Acids, Bases, Salts and Metals

Category : UPSC

 

ACIDS, BASES, SALTS AND METALS

 

 

ACIDS, BASES AND SALTS

 

Some naturally occurring substances that contain acids are given in the following table:

 

Substances

Acid present

Orange, lemon

Citric acid, ascorbic acid (vitamin c)

Apple

Malic acid

Tamarind

Tartaric acid

Vinegar

Acetic acid

Curd

Lactic acid

Tomato

Oxalic acid

Gastric juice

Hydrochloric acid

Tea

Tannic acid

Red ants

Formic acid

 

  • Robert Boyle was the first scientist to term substances as acid and base based on different characteristics of substances.

Swedish scientist Svante Arrhenius in the late nineteenth century first attempted to explain the behaviour of acids and bases from their chemical structure.

 

CONCEPTS OF ACIDS AND BASES

Arrhenius Concept

According to Arrhenius, an acid is a compound that releases \[{{H}^{+}}\]ions in water; and a base is a compound that releases \[O{{H}^{-}}\] ions in water.

 

Bronsted-Lowry Concept

In 1923 J.N. Bronsted and J.M. Lowry independently proposed a broader concept of acids and bases. According to this concept,

  • An acid is any molecule or ion that can donate a proton (\[{{H}^{+}}\])
  • A base is any molecule or ion that can accept a proton
  • An acid is a proton donor while a base is a proton acceptor.
  • Water that accepts a proton is a Bronsted base.

 

Conjugate Acid-Base pairs

An important concept that emanates from Bronsted-Lowry concept is conjugate acid-base pairs. In an acid-base reaction the acid (HA) gives up its proton (\[{{H}^{+}}\]) and produces a new base (\[{{A}^{-}}\]).

  • The new base that is related to the original acid is called a conjugate (meaning related)
  • Similarly the original base (\[{{B}^{-}}\]) after accepting a proton (\[{{H}^{+}}\]) gives a new acid (HB) which is called a conjugate acid.

 

Classes of Bronsted Acids and Bases

Bronsted acids can be classified as per its capacity to furnish protons as follows:

  • Monoprotic acids are capable of donating one proton only,
  • Polyprotic acids are capable of donating two or more protons, e.g. \[{{H}_{2}}S{{O}_{4}},{{H}_{3}}P{{O}_{4}},\]carbonic acid (\[{{H}_{2}}C{{O}_{3}}\]), hydro sulpHuric acid, etc.
  • Monoprotic bases can accept one proton.
  • Polyprotic bases can accept two or more protons, e.g. anions of diprotic and triprotic acids.

 

Strength of Bronsted Acids and Bases

The strength of a Bronsted acid depends upon its tendency to donate a proton. The strength of a Bronsted base depends on its ability to accept a proton.

For example, HCl is nearly 100% ionised in water. Its reaction with water can be depicted by the equation:

 

 

The above reaction has proceeded almost completely to the right; that indicates that HCl has a strong tendency to lose a proton. Also, the base \[{{H}_{2}}O\] has a strong ability to accept a proton. The overall situation is that the acid and base on the left are each stronger than the conjugate acid and conjugate base on the right. That is why the equilibrium is displaced to the right. Thus it may be stated that:

  • a strong acid has a weak conjugate base
  • a strong base has a weak conjugate acid

Strong Acids: \[HCl,HBr,Hl,HN{{O}_{3}},{{H}_{2}}S{{O}_{4}},HCI{{O}_{3}},HCI{{O}_{4}}\]

Strong Bases: Alkali metal hydroxides, \[Ba{{(OH)}_{2}},Sr{{(OH)}_{2}},Ca{{(OH)}_{2}},Mg{{(OH)}_{2}}\]

 

Lewis Concept of Acids and Bases

In the early 1930s, G.N. Lewis gave an even a more general model of acids and bases. According to Lewis theory:

  • an acid is an electron-pair acceptor
  • a base is an electron-pair donor

Lewis visualized an acid and base as sharing the electron pair provided by the base. As a result a covalent bond (or coordinate bond) between the Lewis acid and the Lewis base is formed. The resulting combination is called a Complex.

