CBSE Class 10 Science Metals And Non Metals Notes

METALS AND NON-METALS       

7.1 INTRODUCTION:

There are 118 chemical elements known at present. One the basis of their properties, all these elements can be broadly divided into two main groups: Metals and Non-Metals. A majority of the known elements are metals. All the metals are solids, except mercury, which is a liquid metal. There are 22 non-metals, out of which, 10 non-metals are solids, one non-metal (bromine) is liquid and the remaining 11 non-metals are gases.

7.2 POSITION OF METALS AND NON-METALS IN THE PERIODIC TABLE :

The metals are placed on the left hand side and in the centre of the periodic table. One the other hand, the non-metals are placed on the right hand side of the periodic table. This has been shown in the figure. It may be noted that hydrogen (H) is an exception because it is non-metal but is placed on the left hand side of the periodic table.

Metals and non-metals are separated from each other in the periodic table by a zig-zag line. The elements close to zig-zag line show properties of both the metals and the non-metals. They show some properties of metals and some properties of non-metals. These are called metalloids. The common examples of metalloids are boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te) and polonium (Pi)

In general, the metallic character decreases on going from left to right side in the periodic table. However, on going down the group, the metallic character increases. 

The elements at the extreme left of the periodic table are most metallic and those on the right are least metallic or non-metallic.

7.3  GENERAL PROPERTIES OF METALS AND NON-METALS :

7.3 (a) Electronic Configuration of Metals :

The atoms of metals have 1 to 3 electrons in their outermost shells. For example, all the alkali metals have one electron in their outermost shells (lithium 2, 1; sodium 2,8,1: potassium 2,8,8,1 etc.)

Sodium, magnesium and aluminium are metals having 1,2 and 3 electrons respectively in their valence shells. Similarly, other metals have 1 to 3 electron in their outermost shells.

It may be noted that hydrogen and helium are exception because hydrogen is a non-metal having only  electron in the outermost shell (K shell) of its atom and helium is also a non-metal having 2 electron in the outermost shell (K shell).

7.3 (b) Physical Properties of Metals:

The important physical properties of metals are discussed below:

(i) Metals are solids at room temperature: All metals (except mercury) are solids at room temperature.

MERCURY is a liquid at room temperature.

(ii) Metals are malleable: metals are generally malleable. Malleability means that the metals can be beaten with a hammer into very thin sheets without breaking. Gold and silver are among the best malleable metals. Aluminium and copper re also highly malleable metals.

(iii) Metals are ductile : It means that metals can be drawn (stretched) into this wires. Gold and silver are the most ductile metals. Copper and aluminium are also very ductile, and therefore, these can be drawn into this wires which are used in electrical wiring.

(iv) Metals are good conductors of heat and electricity: All metals are good conductors of heat. The conduction of heat is called thermal conductivity. Silver is the best conductor of heat. Copper and aluminium are also good conductors of heat and therefore, they are used for making household utensil. Lead is the poorest conductor of heat. Mercury metal is also poor conductor of heat.

Metals are also good conductors of electricity. The electrical and thermal conductivities of metals are due to the presence of free electrons in them. Among all the metals, silver is the best conductor electricity. Copper and aluminium are the next best conductors of electricity. Since silver is expensive, therefore, copper and aluminium are commonly used for making electric wires.

(v) Metals are lustrous and can be polished: Most of the metals have shine and they can be polished. The shining appearance of metals is also known as metallic lustre. For example, gold, silver and copper metals have metallic lustre.

SILVER is best conductor of heat and electricity.

(vi) Metals have high densities : Most of the metals are heavy and have high densities. For example, the density of mercury metal if very high (13.6 g cm^-3). However, there are some exceptions. Sodium, potassium, magnesium and aluminium have low densities. Densities of metals are generally proportional to their atomic masses. The smaller the metal atom, the smaller it its density.

(vii) Metals are hard : Most of the metals are hard. But all metals are not equally hard. Metals like iron, copper, aluminium etc. are quite hard. They cannot be cut with a knife. Sodium and potassium are common exceptions which are soft and can be easily cut with a knife.

(viii) Metals have high melting and boiling points : Most of the metals (except sodium and potassium) have high melting and boiling points.

Notes: Tungsten has highest melting point (3410⁰C) among all the metals.

(ix) Metals are rigid : Most of the metals are rigid and they have high tensile strength.

(x) Metals are sonorous: Most of the metals are sonorous i.e., they make sound when hit with an object.

7.3 (c) Electronic Configuration of Non-Metals :

The atoms and non-metals have usually 4 to 8 electrons in their outermost shells. For example, Carbon (At. No6), Nitrogen (At. No. 7), Oxygen (At. No.8 ), Fluorine (At. No. 9) and Neon (At. No. 10) have respectively 4,5,6,7,8 electrons in their outermost shells.

7.3 (d) Physical Properties Of Non-Metals:

The important physical properties of non-metals are listed below :

 (i)  Non-metals are brittle.

(ii) Non- metals are not ductile.

(iii) Non-metals are bad conductor of heat and electricity. (Exception: Graphite is a good conductor because of the presence of free electrons.)

(iv) Non-metals are not lustrous and cannot be polished. (Exception: Graphite and Iodine are lustrous non- metals.)     

(v) Non-metals may be solid, liquid, or gases at room temperature. 

(vi) Non-metals are generally soft. (Exception: Diamond, an allotropic from of non-metal Carbon, is the hardest natural substance known). 

(vii) Non-metals have generally low melting and boiling points. (Exception: Graphite another allotropic form of Carbon, has a melting point of about 37300C).           

