CLASS 10 Science
CHAPTER 3
Metals And Non-Metals (NCERT Notes)
PHYSICAL PROPERTIES
Metals
Metals, in their pure state, have a shining surface. This property is called metallic lustre.
They are generally hard. The hardness varies from metal to metal.
Metals which can be beaten into thin sheets. This property is called malleability. For example gold and silver are the most malleable metals.
The ability of metals to be drawn into thin wires is called ductility. Gold is the most ductile metal. Wire of about 2 km length can be drawn from one gram of gold. It is because of their malleability and ductility that metals can be given different shapes according to our needs.
Metals are good conductors of heat, electricity and have high melting points. The best conductors of heat are silver and copper. Lead and mercury are comparatively poor conductors of heat.
The metals that produce a sound on striking a hard surface are said to be sonorous. For example school bells are made of metals.
Non-metals
Some of the examples of non-metals are carbon, sulphur, iodine, oxygen, hydrogen, etc. The non-metals are either solids or gases except bromine which is a liquid.
All metals except mercury exist as solids at room temperature.
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.
Iodine is a non-metal but it is lustrous.
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 known and has a very high melting and boiling point. Graphite, another allotrope of carbon, is a conductor of electricity.
Alkali metals (lithium, sodium, potassium) are so soft that they can be cut with a knife. They have low densities and low melting points.
CHEMICAL PROPERTIES OF METALS
When Metals are burnt in Air
Almost all metals combine with oxygen to form metal oxides.
Metal + Oxygen → Metal oxide
When copper is heated in air, it combines with oxygen to form copper(II) oxide, a black oxide.
2Cu + O2 → 2CuO
(Copper) (Copper(II) oxide)
Aluminium forms aluminium oxide.
4Al + 3O2 → 2Al2O3
(Aluminium) (Aluminium oxide)
Metal oxides are basic in nature. But some metal oxides, such as aluminium oxide, zinc oxide, etc., show both acidic as well as basic behaviour. Such metal oxides which react with both acids as well as bases to produce salts and water are known as amphoteric oxides.
Aluminium oxide reacts in the following manner with acids and bases –
Al2O3 + 6HCl → 2AlCl3 + 3H2O
Al2O3 + 2NaOH → 2NaAlO2 + H2O
(Sodium aluminate)
Most metal oxides are insoluble in water but some of these dissolve in water to form alkalis. Sodium oxide and potassium oxide dissolve in water to produce alkalis as follows –
Na2O(s) + H2O(l) → 2NaOH(aq)
K2O(s) + H2O(l) → 2KOH(aq)
All metals do not react with oxygen at the same rate. Different metals show different reactivities towards 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. 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. Iron does not burn on heating but iron filings burn vigorously when sprinkled in the flame of the burner. Copper does not burn, but the hot metal is coated with a black coloured layer of copper(II) oxide. Silver and gold do not react with oxygen even at high temperatures.
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 easily to give aluminium articles an attractive finish.
When Metals react with Water
Metals react with water and produce a metal oxide and hydrogen gas. Metal oxides that are soluble in water dissolve in it to further form metal hydroxide. But all metals do not react with water.
Metal + Water → Metal oxide + Hydrogen
Metal oxide + Water → Metal hydroxide
Metals like potassium and sodium react violently with cold water. In case of sodium and potassium, the reaction is so violent and exothermic that the evolved hydrogen immediately catches fire.
2K(s) + 2H2O(l) → 2KOH(aq) + H2(g) + heat energy
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) + heat energy
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.
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.
Metals such as lead, copper, silver and gold do not react with water at all.
When Metals react with Acids
Metals react with acids to give a salt and hydrogen gas.
Metal + Dilute acid → Salt + Hydrogen
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 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 rate of formation of bubbles was the fastest in the case of magnesium. The reaction was also the most exothermic . The reactivity decreases in the order Mg > Al > Zn > Fe. In the case of copper, no bubbles were seen and the temperature also remained unchanged. This shows that copper does
not react with dilute HCl.
Aqua regia, (Latin for ‘royal water’) is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3:1. It can dissolve gold, even though neither of these acids can do so alone. Aqua regia is a highly corrosive, fuming liquid. It is one of the few reagents that is able to dissolve gold and platinum.
Metals react with Solutions of other Metal Salts
Reactive metals can displace less reactive metals from their compounds in solution or molten form.
For example: 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
The reactivity series is a list of metals arranged in the order of their decreasing activities.