 

pH SCALE-THE MEASURE OF ACIDITY

The concentrations of \[{{H}^{+}}\] and \[O{{H}^{-}}\] ions in aqueous solutions are frequently very small and hence not convenient to work with.

  • It was Danish chemist Soren Sorensen who in 1909 proposed a more useful quantity called\[pH\]. The \[pH\] of a solution is defined as the negative logarithm of the hydrogen ion concentration (in mol/L):

\[pH=-\log [{{H}_{3}}{{O}^{+}}]\]              or \[pH=-\log [H]\]

  • The above equation gives a convenient numbers to work with. The negative logarithm gives a positive number for \[pH=-\log [H]\], which otherwise would be negative due to the small value of [\[{{H}^{+}}\]].The\[pH\]of a solution is a dimensionless quantity.
  • \[pH\]is simply a way to express hydrogen ion concentration, acidic and basic solutions at \[25{}^\circ C\]can be distinguished by their \[pH\]values, as follows:
  • A \[pOH\] scale analogous to the \[pH\] scale can be devised using the negative logarithm of the hydroxide ion concentration of a solution. Thus, we define \[pOH\] as:

\[pOH=\log [O{{H}^{-}}]\]

  • In general, from the definitions it follows that

\[p{{H}^{+}}pOH=14.00\]

  • The pH scale ranges from 0 to 14 on this scale. Is considered neutral, below 7 acidic and above 7 basic. Farther from 7, more acidic or basic the solution is.

 

pH in Humans and Animals

Most of the biochemical reactions taking place in our body are m a narrow \[pH\] range of 7.0 to 7.8. Even a small change in \[pH\]bumpers the processes. Any condition in which blood \[pH\]drops below 7.35 is known as acidosis, if \[pH\] rises above 7.45-then it is called alkalosis.

 

Acid Rain

When the \[pH\] of rain water goes below 5.6, it is called acid rain.

Acid rain is a major environmental disaster.

 

pH in Plants

Soils need to be of optimum\[pH\]for plants to have an adequate growth .It should be neither highly alkaline nor highly acidic.

 

In digestive system

pH plays an important part in the digestion of food. Our stomach produces hydrochloric acid (formic acid) which helps in digestion of food. When we eat spicy food, stomach produces too much of acid which causes ‘acidity’ i.e. irritation and sometimes pain too. In order to get cured from this we use ‘antacids’ which are bases like ‘milk of magnesia’ (suspension of magnesium hydroxide in water).

 

Self-defence of Animals and Plants

Bee sting causes severe pain and burning sensation. It is due to the presence of methanoic acid (formic acid) in it. Use of a mild base like baking soda can provides relief from pain. Some plants like ‘nettle plant’ have fine stinging hair which inject methanoic acid into the body of any animal or human being that comes in its contact.

 

BUFFER SOLUTIONS

Generally pH of an aqueous solution decreases on addition of a small amount of HCl because of the increase in the concentration of \[{{H}^{+}}\]ions. On the other hand, if a small amount of NaOH is added, the pH of the solution increases. However, there are some solutions which resist the change in pH on addition of small amount of strong acid or alkali. Such solutions are called buffer solutions. For example, solution of ammonium acetate, blood, a equimolar mixture of \[N{{H}_{4}}OH+N{{H}_{4}}Cl,C{{H}_{3}}COOH+C{{H}_{3}}COONa,\]etc.

 

Types of Buffer Solution

  • Acidic buffer: Acidic buffer solution contains equimolar quantities of a weak acid and its salt with strong base. For example, acetic acid (\[C{{H}_{3}}COOH\]) and sodium acetate (\[C{{H}_{3}}COONa\]). A solution containing equimolar quantities of acetic acid and sodium acetate maintains its pH value around 4.74.
  • Basic buffer: Basic buffer solution contains equimolar quantities of a weak base and its salt with a strong acid. For example, ammonium hydroxide (\[N{{H}_{4}}OH\]) and ammonium chloride (\[N{{H}_{4}}Cl\]).

