CBSE NOTES CLASS 10 SCIENCE CHAPTER 3

METALS AND NON-METALS

Physical Properties of Metals and Non Metals

 Property Metals Non Metals Lustre (shine) Lustrous (shiny) Have no luster, except Iodine. Colour Usually grey or white All colours State of matter at room temperature Usually solid (exception is mercury) Exist as solid, liquid or gas. Hardness Metals are generally hard. Exception Na, K Not as hard as metals. Exception diamond. Melting Point and Boling Point Have high MP and BP. Exception Gallium and Caesium. Low melting and boiling points. Exception diamond. Ductility Ductile Not ductile Malleability Malleable Not malleable Conductivity Good conductors of heat and electricity. Lead and mercury are poor conductors. Bad conductors of heat and electricity. Exception graphite. Sonority Sonorous Not sonorous

Ductility

Property of metals by virtue of which they can be drawn into thin wires is called ductility.

Malleability

Property of metals by virtue of which they can be beaten into thin sheet is called malleability.

Sonority

Property of metals, by virtue of which they make a bell-like sound when struck.

Exceptions

(i) All metals except mercury exist as solids at room temperature.

(ii) Usually metals have high melting points but gallium and caesium have very low melting points. These two metals will melt if you keep them on your palm.

(iii) Iodine and carbon (graphite) are a non-metal but they are lustrous.

(iv) Carbon is a non-metal that can exist in different forms called allotropes. Diamond, an allotrope of carbon, is the hardest natural substance known and has a very high melting and boiling point. Graphite, another allotrope of carbon, is a conductor of electricity.

(v) Alkali metals (lithium, sodium, potassium) are so soft that they can be cut with a knife. They have low densities and low melting points.

(vi) Bromine is the only liquid non metal.

CHEMICAL PROPERTIES OF METALS

Reaction of Metals with Oxygen (Burning)

Almost all metals combine with oxygen to form metal oxides.

Metal + Oxygen → Metal oxide

Depending on the reactivity of the metal, they react differently with oxygen.

• Metals such as potassium and sodium react so vigorously that they catch fire if kept in the open. Hence, to protect them and to prevent accidental fires, they are kept immersed in kerosene oil.

4Na(s) + O2(g) → 2Na2O(s)

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

4Al + 3O2 → 2Al2O3 (Aluminium oxide)

• Iron does not burn on heating but iron filings burn vigorously when sprinkled in the flame of the burner.

4Fe + 3O2 → 2Fe2O3 (Iron(III) oxide)

• When copper is heated in air, it combines with oxygen to form copper(II) oxide, a black oxide.

2Cu + O2 → 2CuO (Copper(II) oxide)

• Silver, gold and platinum do not react with oxygen even at high temperatures.

Behaviour of Metal Oxides

• Most metal oxides are basic in nature.

Na2O + 2HCl → NaCl + H2O

• Amphoteric Oxides

Some metal oxides, such as aluminium oxide, zinc oxide, etc., show both acidic as well as basic behaviour.

Metal oxides which react with both acids as well as bases to produce salts and water, are known as amphoteric oxides.

Al2O3 + 6HCl → 2AlCl3 + 3H2O

ZnO + 2HCl → ZnCl2 + H2O

• Most metal oxides are insoluble in water but sodium oxide, potassium oxide and calcium oxides dissolve in water to produce alkalis

Na2O(s) + H2O(l) → 2NaOH(aq)

K2O(s) + H2O(l) → 2KOH(aq)

CaO(s) + 2H2O(l) → Ca(OH)2(aq)

Anodising

Anodising is a process of forming a thick oxide layer of aluminium. Aluminium develops a thin oxide layer when exposed to air. This aluminium oxide coat makes it resistant to further corrosion. The resistance can be improved further by making the oxide layer thicker. During anodising, a clean aluminium article is made the anode and is electrolysed with dilute sulphuric acid. The oxygen gas evolved at the anode reacts with aluminium to make a thicker protective oxide layer. This oxide layer can be dyed to give aluminium articles with an attractive finish.

Reactions of Metals with Water

• Metals react with water and produce a metal oxide and hydrogen gas.