(viii) Non-metals have low densities. (Exception : Iodine has high density).          

Graphite is a good conductor of electricity, lustrous and has very high melting point.    

7.3 (e) Chemical Properties of Metals :

The atoms of the metals have usually 1, 2 or 3 electrons in their outermost shells. These outermost electrons are loosely held by their nuclei. Therefore, the metal atoms can easily lose their outermost electrons to from positively charged ions. For example, sodium metal can lose outermost one electron to form positively charged ions, Na+. After losing the outermost electron, it gets stable electronic configuration of the noble gas (Ne : 2, 8 ), Similarly, magnesium can lose two outermost electron to from Mg²⁺ ions and aluminium can lose its three outermost electrons to from Al³⁺ ions.

The metal atoms lose electrons and form positively charged ions, therefore, the metals are called electropositive elements.

Some of the important chemical properties of metals are discussed below :

(i) Reaction with oxygen : Metals react with oxygen to from oxides. These oxides are basic in nature. For example, sodium metal reacts with oxygen of the air and form sodium oxide.

Sodium oxide reacts with water to form and alkali called sodium hydroxide. Therefore, sodium oxide is a basic oxide.

Due to the formation of sodium hydroxide (which is an alkali), the solution of sodium oxide in water turns red litmus blue (common property of all alkaline solutions).

When metal oxides are dissolved in water, they give alkaline solutions.

Similarly, magnesium is a metal and it reacts with oxygen to form magnesium oxide. However, magnesium is less reactive than sodium and therefore, heat is required for the reaction.

Thus, when a metal combines with oxygen, it loses its valence electrons and forms positively charged metal ions. We can say that oxidation of metal takes palace.

Reactivity of metals with oxygen:

All metals do not react with oxygen with equal ease. The reactivity of oxygen depends upon the nature of the metal. Some metals react with oxygen even at room temperature, some react on on heating  while still others react only on strong heating.

For example :

(A) Metals like sodium, potassium and calcium react with oxygen even at room temperature to form their oxides.

(B) Metals like magnesium and zinc do not react with oxygen at room temperature. They burn in air only on strong heating to from corresponding oxides.

(C) Metals like iron and copper do not burn in air even on strong heating. However, they react with oxygen only on prolonged heating.

(ii) Reaction with water : Metals react with water to form metal oxide or metals hydroxide and hydrogen. The reactivity of metals towards water depends upon the nature of the metals. Some metals react even with cold water, some react with water only on heating while there are some metals do not react even with steam. For example,

(A) Sodium and potassium metals react vigorously with cold water to form sodium hydroxide and hydrogen gas is liberated.

The reaction between sodium and water is so violent that the hydrogen evolved catches fire.

(B) Calcium reacts with cold water to form calcium hydroxide and hydrogen gas. The reaction is less violent.

 (C) Magnesium reacts very slowly with cold water but reacts rapidly with hot boiling water forming magnesium oxide and hydrogen.

 (D) Metals like zinc and aluminum react only with steam to form their corresponding oxides oxide hydrogen.

(E) Iron metal does not react with water under ordinary conditions. The reactions occurs only when steam is passed over red hot iron and the products are iron (II, III) oxide and hydrogen. 

(F) Metals like copper, silver and gold do not react with water even under strong conditions. The order of reactivities of different metals with water is :

Na >Mg >  AI > Zn   > Fe    > Cu

Rectivity with water decreases

(iii) Reaction with dilute acids : Many metals react with dilute acids and liberate hydrogen gas. Only less reactive metals such as copper, silver, gold etc. do into liberate hydrogen from dilute acids. The reactions of metals with dilute hydrochloric acid (HCI) and dilute sulphuric acid (H2SO4) are similar. With dil. HCI, they given metal chlorides and hydrogen whereas with dil. H2SO4, they give metal sulphates and hydrogen

Dilute nitric acid (HNO3) is an oxidising agent which oxidises metals, but does not produce hydrogen.

The reactivity of different metals is different with the same acid. For example:

(A) Sodium, magnesium and calcium react violently with dilute hydrochloric acid (HCI) or dilute sulphuric acid (H2SO4) liberating hydrogen gas and corresponding metal salt.

Therefore copper is even less reactive that iron

The order of reactivity of different metals with dilute acid:

Na       >          Mg       >          AI        >          Zn        >          Fe        >          Cu

Reactivity with dill acids decreases from sodium to copper.

(iv) Reactions of metals with salt solutions: When a more reactive metals is placed in a salt solution of less reactive metal, then the more reactive metal displaces the less reactive metal from its salt solution. For example, we will take a solution of copper sulphate (blue coloured solution) and put a strip of zinc metal in the solution. It is observed that the blue colour of copper sulphate fades gradually and copper metals are deposited on the zinc strip. this means that the following reaction occurs :

Here, zinc displaces copper from its salt solution.

However, if we take zinc sulphate solution and put a string of copper metal in this solution, no reaction occurs.

This means that copper cannot displace zinc metal from its solution. Thus, we can conclude that zinc is more reactive than copper. However, if we put gold or platinum strip in the copper sulphate solution, then copper is not displaced by gold or platinum. Thus, gold and platinum are less reactive than copper.

7.4 REACTIVITY SERIES OF METALS:

7.4 (a) Introduction:

We have learnt that some metals are chemically very reactive while others are less reactive or do not react at all.

On the basis of reactivity of different metals with oxygen, water acids as well as displacement reactions, the metals have been arranged in the decreasing order of their reactivities.

The arrangement of metals in order of decreasing reactivities is called reactivity series or activity series of metals.

