Metals and Non-metals

"High resistance to corrosion, high melting point and high strength makes titanium and zirconium perfect to design aircraft frames. They are known as strategic metals."

1.0Introduction

Everything around us is made up of different elements. These elements can be classified into metals or nonmetals on the basis of their properties. A few elements have properties common to both metals and non-metals. These are called semi-metals or metalloids.

2.0Physical properties

Metals

Lustre of metals

Most of the metals, in their pure state, have a shining surface. This property is called metallic lustre. For example, gold is shining yellow, copper is brown, iron, aluminium and zinc are lustrous grey.

Active Chemistry 1

Aim To check that metals have lustre i.e. a shining surface.

Materials required Samples of iron, copper, aluminium and magnesium, sand paper.

Method (i) Take samples of iron, copper, aluminium and magnesium. Note the appearance of each sample.

Observation The surface of the metals is dull because they are covered with a layer of oxide, hydroxide, carbonate, etc. due to the attack of gases present in the air on their surface. On rubbing the surface with sand paper this layer is removed and a shining surface appears.

Conclusion Metals in the pure state (or freshly prepared or cut) have shining surface.

Hardness of metals

Most of the metals are hard, but all metals are not equally hard. The hardness of metal varies from metal to metal.

Active Chemistry 2

Aim To test that metals are hard and hardness varies from metal to metal.

Materials required Small pieces of iron, copper, aluminium and magnesium, knife, tong.

Method (i) Take small piece of iron, copper, aluminium and magnesium. Try to cut these metals with a sharp knife. (ii) Hold a piece of sodium metal with a pair of tongs.

Caution Always handle sodium metal with care. Dry it by pressing between the folds of a filter paper. Put it on a watch glass and try to cut it with a knife.

Observation and conclusion All the four metals ( $\mathrm{Fe},\mathrm{Cu},\mathrm{Al}$ and Mg ) are found to be cut with difficulty. This shows that metals are hard. The ease of cutting is found to be in the order $\mathrm{Mg}>\mathrm{Al}>\mathrm{Cu}>\mathrm{Fe}$. This shows that hardness varies from metal to metal. Sodium can be cut very easily. Hence sodium is soft, i.e., it is an exception.

Malleability of metals

The property according to which metals can be beaten with a hammer into very thin sheets without breaking is called malleability. Gold and silver are the most malleable metals. Aluminium and copper are also highly malleable metals. All of these metals can be beaten with a hammer to form very thin sheets, called foil.

Active Chemistry 3

Aim To test that metals are malleable, i.e. can be hammered into sheets.

Materials required Pieces of iron, zinc, lead and copper, block of iron, hammer.

Method (i) Take piece of iron, zinc, lead and copper. (ii) Place any one metal on the block of iron and strike it four or five times with a hammer. (iii) Repeat with other metals. (iv) Record the change in shape of these metals.

Malleability capacity of different metals can be attributed to the strength of the metallic bond and depends upon the extent of force that is applied.

Observation and conclusion It is observed that metals can be beaten into thin sheets i.e. they are malleable.

Ductility of metals

Ductility is also an important property of metals. The ability of metals to be drawn (stretched) into thin wires is called ductility. Generally, wires are made up of iron, copper and aluminium. For example, 100 mg of silver can be drawn into a thin wire of about 200 metres length.

Copper and aluminium are also very ductile, and therefore, they can be drawn into thin wires which are used in electrical wiring.

Active Chemistry 4

Aim To justify that metals are ductile i.e. can be drawn into wire.

Materials required Metal pieces of iron, copper, aluminium, lead, etc.

Method (i) Consider some metals such as iron, copper, aluminium, lead, etc. (ii) Check which of these metals are available in the form of wire.

Observation and conclusion As wires of iron, copper and aluminium are easily available, this shows that metals can be drawn into wires i.e., they are ductile.

Thermal conductivity of metals

The process in which a metal allows the flow of heat through it is called its thermal conductivity. Most of the metals are good conductors of heat, such as silver, gold, iron, copper and aluminium.

Active Chemistry 5

Aim To test that metals are good conductors of heat and have high melting point.

Materials required Aluminium or copper wire, pins, wax, spirit lamp, candle or a burner.

Method (i) Take an aluminium or copper wire. Clamp the wire on a stand. (ii) Fix a pin to the free end of the wire using wax. (iii) Heat the wire with a spirit lamp, candle or a burner near the place where it is clamped.

Silver is the best conductor of heat and lead is the poorest conductor of heat. Good thermal conductivity of metals makes them suitable for wire and utensil making.

Now answer (i) Does the metal wire melt?

Observation and conclusion We observe that on heating the wire near the clamp, after some time the pin falls down. This shows that heat flows through the wire and melts the wax. Further, the wire does not melt even after heating for a long time. This shows that metals have high melting points.

Electrical conductivity of metals

The property in which metal facilitates the flow of electric current through it is called electrical conductivity. All metals are good conductors of electricity because they contain free or mobile electrons. These free electrons conduct electric current.

Active Chemistry 6

Aim To test that metals are good conductor of electricity. Materials required Electric circuit, metal piece.

Method (i) Set up an electric circuit as shown in figure. (ii) Place the metal to be tested in the circuit between terminals $A$ and $B$ as shown in the figure.

Now answer Does the bulb glow? What does this indicate?

Observation and discussion The bulb glows. This shows that electric current flows through the metal.

Conclusion Metals are good conductor of electricity.

Sonorous

The property of metals in which metals produce sound when they strike a hard object or other surface is called sonority or sonorisity. Some metals like copper, silver, gold, aluminium give musical sound when they are struck by themselves or any other object.

Note: Difference in properties of different materials is based on their internal structures.

Non-metals

Among the total known elements, there are only 22 non-metals, out of which 11 are gases like oxygen, nitrogen, hydrogen, one is a liquid (bromine) and the rest 10 are solids such as sulphur, phosphorus, iodine and the allotropes of carbon (diamond and graphite). Do non-metals also have physical properties similar to that of metals? Let us find out.

Active Chemistry 7

Aim To study the physical properties of non-metals.

Materials required Samples of carbon, sulphur, iodine, hammer, knife, battery, etc.

Method (i) Collect samples of carbon (coal or graphite), sulphur and iodine. (ii) Carry out the activities 1 to 6 with these non-metals.

Lead and mercury offer greater resistance to the flow of current. Therefore, they have low electrical conductivities.