Activity series : Relative reactivities of metals
HOW DO METALS AND NON-METALS REACT?
We learnt that noble gases, which have a completely filled valence shell, show little chemical activity. We explain the reactivity of elements as a tendency to attain a completely filled valence shell.
For example 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 us a sodium cation Na+. On the other hand 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 us a chloride anion C1–
. So both these elements can have a give-and-take relation between them as follows: For formation of sodium chloride
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). It should be noted that sodium chloride does not exist as molecules but aggregates of oppositely charged ions.
Formation of one more ionic compound, magnesium chloride.
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
The earth’s crust is the major source of metals. Seawater also contains some soluble salts such as sodium chloride, magnesium chloride, etc. The elements or compounds, which occur naturally in the earth’s crust, are known as minerals. At some places, minerals contain a very high percentage of a particular metal and the metal can be profitably extracted from it. These minerals are called ores.
Enrichment of Ores
Ores mined from the earth are usually contaminated with large amounts of impurities such as soil, sand, etc., called gangue. 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.
Extracting Metals Low in the Activity Series
Metals low in the activity series are very unreactive. The oxides of these metals can be reduced to metals by heating alone. For example, cinnabar (HgS) is an ore of mercury. When it is heated in air, it is first converted into mercuric oxide (HgO). Mercuric oxide is then reduced to mercury on further heating.
Similarly, copper which is found as Cu2S in nature can be obtained from its ore by just heating in air.
Extracting Metals in the Middle of the Activity Series
The metals in the middle of the activity series such as iron, zinc, lead, copper, etc., are moderately reactive. These are usually present as sulphides or carbonates in nature. It is easier to obtain a metal from its oxide, as compared to its sulphides and carbonates. Therefore, prior to reduction, the metal sulphides and carbonates must be converted into metal oxides. The sulphide ores are converted into oxides by heating strongly in the presence of excess air. This process is known as roasting.
The carbonate ores are changed into oxides by heating strongly in limited air. This process is known as calcination. The chemical reaction that takes place during roasting and calcination of zinc ores can be shown as follows –
Roasting
Calcination
The metal oxides are then reduced to the corresponding metals by using suitable reducing agents such as carbon. For example, when zinc oxide is heated with carbon, it is reduced to metallic zinc.
ZnO(s) + C(s) → Zn(s) + CO(g)
Obtaining metals from their compounds is also a reduction process.
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. For example, when manganese dioxide is heated with aluminium powder, the following reaction takes place –
3MnO2(s) + 4Al(s) → 3Mn(l) + 2Al2O3(s) + Heat
These displacement reactions are highly exothermic. The amount of heat evolved is so large that the metals are produced in the molten state. In fact, 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 thermit reaction.
Fe2O3(s) + 2Al(s) → 2Fe(l) + Al2O3(s) + Heat
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. For example, carbon cannot reduce the oxides of sodium, magnesium, calcium, aluminium, etc., to the respective metals. This is because these metals
have more affinity for oxygen than carbon. These metals are obtained by electrolytic reduction. 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 –
At cathode Na+ + e– → Na
At anode 2Cl– → Cl2 + 2e–
Similarly, aluminium is obtained by the electrolytic reduction of aluminium oxide.
Refining of Metals
The metals produced by various reduction processes 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, whereas, the insoluble impurities settle down at the bottom of the anode and are known as anode mud. 
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.
CORROSION
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.
Iron when exposed to moist air for a long time acquires a coating of a brown flaky substance called rust.
Prevention of Corrosion
The rusting of iron can be prevented by painting, oiling, greasing, galvanising, chrome plating, anodising or making alloys.
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.
Alloying is a very good method of improving the properties of a metal. 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. In fact, 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 non- metal. 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.
If one of the metals is mercury, then the alloy is known as an amalgam. The electrical conductivity and melting point of an alloy is less than that of pure metals. For example, brass, an alloy of copper and zinc (Cu and Zn), and bronze, an alloy of copper and tin (Cu and Sn), are not good conductors of electricity whereas copper is used for making electrical circuits. Solder, an alloy of lead and tin (Pb and Sn), has a low melting point and is used for welding electrical wires together.
Pure gold is known as 24 carat gold, is very soft. It is, therefore, not suitable for making jewellery. It is alloyed with either silver or copper to make it hard. Generally, in India, 22 carat gold is used for making ornaments. It means that 22 parts of pure gold is alloyed with 2 parts of either copper or silver.
( From NCERT Book )