The activity series of some common metals is given in Table. In this table, the most reactive metal is placed at the top whereas the least reactive metal is placed at the bottom. As we go down the series the chemical reactivity of metals decreases.

7.4 (b) Reasons for Different Reactivities:

In the activity series of metals, the basis of reactivity is the tendency of metals to lose electrons. If a metals can lose electrons easily to form positive ions, it will react readily with other substances. Therefore, it will be a reactive metal. On the other hand, if a metal loses electrons less rapidly to form a positive ion, it will react slowly with the other substances. Therefore, such a metal will be less reactive. For example, alkali metals such as sodium and potassium lose electrons very readily to from alkali metal ions, therefore, they are very reactive.

7.4 (c) Displacement of Hydrogen from Acids by Metals :

All metals above hydrogen in the reactivity series (i.e. more active than hydrogen) like zinc, magnesium, nickel can liberate hydrogen from acids like HCI and H₂SO₄. These metals have greater tendency to lose electrons than hydrogen. Therefore, the H+ ions in the acids will accept electrons and give hydrogen gas

as :

The metals which are below hydrogen in the reactivity series (i.e. les reactive than hydrogen) like copper, silver, gold cannot liberate hydrogen form acids like HCI, H₂SO₄ etc. These metals have lesser tendency to lose electrons than hydrogen. Therefore, they cannot lose electrons to H⁺ ions.

7.4 (d) Reactivity Series and Displacement Reactions :

The reactivity series can also explain displacement reactions. In general, a more reactive metal (placed higher in the activity series) can displace the less reactive metal from its solution. For example, zinc, displaces copper form its solution.

            Zn (s) + CuSO₄ (aq) → ZnSO₄ (aq) + Cu(s)

7.4 (e) Usefulness of Activity Series:

The activity series is very useful and it gives the following information:

(i) The metal which is higher in the activity series is more reactive than the other. Lithium is the most reactive and platinum is the least reactive.

(ii) The metals which have been placed above hydrogen are more reactive than hydrogen and these can displace hydrogen from its compounds like water and acids to liberate hydrogen gas.

(iii) The metals which are placed below hydrogen are less reactive than hydrogen and these cannot displace hydrogen from its compounds like water and acids.

(iv) A more reactive metal (placed higher in the activity series) can displace the less reactive metal from its solution.

(v) Metals at the top of the series are very reactive and, therefore, they do not occur free in nature. The metals at the bottom of the series are least reactive and, therefore, they normally occur free in nature. For example, gold, present in the reactivity series is found in Free State in nature.

8.1 HOW METALS REACT WITH NON-METALS :

Octet Rule : Octet rule was given by G.N. Lewis and W.Kossel in 1916.

According to actet rule “an atom whose outermost shell contains 8 electrons (octet) is stable.”

This rule, however, does not hold good in case of certain small atoms like helium (He) in which presence of 2 electrons (duplet) in the outermost shell in considered to be the condition of stability.

Examples of elements whose atoms have fully filled or 8 e- in their outermost shell are –

All noble gases contain 8 valence electrons (except He in which 2 valence electrons are present) and are stable. They do not usually form bonds with other elements.

Atoms combine with one another to achieve the inert gas electron arrangement and become stable. Atoms from chemical bonds to achieve stability by acquiring the inert gas configuration or by completing their octet or duplet (in case of small atoms) in outermost shell. An atom can achieve the inert gas electron arrangement in three ways -

(i)  by losing one or more electrons .

(ii) by gaining one or more electrons.

(iii) by sharing one or more electrons.

Noble gases do not usually from bonds with other elements. because they are stable. So. atoms of elements have the tendency to combine with one another to achieve the inert gas configuration.

8.2  CONCEPT OF IONIC BOND :

Except the elements of group 18 of the periodic table all the elements for the remaining group, at normal temperature and pressure, are not stable in independent state. These elements from stable compounds either by combining with the other atoms or with their own atoms. When in gross electronic configuration of the elements there are 8 electrons present then these elements do not take part in the chemical reaction because atoms containing 8 electrons in their outermost shell are associated with extra stability and less energy.

Atoms with other electronic configuration, which do not contain eight electrons in their outermost shell, are unstable and to achieve the stability they chemically combine in such a manner that they achieve eight electrons in their outermost shell. 

Two or more than two types of atoms mutually combine with each other to achieve stable configuration of eight valence electrons. Attempt to achieve eight electrons in the outermost orbit of a element is the reason behind its chemical reactivity or chemical bonding.

8.3  IONIC OR ELECTROVALENT BOND:

This bond is formed by the atoms of electropositive and electronegative elements. Electropositive elements lose electrons in chemical reaction and electronegative elements gain electrons in chemical reaction. When an atom of electropositive element come in contact with that of an electronegative element then the electropositive atom loses electron & becomes positively charged, while the electronegative atom gains the electron to become negatively charged. Electrostatic force of attractions works between the positively and negatively charged ions due to which both ions are bonded with each other. As a result, a chemical bond is produced between the ions, which is known as Ionic or Electrovalent compound.

Number of electrons donated or accepted by any element is called Electrovalence.

In an ionic compound every cation is surrounded by a fixed number of anions and every anion is surrounded by a fixed number of cations and they are bonded in a fixed geometry in a three dimensional structure.

Example : Sodium chloride compound.

Sodium atom (Electropositive element) by losing an electron from its outermost orbit, gets converted into a cation and attains noble gas like stable configuration.

Energy required for this process is called “ionization potential.”

Chlorine atom (Electronegative element) accepts the electron donated by sodium atom in its outermost orbit and forms chloride anion.