 

  • SALTS

A salt is an ionic compound which dissociates to yield a positive ion other than hydrogen ion [\[{{H}^{+}}\]] and a negative ion other than hydroxide ion [\[O{{H}^{-}}\]]

Example:

\[NaCl\to N{{a}^{+}}+C{{l}^{-}}\](fused/ Aq. soln)

 

Classification of Salts

  • Acidic Salt: If a polybasic acid (Example, \[{{H}_{2}}S{{O}_{4}},{{H}_{3}}P{{O}_{4}},{{H}_{2}}S{{O}_{3}}\]etc.) is neutralised partly by a base, the salt formed is acidic.
  • Normal Salt: In case the acid and base neutralise completely the salt formed is a normal salt.
  • Basic Salts: This type of salts are formed by incomplete neutralization of a base with an acid or by partial replacement of hydroxy radicals of a diacids or triacidic base with an acid radical.

Examples:

\[Cu(OH)N{{O}_{3}}\]- Basic copper nitrate.

  • Double Salt - Such a salt is formed by mixing saturated solution of two simple salts followed by crystallisation of the saturated solution.

Example:

\[FeS{{O}_{4}}{{(N{{H}_{4}})}_{2}}S{{O}_{4}}6{{H}_{2}}O\](Mohr’s salt) is a mixture of \[FeS{{O}_{4}}\] (Simple salt) and \[{{(N{{H}_{4}})}_{2}}S{{O}_{4}}\](Simple salt)

  • Mixed Salts - There is no general method for the formation of this type of salt.

Examples: Sodium potassium sulphate \[NaKS{{O}_{4}}\] (two basic radicals,\[N{{a}^{+}}\],\[{{K}^{+}}\])

  • Complex salt - Such a salt is formed by mixing saturated solution of simple salts followed by crystallisation of the solution similar to double salts.

SOME COMMON USEFUL SALTS

A large number of salts are useful for our homes and industry for various purposes. Some are discussed below:

 

Baking Soda

Chemically baking soda is sodium hydrogen carbonate; \[NaHC{{O}_{3}}\].

  • It is an important part of food.
  • Baking soda is manufactured by Solvay's process.
  • It is mainly used for manufacturing washing soda but baking soda is obtained as an intermediate.
  • On heating, sodium hydrogen carbonate is converted into sodium carbonate and carbon dioxide is given off:

\[2NaHC{{O}_{3}}\xrightarrow[{}]{heat}N{{a}_{2}}C{{O}_{3}}+{{H}_{2}}O+C{{O}_{2}}\uparrow \]

Uses:

  • It is used as a component of baking powder.
  • It is used as a tenderizer and leavening agent in baking (In combination with a liquid and acid it releases \[C{{O}_{2}}\])
  • It is used as deodorizer because of its neutralizing action.
  • It is used in laundry work for enhancing the detergents effectiveness because it stabilizes the pH level (acts as a buffer)
  • It is used in fire extinguishers. Baking soda undergoes a chemical reaction that gives off \[C{{O}_{2}}\] that makes it useful in extinguishing small grease or electrical fires.

\[2NaHC{{O}_{3}}\xrightarrow[{}]{heat}N{{a}_{2}}C{{O}_{3}}+{{H}_{2}}O+C{{O}_{2}}\uparrow \]

  • It is used as abrasive cleaner.

 

Washing Soda

  • Washing soda is used for washing of clothes. Chemically, washing soda is sodium carbonate decahydrate, \[N{{a}_{2}}C{{O}_{3}}.10{{H}_{2}}O\].
  • Washing soda is manufactured by Solvay’s process.

 

Uses

  • It is used in the manufacture of caustic soda, glass, soap powders, borax and in paper industry.
  • For removing permanent hardness of water.
  • As a cleansing agent for domestic purpose.

 

Plaster of Paris

Also called POP, chemically, it is \[2CaS{{O}_{4}}.{{H}_{2}}O\]or \[CaS{{O}_{4}}.1/2{{H}_{2}}O\] (calcium sulphate hemi hydrate)

  • Gypsum, (\[CaS{{O}_{4}}.2{{H}_{2}}O\]) is used as the raw material to manufacture POP.
  • The only difference between gypsum (\[CaS{{O}_{4}}.2{{H}_{2}}O\]) and plaster of Paris (\[CaS{{O}_{4}}.1/2{{H}_{2}}O\]) is the less amount of water of crystallization.

