Anomalous Behaviour of Lithium

You are aware of the use of lithium in batteries. Aren’t you? That

means you are aware of what lithium is. No? Well, then let’s learn all

about this element in this chapter. We will see how it behaves

differently than the other group 1 elements. Let’s begin.

Brief about Lithium

Group 1 elements in the periodic table are the Alkali Metals. The first

element of group 1 is lithium. So, why does it behave so differently?

Well, the small size of the element is the reason for its anomalous

behaviour! Let’s see how.

Nature of the Element

Lithium is extremely electropositive in nature. This is the reason why

it can form covalent bonds. The polarization behaviour of its ion

somewhat in the same lines as magnesium ion. Therefore, there exists

a diagonal relationship between magnesium and lithium. There are a

number of reasons why this diagonal relationship exists between them.

Let us look at the various reasons below.

As we move from the top to bottom in a group, the electropositive

nature of the elements increases. However, when we move from left to

right in a period, this nature decreases. It is because of this reason that

we observe similarities in the properties between diagonal elements.

Browse more Topics under The S Block Elements

● Beryllium, Calcium and Magnesium

● Characteristics of the Compounds of the Alkali Earth Metals

● Characteristics of the Compounds of the Alkali Metals

● Group 1 Elements: Alkali Metals

● Group 2 Elements: Alkali Earth Metals

● Some Important Compounds of Sodium and Potassium

We know that the size of the atoms increases from top to bottom in a

group. Because of this, the polarizing power of the element decreases.

However, when we move from the left to right in a period, this

polarizing power keeps on increasing. This is another reason for the

diagonal elements to show similar properties. Thus, we see that

lithium is a strong element that is quite similar to magnesium.

Their ions have comparable boiling and melting points. Due to its

small size, the atom possesses high ionization energy. It reacts with

water and liquid bromine. This forms a highly stable hydride unlike

most of the other alkali metals. It is interesting to note that magnesium

exhibits exactly these similar properties.

(Source: YouTube)

Similarities between Lithium and Magnesium

● Lithium and magnesium both form monoxides.

2Mg + O2 → 2MgO

● They both react with nitrogen to form their nitrides

respectively.

Li(s) + N2(g) → 2Li3N(s)

● They both react slowly with water. They form oxides and

hydroxides which decompose on heating.

Mg(s) + 2H2O(g) → Mg(OH)2(aq) + H2(g)

● Oxides of both the elements do not form super oxides.

● The carbonates of both decompose on heating to form the

oxides and carbon-dioxide.

2 Mg(s) + CO2 → 2MgO(s) + C(s)

● Lithium chloride (LiCl) and magnesium chloride (MgCl2) are

both soluble in ethanol.

● Both the elements are less stable towards heat.

Differences between Lithium and Other Alkali Metals

● Lithium is harder than other alkali metals.

● Melting and boiling point is higher than other alkali metals.

● Out of all the other alkali metals, it is the least reactive metal.

● It is a strong reducing agent compared to other alkali metal.

● It is the only alkali metal that forms its monoxide.

● Compared to other alkali metals, it is not capable of forming

solid hydrogen carbonates.

● It does not react with ethyne to form ethynide. On the other

hand, all other alkali metals form ethynide.

● It reacts slowly with bromine as compared to other alkali

metals.

Let us learn Alkali Metals in

detail.

Solved Example for You

Q: Give some practical applications of lithium.

Ans:

● We use it in the manufacturing of batteries.

● It is used in the glass and ceramic industry.

● It finds its use in the polymers and drug industries as well.

Beryllium, Calcium and Magnesium

Does your mother force you to drink a glass full of milk every day or

your bones will lose their calcium and become weak? Our bones need

calcium to maintain their density. Likewise, beryllium and magnesium

are two other elements that we require for various purposes. Do you

know that these three elements are in the same group of the periodic

table? That means they must have similar properties too. Let’s learn

about these properties of the three elements and their importance.