Metal + Water → Metal oxide + Hydrogen

But all metals do not react the same way with water.

Metal oxides that are soluble in water dissolve in it to further form metal hydroxide.

Metal oxide + Water → Metal hydroxide

• Metals like potassium and sodium react violently with cold water. The reaction is so violent and exothermic that the evolved hydrogen immediately catches fire.

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

2Na(s) + 2H2O(l) → 2NaOH(aq)+H2(g)+ Δ

• The reaction of calcium with water is less violent. The heat evolved is not sufficient for the hydrogen to catch fire.

Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)

Calcium starts floating because the bubbles of hydrogen gas formed stick to the surface of the metal.

• Magnesium does not react with cold water. It reacts with hot water to form magnesium hydroxide and hydrogen. It also starts floating due to the bubbles of hydrogen gas sticking to its surface.

Mg(s) + 2H2O(l) → MgOH)2(aq) + H2(g)

• Metals like aluminium, iron and zinc do not react either with cold or hot water. But they react with steam to form the metal oxide and hydrogen.

2Al(s) + 3H2O(g) → Al2O3(s) + 3H2(g)

3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g)

• Metals such as lead, copper, silver and gold do not react with water at all.

Reactions of Metals with Acids

Metals react with acids to give a salt and hydrogen gas.

2Na + 2HCl → 2NaCl + H2

Ca + 2HCl → CaCl2 + H2

Zn + H2SO4→ ZnSO4 + H2

Hydrogen gas is not evolved when a metal reacts with nitric acid. It is because HNO3 is a strong oxidising agent. It oxidises the H2 produced to give water and itself gets reduced to any of the nitrogen oxides (N2O, NO, NO2). But magnesium (Mg) and manganese (Mn) react with very dilute HNO3 to evolve H2 gas.

The reactivity decreases in the order Mg > Al > Zn > Fe.

This shows that copper does not react with dilute HCl.

Aqua regia (Latin for ‘royal water’):

Aqua regia is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3:1. Aqua regia is a highly corrosive, fuming liquid. It is one of the few reagents that are able to dissolve gold and platinum. It can dissolve gold, even though neither of these acids can do so alone.

Reactions of Metals with Compounds of Other Metals

Reactive metals can displace less reactive metals from their compounds in solution or molten form.

CuSO4 + Zn → ZnSO4 + Cu

The Reactivity Series

• If metal A displaces metal B from its solution, it is more reactive than B.

Metal A + Salt solution of B → Salt solution of A + Metal B

• The reactivity series is a list of metals arranged in the order of their decreasing activities.
 K Potassium Most reactive Na Sodium Ca Calcium Mg Magnesium Al Aluminium Zn Zinc Fe Iron Pb Lead H Hydrogen Cu Copper Hg Mercury Ag Silver Au Gold Least reactive

Reaction of Metals with Non-metals

Elements take part in chemical reaction to attain a completely filled valence shell or a noble gas configuration.

Atoms of metals loose electron to attain the stable octet and give positively charged ions (cation).

Similarly atoms of non-metals gain electron to attain the stable octet and give negatively charged ions (anion).

• Ionic Compounds:The compounds formed by the transfer of electrons from a metal to a non-metal are known as ionic compounds or electrovalent compounds.

Examples: Sodium Chloride, Magnesium Chloride, Sodium Oxide etc.

• A sodium atom has one electron in its outermost shell. If it loses the electron from its M shell then its L shell now becomes the outermost shell and that has a stable octet. The nucleus of this atom still has 11 protons but the number of electrons has become 10, so there is a net positive charge giving a sodium cation Na+.
 Na → Na+ (Sodium cation) + e– 2,8,1 2,8

• Chlorine has seven electrons in its outermost shell and it requires one more electron to complete its octet. If sodium and chlorine were to react, the electron lost by sodium could be taken up by chlorine.

After gaining an electron, the chlorine atom gets a unit negative charge, because its nucleus has 17 protons and there are 18 electrons in its K, L and M shells. This gives a chloride anion Cl.

 Cl + e– → Cl- (Chloride ion) 2,8,7 2,8,8,

• Sodium and chloride ions, being oppositely charged, attract each other and are held by strong electrostatic forces of attraction to exist as sodium chloride (NaCl).