 

Uses

  • In medicine, used for making plaster casts to hold fractured bones in place while they set. It is also used for making casts in dentistry.
  • For making fire proof materials.

 

Bleaching Powder

Bleaching is a process of removing colour from a cloth to make it whiter. Bleaching powder has been used for this purpose since long. Chemically, it is calcium oxy chloride\[CaOC{{l}_{2}}\].

 

Uses:

  • Used in textile industry for bleaching of clothes.
  • In paper industry for bleaching of wood pulp.
  • It makes wool unshrinkable.
  • Used as disinfectant and germicide for sterilization of drinking water and swimming pool water.
  • For the manufacture of chloroform (\[CHC{{l}_{3}}\])
  • Used as an oxidizing agent in chemical industry.

 

Sodium Hydroxide

Also known as Caustic Soda, chemically it is NaOH. Industrial methods of its production are:

  • Causticisation process(Gossage process)
  • Castner Kellner cell
  • Chlor-Alkali process

 

Handy Facts

Some properties of NaOH are:

NaOH (caustic soda) is a white translucent solid available in pellets, flakes and granules as 50% saturated solution.

In solution when exposed to air reacts with atmospheric \[C{{O}_{2}}\]and from a crust of \[N{{a}_{2}}C{{O}_{3}}\]at the surface.

 

Uses:

  • It is used in many industries, mostly as a strong chemical base in the manufacture of pulp and paper, textiles, soaps, dyes, cellulose, detergents etc.
  • It is used in petroleum refining.

 

Sodium Carbonate

Also called soda ash, its chemical formula is \[N{{a}_{2}}C{{O}_{3}}\]

It exists in various forms, namely anhydrous sodium carbonate \[N{{a}_{2}}C{{O}_{3}}\] (Soda-ash). Monohydrate \[N{{a}_{2}}C{{O}_{3}}.{{H}_{2}}O\] (crystal carbonate), heptahydrate. \[N{{a}_{2}}C{{O}_{3}}.7{{H}_{2}}O\] and decahydrate \[N{{a}_{2}}C{{O}_{3}}.10{{H}_{2}}O\] (washing soda or sal soda)

Sodium carbonate is generally prepared now a days by ammonia soda or Solvay process. The ingredients of this process are readily available and inexpensive. These are Salt brine (NaCl) (from sea), ammonia (\[N{{H}_{3}}\]) and lime stone \[CaC{{O}_{3}}\](from mines). The process consists of many sections, \[CaC{{l}_{2}}\] is an important by-product obtained.

 

Uses:

  • It is used for softening of hard water. For this purpose hydrated sodium carbonate, \[N{{a}_{2}}C{{O}_{3}}.10{{H}_{2}}O\] known as washing soda is used.
  • A mixture of sodium carbonate (\[N{{a}_{2}}C{{O}_{3}}\]) and potassium carbonate (\[{{K}_{2}}C{{O}_{3}}\]) is used as fusion mixture.
  • It is used in paper, paint and textile industries

It is used for the manufacture of glass, borax, soap and caustic soda (NaOH)
 

PROPERATIES, EXTRACTION AND USES OF METALS

 


 Out of 118 chemical elements known till date, 103 has been well characterized in terms of their properties. The systematic classification of these 103 elements shows their numbers as:

  • Metals=79
  • Non-metals=17
  • And metalloids=7

 

COMPARISON OF METALS AND NON-METALS

Metals and non-metals differ both in physical and chemical properties.

The characteristic physical properties of metals and non-metals are listed in following table:

 

Serial No.

Metals

Non-metals

1

Metals have luster. They reflect light from polished or freshly cut surface

Non-metals do not have luster( exceptions - Diamond and Iodine )

2

Metals generally have high density.

Non-metals generally have low density.

3

They are good conductors of heat and electricity.

