CBSE Class 10 Science Chapter 3 - Metals & Non-Metals

Class Notes of Ch 3 : Metals & Non-metals 
Class 10th Science

Metals & Non-Metals

      Topics:

• Introduction
• Physical Properties Of Metal
• Physical Properties of Non-Metal
• Reaction of Metal with Oxygen
• Exceptions to metallic oxides being basic in nature
• Reactions of metal with water
• Reactions of metal with acids
• Reactivity series
• Reaction of metals with non-metals
• Ionic Compounds
• Properties of Ionic Compounds
• Extraction of Metals
• Extracting Metals Low in the Activity Series
• Extracting Metals in the Middle of the Activity Series
• Extracting Metals towards the Top of the Activity Series
• Refining of metals
• Corrosion
• Prevention from Corrosion

Introduction

In our day to day life we use many well-known materials like iron, aluminum, copper, gold, silver and many more. We are also familiar with the elements like oxygen, carbon, Sulphur as well. These are all obtained from some naturally occurring substances encompassing one or more elements or their compounds. They are termed as minerals. Minerals form which elements can be profitably extracted are termed as ores.

• There 92 well known naturally occurring minerals of which 70 are metals and rets 20 are the non-metals.
• Some metals possess the characteristics of both metals as well as non-metals. They are termed as metalloids. For instance, boron, silicon, germanium, arsenic, antimony, tellurium, and polonium.
• Elements found in free state includes some metals like gold, silver, platinum etc. and some non-metals like helium, neon, argon etc. as well.
• A major proportion of metals are found in combined states like oxides, sulphides, carbonates, silicates etc. Some non-metals like sulphur, phosphorous too are found to exist in combined state.

Physical Properties of Metal

• Metals are solid except mercury that is found in liquid state at room temperature.
• They are hard and tough except sodium and potassium that can be cut with a knife.
• They are lustrous i.e. they shine in light because metals possess free electrons that vibrates on getting in contact with light.
• They have high melting points except caesium and gallium have very low melting point.
• They are good conductors of heat and electricity. The best conductors of electricity are silver and copper whereas lead and mercury are comparatively poor conductors.
• They have high density except alkali metals like lithium, sodium and potassium.
• They are malleablee. they can be beaten into thin sheets. For instance, the aluminium foil we use to pack food is manufactured by beating the aluminium metal into thin sheets. This a characteristic property of metal.
• They are ductilee. they can be drawn into wires. For instance. We all are familiar with copper and aluminium wires. It is due to their characteristic property of ductility that these metals can be drawn into wires.
• They are sonorouse. they produce a ringing sound when struck against a hard object. For instance, your school bells are made of metal that produces a ringing sound after the period is over when struck hard by a hammer.

Physical Properties of Non-Metal

• Non-metals may be either solids, liquids or gases.
• Solid non-metals are brittle and break down into powdery mass on striking with a hammer except diamond which is the hardest non-metal.
• They have a dull luster but iodine is lustorous.
• They have low melting points except diamond that has very high melting point.
• They are poor conductors of heat and electricity except graphite.
• They are not malleable.
• They are not ductile.
• They are not sonorous.
• Examples are carbon, Sulphur, oxygen, phosphorous and many more.

Reactions of metal with oxygen

Metals reacting with oxygen is observed in our day to day life. We all must have observed rusting of iron, silver jewellery getting tarnished, or copper articles getting covered by green layer. Metals reacts with atmospheric oxygen and produces metal oxides that are basic in nature because they react with water to form bases.

• In case of rusting of iron, the iron reacts with the oxygen present in air and moisture and develops rust (hydrated iron (III) oxide).
• In case of rusting of copper, the metallic copper reacts with oxygen, carbon-dioxide and atmospheric moisture and develops a green coloured coating of copper hydroxide and copper carbonate.
• In case of tarnishing of silver articles, the metallic silver reacts with hydrogen sulphide or sulphur present in air and gets tarnished.
• Copper burns in air to combines with oxygen and form copper (II) oxide, a black oxide.

Exceptions to metallic oxides being basic in nature

Aluminium forms aluminium oxide.

And show both acidic as well as basic actions. Such metal oxides reacting with both acids as well as bases to produce salts and water are known as amphoteric oxides.

• Metallic oxides are generally insoluble in water but some of the metallic oxides like Sodium oxide and potassium oxide dissolves in water to produce alkalis.

• At normal temperature, the surfaces of metals such as magnesium, aluminium, zinc, lead, etc., develops a coating of thin layer of oxide. The protective oxide layer prevents the metal from further oxidation.

Magnesium burns in air to form magnesium oxide.

       2Mg + O2 --> 2MgO

• Iron does not burn on heating but iron filings burn vigorously when scattered
  in flames.