Browse more Topics Under The S Block Elements

● Anomalous Behaviour of Lithium

● Beryllium, Calcium and Magnesium

● Characteristics of the Compounds of the Alkali Earth Metals

● Characteristics of the Compounds of the Alkali Metals

● Group 1 Elements: Alkali Metals

● Group 2 Elements: Alkali Earth Metals

● Some Important Compounds of Sodium and Potassium

Beryllium, Calcium and Magnesium

Naturally occurring elements are quite rare but of extensive use to the

humankind. One of such class of elements occurs in the second

column of the periodic table. These chemical elements are collectively

called as the Alkaline Earth Metals. Beryllium, Calcium and

Magnesium are three of the six elements that fall into this category.

The outer electronic structure of all these elements is similar due to

which they all have similarity in their chemical and physical

properties. They are all shiny, though fairly soft but still harder than

alkali metals. Further, these are usually white or silvery coloured

elements.

They react with water to form hydrogen gas and metal hydroxide and

with oxygen, they form oxides. It will be fascinating to know more

about these 3 elements and their uses and characteristics. So let’s learn

about these elements one by one.

Beryllium(Be)(Z = 4)

Beryllium is the lightest of the entire alkali earth metal family. It was

discovered by a French Chemist Louis-Nicholas Vauquelin

(1763-1829) in 1798. He suggested the name glucinium, meaning

‘sweet tasting’, for the element because the element and some of its

compounds have a sweet taste. The name beryllium was officially

adopted in 1957 after the mineral beryl, which was the form of its

discovery.

By far the greatest use of beryllium metal is in alloys (combination of

two or more metals melted and mixed). Beryllium alloys are popular

because they are tough, stiff, and lighter than similar alloys. Beryllium

finds use in alloys with copper or nickel to make gyroscopes, springs,

electrical contacts, spot-welding electrodes and non-sparking tools.

It occurs in about 30 different mineral species. One of the most

important is Beryl (beryllium aluminium silicate). Emerald and

aquamarine are its precious forms.

s-Block Elements

Calcium (Ca)(Z = 20)

Calcium is a very commonly used element of the alkali earth metal

family. It has been in use for as long as can be dated back even before

its official discovery by an English Chemist Humphry Davy

(1778-1829) in 1808. It had been in use in compound format (one of

the many was limestone) but Humphry discovered the pure Calcium

and named it thus, from the Latin word ‘calx’ meaning lime.

Metallic calcium has relatively few uses. However, calcium

compounds are well known and widely used. The starting point for the

manufacture of most calcium compounds is limestone. Limestone

occurs naturally in large amounts and used in the production of

metals. Another important use of lime is in pollution control. Many

factories release harmful gases into the atmosphere through

smokestacks. Lining a smokestack with lime allows some of these

gases to get absorbed before releasing.

Calcium is essential to both plant and animal life. In humans, it makes

up about two percent of body weight. About 99 percent of the calcium

in a person’s body is found in bones and teeth. Milk is a good source

of calcium.

(Source: Livestrong)

Magnesium (Mg)(Z = 12)

The first person to recognise that magnesium was an element was

Joseph Black at Edinburgh in 1755. He distinguished magnesia

(magnesium oxide, MgO) from lime (calcium oxide, CaO) although

both were produced by heating similar kinds of carbonate rocks,

magnesite and limestone respectively. The name was coined from

Magnesia. A pure, but tiny, amount of the metal was isolated in 1808

by Humphry Davy by the electrolysis of magnesium oxide.

Magnesium is used in products that benefit from being lightweight,

such as car seats, luggage, laptops, cameras and power tools.

Magnesium ignites easily in air and burns with a bright light; hence

it’s used in flares, fireworks and sparklers.

Magnesium oxide is used to make heat-resistant bricks for fireplaces

and furnaces. Also, Chlorophyll contains magnesium at its centre

which enables plants to carry out the process of photosynthesis.

Magnesium is very rarely found in the purest form. It usually occurs in

a combined state in nature.

Solved Example

Q: What are the practical applications of Beryllium, Calcium and

Magnesium?

Ans:

● Beryllium is used mostly for military applications, but there are

other uses of beryllium, as well. In electronics, it is used in

some semiconductors, and beryllium oxide is used as a

high-strength electrical insulator and heat conductor.