 In this process energy is released which is known as “electron affinity.”

Due to the opposite charges on the Na+ and CI- ions, they are bonded by electrostatic force of attraction to from NaCI compound.

Here electrovalent of sodium and chlorine atom is one.

For the formation of ionic bond, it is necessary that the ionization potential of electropositive element should be less and the electron affinity of electronegative element should be high.

8.3 (a) Properties of Ionic Compounds:

(i)  Ionic compounds consist of ions: All ionic compounds consist of positively and negatively charged ions and not molecules. For example, sodium chloride consists of Na+ and CI- ions, magnesium fluoride consists of Mg²⁺ and F- ions and so on.

(ii) Physical nature : Ionic compounds are solid and relatively hard due to strong electrostatic force of attraction between the ions of ionic compound.

(iii) Crystal structure : X-ray studies have shown that ionic compounds do not exist as simple single molecules as Na+CI-, This is due to the fact that the forces of attraction are not restricted to single unit such as Na⁺ and CI^- but due to uniform electric field around and ion, each ion is attracted to a large number of other ions. For example, one Na+ ion will not attract only one CI^- ion but it can attract as many negative charges as it an. Similarly, the CI^- ion will attract several Na^- ions. As a result, there is a regular arrangement of these ions in three dimensions as shown in diagram. Such a regular arrangements is called crystal lattice.

(iv) Melting point and boiling point : Strong electrostatic force of attraction if present between ions of opposite charges. To break the crystal lattice more energy is required so their melting points and boiling points are high.

(v) Solubility : Ionic compounds are generally soluble in polar solvents like water and insoluble in no - polar solvents like carbon tetrachloride, benzene, ether alcohol etc.

(vi) Brittle nature: Ionic compounds on applying external force or pressure are broken into small pieces, such substances are known as brittle and this property is known as brittleness. When external force is applied on the ionic compound, layers of ions slide over one another and particles of the same charge come near to each other as a result due to the strong repulsion force, crystals of compounds are broken.

(vii) Electrical conductivity : Electrical conductivity in any substance is due to the movement of free electrons of ions. In metals electrical conductivity is due to the free movement of valency electrons. As ionic compound exhibits electrical conductivity due to the movement of ions either in the fused state or in the soluble state in the polar solvent. But in the solid state due to strong electrostatic force of attraction free ions are absent so they are insulator in the solid state.

9.1  OCCURRENCE OF METALS :

All metals are present in the earth’s crust either in the free state or in the form of their compounds. Aluminium is the most abundant metal in the earth’s crust. The second most abundant metal is iron and their one is calcium.

9.1 (a) Native and Combined State of Metals :

Metals occur in the crust of earth in the following two states -

(i)  Native state of free state: A metal is said to occur in a free or a native state when it is found in the crust of the earth in the elementary or uncombined form.

The metals which are very uncreative (lying at the bottom of activity series) are found in the free state. These have no tendency to react with oxygen and are not attacked by moisture, carbon dioxide of air or other no-metals. Silver, copper, gold and platinum are some examples of such metals.

(ii) Combines state : A metal is said to occur in a combined state if it is found in nature in the form of its compounds. e.g. Sodium, magnesium etc. Copper and silver are metals which occur in the free state as well as in the combined state.

9.2  MINERALS AND ORES :

The natural substances in which metals or their compounds occur either in native state or combined state are called minerals.

The minerals are not pure and contain different types of other impurities. The impurities associated with minerals are collectively known as gangue or matrix.

The mineral from which the metal can be conveniently and profitably extracted, is called an ore.

For example, aluminium occurs in the earth’s crust in the form of two minerals, bauxite (AI2,O3. 2H2O) and caly (AI₂O₃. 2SiO₂. 2H₂O). Out of these two, aluminium can be conveniently and profitably extracted from bauxite. So, bauxite is an are of aluminium.

 Oxygen is the most abundant element on earth’s crust.

9.2 (a) Types of Ores :

The most common ores of metals are oxides, sulphides, carbonates, sulphates, halides, etc. In general, very uncreative metals (such as gold, silver, platinum etc.) occur in elemental form or Free State.

(i)  Metals which are only slightly reactive occur as sulphides (e.g., CuS, Pbs etc.).

 (ii) Reactive metals occur as oxides (e.g., MnO₂,  AO₂O₃ etc.)

(iii) Most reactive metals occur as salts as carbonates, subparts, halides etc.

SOME COMMON ORES ARE LISTED IN THE TABLE

9.3  METALLURGY:

The process of extracting pure metals from their ores and then refining them for use is called metallurgy.

In other words, the process of metallurgy involves extraction of metals from their ores and then refining them from use. The ores generally contain unwanted impurities such as sand, stone, earthy particles, limestone, mica, etc., these are called gangue or matrix.

The process of metallurgy depends upon the nature of the ore, nature of the metals and they types of impurities present. Therefore, there is not a single method for the extraction of all metals. However, most of the metals can e extracted by a general procedure which involves the following steps.

Various steps involved in metallurgical processes are -

(a) Crushing and grinding of the ore.           

(b) Concentration of the ore or enrichment of the ore.           

(c) Extraction of metal from the concentrated ore.           

(d) Refining or purification of the impure metal.

 These steps are briefly discussed below -

9.3 (a) Crushing and Grinding of Ore :

Most of the ores occur as big rocks in nature. They are broken into small pieces with the help of crusher. These pieces are then reduced to fine powder with the help of a ball mill or a stamp mill. 