Observation

Conclusion Non-metals possess different properties than metals. The important physical properties of non-metals are given below : (1) Non-metals may be solids (such as sulphur, phosphorus and diamond), liquid (bromine), or gases (such as oxygen, nitrogen, hydrogen, neon, argon, etc.) at room temperature. (2) Non-metals are usually brittle and cannot be used to make sheets and wires. (3) Non-metals are non-lustrous and cannot be polished. (Exception : Graphite and iodine are lustrous non-metals). (4) Non-metals are generally bad conductor of heat and electricity.

Exception : Graphite is a good conductor of electricity. Non-metals do not conduct the electric current due to absence of mobile electrons. (5) Non-metals can be easily broken due to their low tensile strength. (6) Non-metals are generally light and have low densities. (7) Unlike metals, non-metals do not produce any ringing sound when struck with an object. (8) Non-metals are soft (Exception : Diamond) (9) Non-metals have low melting and boiling points. (Exception : Graphite has very high melting point $\left(3730^{\circ }\mathrm{C}\right)$ ) On the basis of the above discussion of the physical properties of metals and non-metals, we have concluded that elements cannot be grouped according to the physical properties alone, as there are many exceptions. Note: Allotropes differ in physical properties but have almost similar chemical properties.

Some exceptions (i) All metals except mercury are solid at room temperature. We know that metals have very high melting points but gallium (Ga) and cesium (Cs) have very low melting points. These two metals will melt if we keep them on our palm. (ii) Iodine is a non-metal but it is lustrous. (iii) Alkali metals such as lithium, sodium and potassium are so soft, that they can be easily cut with a knife i.e. they have low densities and low melting points. (iv) Carbon is a non-metal that can exist in different forms. Each form is called an allotrope. Diamond, an allotrope of carbon is the hardest natural substance, which has very high melting and boiling point. Graphite is another allotrope of carbon which is good conductor of electricity.

Nature of oxides

Elements can be more clearly classified as metals and non-metal on the basis of their chemical properties.

Active Chemistry 8

Aim To show that metal oxides are basic while non-metal oxides are acidic.

Materials required Magnesium ribbon, sulphur powder, burner, china dish, red and blue litmus paper, test tubes, water.

Method (i) Take a magnesium ribbon and some sulphur powder. (ii) Burn the magnesium ribbon. Collect the ashes formed and dissolve them in water. (iii) Test the resultant solution with both red and blue litmus paper. (iv) Now burn sulphur powder. Place a test tube over the burning sulphur. (v) Add some water to the above test tube and shake. (vi) Test this solution with blue and red litmus paper.

Now answer (i) Are the products formed on burning magnesium and sulphur acidic or basic? (ii) Can you write equations for these reactions?

Observation and discussion

The solution obtained on dissolving ashes after burning magnesium turns red litmus blue whereas solution obtained on dissolving gases from burning sulphur turns blue litmus red. This shows that magnesium oxide ( MgO ) is basic while oxides of sulphur $\left({\mathrm{SO}}_2,{\mathrm{SO}}_3\right)$ are acidic. The reactions are:

Conclusion Metal oxides are ionic and basic in nature, whereas non-metallic oxides are covalent and acidic in nature.

For example, ${\mathrm{Na}}_2\mathrm{O},{\mathrm{K}}_2\mathrm{O}$ are ionic and ${\mathrm{CO}}_2,{\mathrm{NO}}_2$ are covalent. Some of the non-metal oxides are neutral in nature like ${\mathrm{H}}_2\mathrm{O},\mathrm{CO},{\mathrm{N}}_2\mathrm{O}$ and NO .

3.0Chemical properties of metals

We have studied the physical characteristics of metals. Now let us focus our attention on their chemical properties. Metals in general have tendency to lose one or more electrons present in the valence shells of their atoms to form positive ions. Metals are therefore, regarded as electropositive elements. $\underset{(\text{ metal atom })}{\mathrm{M}}⟶\underset{(\text{ metal ion })}{{\mathrm{M}}^{\mathrm{n}+}}+{\mathrm{ne}}^{-}$

Lithium cannot be stored in kerosene oil because it is the lightest metal and it floats on its surface and reacts with air. Thus, it is kept wrapped in paraffin wax.

The chemical properties of the metals are mostly linked with the electron releasing tendency of their atoms. Greater the tendency, more will be the reactivity of the metal.

(1) Reaction of metal with oxygen

Almost all metals combine with oxygen to form metal oxides. But they possess different reactivity towards oxygen.

Active Chemistry 9

Aim To study the burning of metals in air to form oxides. Materials required Samples of sodium, potassium, magnesium, copper and aluminium, tong, burner, china dish.

Method (i) Take samples of sodium, potassium, magnesium, copper and aluminium. (ii) Hold sample with a pair of tongs and try to burn it over the flame. Repeat with other metal samples. (iii) Collect the products if formed. (iv) Let the products and the metal surface cool down.

Now answer (i) Which metals burn easily? (ii) What flame colour did you observe when the metal burns? (iii) Arrange the metals in the decreasing order of their reactivity towards oxygen. Are the products soluble in water? (iv) How does the metal surface appear after burning?

Observation and discussion

Sodium and potassium react vigorously with oxygen.

$$\begin{gathered} 4\mathrm{Na}(\mathrm{s})+{\mathrm{O}}_2(\mathrm{g})⟶2{\mathrm{Na}}_2\mathrm{O}(\mathrm{s}) \\ 4\mathrm{K}(\mathrm{s})+{\mathrm{O}}_2(\mathrm{g})⟶2{\mathrm{K}}_2\mathrm{O}(\mathrm{s}) \end{gathered}$$

Sodium and potassium burns with a golden yellow and lilac colour flame respectively to form sodium and potassium oxides, which dissolve with water to form alkali called sodium hydroxide and potassium hydroxide.

Magnesium also burns easily, to form magnesium oxide. $2\mathrm{Mg}(\mathrm{s})+{\mathrm{O}}_2(\mathrm{g})⟶2\mathrm{MgO}(\mathrm{s})$

Copper and Aluminium do not burn but on heating in air form black copper (II) oxide and white aluminium oxide ( ${\mathrm{Al}}_2{\mathrm{O}}_3$ ) respectively.

$2\mathrm{Cu}(\mathrm{s})+{\mathrm{O}}_2(\mathrm{g})⟶2\mathrm{CuO}(\mathrm{s})$ Copper Copper (II) oxide (Black)

$4\mathrm{Al}(\mathrm{s})+3{\mathrm{O}}_2(\mathrm{g})⟶2{\mathrm{Al}}_2{\mathrm{O}}_3(\mathrm{s})$ Aluminium Aluminium oxide (White) These metal oxides are found to be insoluble in water.