Sodium chloride does not exist as molecules but aggregates of oppositely charged ions.

Also,

 Mg → Mg2+ (Magnesium cation) + 2e– 2,8,2 2,8 Cl + e– → Cl- (Chloride ion) 2,8,7 2,8,8

General properties of ionic compounds

(i) Physical nature: Ionic compounds are solids and are somewhat hard because of the strong force of attraction between the positive and negative ions. These compounds are generally brittle and break into pieces when pressure is applied.

(ii) Melting and Boiling points: Ionic compounds have high melting and boiling points. This is because a considerable amount of energy is required to break the strong inter-ionic attraction.

(iii) Solubility: Electrovalent compounds are generally soluble in water and insoluble in solvents such as kerosene, petrol, etc.

(iv) Conduction of Electricity: The conduction of electricity through a solution involves the movement of charged particles. A solution of an ionic compound in water contains ions, which move to the opposite electrodes when electricity is passed through the solution. Ionic compounds in the solid state do not conduct electricity because movement of ions in the solid is not possible due to their rigid structure. But ionic compounds conduct electricity in the molten state. This is possible in the molten state since the elecrostatic forces of attraction between the oppositely charged ions are overcome due to the heat. Thus, the ions move freely and conduct electricity.

Occurrence of Metals

Minerals

The elements or compounds, which occur naturally in the earth’s crust, are known as minerals.

Ores

Minerals which contain a very high percentage of a particular metal and the metal can be profitably extracted from them, are called ores.

Gangue

The impurities, like sand, in an ore are called gangue.

Different techniques 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.

The process of removal of gangue from the ore is called concentration of ore.

EXTRACTION OF METALS

On the basis of reactivity, we can group the metals into the following three categories

(i) Metals of low reactivity;

(ii) Metals of medium reactivity;

(iii) Metals of high reactivity.

Different techniques are to be used for obtaining the metals falling in each category.

 $\begin{array}{c}\mathrm{Potassium}\\ \mathrm{Sodium}\\ \mathrm{Calcium}\\ \mathrm{Aluminium}\\ \mathrm{Magnesium}\end{array}}$ Electrolysis Reduction with C $\begin{array}{c}\mathrm{Copper}\\ \mathrm{Mercury}\end{array}}$ Roasting $\begin{array}{c}\mathrm{Silver}\\ \mathrm{Gold}\end{array}}$ Free state

Extracting Metals Low in the Activity Series

The oxides of these metals can be reduced to metals by heating. For example, cinnabar (HgS) is an ore of mercury. When it is heated in air, it is first converted into mercuric oxide (HgO) and then to mercury on further heating.

2HgO(s) 2Hg(l) + O2 (g)

Similarly, copper which is found as Cu2S in nature can be obtained from its ore by just heating in air.

2Cu2S + 3O2 (g) 2Cu2O(s) + 2SO2 (g)

2Cu2O + Cu2S 6Cu(s) + SO2 (g)

Extracting Metals in the Middle of the Activity Series

• The metals in the middle of the activity series such as iron, zinc, lead, etc., are moderately reactive. These are present as sulphides or carbonates in nature.

• It is easier to obtain a metal from its oxide, as compared to its sulphides and carbonates.

• Prior to reduction, the metal sulphides and carbonates must be converted into metal oxides.

• Roasting: The sulphide ores are converted into oxides by heating strongly in the presence of excess air. This process is known as roasting.

2ZnS(s) + 3O2(g) 2ZnO(s) + 2SO2(g)

• Calcination: The carbonate ores are changed into oxides by heating strongly in limited air. This process is known as calcination.

ZnCO3(s) ZnO(s) + CO2(g)

• Reduction: The metal oxides are then reduced to the corresponding metals by using suitable reducing agents such as carbon.

ZnO(s) + C(s) → Zn(s) + CO(g)

• Obtaining metals from their compounds is a reduction process.

• The highly reactive metals such as sodium, calcium, aluminium, etc., can be used as reducing agents because they can displace metals of lower reactivity from their compounds.