● Magnesium has many different uses. One of its most common

use was in industry, where it is often alloyed with aluminium or

zinc to form materials with more desirable properties than any

pure metal. Also, it is used in the production of iron, steel and

titanium.

● Calcium can be used as a reducing agent in the separation of

other metals from ore, like uranium. It is also used in the

production of the alloys of many metals, such as aluminium

and copper alloys, and is also used to deoxidize alloys as well.

Calcium is also used in the making of cheese, mortars, and

cement.

Characteristics of the Compounds of the Alkali Earth Metals

Every object around us displays some unique characteristics, which

sets aside that object from the rest. These properties become the

reason why an element behaves a certain way and how we are able to

predict its behaviour. We will learn about the characteristics of the

compounds of the alkali earth metals in this topic. How are the

compounds of the alkali earth metals any different from the

compounds of other elements in the periodic table? Let’s find out.

What are Alkali Earth Metals?

If we consider the periodic table, the elements that would fall in the

group 2 of the table are usually known as alkali earth metals. Included

in these metals are beryllium(Be), magnesium(Mg), strontium(Sr),

barium(Ba) and radium(Ra). Each of these elements contains two

electrons in their outermost shell. Let us discuss the characteristics of

the compounds of the alkali earth metals.

Browse more Topics under The S Block Elements

● Anomalous Behaviour of Lithium

● Beryllium, Calcium and Magnesium

● Characteristics of the Compounds of the Alkali Metals

● Group 1 Elements: Alkali Metals

● Group 2 Elements: Alkali Earth Metals

● Some Important Compounds of Sodium and Potassium

Physical Properties of the Compounds of Alkali Earth Metals

● They are silverish, white, and hard metals. They are soft but

harder than alkali metals in comparison.

● Some of them appear whitish but beryllium and magnesium

appear greyish in colour.

● Their melting and boiling points are higher compared to the

alkali metals.

● These metals are strongly electropositive in nature. Alkaline

earth metals give different colour with the flame test such as

calcium gives brick red colour, strontium gives crimson colour

and barium gives apple green colour, all of which are different

for different metals.

Chemical Properties of the Compounds of Alkali Earth Metals

● All alkaline earth metals tend to form monoxide except the

metal, beryllium.

● They usually have high electrical and thermal conductivities as

they have a metallic bonding.

● The oxides of alkaline earth metals are basic but less basic in

comparison to alkali metals.

● Hydroxides of alkaline earth metals are basic in nature except

for beryllium hydroxide.

● Alkali earth metals form solid carbonates. As one moves from

beryllium to barium thermal stability of carbonates usually

increases.

● Alkaline earth metals are also capable to form sulphates such as

BeSO4, and MgSO4. Beryllium sulphate and magnesium

sulphate is soluble in water as compared to other sulphates of

alkaline earth metals.

● Group 2 elements also form hydrated, crystallized nitrates.

Heating of nitrates forms oxides. Barium nitrate crystallizes to

form an anhydrous salt of barium oxide whereas magnesium

nitrate crystallizes with six molecules of water.

2Ba(NO3)2 + heat → 2BaO + 4NO2 + O2

● Alkaline earth metals form halides only after reacting with

halogens. Beryllium chloride polymerizes in the solid phase.

● All beryllium halides are essentially covalent and are soluble in

organic solvents. They are hygroscopic, and fume in the air due

to hydrolysis. On hydrolysis, they produce an acidic solution.

● The halides of all remaining alkaline earth metals are ionic in

nature. Their ionic character increases as the size of the metal

Ion increases.

● As the ionic character increases or the covalent character

decreases, their tendency towards undergoing hydrolysis

decreases.

● The hydrated chlorides, bromides and iodides of Ca, Ba and Sr

can be dehydrated on heating but those of Be and Mg undergo

hydrolysis.

● BeF2 in highly soluble in water due to the high hydration

enthalpy of the small Be2+ ion. The other fluorides are almost

insoluble in water.

● The chlorides, bromides and iodides of all of the elements i.e.

Mg, Ca, Ba, Sr are ionic, have a lower melting point than the

fluorides and are readily soluble in water. The solubility

decreases somewhat with increasing atomic number.