They are usually bad conductors of heat  (exception - carbon in the form of gas carbon and graphite )

4

Metals are malleable and ductile. They can be beaten into thin sheets and drawn into wires.

Non-metals are not malleable and ductile. They can be crushed into powder.

5

They have a three dimensional crystal structure with metallic bonds

They have different types of structures with covalent and van-der-Waals bonds

6

Metals are generally hard.

Non- metals are generally soft (Exception: Diamond)

7

Metals generally have 1 to 3 electrons in their outermost shell

Non-metals generally have 4 to 8 in their outermost shell of the atom.

8

They show valency 1 to 4.

They show valency 1 to 7.

9

They are electropositive in nature.

They are electronegative in nature. j

10

They generally form basic oxides.

They generally form acidic oxides.

11

They act as reducing agents.

They act as oxidizing agents.

12

Active metals react with cold and hot water.

Non-metals usually do not react with cold or hot water.

13

Active metals react with non- oxidizing acids to form their oxides or oxy acids.

Solid non-metals react with oxidizing acids to form hydrogen gas.

14

They react with non-metals under different conditions to form salts.

They react with metals as well as non-metals under different conditions to form salts.

 

CHEMICAL PROPERTIES OF METAL

The chemical reactions undergone by metals are furnished below:

 

Reaction with Oxygen

Most of the metals react with oxygen and form oxides. The reaction may take place without heating as in sodium, calcium or potassium, while some metals react with oxygen on heating to form oxides.

  • Oxides of metals are basic in nature as they react with water and form bases, e.g. \[N{{a}_{2}}O,CaO,MgO,{{K}_{2}}O,\], etc.
  • Oxides of aluminium (\[A{{l}_{2}}{{O}_{3}}\]), zinc (ZnO), tin (SnO) and iron (\[F{{e}_{2}}{{O}_{3}}\]) are amphoteric in nature as they react with acids as well as with bases.

 

Reaction of Metals with Acids

Metals react with common acids like dilute HCl and dilute \[{{H}_{2}}S{{O}_{4}}\] with evolution of\[{{H}_{2}}\].

 

Reaction of Metals with Water

Many metals react with water to form hydroxides. Hydroxides are basic in nature. Sodium and potassium react with cold water.

  • Magnesium reacts with hot water

\[Mg(s)+{{H}_{2}}O(l)\to Mg{{(OH)}_{2}}(aq)+{{H}_{2}}(g)\]

  • Metals like Al or Fe react on heating with water or with steam. In these conditions metals form metal oxides.

 

Reaction of Metals with Common Bases

Some metals like aluminium and zinc react with common bases.

\[Sn(s)+2NaOH(aq)+{{H}_{2}}O(l)\to N{{a}_{2}}Sn{{O}_{3}}\]

sodium stannate

\[Zn(s)+2NaOH(aq)\to N{{a}_{2}}Zn{{O}_{2}}\]                       

Sodium zincate

REACTIVITY SERIES

A metal that can lose electrons more easily in comparison to other metal is more electropositive and would be more active in nature then other metal. Such a metal when dipped in a solution of salt of a less active metal would displace it.

The arrangement of metals in the decreasing order of their activity is known as activity or reactivity series. It is also known as electrochemical series.

 

EXTRACTION OF METALS-METALLURGY

The process of extraction of metal from its ore is called metallurgy.

  • Minerals are naturally occurring compounds of metals.
  • Ores are those minerals from which metals can be economically extracted.

Examples

  • Aluminium is the most common metal in the Earth's crust, occurring in all sorts of minerals. However, it isn’t economically worthwhile to extract it from most of these minerals. Instead, the usual ore of aluminium is bauxite - which contains from 50 - 70% of aluminium oxide.