9.3 (b) Concentration of Ore or Enrichment of Ore :

The process of removal of unwanted impurities (gangue) from the ore is called ore concentration or ore enrichment.

(i) Hydraulic washing (washing with water) :

Principle: This method is based upon the difference in the densities of the ore particles and the impurities (gangue).

Ores of iron, tin and lead are very heady and, therefore, they are concentrated by this method.

(ii) Front floatation process:

Principle: this method is based on the principle of difference in the wetting properties of the ore and gangue particles with water and oil.

 This method is commonly used for sulphide ores.

(iii) Magnetic separation :

Principle: This method depends upon the difference in the magnetic properties of the ores and gangue.

This method is used for the concentration of haematite, an ore of iron.

The froth floatation process if commonly used for the sulphide ores copper, zinc, lead etc.

10.1 EXTRACTION OF THE METAL FROM THE CONCENTRATED ORE :

The metal is extracted from the concentrated ore by the following steps :

(a) Conversion of the concentrated ore into its oxide : The production of metal from the concentrated ore mainly involves reduction process. This can be usually done by two processes known as calcination and roasting process. The method depends upon the nature of the ore.

(b) Conversion of oxide to metal y reduction process

10.1 (a) Conversion of Ore into Metal Oxide :

 These are briefly discussed below :

(i) It is the process of heating the concentrated ore in the absence of air.

The calcination process is used for the following changes :

1.  to convert carbonate ores into metal oxide.
2.  to remove water from the hydrated ores.
3.  to remove volatile impurities from the ore.

(ii) Roasting : It is the process of heating the concentrated ore strongly in the presence of excess air.

This process is used for converting sulphide ores to metal oxide. In this process, the following changes take place :

1. the sulphide ores undergo oxidation to their oxides.
2. moisture is removed
3. volatile impurities are removed.          

For example :

Calcination is used for hydrated and carbonate ores and roasting is used for sulphide ores.

10.1 (b) Conversion of Metal Oxide to Metal:

The metal oxide formed after calcination or roasting is converted into metal by reduction. The method used for reduction of metal oxide depends upon the nature and chemical reactivity of metal.

The metals can be grouped into the following three categories on the basis for their reactivity:

1. Metals of low reactivity.
2. Metals of medium reactivity.
3. Metals of high reactivity.

These different categories of metals are extracted by different technique. the different steps involved in separation are as follows :

(i) Reduction by heating : Metals placed low in the reactivity series are very less reactive. They can be obtained from their oxides by simple heating in air.

(ii) Chemical Reduction (For metals in the middle of the reactivity series):

The metals n the middle of the reactivity series, such as iron, zinc, lead, copper etc. are moderately reactive. These are usually present as sulphides or carbonates. Therefore, before reduction the metal sulphides and carbonates must be converted to oxides. This is done by roasting and calcination. The oxides of these metals cannot be reduced by heating alone. Therefore, these metal oxides are reduced to free metal by using chemical agents like carbon, aluminium, sodium or calcium.

(A) Reduction with carbon : The oxides of moderately reactive metals (occurring in the meddle of reactivity series) like zinc, copper, nickel, tin, lead etc. can be reduced by using carbon as reducing agent.

One disadvantage of using carbon as reducing agent is that small traces of carbon are added to metal as impurity. Therefore, it contaminates the metals.

Coke is very commonly used as a reducing agent because it is cheap.

(B) Reduction with carbon monoxide: Metals can be obtained from oxides by reduction with carbon monoxide in the furnace.

(C) Reduction with aluminium : Certain metal oxides are reduced by aluminium to metals.

Reduction of metals oxides with aluminium is known as aluminothermy or thermite process.

(iii) Reduction of electrolysis or electrolytic reduction : The oxide of active metals (which are high up in the activity series) are very stable and cannot be reduced by carbon or aluminium. These metals are commonly extracted by the electrolysis of their fused salts using suitable electrodes. This is also called electrolytic reduction i.e. reduction by electrolysis.

For example, aluminium oxide is very stable and aluminium cannot be prepared by reduction with carbon. It is prepared by the electrolysis of molten alumina (AI₂O₃).

It may be noted that during electrolytic reduction of molten salts, the metals are always obtained at the cathode (negative electrode).

The process of extraction of metals by electrolysis process is called electrometallurgy.

10.2  PURIFICATION OR REFINING OF METALS :

The metal obtained any of the above methods is usually impure and is known as crude metal. The process of purifying the crude metal is called refining.

10.2 (a) Liquation :

This method is use for refining the metals having low melting points, such as tin, lead, bismuth etc. This is based on the principle that the metal to be refined is easily fusible (melt easily (but the impurities do not fuse easily.

10.2 (b) Distillation:

This method is used for the purification of volatile metals (which form vapours readily) such as mercury and zinc.

10.2 (c) Electrolytic Refining :

This is most general and widely used method for the refining of impure metals. Many metals such as copper, zinc, tin, nickel, silver, gold etc. are refined electrolytic ally. It is based upon the phenomenon of electrolysis. In this method, the crude metal is cast into thick rods and are made as anodes, while the thin sheets of pure metal are made as cathodes, An aqueous solution of some salt of the metal is used as an electrolyte. On passing 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 in the solution whereas the insoluble impurities settle down at the bottom of the node and are known as anode mud. In this way, the pure metal from anode goes into electrolyte and from electrolyte it goes to the cathode

In electrolytic refining impure metal is made anode and pure metal is made cathode.

Zone refining and Van Arkel method are used for obtaining metals (Si, Ge etc.) of very high purity for certain specific applications.