The order of reactivity with oxygen is : $\mathrm{K}>\mathrm{Na}>\mathrm{Mg}>\mathrm{Al}>\mathrm{Cu}$

At ordinary temperature, the surfaces of metals such as magnesium, zinc and lead, etc. are covered with a thin layer of the oxide. The protective layer of the oxide prevents the metal from further oxidation.

Conclusion Almost all metals combine with oxygen to form metal oxides. Metal + Oxygen $⟶$ Metal oxide

Nature of metallic oxide

Generally, metallic oxides are basic in nature except aluminium and zinc oxides which are amphoteric in nature. That means these oxides $\left({\mathrm{Al}}_2{\mathrm{O}}_3,\mathrm{ZnO}\right)$ react with base as well as acid. The basic oxide of metals react with acid to give salt. For example,

Some oxides of metals dissolve in water and form alkalis. For example,

Reactions showing amphoteric nature of ${\mathrm{Al}}_2{\mathrm{O}}_3$ and ZnO .

Nature of oxides

(2) Reaction of metals with water

Metal reacts with water and produce a metal oxide and hydrogen gas. Metal oxides that are soluble in water dissolve in it to form metal hydroxide. But all metals do not react with water. $\begin{array}{llll}\text{ Metal }& +& \begin{array}{c}\text{ (Steam) }\\ \text{ Water }\end{array}& ⟶\end{array}$ Metal oxide + Hydrogen gas $=\begin{array}{c}\text{ Water }\\ \text{ (Cold/Boiling) }\end{array})⟶$ Metal hydroxide + Hydrogen gas

The following activity clears this reaction properly.

Active Chemistry 10

Aim To study the reactivity of metals with water.

• Anodising is a process of forming a thick oxide layer of aluminium. This oxide coat of aluminium (Al) makes it resistant to further corrosion.

Materials required Samples of sodium, potassium, calcium, magnesium, zinc and copper, beaker, cold water, hot water, burner, glass wool, test tube.

Method (i) Collect the samples of sodium, potassium, calcium, magnesium, zinc and copper. (ii) Put small piece of the samples separately in beakers halffilled with cold water. (iii) Put the metals that do not react with cold water in beaker half-filled with hot water. (iv) For the metals that do not react with hot water, arrange the apparatus (to produce steam) and observe their reaction with steam.

Warning sign displayed on containers containing concentrated acids and bases.

Now answer (i) Which metals reacted with cold water? Arrange them in the increasing order of their reactivity with cold water. (ii) Did any metal produce fire on water? (iii) Does any metal start floating after some time? (iv) Which metals reacted with steam? (v) Which metals did not react even with steam?

Observation and discussion

(i) Na and K metals react vigorously with cold water to form NaOH and KOH respectively and ${\mathrm{H}}_2$ gas is liberated.

(ii)

These reactions are so violent and exothermic that the ${\mathrm{H}}_2$ gas evolved, catches fire. Calcium reacts with cold water to form $\mathrm{Ca}(\mathrm{OH}{)}_2$ and ${\mathrm{H}}_2$ gas. It is less violent.

$Ca(s)+2H_{2}O(l)\to Ca(OH{)}_2(aq)+H_2(g)$ Calcium + Cold water → Calcium hydroxide + Hydrogen gas

(iii) Magnesium reacts with hot boiling water to form $\mathrm{Mg}(\mathrm{OH}{)}_2$ and ${\mathrm{H}}_2$ gas.

$Mg(s)+2H_{2}O(l)\to Mg(OH{)}_2(aq)+H_2(g)$ Magnesium + Boiling water → Magnesium hydroxide + Hydrogen gas

Magnesium and calcium floats on surface of water due to bubbles of ${\mathrm{H}}_2$ gas sticking on its surface.

(iv) Aluminium does not react either with cold or hot water. But it reacts only with steam to form aluminium oxide and hydrogen gas.

$2Al(s)+3H_{2}O(g)\to A{l}_2{O}_3(s)+3H_2(g)$ Aluminium + Steam → Aluminium oxide + Hydrogen gas

Similarly, zinc reacts with steam to form zinc oxide and ${\mathrm{H}}_2$ gas.

$Zn(s)+H_{2}O(g)\to ZnO(s)+H_2(g)$ Zinc + Steam → Zinc oxide + Hydrogen gas

(v) Copper do not react with water even under strong conditions. The above reactions indicate that sodium and potassium are the most reactive metals while copper is less reactive, metals such as lead, copper, silver and gold do not react with water at all.

Conclusion The reactivity order of these metals with water are $\mathrm{K}>\mathrm{Na}>\mathrm{Ca}>\mathrm{Mg}>\mathrm{Al}>\mathrm{Zn}>\mathrm{Fe}>\mathrm{Cu}$ Reactivity with water decreases $⟶$

The reaction (a) Potassium metal (stored in mineral oil to prevent oxidation) and (b) water. (c) The reaction of potassium with water. The flame occurs because of the produced hydrogen gas. $({\mathrm{H}}_2(\mathrm{g})$ burns in air which reacts with ${\mathrm{O}}_2(\mathrm{g})$, at the high temperature)

(3) Reaction of metals with acids

The highly reactive metals react with dilute acid to displace hydrogen from acid and give a salt. Metal + Dilute acid $⟶$ Salt + Hydrogen The reactivity of different metals is different with same acid. This is made clear by the following activity.

Active Chemistry 11

Aim To study the reaction of metals with acids.

• Do not take sodium and potassium in this activity as they react vigorously even with cold water.

Materials required Samples of magnesium, aluminium, iron, copper, test tubes, dilute hydrochloric acid, thermometers.

Method (i) Put the samples of $\mathrm{Mg},\mathrm{Al},\mathrm{Fe}$ and Cu separately in test tubes containing dilute hydrochloric acid. (ii) Suspend thermometers in the test tubes so that their bulbs are dipped in the acid. (iii) Observe the rate of formation of bubbles carefully.

Now answer (i) Which metals reacted vigorously with dilute hydrochloric acid? (ii) With which metal did you record the highest temperature? (iii) Arrange the metals in the decreasing order of reactivity with dilute acids.

Observation and discussion In the test tube which contains Mg , the hydrogen bubbles appear very rapidly.