● Except for BeCl2 and MgCl2, the other chlorides of alkaline

earth metal impart characteristic colours to flame.

A Solved Question for You

Q: Discuss the solubility and thermal stability of sulphates of alkali

earth metals.

Ans: The sulphates of alkaline earth metal are prepared by the action

of sulphuric acid on metals, metal oxides, hydroxides and carbonates.

The sulphates of alkaline earth metal are all white solids. Beryllium,

magnesium and calcium sulphate crystallise in the hydrated form i.e.

BeSO4·4H2O, MgSO4·7H2O, CaSO4·2H2O but sulphates of Strontium

and barium crystallise without water of crystallisation.

The solubility of sulphates in water decreases down the group. The

magnitude of the lattice enthalpy remains almost constant as the

sulphate ion is so big that small increase in the size of cation from Be

to Ba does not make any difference. The hydration enthalpy decreases

from Be2+to Ba2+ as the size of the cation increases down the group.

Hence the solubility of sulphates of alkaline earth metal decreases

down the group mainly due to decreasing hydration enthalpy from

Be2+ to Ba2+. The high solubility of BeSO4 and MgSO4 is due to the

high hydration enthalpy because of the smaller size of Be2+ and Mg2+

ions. The reason for this is because the sulphates of alkaline earth

metals decompose on heating giving their corresponding oxides and

SO3.

Characteristics of the Compounds of the Alkali Metals

Science manifests itself in strange, and sometimes, surprising ways.

Elements show various properties in ways that are hard to predict and

imagine at times. Among the elements to show amazing and surprising

properties are Alkali metals and the compounds of the alkali metals.

The study of such alkali metals is quite amazing. Let us get to know

more about the compounds of the Alkali Metals in the following

sections.

Characteristics of the Compounds of the Alkali Metals

All alkali metals form various oxides, hydroxides, carbonates and

nitrates. They are hence, known as the most reactive elements as they

have the weakest nuclear charge in the respective period. They have

the tendency to lose their one valence electron in the last shell and

form strong ionic bonds with their anions. Let’s, one by one study the

characteristics of the compound of the alkali metals.

Forming Oxides and Hydroxides

The property of alkali metals allows their oxides, their peroxides and

their super-oxides to dissolve in water quite readily. Such dissolving

in water produces corresponding hydroxides which are basically very

strong alkalis. Certain examples in the form of a chemical equation

that display such phenomenon are as follows:

● 2Na + 2H2O → 2NaOH + H2

● Na2O + 2H2O→ 2NaOH

● Na2O2+ 2H2O → 2NaOH + H2O2

● 2KO2+ 2H2O → 2KOH + H2O2 + O2

It will be right to say that peroxides and super-oxides act as oxidising

agents since they are able to react with water to form hydrogen oxide

and oxygen easily. For any hydroxide of an alkali metal, one can

observe whitish crystalline solids. As a base, they are very strong and

can easily dissolve in water, also emitting a large quantity of heat in

the process.

Moving down the periodic table, we can observe that the basic

strength of such hydroxides, tends to increase. The hydroxides of

alkali metals usually behave as strong bases owing to their low

ionization energies which go down in the group. This decrease in

ionization energies usually leads to the weakening of the bond

between the metal and hydroxide ions and M – O bond in M – O – H

can easily break, hence giving M+ and OH–

This ultimately results in increased concentration of hydroxyl ions in

the solution which can be defined by the increase in basic characters.

All such hydroxides are very soluble in water and thermally stable

with the exception of lithium hydroxide. Alkali metals with their

hydroxides, being strongly basic, tend to react with all acids, leading

to the formation of salts.

Forming Halides

The alkali metals tend to combine directly with different halogens

under appropriate conditions, thus forming halides of the general

formula MX. Examples that demonstrate the formation of such halides

are as follows:

● M2O + 2HX → 2MX + H2O

● MOH + HX → MX + H2O

● M2CO3+ 2HX → 2MX + CO2 + H2O

(where M = Li, Na, K, Rb or Cs and X = F, Cl, Br or I). All of these

halides are usually colourless, are high melting crystalline solids that

have high negative enthalpies of formation.