• Copper does not burn on heating , but the hot metal develops a coating of black coloured copper(II) oxide.

        2Cu + O2 ---> 2CuO

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

Reactions of metal with water

• Sodium being very reactive reacts vigorously with water leading to the production of sodium hydroxide and hydrogen. Therefore it is stored in kerosene.

          2Na + 2H2O à 2NaOH + H2

• Magnesium does not react with cold water. Magnesium undergoes reaction with hot water to form magnesium hydroxide and hydrogen.

        Mg +H2O à Mg(OH)2 + H2

  

• The reaction of calcium with water evolves heat that is not sufficient for the hydrogen to catch fire.

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

• Whereas metals like aluminium, iron and zinc reacts neither with cold nor
  hot water. Instead they react with steam to form the metal oxide and hydrogen.

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

Reactions of metal with acids

Metals react with acids to produce hydrogen gas. If a matchstick is brought near the mouth of the tube containing the product of the reaction then we hear a pop sound. It is this hydrogen gas that burns with a pop sound.

• For instance, Magnesium reacts with dilute hydrochloric acid to form magnesium chloride and hydrogen.

Mg + 2HCl --> MgCl2 + H2

• Copper does not react with hydrochloric acid but reacts with sulphuric acid.

CuO + H2SO4 --> CuSO4 + H2O

 Hydrogen gas is not evolved when a metal reacts with nitric acid. This is due to the fact that HNO3 is a strong oxidising agent. It oxidises the H2 produced to water and 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.

Reactions of metal with other metals: Reactivity series

A metal reacts with another metals and displaces them from their solution. This is known as displacement reaction. In displacement reaction a more reactive metal displaces a less reactive metal but a less reactive metal cannot displace a more reactive metal.

Reaction of metals with non-metals

Metals are electropositive because they have tendency to lose electrons. E.g. Na+ whereas non-metals are electronegative since they have a tendency to gain electrons. E.g. Cl-.

Atomic number of sodium is 11. So the electronic configuration stands out to be 2,8,1 i.e. there is one electron in the outermost shell. In order to gain inert gas configuration it is better for sodium to lose one electron and achieve the nearest noble gas configuration of neon with atomic number 10. Sodium has 1 electron in its outermost shell and hence valency of sodium is 1.

On the other hand atomic number of chlorine is 17. So electronic configuration stands out to be 2,8,7. In order to achieve noble gas configuration to become stable it requires one electron then it will acquire the configuration of neon (noble gas). Therefore valency of chlorine is 1.

That means the every element tries attain stability by acquiring noble gas configuration for which it tries to either gain electron or donate electron. Na donates 1 electron in its outermost shell to attain noble gas configuration whereas chlorine acquires 1 electron in its outermost shell to acquire noble gas configuration. Hence formation of NaCl takes place in the following manner.

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

Ionic Compounds

Metals and non-metals leads to the formation of Compounds that possess charged species. These charged species are called ions. These oppositely charged positive and negative ions hold together in an ionic bonds, forms ionic compounds i.e. compounds made of ions.

These charged species can be either positively charged called cation or negatively charged called anion. For instance, In sodium chloride (NaCl), Na exist as cation Na+ whereas Cl exist as anion Cl-. Hence it is an ionic compound.

Properties of Ionic Compounds

Properties of ionic compounds are as follows.

• Ionic compounds are solids and hard due to the strong attracting force between the positive and negative ions. These compounds are generally brittle and break into pieces on application of pressure.
• Ionic compounds have high melting and boiling points because a application of significant amount of energy can break the strong inter-ionic attraction.
• Ionic compounds are soluble in water but insoluble in solvents like kerosene, petrol, etc.
• Conduction of electricity through a solution is possible when there is movement of charged particles. 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.
• 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 conduct electricity in the molten state as in the molten state the electrostatic forces of attraction between the oppositely charged ions overcome due to the heat. Thus, the ions move freely and conduct electricity.

Extraction of Metals

• The metals at the top of the activity series (K, Na, Ca, Mg and Al) are so reactive that they are never found in nature as free elements.
• The metals in the middle of the activity series (Zn, Fe, Pb, etc.) are moderately reactive. They are found in the earth’s crust mainly as oxides, sulphides or carbonates.
• Below copper lies gold and silver that are found in Free State. These metals have low reactivity.

Ores that are mined out from the earth contains contaminants with large amounts of impurities including soil, sand, etc. These are known as gangue which must be removed from the ore before extraction of the metal. The processes involved in removing the gangue from the ore are based on the alterations between the physical or chemical properties of the gangue and the ore.