$Mg(s)+2HCl(aq)\to MgC{l}_2(aq)+H_2(g)$ Magnesium + Hydrochloric acid → Magnesium chloride + Hydrogen

In the test tubes containing Al and Zn , reaction with acid is fast.

$2\mathrm{Al}(\mathrm{s})+6\mathrm{HCl}(\mathrm{aq})⟶2{\mathrm{AlCl}}_3(\mathrm{aq})+3{\mathrm{H}}_2(\mathrm{g})$ Aluminium + Hydrochloric acid $⟶$ Aluminium chloride + Hydrogen

The reaction between Fe and acid is slow. $Fe(s)+2HCl(aq)\to FeC{l}_2(aq)+H_2(g)$ Iron + Hydrochloric acid → Ferrous chloride + Hydrogen

No reaction is observed in the test tube which contain Cu and dil HCl . $\mathrm{Cu}(\mathrm{s})+\mathrm{HCl}(\mathrm{aq})⟶$ No reaction $(x)$ Copper Hydrochloric acid Temperature was found to rise in case of all the metals that reacted with dilute acid showing that reaction is exothermic. The rise in temperature is maximum in case of magnesium.

Conclusion The order of reactivity with dilute HCl is: $\mathrm{Mg}>\mathrm{Al}>\mathrm{Fe}>\mathrm{Cu}$

Important information Hydrogen gas is not evolved when metals such as $\mathrm{Zn},\mathrm{Fe},\mathrm{Cu}$ and Al react with nitric acid. Because ${\mathrm{HNO}}_3$ is strong oxidising agent. It oxidises ${\mathrm{H}}_2$ gas to water and is itself reduced to oxides of nitrogen ( $\mathrm{NO},{\mathrm{N}}_2\mathrm{O}$ and ${\mathrm{NO}}_2$ ).

But copper reacts with hot concentrated sulphuric acid $\left({\mathrm{H}}_2{\mathrm{SO}}_4\right)$ to produce copper sulphate, sulphur dioxide and water.

$Cu(s)+2H_{2}S{O}_4(aq)\to CuS{O}_4(aq)+S{O}_2(g)+2H_{2}O(l)$ Copper + Sulphuric acid → Copper sulphate + Sulphur dioxide + Water

Mg reacts with very dilute ${\mathrm{HNO}}_3$ to evolve ${\mathrm{H}}_2$ gas.

$Mg(s)+2HN{O}_3(aq)\to Mg(N{O}_3{)}_2(aq)+H_2(g)$ Magnesium + Nitric acid (dil) → Magnesium nitrate + Hydrogen

Fe reacts with dil ${\mathrm{H}}_2{\mathrm{SO}}_4$ to evolve ${\mathrm{H}}_2$.

$Fe(s)+dil{H}_{2}S{O}_4\to FeS{O}_4(aq)+H_2(g)$ Iron + Sulphuric acid → Ferrous sulphate + Hydrogen

• Aqua Regia (Royal water): Aqua regia is a Latin word it means "royal water". It is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of $3:1$. It is a highly corrosive, fuming liquid and is used to dissolve gold and platinum.

$\mathrm{HCl}:{\mathrm{HNO}}_3-3$ : 1

(4) Reaction of metal with solutions of other metal salts

When a more reactive metal is placed in a salt solution of less reactive metal, then the more reactive metal displaces the less reactive metal from its salt solution. This reaction is also known as displacement reaction.

Let us learn it with the help of following activity.

Active Chemistry 12

Aim To study the reaction of metals with solutions of other metal salts.

Materials required Clean wire of copper, iron nail, solution of iron sulphate, solution of copper sulphate, test tubes.

Method (i) Take a clean wire of copper and an iron nail. (ii) Put the copper wire in a solution of iron sulphate and the iron nail in a solution of copper sulphate taken in test tubes. (iii) Record your observations after 20 minutes.

It may be noted that hydrogen is not a metal but even then it has been placed in the reactivity series because hydrogen can also lose electron and form positive ion ${\mathrm{H}}^{+}$ like metals.

Now answer (i) In which test tube did you find that a reaction has occurred? (ii) On what basis can you say that a reaction has actually taken place? (iii) Can you correlate observations for the Activities 9, 10 and 11? (iv) Write a balanced chemical equation for the reaction that has taken place. Name the type of reaction.

Observation and discussion Reaction is found to occur in the tube containing iron nail dipped in copper sulphate solution. This is because in this tube, blue colour of copper sulphate solution fades and light green colour due to formation of iron (II) sulphate appears. Moreover, a brown deposit of copper takes place on iron nail. Thus, the following reaction takes place. $\mathrm{Fe}(\mathrm{s})+{\mathrm{CuSO}}_4(\mathrm{aq})⟶{\mathrm{FeSO}}_4(\mathrm{aq})+\mathrm{Cu}(\mathrm{s})$ Iron Copper sulphate Iron (II) sulphate Copper (Blue) (Green) (Brown)

Conclusion Iron is more reactive than copper and displaces copper from copper sulphate solution. But, as copper is less reactive than iron, so it will not displace iron from iron sulphate solution. In general, a more reactive metal displaces a less reactive metal from its salt in the solution.

• Metal A + salt solution of metal B $\to$ salt solution of metal A + metal B. (only if metal A is more reactive than metal B ) Let us discuss one more example of displacement reaction.

Reaction of copper with silver nitrate solution When a strip of copper metal is placed in a solution of ${\mathrm{AgNO}}_3$, the solution becomes gradually blue due to the formation of copper nitrate and a shining coating of silver metal gets deposited on the copper strip. The reaction may be written as :

$2AgN{O}_3(aq)+Cu(s)\to Cu(N{O}_3{)}_2(aq)+2Ag(s)$ Silver nitrate (colourless solution) + Copper → Copper nitrate (blue colour) + Silver

However, if we place silver wire in a copper sulphate solution no reaction occurs. This means copper can displace silver from its salt solution but silver cannot displace copper from its solution. i.e. copper is more reactive metal than silver.

The reactivity series

The arrangement of metals in order of decreasing reactivities is called reactivity series or activity series of metals. After performing displacement experiments the following series has been developed.

Reactivity series of metals

Here is a tip to memorise the activity series of some metals which can help you to perform displacement reaction.

Reactivity series of some metals

In earlier standards, we have seen how elements combine to form stable molecules. $\star$ We have also learnt that the driving force behind formation of a chemical bond is to attain an electronic configuration with a complete octet. $\star$ The atoms attain a complete octet by giving, taking or sharing of electrons with each other. $\star$ In previous class, we have learnt about atomic number, electronic configuration and number of electrons in different shells like - K, L, M, N.