11.1 CORROSION OF METALS :

Surface of many metals is easily attacked when exposed to atmosphere. The react with air or water  present in the environment and form undesirable compounds  on their surface. These undesirable compounds are generally  oxides.

Thus, corrosion is a process of deterioration of metal as a result of its reaction with air or water (present in environment) surrounding it.

11.1 (a) Corrosion of Iron:

iron corrodes readily when exposed to moisture and gets covered with a brown flaky substance called rust.  This is also called Rusting of Iron. Chemically, the rust if hydrated iron (III) oxide,  Fe2O3.XH2O. Rusting is an oxidation process  in which iron metal is slowly oxidized by the action of air (in presence of water). Therefore, rusting of iron takes place under the following conditions :

1. Presence of air (or oxygen)
2. Presence of water (moisture)           

More the reactivity of the metal, the more will be the possibility of the metal getting corroded.

(i) Experiment to show that rusting of iron requires both air and water -

We take three test tubes and put one clean iron nail in each of the three test tubes :

(A) In the first test tube containing iron nail, we put some anhydrous calcium chloride to absorb water (or moisture) from the damp air present in the test tube and make it dry.

(B) In the second test tube containing iron nail, we put boiled water because boiled water does not contain any dissolved air or oxygen in it. A layer of oil is put over boiled water in the test tube to prevent the outside air from mixing with boiled water.

(C) In the third test tube containing an iron nail, we put unboiled water so that about two-third of the nail is immersed in water and the rest is above water exposed to damp air.

After one week, we observe the iron nails kept in all the three test tubes.

(ii) We will obtain the following observations from the experiment :

(A) No rust in seen on the surface of iron nail kept in dry air in the first test tube. This tells us that  rusting of iron does not takes place in air alone.

(B) No rust is seen on the surface of iron nail kept in air free boiled water in the second test tube, This tells us that rusting of iron does not take place in water alone.

(C) Red brown rust is seen on the surface of iron nail kept in the presence of the air and water in the third test tube. This tells us that rusting of iron takes place in the presence of both air and water together.

(iii) Prevention of rusting

(A) Corrosion of metals can be prevented by coating the metal surface with a thin layer of pant, varnish or grease.

(B) Iron is protected from rusting by coating it with a thin layer of another metal which is more reactive that iron. This prevents the loss of electrons from iron because the active metal loses electrons in process of covering iron with zinc is called galvanization.  Iron is also coated with other metals such as tin known as tin coating.

(C) By alloying: Some metals when alloyed with other metals become more resistant to corrosion. For example, when iron is alloyed with chromium and nickel, it form stainless steel. This is resistant to corrosion and does not rust at all.

(D) To decrease rusting of iron, certain antirust solutions are used. For example, solutions of alkaline phosphates are used as antirust solutions.

11.1 (b) Corrosion of Aluminum :

Due to the formation of a dull layer of aluminum oxide when exposed to moist air, the aluminum metal loses its shine very soon after use. This aluminum oxide layer is very tough and prevents the metal underneath from further corrosion (because moist air is not able to pass through this aluminum oxide layer). This means sometimes corrosion is useful.

11.1 (c) Corrosion of Copper

When a copper object remains in damp air for a considerable time, then copper reacts slowly with carbon dioxide and water of air to form a green coating of basic copper carbonate [CuCO₃, Cu(OH(₂] on the surface of the object. Since copper metal is low in the reactivity series, the corrosion of copper metal is very, very slow.

11.1 (d) Corrosion of Silver :

Silver is a highly uncreative metal, so it does not reacts with oxygen of air easily. But, air usually contains a little of sulphur compounds such as hydrogen sulphide gas (H₂S), which react slowly with silver to form a black coating of silver sulphide (Ag₂S). Silver ornaments gradually turn black due to the formation of a thin silver sulphide layer on their surface and silver is said to be tarnished.

11.2 ALLOYS :

An alloy is a homogenous mixture of two or more metals or a metal and a non-metal. For example, iron is the most widely used metal. But it is never used in the pure form. this is because iron is very soft and stretches easily when not. But when it is mixed with a small amount of carbon (about 0.05%), it becomes hard and strong. The new form of iron is called steel.

11.2 (a) Objective of Ally Making :

Alloys are generally prepared to have certain specified properties which are not possessed by the contituent metals. The main objects of ally-making are :

(i) To increase resistance to corrosion : For example, stainless steel is prepared which has more resistant to corrosion than iron.

(ii) To modify chemical reactivity : The chemical reactivity of sodium is decreased by making an alloy with mercury which is known as sodium amalgam.

(iii) To increase the hardness: Steel, an alloy of iron and carbon is harder than iron.

(iv) To increase tensile strength: Magnesium is an alloy or magnesium and aluminum.

It has greater tensile strength as compared to magnesium and aluminum.

(v) To produce good casting: Type metal is an alloy of lead, tin and mercury.

(vi) To lower the melting point: For example, solder is an alloy of lead and tin (50) Pb and 50% Sn). It has a low melting point and is used for welding electrical wires together.

11.2 (b) Some Important Alloys :

 The approximate composition and used of some important alloys are given below :

(i) Steel :  Steel is an alloy of iron and carbon containing 0.1 to 1.5% carbon. Steel is very hard, tough and strong. It is used for making rails, screws, girders, bridges, railway lines etc. Steel can also be used for the contraction for building, vehicles, ships etc.

(ii) Allow Steels : Steel obtained by the addition of some other elements such as chromium, vanadium, titanium, molybdenum, manganese, cobalt or nickel to carbon steel are called Ally Steel.