Other Compounds of the Alkali Metals

1) Sodium Bicarbonate

A concentrated solution of sodium carbonate can absorb Carbon

Dioxide to give sparingly soluble sodium bicarbonate. We can

demonstrate this by the following chemical equation:

Na2CO3 + CO2 + H2O → 2NaHCO3

Quite sparingly, this happens to be soluble in water. When heated

between 250°C and 300°C, it gets converted into pure anhydrous

sodium carbonate which can be later used for standardising acids.

2) Sodium Chloride

We usually call it as ‘common salt’ that occurs abundantly in nature as

a rock salt or halite. The most abundant source of Sodium Chloride is

sea water where sodium chloride occurs to the extent of 3 percent. The

sea water is exposed to the sun and air in large shallow pits.

The gradual evaporation of water leads to the crystallization of the

salt. The solution is later saturated with a current of dry hydrogen

chloride whereby crystals of pure sodium chloride separate out in the

process. Sodium Chloride is a colourless crystalline salt that is almost

insoluble in alcohol but highly soluble in water.

3) Potassium Chloride

It is prepared from fused carnallite that is nearly pure Potassium

Chloride, separated from the melt, leaving fused MgCl2 behind. The

salt is colourless in cubic crystal-like solid soluble in water. Its

solubility increases in linear proportion with the temperature.

4) Potassium Sulphate

We can obtain it by strongly heating potassium chloride with

concentrated Sulphuric acid. It is a colourless crystalline salt, and it is

less soluble in water than sodium sulphate.

A Solved Question for You

Q: Why does lithium form only lithium oxide and not peroxide or

superoxide?

Ans: Due to the small size of lithium compounds, the element has a

strong positive field around it. When combined with the oxide anion

(O2–), the positive field around the lithium-ion restricts the spreading

movement of the negative charge towards another oxygen atom and

thus prevents the formation of higher oxides. This is why lithium does

not form lithium peroxide or lithium superoxide and gets restricted to

only lithium oxide.

Group 1 Elements: Alkali Metals

Did you know that the elements in the periodic table are further

classified on the basis of their properties? Well, depending on the

nature of reaction that the metals display, some of the metals are

called Alkali metals. Making such classification is sometimes vital to

the understanding of different elements.

It also helps us in determining the expected outcome of a given

element, based on its placement in the periodic table. Hence, studying

alkali metals and their properties is very interesting and stimulating to

the overall knowledge of the periodic table. Let us understand more

about these metals.

Alkali Metals – Group 1 Elements

(Source: Image40)

Included in the Group 1 of the periodic table are the following

elements:

● Lithium

● Sodium

● Potassium

● Rubidium

● Caesium

The general electronic configuration of Group 1 elements is ns1. They

have a strong tendency to donate their valence electron in the last shell

to form strong ionic bonds. They have the least nuclear charge in their

respective periods. As we move down the group, the atomic radius

increases. Therefore, the nuclear charge decreases. Caesium is the

most metallic element in the group.

In order to prevent the elements from coming in contact with oxygen,

they are stored in jars that contain oil. The melting points of these

elements are quite low, which is 180° Celsius in the case of Lithium,

while it is 39° Celsius in the case of Rubidium. When it comes to the

density of the metal, group one elements display a very low level of

density of up to 1 gcm-3 which means that they can easily float on the

surface of the water.

Therefore, if we decide to cut these metals, we will be able to do so,

without much trouble. Upon being cut into two halves, we can observe

that their surface is as shiny as any other metal but even after they are

stored in oil, they undergo tarnishing. As a conductor of heat and

electricity, they are excellent.

Learn more about Group 14 Elements here.

The Reaction of Alkali Metals with Water

Alkali metals derive their classification because of the results of their

reaction with water. It is known upon the reaction with water that

alkali metals produce an alkaline solution, along with the release of

hydrogen gas. The following chemical equations demonstrate how

various metals react with water:

lithium + water → lithium hydroxide +

hydrogen

2 Li(s) + 2 H2O(l) → 2 LiOH (aq) + H2(g)

sodium + water → sodium hydroxide + hydrogen

2 Na(s) + 2 H2O(l) → 2 NaOH (aq) + H2(g)

potassium + water → potassium hydroxide + hydrogen

2 K(s) + 2 H2O(l) → 2 KOH (aq) + H2(g)

All elements in a particular group react in an analogous manner.