4.0Reaction of metals with non-metals

Atoms of elements combine to form stable molecules. The combining power of an atom is expressed as valency. Each atom has a tendency to attain a completely filled valence shell. The noble gases, which have a completely filled valence shell or outermost shell, are very stable.

• The electronic configuration of noble gases and some metals and non-metals are given in the following table.

Electronic configuration of some elements

It is clear from the above table that except helium, all other noble gases have 8 electrons (octet) in their outermost shell, which represent a highly stable electronic configuration. Due to this stable configuration, the noble gases have no tendency to lose or gain electrons. So they exist in monoatomic form. However, metals and non-metals which do not have complete octet will try to attain stability either by gaining or loosing electrons.

It should be noted that ionic compounds do not exist as molecules but aggregates of oppositely charged ions. These compounds do not bear any charge i.e. they are neutral.

Lets discuss formation of sodium chloride ( NaCl ). Sodium atom has one electron in its outermost shell. If it loses the electron from its M shell then its $L$ shell becomes the outermost shell, which has stable octet like noble gases. The nucleus of this atom still has 11 protons but the number of electrons becomes 10. Therefore, it becomes positively charged sodium ion or cation ( ${\mathrm{Na}}^{+}$).

On the other hand, chlorine has seven electrons in its outer most shell and it require one more electron to complete its octet. The nucleus of chlorine atom has 17 protons and the number of electrons become 18. Therefore, it becomes negatively charged chloride ion ( ${\mathrm{Cl}}^{-}$) or anion.

So, ${\mathrm{Na}}^{+}$and ${\mathrm{Cl}}^{-}$ions being oppositely charged attract to each other and are held by strong electrostatic forces of attraction to exist as NaCl . In other words, ${\mathrm{Na}}^{+}$and ${\mathrm{Cl}}^{-}$ions are held together by electrovalent or ionic bond.

Electrovalent bond or ionic bond may be defined as the electrostatic force of attraction which holds the oppositely charged ions together or it may also be defined as a chemical bond formed between two atoms by complete transfer of electrons from one atom to another so as to complete their octet and hence acquire the stable nearest noble gas configuration. The number of electrons lost or gained by the atom is called its electrovalency. Let us discuss the formation of one more ionic compound magnesium chloride $\left({\mathrm{MgCl}}_2\right)$. The electronic configuration of magnesium ( Mg ) and chlorine ( Cl ) atoms are

Magnesium atom has two electrons in its valence shell, so has a tendency to lose both of its electrons to attain the nearest noble gas configuration i.e. Ne.

On the other hand, chlorine has only one electron less than the nearest noble gas (i.e. Ar) configuration. Magnesium transfers both its valence electrons to two chlorine atoms, each of which needs one electron to form ${\mathrm{Cl}}^{-}$ion.

$Mg\to M{g}^{2+}+2e^{-}$ (2, 8, 2) (2, 8)

$2Cl+2e^{-}\to 2C{l}^{-}$ (2, 8, 7) (2, 8, 8)

Some common ionic compounds are:

Properties of ionic compounds

To learn about the properties of ionic compounds, let us perform the following activity.

Active Chemistry 13

Aim To study the properties of ionic compounds.

Materials required Samples of sodium chloride, potassium iodide, barium chloride, metal spatula, burner, water, kerosene, battery, beaker, bulb, switch, electrodes of graphite.

Method (i) Take samples of sodium chloride, potassium iodide, barium chloride or any other salt from the science laboratory. (ii) Take a small amount of sample on a metal spatula and heat directly on the flame. Repeat with other samples. (iii) Try to dissolve the compound in water and kerosene. (iv) Make a circuit as shown in figure and insert the electrodes into a solution of one salt. Test the other salt samples too in this manner.

Now answer (i) What is the physical state of the salt taken? (ii) Did the samples impart any colour to the flame on heating? (iii) Did the compounds melt on heating? (iv) Are the compounds soluble in water or kerosene? (v) Did the electric bulb glow on passing electric current? (vi) What is your inference about the nature of these compounds?

Observation and discussion (i) All the salts taken are solids. (ii) Each salt imparted a particular colour to the flame. (iii) The compounds did not melt on heating.

Conclusion (i) Ionic compounds are generally solids. (ii) They impart a characteristic colour to the flame. (iii) No, they have high melting point. (iv) They are soluble in a polar solvent like water and insoluble in non-polar solvents like kerosene, petrol, etc. (v) Their molten or aqueous solution conduct electricity. (iv) Following are the general properties of ionic compounds.

(a) Physical state Ionic compounds are solids and relatively hard because of the strong force of attraction between the positive and negative ions. This force of attraction is also known as strong electrostatic force of attraction. These compounds are generally brittle and break into pieces when pressure is applied.

(b) Solubility Electrovalent compounds are generally soluble in water (because of their polar nature) and insoluble in solvents such as kerosene, petrol, etc.

(c) Melting and boiling points Ionic compounds have high melting and boiling points, due to the strong electrostatic force of attraction between the oppositely charged ions. Therefore, large amount of energy is needed to break these bonds.

(d) Conduction of electricity Ionic compounds in the solid state do not conduct electricity because movement of ions in the solid state is not possible due to their rigid structure. But they can conduct electricity in molten or aqueous state.

(e) Colour to the flame Most of the salts when brought into the flame, impart characteristic colour to the flame. Flame colours are produced from the movement of the electrons in the metal ions present in the compounds.

• Here are some important physical properties of metals and nonmetals.

• Here are some important chemical properties of metals and nonmetals.

Why sodium chloride has high melting point? Explanation: Sodium chloride consist of ${\mathrm{Na}}^{+}$and ${\mathrm{Cl}}^{-}$ions. These oppositely charged ions are strongly attracted towards each other. To break these strong forces of attraction, a large amount of energy is needed and hence sodium chloride has a high melting point.

5.0Occurrence of metals

The earth's crust is the major source of metals. They are present in nature in the free state as well as in combined state. Oxygen & silicon are the main elements present in earth's crust. Both are non-metals. Seawater also contains some soluble salts such as $\mathrm{NaCl},{\mathrm{MgCl}}_2$, etc.

Native and combined states of metals

Metals occur in the crust of earth in two states : native state and combined state.