 

Examples of Ores

Some important ores and the metals present in these ores are listed in the following table:

 

 

Type of ore

Metals

Native Metals (Found in Free State)

Gold (Au), silver (Ag)

Oxide ores

Iron (Haematite, \[F{{e}_{2}}{{O}_{3}}\]); Aluminum (Bauxite, \[A{{l}_{2}}{{O}_{3}}.2{{H}_{2}}O\]); Tin(Cassiterite, \[Sn{{O}_{2}}\]); Copper Cuprite, \[C{{u}_{2}}O\]); Zinc (Zincite, ZnO); Titanium (llmenite, \[FeTi{{O}_{3}}\], Rutile, \[Ti{{O}_{2}}\])

Sulphide ores

Zinc (Zinc blende, ZnS); Lead (Galena, PbS); Copper (Copper glance, \[C{{u}_{2}}S\]); Silver (Silver glance or Argentite, \[A{{g}_{2}}S\]); Iron(Ironpyrites, \[Fe{{S}_{2}}\])

Carbonate ores

Iron (Siferite, \[FeC{{O}_{3}}\]); Zinc (Calamine, \[ZnC{{O}_{3}}\]), Lead (Cerrusite, \[PbC{{O}_{3}}\])

Sulphate ores

Lead(Anglesite, \[PbS{{O}_{4}}\])

Halide ores

Silver (Horn silver, AgCl); Sodium (Common salt or Rock salt. NaCl); Aluminum (Cryolite, \[N{{a}_{3}}AI{{F}_{6}}\])

Silicate ores

Zinc (Hemimorphite, \[2ZnO.Si{{O}_{2}}.{{H}_{2}}O\])

 

The method used to extract metals from the ore in which they are found depends on their reactivity.

 

Steps Involved in the Extraction of Metals from Ores

Metallurgical Steps

Few steps involved in extracting the metal as shown above are discussed below:

 

Concentrating the Ore

This means getting rid of as much of the unwanted rocky material (called gangue or matrix) as possible before the ore is converted into the metal.

Some methods are:

Froth floatation: The ore is first crushed and then treated with something which will bind to the particles of the metal compound that are required and make those particles hydrophobic. “Hydrophobic” stands for “water fearing”.

Magnetic separation: Magnetic ores like pyrolusite (\[Mn{{O}_{2}}\]) and chromite (\[FeO.C{{r}_{2}}{{O}_{3}}\]) are enriched by this method by making use of the difference in the magnetic properties of the ore and gangue particles.

 

Conversion of concentrated ore to oxide

It is easier to obtain a metal from its oxide form as compared to its sulphide, carbonate or any other form. Therefore, prior to reduction usually the metal is converted to its oxide form.

Following methods are used to convert the concentrated ore to its oxide form.

Calcination: It is a process in which the ore is heated strongly in absence of air. The ore is heated at a temperature well below its melting point.

This method is generally used for carbonate and hydrated ores.

Roasting: It is a process wherein the ore is heated either alone or with some other material in excess of air below the fusion point of the ore.

  • Usually, this method is used for sulphide ores.
  • Ores of metals like zinc, lead, copper and nickel, when roasted in air, are converted to their oxides.
  • The purpose of roasting is to convert the ore in a form suitable to reduce. The gaseous product of sulphide roasting, sulphur dioxide is often used to produce sulphuric acid.

The differences between roasting and calcinations are given in the following table:

Reducing the metal compound to the metal

  • Carbon reduction

Carbon (as coke or charcoal) is cheap. It not only acts as a reducing agent, but it also acts as the fuel to provide heat for the process.

  • Reduction using a more reactive metal Titanium is produced by reducing titanium chloride using a more reactive metal such as sodium or magnesium. This is the only way of producing high purity metal.
  • Reduction by electrolysis

This is a common extraction process for the more reactive metals - for example, for aluminium and metals above it in the electrochemical series.

An advantage is that it can produce very pure metals.

 

Refining of metals

Refining of metals is done to obtain metals in very pure form by removing impurities present in it. Electrolytic refining is widely used method for this purpose.

 

ALLOYS

Alloys are metallic materials prepared by mixing two or more molten metals.

These are used for many purposes, such as construction, and are central to the transportation and electrical industries. Composition of some common alloys and their uses are given in the following table.

 

 

Alloy

Composition

Uses

1.

Brass

Cu = 80%, Zn = 20%

For making utensils and cartridges.

2.