How Do Alkali Metals React Otherwise?

Labelled as the most reactive group of metals in the periodic table,

each of the alkali metals is capable of reacting with different elements

to produce different results. Following chemical equations

demonstrate some of the ways in which they react:

potassium + oxygen → potassium oxide

4 K(s) + O2(g) → 2 K2O(s)

sodium + chlorine → sodium chloride

2 Na(s) + Cl2(g) → 2 NaCl(s)

potassium + chlorine → potassium chloride

2 K(s) + Cl2(g) → 2 KCl (s)

All alkalis manifest themselves as a white solid in their compound

form that is capable of being dissolved in water. Most of these

compounds are ionic in nature.

Learn about Group 16 Elements here.

The History of Alkali Metals

Alkali metal salts were known to the ancients through the Old

Testament which refers to a salt called ‘Neter’ (sodium carbonate),

extracted from the ash of vegetable matter. Saltpetre (potassium

nitrate) was used in gunpowder, which was invented in China around

about the 9th century AD and had been introduced into Europe by the

13th century.

In October in the year 1807, the English chemist Sir Humphry Davy

isolated potassium and later sodium. The name sodium comes from

the Italian soda, a term applied in the Middle Ages to all alkalis,

potassium comes from the French ‘potasse’, a name used for the

residue left in the evaporation of aqueous solutions derived from wood

ashes.

In the year 1817 Swedish chemist Johan August Arfwedson

discovered Lithium while analyzing the mineral petalite. The name

lithium comes from lithos, the Greek word for stony. The element was

not isolated in pure form until Davy produced a minute quantity by the

electrolysis of lithium chloride.

Learn more about Group 17 Elements here.

A Solved Question for You

Q: What are the properties of alkali metals?

Ans: Alkali metals are highly reactive in nature, which is why they

manifest themselves in combination with other elements, in nature.

Most of these metals are easily soluble in water, which makes their

extraction quite easy. All alkali metals show a silver-like lustre, which

makes them appear shiny. They are highly ductile and conduct

electricity without any trouble.

All alkali metals have a very low melting point and the alloys of such

alkali metals display even lower melting points. They react most

easily with the oxygen in the atmosphere and water vapor. They are

also capable of reacting quite vigorously to form hydrogen gas and

strong caustic solutions.

Group 2 Elements: Alkali Earth Metals

Possibly the neighbours to the most reactive elements in the group,

Alkali earth metals belong to the group 2 of the periodic table.

Somehow, they are very similar to their neighbouring elements of the

table. Yet they manage to be quite different from them. These metals

display a fair share of interesting properties which are absolutely fun

to study about. So let’s learn about the Alkali Earth Metals.

Alkali Earth Metals – Group 2 Elements

Included in the group two elements are Beryllium(Be),

Magnesium(Mg), Calcium(Ca), Strontium(Sr), and Barium(Ba).

Usually, there is no need to store these elements in oil, unlike the

group one elements. For a metal, alkali earth metals tend to have low

melting points and low densities. Being a metal, they are obviously

good conductors of heat and electricity.

The general electronic configuration of Group 2 elements is ns2.

Alkali earth metals have the capability to lose the two electrons in

their outer shell. Thus, they react with other elements and form ionic

compounds. Let’s take some examples to understand the reactions of

such metals.

● The reaction of magnesium with water takes place very slowly,

wherein, the release of hydrogen gas is also very slow.

However, upon reaction, calcium tends to frizz away quite

quickly. As a result, an alkaline solution is formed, which can

be understood better by the following equation:

calcium + water → calcium hydroxide + hydrogen

i.e.

Ca (s) + 2 H2O (l) → Ca (OH)2 (aq) + H2 (g)

● Strontium tends to give off the hydrogen gas, much more

easily.

● Barium also reacts very quickly with water.