(1) Native state A metal is said to occur in native or free state when a metal is found in nature in the elementary or metallic state. The metals at the bottom of the activity series are least reactive. They are often found in free state. For example, Gold, silver, copper and platinum are found in free state because they are very unreactive metals. So, they have no tendency to react with oxygen and they do not react with moisture, ${\mathrm{CO}}_2$ of air or any other non-metal.

(2) Combined state The metals at the top of reactivity series are not expected to occur in free state due to their reactive nature. They exist in combination with other elements as oxides, carbonates, halides, sulphates, sulphides, etc. For example, Sodium, potassium, calcium, aluminium, magnesium, etc. are very reactive metals. All of these are lying at the top of activity series. These are never found in the free state. The metals in the middle of the activity series such as zinc, iron, lead, etc. are moderately reactive. They are found in the earth crust mainly as oxides, sulphide or carbonates.

• On the basis of reactivity, metals are divided into the following three categories.
  
  Occurrences of metals in the activity series

Minerals and ores

The elementary state of the compounds in the form of which the metals occur in nature are called minerals. The minerals from where metals can be conveniently and profitably extracted are called ores. For example, Copper occurs in nature in the form of several mineral like copper pyrites $\left({\mathrm{CuFeS}}_2\right)$, copper glance ( ${\mathrm{Cu}}_2\mathrm{S}$ ) and cuprite ( ${\mathrm{Cu}}_2\mathrm{O}$ ). We obtain copper metal profitably from copper pyrites mineral, so it is called ore of copper.

6.0Extraction of metals: Metallurgy

The process through which a pure metal is extracted from its ores is known as extraction of metals. The series of various processes involved in the extraction of metals from their ores, followed by refining of the metal is known as metallurgy.

Various steps involved in the extraction of metals or metallurgical process:

• Crushing and grinding of the ore.
• Concentration of the ore or enrichment of the ore.
• Extraction of metal from the concentrated ore.
• Refining or purification of the impure metal.

Crushing and grinding of the ore

Most of the ores in nature occur as big rocks. They are broken into small pieces with the help of crushers. These pieces are then reduced to fine powder with the help of a ball mill or a stamp mill. This process is known as pulverization of the ore.

Enrichment of ore or concentration of ore

The ores mined from the earth's crust contain a number of impurities, such as soil, sand, etc. called gangue or matrix. The process of removal of impurities (gangue) from the ore is called enrichment of ore or concentration of ore. The impurities must be removed from the ore prior to the extraction of the metal. The processes used for removing the gangue from the ore are based on the differences between the physical or chemical properties of the gangue and the ore. Different separation techniques are accordingly employed.

Extraction of metal from the enriched ore

The method used for extraction of the metal from the concentrated ore depends upon the nature of metal.

Based on the reactivity, the metals have been grouped into the following three categories: (I) Metals of low reactivity. (Low in the activity series) (II) Metals of medium reactivity. (In the middle of the activity series) (III) Metals of high reactivity. (At the top of the activity series) (I) Extraction of metals low in the activity series ( $\mathbf{Cu},\mathbf{Hg},\mathbf{Ag},\mathrm{Pt},\mathrm{Au}$ )

As these metals are very less reactive, they are either found in native state or in the form of sulphide ores. For example, Cinnabar (HgS)

These sulphide ores can be converted to oxide ores on heating in the presence of excess of air called Roasting.

$2HgS(s)+3O_2(g)\stackrel{\text{Heat}}{\to }2HgO(s)+2S{O}_2(g)$ Mercuric sulphide (Cinnabar) + Oxygen $\stackrel{\text{Heat}}{\to }$ Mercuric oxide + Sulphur dioxide

This oxide can be reduced to metal by further heating.

$2HgO(s)\stackrel{\text{Heat}}{\to }2Hg(l)+O_2(g)$ Mercuric oxide $\stackrel{\text{Heat}}{\to }$ Mercury + Oxygen

Similarly, when copper glance ( Cu 2 S ) an ore of copper, is subjected to roasting, it directly gives copper.

(II) Extraction of metals in the middle of the activity series ( $\mathrm{Fe},\mathrm{Zn},\mathrm{Pb}$, etc.)

These metals are found in the form of their oxides, sulphides and carbonates. For easy extraction, sulphide and carbonate ores are first converted into the oxide. (a) Conversion into metal oxide (i) Calcination : for the conversion of carbonate ores into oxides. It is the process of heating the ore strongly in the absence of air. For example, ${\mathrm{ZnCO}}_3(\mathrm{s})\underset{\text{ (Absence of air) }}{\overset{\text{ Heat }}{\to }}\mathrm{ZnO}(\mathrm{s})+{\mathrm{CO}}_2(\mathrm{g})$ (ii) Roasting : for the conversion of sulphide ores into oxides. It is the process of heating the ore strongly in the presence of excess of air. $2\mathrm{ZnS}(\mathrm{s})+3{\mathrm{O}}_2(\mathrm{g})\underset{\text{ Presence of excess of air }}{\overset{\text{ Heat }}{\to }}2\mathrm{ZnO}(\mathrm{s})+2{\mathrm{SO}}_2(\mathrm{g})$ Zinc sulphide Zinc Sulphur (Zinc blende-ore of Zn ) oxide dioxide

(b) Reduction of the metal oxide to metal For reduction suitable reducing agents are used, like carbon, carbon monoxide, aluminium, sodium or calcium. (i) Reduction by heating with carbon (coke) - When zinc oxide is heated with carbon, zinc metal is produced. $ZnO(s)+C(s)⟶Zn(s)+CO(g)$ Reduction by carbon is also known as smelting. Similarly, iron and lead are also obtained from their oxides by heating with carbon. $F{e}_2{O}_3(s)+3C(s)\stackrel{\text{Heat}}{\to }2Fe(s)+3CO(g)$ $PbO(s)+C(s)\stackrel{\text{Heat}}{\to }Pb(s)+CO(g)$

(ii) Reduction with CO Iron is obtained from ferric oxide by heating with CO. ${\mathrm{Fe}}_2{\mathrm{O}}_3(\mathrm{s})+3\mathrm{CO}(\mathrm{g})\stackrel{\text{ Heat }}{\to }2\mathrm{Fe}(\mathrm{s})+3{\mathrm{CO}}_2(\mathrm{g})$

(iii) Reduction by aluminium Besides using carbon (coke) to reduce metal oxides to metals, sometimes displacement reactions can also be used. The highly reactive metals such as sodium, calcium, aluminium, etc. are used as reducing agents because they can displace metals of lower reactivity from their compounds.

