metals and nonmetals class 10 notes

[h]Physical Properties of Metals and Nonmetals [/h]

1. Metals

  • Metals, in their pure state, have a shining surface. This property is called metallic lustre.
  • Metals can be beaten into thin sheets. This property is called malleability.
  • The ability of metals to be drawn into thin wires is called ductility. Gold is the most ductile metal. 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 and have high melting points.
  • The metals that produce a sound on striking a hard surface are said to be sonorous.
  • Some Example of Metals are iron, copper, aluminium, magnesium, sodium, lead, zinc

 

2. 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.
  • Most non-metals produce acidic oxides when dissolve in water.
  • Most metals, give rise to basic oxides.

 

Important facts about Metals and non-metals

  •  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.

 

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

 

[h]CHEMICAL PROPERTIES OF METALS[/h]

Metals Reaction with oxygen in Air

Almost all metals combine with oxygen to form metal oxides.

Metal + Oxygen → Metal oxide

For example, when copper is heated in air, it combines with oxygen to form copper(II) oxide, a black oxide.

Similarly, aluminium forms aluminium oxide

  • We have learnt that 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.
  • 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 –
  • Metals such as potassium and sodium react so vigorously with the oxygen in air 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.

 

Reaction of Metals 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.
  • 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.

 

Reaction of Metals 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.

 

Reaction of Metals with the Solutions of other Metal Salts

  • Reactive metals can displace less reactive metals from their compounds in solution or molten form.
  • It is simple and easy 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

Defination : The reactivity series is a list of metals arranged in the order of their decreasing activities.

Activity series : relative reactivities of metals
Activity series : relative reactivities of metals

 

Reaction of Metals and non-metals

  • A completely filled valence shell, show little chemical activity. We, therefore, explain the reactivity of elements as a tendency to attain a completely filled valence shell.
  • Example 1: We can see from Table 3.3 that 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 hasbecome 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

    Formation of sodium chloride
    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.

Example 2

formation of  ionic compound, magnesium chloride

Formation of magnesium chloride
Formation of magnesium chloride

 

  • The compounds formed in this manner by the transfer of electrons from a metal to a non-metal are known as ionic compounds or electrovalent compounds.

 

Properties of ionic compounds

You may have observed the following general properties for 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. 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.

 

[h] Occurrence of Metals[/h]

  • 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.

Extraction of Metals

  • Some metals are found in the earth’s crust in the free state. Some are found in the form of their compounds.
  • The metals at the bottom of the activity series are the least reactive. They are often found in a free state. For example, gold, silver, platinum and copper are found in the free state. Copper and silver are also found in the combined state as their sulphide or oxide ores.
  • 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.
  •  ores of many metals are oxides. This is because oxygen is a very reactive element and is very abundant on the earth
Metals Activity series and related metallurgy
Metals Activity series and related metallurgy

 

  • Several steps are involved in the extraction of pure metal from ores.
    Steps involved in the extraction of metals from ores
    Steps involved in the extraction of metals from 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 –

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)

 

  • 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 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 .

Electrolytic refining of copper
Electrolytic refining of copper

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.

 

CORROSION

Corrosion is the slow eating up off Metal by the action of moisture , etc.

  • 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.

Alloying is a very good method of improving the properties of a metal. We can get the desired properties by this method.

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.

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

brass, an alloy of copper and zinc (Cu and Zn), and bronze, an alloy of copper and tin (Cu and Sn),

 

 

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