Extracting Metals Low in the Activity Series

• These metals have low reactivity.
• The oxides of these metals can be reduced to metals by heating alone.
• Cinnabar (HgS) is an ore of mercury which on exposure to heat first converts into mercuric oxide (HgO) and on further heating gets reduced to mercury.

• Another instance is reduction of Cu2S (ore of copper) to copper by heating.

Extracting Metals in the Middle of the Activity Series

• The metals in the middle of the activity series (Zn, Fe, Pb, etc.) are moderately reactive.
• They are found in the earth’s crust mainly as oxides, sulphides or carbonates.
• It is easy to extract metal from their oxide, as compared to their sulphides and carbonates. Therefore before extracting metals form their respective ores, the metal sulphides and carbonates must be converted into metal oxides. For instance, sulphide and carbonate ores of zinc are heated to convert into their oxides.

• This conversion of sulphide ores into their oxides by heating strongly in the presence of excess air is known as roasting.

• This conversion of carbonate ores into their oxides by heating strongly in limited air in the presence of suitable reducing agents such as carbon is known as calcination.
• Sometimes displacement reaction is also used to derive the metal from their ores. Highly reactive metals displace the low reactive metals from their compounds. For instance, manganese dioxide is heated with aluminium powder.

The displacement reactions are highly exothermic and hence the metals are produced in the molten state. In fact, the reaction of iron (III) oxide (Fe2O3) with aluminium is used for joining railway tracks or cracked machinery parts. This reaction is known as the Thermit reaction.

Extracting Metals towards the Top of the Activity Series

• The metals at the top of the activity series are highly reactive.
• They cannot be obtained from their compounds by heating with carbon.
• These metals have more affinity for oxygen.
• These metals are obtained by electrolytic reduction.
• Sodium ions migrate to the cathode, where electrons enter the melt and are reduced to sodium metal.
• Chloride ions migrate the other way, toward the anode. They give up their electrons to the anode and are oxidized to chlorine gas.

Refining of metals

• A wide range of metals, such as copper, zinc, tin, nickel, silver, gold, etc., are refined electrolytically after obtaining them from their ores.
• An impure metal is made the anode.
• Thin strip of pure metal is made the cathode.
• A solution of the respective metal salt is used as electrolyte.
• Electric current is passed through the solution.
• Passing current through the electrolyte, dissolves the pure metal from the anode into the electrolyte.
• On the other hand an equal amount of pure metal from the electrolyte gets deposited on the cathode.
• The soluble impurities go into the electrolyte.

Whereas, the insoluble impurities settle down at the bottom of the anode and are termed as anode mud.

Corrosion

• In case of rusting of iron, the iron reacts with the oxygen present in air and moisture and develops rust (hydrated iron (III) oxide).
• In case of rusting of copper, the metallic copper reacts with oxygen, carbon-dioxide and atmospheric moisture and develops a green coloured coating of copper hydroxide and copper carbonate.
• In case of tarnishing of silver articles, the metallic silver reacts with hydrogen sulphide or sulphur present in air and gets tarnished.

Prevention from Corrosion

Following methods can be used to prevent corrosion:

• Galvanisation

Method to protect steel and iron articles from rusting by coating them with a thin layer of zinc.

• Alloying

Homogeneous mixture of two or more metals, or a metal and a non-metal to change the properties of the pure state metals and protect them from rust.

1. For instance, Iron is used widely for different purposes but it is never used in pure state. It is due to the fact that pure iron is very soft and stretches in hot state. But mixing with small amount of carbon (about 0.05%) makes it strong and tough. Mixing nickel and chromium with iron gives stainless steel, which is hard and does not rust.
2. 
   
3. Another vital property of alloy is that electrical conductivity and melting point of an alloy is less than that of pure metals. For instance, brass is an alloy of copper and zinc (Cu and Zn), and bronze is an alloy of copper and tin (Cu and Sn), possessing poor conductivity towards electricity. On the other hand copper in pure state is used for making electrical circuits.

1. Solder is an alloy of lead and tin (Pb and Sn) possessing low melting point and is used for welding electrical wires together.

SCIENCE Revision Notes

Chapter:01  Chemical Reaction & Equation
Chapter:02  Acid Base & Salt
Chapter:03  Metals & Non Metals
Chapter:04  Carbon & its Components
Chapter:05  Periodic Classification of Elements
Chapter:06  Life Processes
Chapter:07  Control & Coordinates
Chapter:08  How do Organisms Reproduce
Chapter:09  Heridity & Evolution
Chapter:10  Light Reflection Refraction
Chapter:11  The Human Eye & the Colourful World
Chapter:12  Electricity
Chapter:13  Magnetic Effect of Electric Current
Chapter:14  Source Of Energy
Chapter:15  Our Environment
Chapter:16  Management of Natural Resource

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