(iii) Alloys of Aluminum : This common alleys of aluminum are :

(A) Duralumin. It is an alloy containing aluminum, copper and traces of magnesium and manganese. Its

percentage composition is - AO\I  95%, Cu = 4%, Mg = 0.5 % Mn = 0.5 % It is stronger than pure aluminum, Since duralumin is light and yet strong, it is used for making bodies of aircrafts, helicopter, jets and kitchenware’s like pressure cookers etc.

(B) Magnesium.  It is an alloy of aluminum and magnesium having the composition: Al - 95%, Mg = 5% It is very light and hard. It is more hard than pure aluminum. It is used for making light instruments, balance beams, pressure cookers etc.

(C) Alnico . It is an alloy containing aluminum, iron nickel, and cobalt. It is highly magnetic in nature and can be used for making powerful magnets.

(iv) Alloys of Copper:  The important alloys of copper are Brass and Bronze.

(A) Brass - It is an alloy of copper and zinc having the composition = Cu = 80% Zn = 20% Brass is more malleable and more strong than pure copper. It is used for making cooking utensils, condenser sheets, pipes, hardware, nuts, bolts, screws, springs etc.

(B) Bronze - It is an alloy of copper and tin having the composition : Cu = 90% Sn = 10% Bronze is very though and highly resistant to corrosion. It is used for making utensils, statues, cooling pipes, coins, hardware etc.

(C) German Silver - It is an alloy of copper, zinc and nickel having the composition: Cu = 60%, Zn = 20%, Ni = 20%. It is used for making silverware, utensils and for electroplating.

(v) Alloying of Gold : Pure gold is very soft and cannot be used as such for jewellery. Therefore, it is generally alloyed with other metals commonly copper or silver to make it harder and modify its colour. The purity of gold is expressed as carats. Pure gold is of 24 carat. A 18 carat gold means that is contains 18 parts of gold is 24 parts by weight of alloy. Most of the jewellery is made of 22 carat gold.

Amalgams are homogenous mixtures of a metal and mercury. For example, sodium amalgam contains sodium and mercury.

Different amalgams are prepared according to their used. For example,

(i) Sodium amalgam is produced to decrease the chemical reactivity of sodium metal. It is also used as a good reducing agent.

(ii) Tin amalgam is used for silvering cheap mirrors.

(iii) The process of amalgamation is used for the extraction of metals like godl or silver from their native ores.

GIST OF THE LESSON

Elements are classified broadly into two categories on the basis of properties:

Metals: Iron, Zinc, Copper, Aluminium etc. Non – metals: Chlorine, Nitrogen, Hydrogen, Oxygen, Sulphur etc. Apart from metals and non-metals some elements show properties of both metals and non – metals, e.g. Silicon, Arsenic, Germanium .They are called metalloids

Properties of ionic compounds
1. Physical nature:solid and hard due to strong force of attraction. (generally brittle)
2. Melting point and boiling point:have high M.P and B.P, as large amount of heat energy is required to break strong ionic attraction.
3. Solubility: soluble in water and insoluble in kerosene and pertrol.
4. Conduction of electricity:ionic compounds in solid state-----does not conduct electricity.
Reason—Ions can not move due to rigid solid structure. Ionic compounds conduct electricity in molten state.
Reason-- Ions can move freely since the electrostatic forces of attraction between the oppositely charged ions are overcome due to heat.

METALS AND NON-METALS

About 118 elements are known today. There are more than 90 metals, 22 non metals and a few metalloids.

Sodium (Na), potassium (K), magnesium(Mg), aluminium(Al), calcium(Ca), Iron(Fe), Barium(Ba) are some metals. Oxygen(O), hydrogen(H), nitrogen(N), sulphur(S), phosphorus(P), fluorine(F), chlorine(Cl), bromine(Br), iodine(l) are some non-metals

Physical properties of metals:

Solid at room temperature except mercury

Ductile (drawn into wires)

Malleable (beaten into thin sheets)

Sonorous(produce sound)

Lustrous(natural shine)

Have high melting point. Cesiumand galliumhave very low melting point.

Generally good conductor of heat and electricity, except lead and mercury which are comparatively poor conductors. Silver and copper are best conductors.

Have high density. Sodium and potassium can be cut with knife, they have low density.

Physical properties of non-metals:
• Occur as solid or gas. Bromine is liquid.
• Generally bad conductors of heat and electricity. Graphite a natural form of carbon is a good conductor.
• Non-lustrous, only iodine has lustre.
• Metals form basic oxides like Magnesium oxide(MgO), while non-metals form acidic oxides (as in acid rain).
*Chemical properties of metals:
1. Reaction with air
Metals can burn in air, react or don't react with air.
Metal + oxygen → Metal Oxide
• Some metals like Na and K are kept immersed in kerosene oil as they react vigorously with air and catch fire.
• Some metals likeMg,Al, Zn, Pb react slowly with air and form a protective layer.
• Mg can also burn in air with a white dazzling light to form its oxide
• Fe and Cu don't burn in air but combine with oxygen to form oxide.When heated iron filings burn when sprinkled over flame.