Bronze

Cu = 90%, Sn = 10%

For making statues, medals, ships, coins and machines

3.

Solder

Sn = 50%, Pb = 50%

For joining metals, solding wire and electronic components etc.

4.

Duraluminum

Al = 95.5%, Cu = 3%,

Used in bodies of aircrafts, kitchen parts etc. ware and automobile

 

 

Mn = 1.0%, Mg = 0.5%

 

 5.

German Silver

Cu = 60%, Zn = 20%, Ni = 20%

For making utensils and ornaments

6.

Gun metal

Cu = 90%, Sn = 10%

For Gears and castings etc.

7.

Bell metal

Cu = 80%, Sn = 20%

For bells, gangs etc.

8.

Magnalium

Al = 90%,Mg=10%

For balance beams, light instruments.

9.

Type metal

Pb = 82%, Sb = 15%, Sn = 3%

For casting type

10.

Stainless steel

Fe, Ni, Cr, C

For utensils, cutlery etc.

 

  • In homogeneous alloys, atoms of the different elements are distributed uniformly. Examples include brass, bronze, and the coinage alloys.
  • Heterogeneous alloys, such as tin-lead solder and the mercury amalgam sometimes used to fill teeth, consist of a mixture of crystalline phases with different compositions.

An amalgam is an alloy of mercury with one or more metals. Most of the metals form amalgams with mercury except iron and platinum. Amalgams of sodium and aluminium are good reducing agents. Amalgam of silver, tin, cadmium and copper have been utilized as dental fillings.

 

SOME IMPORTANT METALS AND THEIR USES

Iron (Fe)

Iron is the most abundant metal which occurs in the earth’s crust. The most commonly used iron ores are haematite, \[F{{e}_{2}}{{O}_{3}}\], and magnetite, \[F{{e}_{3}}{{O}_{4}}\].

 

Cast Iron:

The molten iron from the bottom of the furnace can be used as cast iron. However, it is very impure, containing about 4% of carbon. This carbon makes it very hard, but also very brittle. If you hit it hard, it tends to shatter rather than bend or dent.

Use: Cast iron is used for things like manhole covers, cast iron pipes, valves and pump bodies in the water industry, guttering and drainpipes, cylinder blocks in car engines,

Aga-type cookers, and very expensive and very heavy cookware.

Wrought iron

If all the carbon is removed from the iron to give high purity iron, it is known as wrought iron. Wrought iron is quite soft and easily worked and has little structural strength.

Use: It was once used to make decorative gates and railings, but these days mild steel is normally used instead.

 

Mild steel

Mild steel is iron containing up to about 0.25% of carbon. The presence of the carbon makes the steel stronger and harder than pure iron. The higher the percentage of carbon, the harder the steel becomes.

Use: Mild steel is used for lots of things - nails, wire, car bodies, ship building, girders and bridges amongst others.

 

High carbon steel

High carbon steel contains up to about 1.5% of carbon. The presence of the extra carbon makes it very hard, but it also makes it more brittle.

Use: Used for cutting tools and masonry nails (nails designed to be driven into concrete blocks or brickwork without bending). It tends to fracture rather than bend if you mistreat it.

 

Special steels

These are iron alloyed with other metals. Examples are given in the following table:

 

Variety

Iron mixed with

Special Properties

Uses

Painless steel

chromium and nickel

resists corrosion

cutlery, cooking utensils, kitchen sinks, industrial equipment for food and drink processing

titanium steel

titanium

withstands high temperatures

gas turbines, spacecraft

manganese steel

manganese

very hard

rock-breaking machinery, some railway track (e.g. points), military helmets

 

 

Handy Facts

The hardness and elasticity of steel can be controlled by heat treatment. The steel is heated to a temperature below redness. It is then cooled slowly. The process is called tempering of steel. It is used to bring the steel to a suitable state of hardness and elasticity. Hard steel can be softened by heating it to a high temperature and then allowing it to cool down slowly. This process is called annealing.