Certain metal oxides are reduced by aluminium to metals. This method is known as aluminothermy or thermite process. For example, Chromium, manganese, vanadium metals are obtained by the reduction of their oxides with Al powder. The following reaction takes place. $3{\mathrm{MnO}}_2(\mathrm{s})+4\mathrm{Al}(\mathrm{s})\stackrel{\text{ Heat }}{\to }3\mathrm{Mn}(\ell )+2{\mathrm{Al}}_2{\mathrm{O}}_3(\mathrm{s})+$ Heat ${\mathrm{Cr}}_2{\mathrm{O}}_3(\mathrm{s})+2\mathrm{Al}(\mathrm{s})\stackrel{\text{ Heat }}{\to }2\mathrm{Cr}(\ell )+{\mathrm{Al}}_2{\mathrm{O}}_3(\mathrm{s})+$ Heat These displacement reactions are highly exothermic, so a large amount of heat is evolved and metals are produced in the molten state. In fact the reaction of iron(III) oxide $\left({\mathrm{Fe}}_2{\mathrm{O}}_3\right)$ with aluminium, is used to weld railway tracks or cracked machine parts. This reaction is known as thermite reaction. The mixture of iron oxide and aluminium powder is called thermite. ${\mathrm{Fe}}_2{\mathrm{O}}_3(\mathrm{s})+2\mathrm{Al}(\mathrm{s})\stackrel{\text{ Heat }}{\to }2\mathrm{Fe}(\ell )+{\mathrm{Al}}_2{\mathrm{O}}_3(\mathrm{s})+$ Heat

(III) Extraction of metals high up in the activity series ( $\mathrm{K},\mathrm{Ca},\mathrm{Na},\mathrm{Mg}$ and Al ) The highly reactive metals such as $\mathrm{K},\mathrm{Na},\mathrm{Mg}$ have strong affinity for oxygen, so they can not be reduced with the help of carbon. Hence these metals are obtained by electrolysis of their molten or fused oxides or chlorides, this method is called electrolytic reduction. On electrolysis, metal ions, being positive, are liberated at the cathode (negative electrode) where they gain electrons and convert in the metal atoms. For examples, (i) Sodium metal is obtained by electrolysis of molten sodium chloride. $\mathrm{NaCl}(\mathrm{s})\underset{\text{ to melt }}{\overset{\text{ Heat }}{\to }}{\mathrm{Na}}^{+}(\ell )+\mathrm{Cl}-(\ell )$ At Cathode: ${\mathrm{Na}}^{+}(\ell )+{\mathrm{e}}^{-}⟶\mathrm{Na}(\mathrm{s})$ (Reduction) Sodium ion electron Sodium metal At Anode : $C{l}^{-}(l)⟶Cl(g)+e^{-}$ $Cl(g)+Cl(g)⟶C{l}_2(g)$ Thus, sodium metal is obtained at cathode whereas chlorine gas is obtained at the anode.

(ii) Aluminium oxide is reduced to aluminium by the electrolysis of molten aluminium oxide. ${\mathrm{Al}}_2{\mathrm{O}}_3(\mathrm{s})\underset{\text{ to melt }}{\overset{\text{ Heat }}{\to }}2{\mathrm{Al}}^{3+}(\ell )+3{\mathrm{O}}^{2-}(\ell )$ At Cathode: ${\mathrm{Al}}^{3+}(\ell )+3\mathrm{e}⟶\mathrm{Al}(\mathrm{s})$ (Reduction) At Anode : ${\mathrm{O}}^{2-}(\ell )⟶\mathrm{O}(\mathrm{g})+2{\mathrm{e}}^{-}$(Oxidation) Oxide ions Oxygen atom $\mathrm{O}(\mathrm{g})+\mathrm{O}(\mathrm{g})⟶{\mathrm{O}}_2(\mathrm{g})$ Oxygen atoms Oxygen gas

Refining of impure metals

The metals produced by various reduction processes described above are not very pure. They contain impurities, which must be removed to obtain pure metals. The most widely used method for refining of impure metals is electrolytic refining.

Process (i) In this process, the impure metal is made the anode and a thin strip of pure metal is made the cathode. (ii) A solution of the metal salt is used as an electrolyte. On passing the electric current through the electrolyte, the pure metal from the anode dissolves into the electrolyte. (iii) An equivalent amount of pure metal from the electrolyte gets deposited on the cathode. The soluble impurities go into the solution, leaving the insoluble impurities which settle down at the bottom of the anode. At Anode: $M(s)⟶M^{n+}(aq)+n{e}^{-}$ At Cathode: $M^{n+}(aq)+n{e}^{-}⟶M(s)$

Anode mud / anode sludge

The soluble impurities present in the impure metal pass into solution whereas insoluble impurities fall below the anode as anode mud. For example, Electrolytic refining of copper.

Electrolytic refining of copper. The electrolyte is a solution of acidified copper sulphate. The anode is impure copper, whereas the cathode is a strip of pure copper. On passing electric current, pure copper is deposited on the cathode.

Explain why carbon can reduce copper oxide to copper but not sodium oxide to sodium? Explanation: Carbon is a strong reducing agent. Hence, it can reduce copper oxide to copper as follows. $\mathrm{CuO}(\mathrm{s})+\mathrm{C}(\mathrm{s})⟶\mathrm{Cu}(\mathrm{s})+\mathrm{CO}(\mathrm{g})$ Sodium is much more reactive than copper. It has greater affinity for oxygen than the affinity for carbon. Moreover, at high temperature, sodium can combine with carbon to form sodium carbide.

• Process of metallurgy
  

7.0Corrosion

When the surface of a metal is attacked by the gases and water vapour present in the air, it is said to corrode and the phenomenon is called corrosion. Thus, corrosion may be defined as follows: The process of slowly eating up of metals due to their conversion into oxides, carbonates, sulphide, sulphates, etc. by the action of atmospheric gases and moisture is called corrosion. In case iron is the metal involved in the chemical process, then corrosion is called rusting.

Factors which promote corrosion (a) Position of metal in the reactivity series: Active metals placed above hydrogen are easily corroded as compared to metals which are placed below hydrogen. (b) Presence of water vapours and gases like ${\mathrm{CO}}_2,{\mathrm{SO}}_2$, etc. in the air. (c) Presence of salts or electrolyte in water promotes corrosion. e.g. Rusting of iron is faster in sea water than in ordinary or distilled water.

Active Chemistry 14

Aim To prove that air and moisture are essential for rusting of iron.

Materials required Three test tubes, iron nails, water, oil, anhydrous calcium chloride (drying agent), corks.