Since the reactive ability of group two elements is quite less in

comparison to group one elements, they are used to be added in acids,

in order to dilute them. For example, magnesium and calcium, added

in hydrochloric acid would produce the following output:

magnesium + hydrochloric acid → magnesium chloride +

hydrogen

Mg (s) + 2 HCl (aq) → MgCl2 (aq) + H2 (g)

Uses of the Alkali Earth Metals

As far as the uses of the group two elements and their compounds are

concerned, there is a lot to be understood on that front.

Magnesium usually burns with a bright whitish flame and this has

allowed it to be used in fireworks and rescue flares, along with the

other type of such variety. A unique use of the metal is in the

manufacture and production of high-performance car engines.

Let us take, for example, the Volkswagen ‘Beetle’ has a magnesium

crankcase and other engine parts. The Porsche 911 contains more than

50 kilogram of magnesium. It is brought into use because of its low

density, thereby reducing fuel consumption and reducing the emitting

of pollution from these cars. Magnesium compounds are useful as

well. For example:

● The active ingredient Magnesium hydroxide is used some

indigestion remedies. It neutralizes the excess acid that causes

heartburn in humans.

● Magnesium oxide has a very high melting point hence used as

a lining inside furnaces.

● Epsom salt, which is a laxative, has Magnesium sulphate in it.

Strontium compounds find their use in fireworks to produce a crimson

red colour. Barium compounds are very poisonous. Rat poison has

barium carbonate in it. However, you might have heard of ‘barium

meals’ in hospitals. Patients swallow a white substance that shows up

their digestive tract when X-rayed. This contains barium sulphate,

which is insoluble in water and so just passes through your body

without doing any harm.

A Solved Question for You

Q: Discuss the physical properties of the group II elements.

Ans: The atomic radii, as well as ionic radii of the members of the

family of group II elements, are smaller than the corresponding

members of alkali metals. The alkaline earth metals, owing to their

large size of atoms have fairly low values of ionization energies as

compared to the p – block elements. However, within the group, the

ionization energy decreases as the atomic number increases.

It is because of increase in atomic size due to the addition of new

shells and increase in the magnitude of screening effect of the

electrons in inner shells. Because their (IE) 1 is larger than that of

their alkali metal neighbours, the group IIA metals trend to the

somewhat less reactive than alkali metals.

Atomic weight increases from Be to Ba in a group and volume also

increases, but increase in atomic weight is more as compare to atomic

volume. Therefore the density increases from Be to Ba. The alkaline

earth metals have higher melting and boiling points as compared to

those of alkali metals mainly attributed to their small size and more

closely packed crystal lattice as compared to alkali metals and

presence of two valence electrons.

Since the alkaline earth metals (except Be) tend to lose their valence

electrons readily, they act as strong reducing agents as indicated by E0

red values. The less negative value for Beryllium arises from the large

hydration energy associated with the small size of Be2+ and the

relatively large value of the heat of sublimation.

Some Important Compounds of Sodium and Potassium

Sodium and potassium salts are the better forms of ‘common salt’. As

much as it is difficult to imagine food without salts, these two

elements find use in several industries. This makes them highly

desired and useful elements found on planet earth. Let us get to know

more about the compounds of sodium and potassium.

Compounds of Sodium and Potassium

The compounds of sodium and potassium are as useful as the

elements. They are very useful in industries. They have their unique

properties. Before learning about the compounds of sodium and

potassium. let’s learn about the elements first.

Sodium (Na)

An alkali metal, Sodium belongs to Group 1. Sodium has an atomic

number of 11. It usually manifests itself as a soft, white and highly

reactive alkali metal. It has one electron in the outermost shell, upon

losing which, it tends to form a sodium ion. Sodium does not occur

freely in nature as it is a highly reactive metal. You can store it in

kerosene oil to prevent its reaction in the atmosphere and with the air.

It mainly consists of three minerals such as sodalite, feldspar, and rock

salt.

Properties of Sodium

We know the following properties of sodium:

● It is a highly reactive alkali metal.

● Sodium appears yellow in the flame test.

● Sodium has a soft texture, hence, it can be easily cut with a

knife.

● The melting and boiling points of sodium are lower than that of

lithium.

● Sodium possesses metallic bonding. It is also conducting in

nature as it has one free electron.