Steel produced in this way is known as quenched steel and the process of making such is known as quenching or hardening of steel

 

Some important compounds of iron:

Green Vitriol: Iron sulphate or ferrous sulphate is the chemical compound with the formula (\[FeS{{O}_{4}}\]).\[FeS{{O}_{4}}\]. \[7{{H}_{2}}O\]is called Green vitriol.

It is used as:

  • A reducing agent, mostly for the reduction of chromate in cement.
  • As nutritional supplement to patients suffering from iron deficiency.
    • As a colorant.

Mohr’s Salt: Ammonium iron sulphate, or Mohr’s Salt, is a double salt of iron sulphate and ammonium sulphate, with the formula\[{{[N{{H}_{4}}]}_{2}}[Fe]{{[S{{O}_{4}}]}_{2}}.6{{H}_{2}}O\]. It is used in the laboratory.

 

Copper (Cu)

Copper is malleable and ductile and is a good conductor of both heat and electricity. The important ores of copper is chalcopyrite, \[CuFe{{S}_{2}}\](also known as copper pyrites)

Copper has multifarious uses like:

  • Making brass: Brass is a copper-zinc alloy.

Alloying produces a metal harder than either copper or zinc individually. Bronze is another copper alloy – this time with tin.

  • Coinage: Used for masking coins in its alloy forms.

 

Some important compounds of Copper

  • Blue vitriol

Copper sulphate is the compound with the formula\[CuS{{O}_{4}}\]. This salt exists as a series of compounds that differ in their degree of hydration. The anhydrous form is a pale green or gray – white powder, whereas the penta hydrate (\[CuS{{O}_{4}}.5{{H}_{2}}O\]), the most commonly encountered salt, is bright blue.

  • Black oxide of copper

Copper oxide or cupric oxide (CuO) is the higher oxide of copper. As a mineral, it is known as tenorite.

Uses:

  • Cupric oxide is used as a pigment in ceramics to produce blue, red, and green (and sometimes gray, pink, or black) glazes.
  • To produce cup ammonium hydroxide solutions, used to make rayon.
  • Occasionally used as a dietary supplement in animals, against copper deficiency.
  • Copper oxide has application as a p-type semiconductor, because it has a narrow band gap.
    • Red oxide of copper
  • Copper oxide or cuprous oxide is the inorganic compound with the formula-\[C{{u}_{2}}O\]. It is one of the principal oxides of copper.

Uses

Cuprous oxide is commonly used as a pigment, a fungicide, and an antifouling agent for marine paints.

Rectifier based on this material have been used industrially.

 

Aluminium

The usual aluminium ore is bauxite. Bauxite is essentially an impure aluminium oxide. The major impurities include iron oxides, silicon dioxide and titanium dioxide. Bauxite actually contains one of a variety of hydrated aluminium oxides some of which can be written as\[A{{l}_{2}}{{O}_{3}},x{{H}_{2}}O\].

Uses: Aluminium is used in aircraft making, container vehicle bodies, tube trains, in making saucepans, etc.

 

Science in Action

Anodizing of Aluminium: In anodising the aluminium is first etched with sodium hydroxide solution to remove the existing oxide layer, and then making the aluminium article the anode in electrolysis of dilute sulphuric acid. The oxygen given off at a film of oxide up to about 0.02 mm thick.

 

Zinc (Zn)

It is extracted from ores like zinc blende (ZnS).

  • Zinc is an essential trace element, necessary for plants, animals, and microorganisms.
  • It is typically the second most abundant transition metal in organisms after iron and it is the only metal which appears in all enzyme classes.
  • Zinc is used in galvanizing. Galvanization is a metal coating process in which a ferrous part is coated with a thin layer of zinc.

 

Handy Facts

Important uses of some other metals are also listed below:

  • Tin is used for soldering, for preparing foils, for metal coatings to prevent chemical action and corrosion, for panel lighting etc.
  • Lead is used in making water pipes, in pigments, batteries, in alloys, etc.
  • Titanium finds extensive use in aircraft industries.
  • Pure metals, which display zero resistance to electrical currents, are called Mercury (Hg), Niobium (Nb) are examples of superconductors. They become superconductors below a critical temperature of 4.2 K and 9.2 K respectively.



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