of the hydrides increases from Li to Cs and their stability decreases. The hydrides behave as strong reducing agents and their reducing nature increases down the group.

Reaction With Halogen

Alkali metals combine readily with halogens to form ionic halides MX. Reactivity of alkali metals with halogens increases down the group because of corresponding decrease in ionisation enthalpy.

2M + X₂ → 2 MX
(M = Li, Na, K, Rb, Cs) (X = F, Cl, Br, I)

All metal halides are ionic crystals. However Lithium iodide shows covalent character, as it is the smallest cation that exerts high polarising power on the iodide anion. Additionally, the iodide ion being the largest can be polarised to a greater extent by Li⁺ ion.

Reaction With Liquid Ammonia:

Alkali metals dissolve in liquid ammonia to give deep blue solutions that are conducting in nature. The conductivity is similar to that of pure metals (The specific conductivity of Hg is 10⁴ Ω^-1 and for sodium in liquid ammonia is 0.5 × 10⁴Ω^-1).

This happens because the alkali metal atom readily loses its valence electron in ammonia solution. Both the cation and the electron are ammoniated to give ammoniated cation and ammoniated electron.

M + (x + y)NH₃ → [M(NH_3x]⁺] + [e(NH₃)_y]^–

The blue colour of the solution is due to the ammoniated electron which absorbs energy in the visible region of light and thus imparts blue colour to the solution. The solutions are paramagnetic and on standing slowly liberate hydrogen resulting in the formation of an amide.

M⁺ + e^– + NH₃ → MNH₂ + ½H₂

In concentrated solution, the blue colour changes to bronze colour and become diamagnetic.

Reaction With Water:

Alkali metals react with water to give corresponding hydroxides with theliberation of hydrogen.

2 Li + 2 H₂O → 2 LiOH + H₂

They also react with alcohol, and alkynes which contain active hydrogens.

Reducing Activity:

Alkali metals can lose their valence electron readily hence they act as good reducing agents.

M_(s) → M⁺_(g) + e^–

Reaction With Carbon:

Lithium directly reacts with carbon to form the ionic compound, lithium carbide. Other metals do not react with carbon directly. However, when they are treated with compounds like acetylene they form acetelydes.

2Li + 2C → Li₂C₂

Uses of Alkali Metals:

1. Lithium metal is used to make useful alloys. For example with lead it is used to make ‘white metal’ bearings for motor engines, with aluminium to make aircraft parts, and with magnesium to make armour plates. It is used in thermonuclear reactions.

2. Lithium is also used to make electrochemical cells.

3. Sodium is used to make Na/Pb alloy needed to make Pb(Et)₄ and Pb(Me)₄. These organolead compounds
were earlier used as anti-knock additives to petrol, but nowadays lead-free petrol in use.

4. Liquid sodium metal is used as a coolant in fast breeder nuclear reactors.

5. Potassium has a vital role in biological systems. Potassium chloride is used as a fertilizer.

6. Potassium hydroxide is used in the manufacture of soft soap. It is also used as an excellent absorbent of carbon dioxide.

7. Caesium is used in devising photoelectric cells.