Method (i) Take three test tubes and place clean iron nails in each of them. (ii) Label these test tubes A, B and C. Pour some water in test tube A and cork it. (iii) Pour boiled distilled water in test tube B and add about 1 ml of oil and cork it. The oil will float on water and prevent the air from dissolving in the water. (iv) Put some anhydrous calcium chloride in test tube C and cork it. Anhydrous calcium chloride will absorb the moisture, if any, from the air. Leave these test tubes for a few days and then observe.

Observation It is observed that iron nails rust in the tube A but they do not rust in test tubes B and C . In test tube A , the nails are exposed to both air and water. In test tube $B$, the nails are exposed to water only and the nails in test tube C are exposed to dry air free from water vapour.

Conclusion Presence of both air and moisture is essential for rusting to take place.

Rusting is the term used only in case of iron. If the rusted surface of iron is rubbed with a sand paper, the rust will appear again in a few days. This shows that the rust is formed by a chemical reaction and not by a physical process.

Example of corrosion: (i) When iron is exposed to moist air for a long time, its surface acquires a brown flaky substance called rust and the process is known as rusting. Rust is a mixture of ${\mathrm{Fe}}_2{\mathrm{O}}_3$ and $\mathrm{Fe}(\mathrm{OH}{)}_3$. (ii) Copper reacts with ${\mathrm{CO}}_2$ in the air and slowly loses its shiny brown surface and acquires a green coating of basic copper carbonate in moist air. $2Cu(s)+C{O}_2(g)+O_2(g)+H_{2}O(l)⟶CuC{O}_3.Cu(OH{)}_2$ (iii) Silver articles becomes black after sometime when exposed to air. [Due to reaction with ${\mathrm{H}}_2\mathrm{S}$ in the air to form a black coating of silver sulphide( ${\mathrm{Ag}}_2\mathrm{S})$.] (iv) Lead or stainless steel lose their lusture due to corrosion.

Prevention of corrosion

(i) By painting : The corrosion of a metal can be prevented simply by painting the metal surface by grease or varnish that forms a protective layer on the surface of the metal which protect the metal from moisture and air.

(ii) Self prevention : Some metals form protective layers. For example, When zinc is left exposed to the atmosphere, it combines with the oxygen of air to form a layer of zinc oxide over its surface. The oxides layer does not allow air to go inside the metal. Thus, zinc is protected from corrosion by its own protective layer. Similarly, aluminium combines with oxygen to form a dull layer of aluminium oxide on its surface which protects aluminium from further corrosion.

(iii) Cathodic protection : In this method, the more reactive metal which is more corrosionprone is connected to a bar of another metal which is less reactive and to be protected. In this process, electron flows from more reactive metal to the less reactive metal. The metal to be protected becomes the cathode and the more reactive metal becomes the anode. In this way, the two metals form an electrochemical cell and oxidation of the metal is prevented. For example, The pipelines (iron) under the surface of the earth are protected from corrosion by connecting them to a more reactive metal (magnesium or Zn ) which is buried in the earth and connected to the pipelines by a wire.

(iv) Oiling and greasing : Both protect the surface of metal against moisture and chemicals, etc. In addition the oil and grease prevent the surface from getting scratched.

(v) Electroplating : It is a very common and effective method to check corrosion. The surface of metal is coated with chromium, nickel or aluminium, etc. by electrolysis also called electroplating. They are quite resistant to the attack by both air and water. If the surface of metal is electroplated by zinc, it is known as galvanisation and in case tin metal is used, then the process is called tinning.

(vi) By alloying : It is a very good method of improving the properties of a metal. An alloy is a homogeneous mixture of two or more metals or non-metal. It can be prepared by first melting the metal and then dissolving the other elements (metal or non-metal) in proper proportions. The physical properties of an alloy are different from the constituent metals (from which it is made).

Some of the common alloys are

(i) Steel : Iron is the most widely used metal. But it is never used in its pure state. This is because pure iron is very soft and stretches easily when hot. But, if it is mixed with a small amount of carbon (about 0.05%) it becomes hard and strong. When iron is mixed with nickel and chromium to form stainless steel which is hard and does not rust.

(ii) Amalgam : An alloy of mercury and one or more other metals is known as an amalgam. It may be solid or liquid. A solution of sodium metal in liquid mercury metal is called sodium amalgam, which is used as a reducing agent. Amalgam of silver, tin and zinc is used by dentists for filling in teeth.

(iii) Brass : Brass is an alloy of copper ( Cu ) and Zn . It contains $80\%$ copper and $20\%$ zinc. It is more malleable and more stronger than pure copper. Brass is used for making cooking utensils, condenser sheets, pipe, screws, bolts, wire, scientific instruments, ornaments, etc.

(iv) Bronze : It is also the alloy of copper. It contain $90\%$ of copper and $10\%$ tin. It is highly resistant to corrosion and

(v) Solder : It is an alloy of lead (50%) and tin (50%). It is used for soldering (or welding) electrical wires together as it melts at a low temperature.

(vi) Alloys of Gold : The purity of gold is expressed is terms of 'carats'. Pure gold is known as 24 carats gold. It is very soft due to which, it is not suitable for making jewellery. It is alloyed with either silver or copper to make it hard and more suitable for making ornaments. In India, gold ornaments are usually made of 22 carats gold. It is an alloy of gold with silver or copper.

The wonder of ancient Indian metallurgy

The iron pillar near the Qutub-Minar in Delhi was made around 400 BC by the iron workers of India. They had developed a process which prevented the wrought iron pillar(a type of iron) from rusting even after thousands of years. This is likely because of formation of a thin film of magnetic oxide $\left({\mathrm{Fe}}_3{\mathrm{O}}_4\right)$ on the surface as a result of finishing treatment given to the pillar, painting it with a mixture of different salts then heating and quenching (rapid cooling). The iron pillar is 8

• Some important alloys are mentioned below -

8.0Basic terminology

• Hammered - To hit something several times.
• Resistance - Trying to stop something from happening.
• Lustre - Shine
• Clamp - A tool that is used to hold two thing tightly.
• Brittle - To break into particles.
• Violent - Very strong and impossible to control.
• Affinity - Tendency to combine.
• Malleability - To form sheets
• Ductility - To form wires.
• Conductivity - Property to allow heat or electricity to go through something.
• Sonorous - Production of sound.
• Enrichment - To increase the percentage.
• Refining - To purify.
• Lilac - Light purple or pale violet tone.

9.0Mindmap