Try Balancing these Chemical equations yourself

3. Reaction with dilute acids:
Metal + dilute acid → Salt + Hydrogen gas

Metals react with dilute hydrochloric acid and dilute sulphuric acid to form
salt and hydrogen gas.
Fe + 2HCl → FeCl₂ + H₂
Mg + 2HCl→ MgCl₂ + H₂
Zn + 2HCl → ZnCl₂ + H₂
2Al + 6HCl → 2AlCl₃ + 3H₂

Copper, mercury and silver don’t react with dilute acids.
Hydrogen gas produced is oxidised to water when metals react with nitric acid. But Mg and Mn, react with very dilute nitric acid to evolve hydrogen gas.
Mg + 2HNO₃ → Mg(NO₃)2 + H₂

4. Reaction of metals with other metal salts :

                 Salt           Salt
Metal A + solution → solution + Metal B
                of B             of A
All metals are not equally reactive. Reactive metals can displace less reactive metals from their compounds in solution. This forms the basis of reactivity series of metals.
Reactivity series is a list of metals arranged in order of their decreasing activities. 

Reaction between Metals and Non-Metals :

– Reactivity of elements can be understood as a tendency to attain a completely filled valence shell.
– Atom of metals can lose electrons from valence shells to form cations (+ve ions).
– Atom of non-metals gain electrons in valence shell to form anions (–ve ions).
– Oppositely charged ions attract each other and are held by strong electrostatic forces of attraction forming ionic compounds.

Properties of Ionic Compounds :
– Are solid and mostly brittle.
– Have high melting and boiling points. More energy is required to break the strong inter-ionic attraction.
– Generally soluble in water and insoluble in kerosene, petrol.
– Conduct electricity in solution and in molten state. In both cases, free ions are formed and conduct electricity.

Occurance of Metals
Minerals : elements of compounds occuring naturally are minerals.
ORES : mineral fromwhichmetal can be profitably extracted is an ore. For example, sulphide ore, oxide ore, carbonate ore.
– Metals at the bottom of activity series like gold, platinum, silver, copper generally occur in free state. But copper and silver also occur in sulphide and oxide ores.
– Metals ofmediumreactivity (Zn, Fe, Pb etc.) occurmainly as oxides, sulphides or carbonates.
– Metals of high reactivity (K, Na, Ca, Mg andAl) are very reactive and thus found in combined state.
GANGUE : ores are naturally found mixed impurities like soil, sand, etc. called gangue. The gangue is removed from the ore.
METALLURGY : step-wise process of obtaining metal from its ore.
*Enrichment of ore
*Obtaining metal from enriched ore.

Extracting Metals Low in the Activity Series :
By heating the ores in air at high temperature.
*Mercury from cinnabar 

– In the above reaction molten iron is formed and is used to join railway tracks.
This is called thermit reaction.

Extracting Metals at the Top of Activity Series :
These metals
– have more affinity for oxygen than carbon.
– are obtained by electrolytic reduction. Sodium is obtained by electrolysis of
its molten chloride NaCl  → Na⁺ + Cl^–
As electricity is passed through the solution metal gets deposited at cathode and non-metal at anode.
– At cathode :
Na⁺ + e^– → Na
– At anode :
2Cl^– → Cl₂ + 2e–

Refining of Metals :

– Impurities present in the obtainedmetal can be removed by electrolytic refining.
Copper is obtained using this method. Following are present inside the electrolytic tank.
– Anode – slab of impure copper
– Cathode – slab of pure copper
– Solution – aqueous solution of copper sulphate with some dilute sulphuric acid
– From anode copper ions are released in the solution and equivalent amount of copper from solution is deposited at cathode.
– Impurities containing silver and gold gets deposited at the bottom of anode as anode mud.

Corrosion :
– Metals are attacked by substances in surroundings like moisture and acids.
– Silver - it reacts with sulphur in air to form silver sulphide and articles become black.
– Copper - reacts with moist carbon dioxide in air and gains a green coat of copper carbonate.
– Iron-acquires a coating of a brown flaky substance called rust. Both air and moisture are necessary for rusting of iron.

Prevention of corrosion:
– Rusting of iron is prevented by painting, oiling, greasing, galvanizing, chrome plating, anodising and making alloys.
– In galvanization, iron or steel is coated with a layer of zinc because zinc is preferably oxidized than iron.

Alloys :

– These are mixture of metals with metals or non-metals  Adding small amount of carbon makes iron hard and strong.
– Stainless steel is obtained bymixing iron with nickel and chromium. It is hard and doesn’t rust.
– Mercury is added to other metals to make amalgam.
Brass : alloy of copper and zinc.
Bronze : alloy of copper and tin.
– In brass and bronze, melting point and electrical conductivity is lower than that of pure metal.
Solder : alloy of lead and tin has low melting point and is used for welding electrical wires.

METALS AND NON-METALS IN BRIEF

– Metals are generally solid, sonorous, lustrous, good conductor of heat and electricity, malleable, ductile, high melting point, high densities, form basic oxides, form +vely charged ion.
– Non-metals are generally solid or gas, non-lustrous, non-sonorous, bad conductor of heat and electricity, have low melting point, form acidic oxides and form -vely charged ions.
– Metals like Na, K and Ca are highly reactive, while others likeMagnesium, Aluminium, Zinc and Lead are less reactive and some others are least reactive like silver, gold and platinum.
– Metals generally displace hydrogen from acids.
– Reactivity series is based on displacement capability of metals and is a series of metals in the order of their decreasing reactivity.
– Metals and non-metals react to form ionic compounds which are soluble in water, have high melting point and are good conductor of electricity in their aqueous solution or molten state.
– Ores are minerals from which a metal can be profitably extracted.
– Metals are extracted from their ores according to their reactivity.
– Sulphide and chloride ores are roasted while carbonate ores are roasted.
– Pure metals can be obtained using electrolytic refining process.
– Metals are generally attacked by air and corrode. To alter the properties of metals alloys are made.
– Steel, stainless steel, amalgams, brass, bronze and solder are some alloys.

METALS AND NON-METALS