● It has lower first ionization energy.

● The common oxidation state of sodium atom is +1

Uses of Sodium

Among the most common uses of Sodium in the human body are that

it regulates the flow of water across the membrane and helps in

transporting sugars and amino acids into various cells.

Potassium (K)

Potassium is an s-block element occurring in Group 1 below Sodium.

It has an atomic number of 19. Hence its electronic configuration is

ns1. It has one valence electron which it readily donates to accepting

atoms. Hence, it forms strong ionic bonds and becomes a cation. The

size of the cation of potassium is smaller than its atom since it loses its

electron.

Potassium, like sodium, is a soft metal. It can be cut with a knife, It

forms various compounds like salts, oxides, hydroxides, etc. Let’s

study the compounds of sodium and potassium respectively now.

Important Compounds of Sodium

Some of the important compounds of Sodium are Sodium Carbonate,

Sodium Chloride, Sodium Hydroxide, and Sodium Hydrogen

Carbonate.

● Sodium Carbonate: Sodium Carbonate, also commonly known

as washing soda, has a molecular formula Na2CO3.10H2O. It is

readily soluble in water. Heating of sodium carbonate

decahydrate leads to the formation of sodium carbonate

monohydrate. On further heating, monohydrate converts into

an anhydrous form of sodium carbonate. It found uses in

cleaning, softening, and laundering. It has found uses in the

textiles industry. Sodium carbonate also supports the

manufacture of glass, borax, soap, and caustic soda.

● Sodium Chloride: As for Sodium Chloride, its main source is

sea water. Crude sodium chloride is obtained by crystallization

of brine solution, containing sodium sulphate, calcium

sulphate, calcium chloride and magnesium chloride.

● Sodium hydroxide: We commonly know caustic as sodium

hydroxide. Among its many uses are manufacturing of soap,

paper, artificial silk etc., used in textiles industries such as

cotton industries, used as a precipitating agent in the

laboratories and more.

Important Compounds of Potassium

The two most important compounds of Potassium are Potassium

Fluoride and Potassium chlorides. Some other compounds of

Potassium are as follows:

● Potassium Permanganate: It is a dark purple crystal at room

temperature. It is soluble in water. Its melting point is 240°C

and density is 2.7 g/cm3. KMnO4 is a strong oxidizing agent. It

is widely used to prevent infection, water purification etc.

● Potassium Hypochlorite: Its formula is KClO and molar mass

is 90.5507. We use it as a disinfectant.

● Potassium Phosphate: It is a white powder at room temperature.

It is soluble in water. Its melting point is 1380°C and density is

2.564 g/cm3. K3PO4 can be used as fertilizer or food additive.

● Potassium Oxalate: It is a white crystal at room temperature. It

is soluble in water. We use K2C2O4 mainly in the medical field

e.g. as an anti-coagulant. It can also be used as a bleaching

agent.

● Potassium Chromate: It is a yellow powder at room

temperature. It is soluble in water. Its melting point is 968 °C

and density is 2.732 g/cm3. K2CrO4 can be used as an

oxidizing agent.

● Potassium Hydrogen Phthalate: It is a white solid at room

temperature. It is soluble in water. Its melting point is 295 °C

(563 °F) and density is 1.636 g/cm3. KHC8H4O4 can be used

for pH meter calibration or as a buffering agent.

● Potassium Hydrogen Carbonate: It is a white crystal at room

temperature. It is soluble in water. Its melting point is 292 °C

and density is 2.17 g/cm3. KHCO3 can be used in baking

similar to soda. We can also use it as a pH regulator.

Question for You

Q: Discuss the properties of Sodium Hydrogen Carbonate and its uses.

Ans: Also known as baking soda, the molecular formula of Sodium

Hydrogen Carbonate is NaHCO3. Upon decomposition, it leads to the

formation of carbon dioxide, which you can understand through the

equation:

2NaHCO3 (s) CO2 (g) + H2O (g) + Na2CO3 (s)

You can use it as an antiseptic during a skin infection or in fire

extinguishers. Most bakeries also make use of this compound to

prepare and preserve pasties and cakes. It has a wide use across

various industries.