3MnO2(s) + 4Al(s) → 3Mn(l) + 2Al2O3(s) + Δ

• These displacement reactions are highly exothermic. The amount of heat evolved is so large that the metals are produced in the molten state.

• Thermite reaction: The reaction of iron(III) oxide (Fe2O3) with aluminium is used to join railway tracks or cracked machine parts. This reaction is known as the thermite reaction.

Fe2O3(s) + 2Al(s) → 2Fe(l) + Al2O3(s) + Δ

Extracting Metals towards the Top of the Activity Series

• The metals high up in the reactivity series are very reactive. They cannot be obtained from their compounds by heating with carbon.

This is because these metals have more affinity for oxygen than carbon.

• These metals are obtained by electrolytic reduction of their molten salts. For example, sodium, magnesium and calcium are obtained by the electrolysis of their molten chlorides.

• The metals are deposited at the cathode (the negatively charged electrode), whereas, chlorine is liberated at the anode (the positively charged electrode).

• The reactions are,

In the solution: Na ↔ Na+ + Cl-

At cathode: Na+ + e → Na

At anode: 2Cl → Cl2 + 2e

• Similarly, aluminium is obtained by the electrolytic reduction of aluminium oxide.

In the solution: Al2O3 ↔ 2Al3+ + 3O2-

At cathode: Al3+ + 3e → Al

At anode: 2O2– → O2 + 4e

Refining of 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 impure metals is electrolytic refining.

• Electrolytic Refining: Many metals, such as copper, zinc, tin, nickel, silver, gold, etc., are refined electrolytically.

In this process, the impure metal is made the anode and a thin strip of pure metal is made the cathode. A solution of the metal salt is used as an electrolyte.

On passing the current through the electrolyte, the pure metal from the anode dissolves into the electrolyte. An equivalent amount of pure metal from the electrolyte is deposited on the cathode.

The soluble impurities go into the solution.

Anode mud: The insoluble impurities settle down at the bottom of the anode and are known as anode mud.

Corrosion

It is the process in which metals are slowly eaten up by the action of air, moisture or chemicals.

• For example rusting is a form of corrosion in which iron is eaten up by the action of air and moisture and a reddish brown coating of iron oxide, called rust, is formed as shown in the following chemical reaction.

3Fe + 4H2O → F3O4 + H2

• Silver articles become black after some time when exposed to air. This is because it reacts with sulphur in the air to form a coating of silver sulphide.

• Copper reacts with moist carbon dioxide in the air and slowly loses its shiny brown surface and gains a green coat. This green substance is copper carbonate.

• Corrosion causes damage to car bodies, bridges, iron railings, ships and to all objects made of metals, especially those of iron. Corrosion of iron is a serious problem. Every year an enormous amount of money is spent to replace damaged iron.

Methods of protecting the metals from corrosion

• Painting, oiling, greasing, galvanising, Electroplating (e.g., chrome plating), anodising or making alloys are methods of protecting metals from corrosion.

• Galvanisation is a method of protecting steel and iron from rusting by coating them with a thin layer of zinc.

The galvanised article is protected against rusting even if the zinc coating is broken, how?

The outer layer of zinc of any galvanized material reacts with the atmospheric oxygen to form Zinc Oxide (ZnO), which is stronger than Zinc. Thus, even if the outer layer of zinc undergoes corrosion, the material is getting coated with a stronger substance (ZnO), and thus is better able to resist corrosion.

• Alloying is a very good method of improving the properties of a metal.

The properties of any metal can be changed if it is mixed with some other substance.

The substance added may be a metal or a non-metal.

An alloy is a homogeneous mixture of two or more metals, or a metal and a nonmetal.

It is prepared by first melting the primary metal, and then, dissolving the other elements in it in definite proportions. It is then cooled to room temperature.

For example, 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, we get stainless steel, which is hard and does not rust. Thus, if iron is mixed with some other substance, its properties change.

• The wonder of ancient Indian metallurgy

The iron pillar near the Qutub Minar in Delhi was built more than 1600 years ago by the iron workers of India. They had developed a process which prevented iron from rusting. For its quality of rust resistance it has been examined by scientists from all parts of the world. The iron pillar is 8 m high and weighs 6 tonnes (6000 kg).