Metals and Non-metals Notes – Class 10 Science

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Metals and Non-metals

 

Elements are divided mainly into two groups on the basis of physical and chemical properties – Metal and Non-metal.

 

Physical Properties of Metals:

 

Hardness: Most of the metals are hard, except alkali metals, such as sodium, potassium, lithium, etc. Sodium, potassium, lithium etc. are very soft metals, these can be cut using knife.

Strength: Most of the metals are strong and have high tensile strength. Because of this big structures are made using metals, such as copper and iron.

State: Metals are solid at room temperature except mercury.

Sound: Metals produce ringing sound, so, metals are called sonorous. Sound of metals is also known as metallic sound. This is the cause that metal wires are used in making musical instruments.

Conduction: Metals are good conductor of heat and electricity. This is the cause that electric wires are made of metals like copper and aluminium.

Malleability: Metals are malleable. This means metals can be beaten into thin sheet. Because of this property iron is used in making big ships.

Ductility: Metals are ductile. This means metals can be drawn into thin wire. Because of this property wires are made of metals.

Melting and boiling point: Metals have generally high melting and boiling points.

Density: Most of the metals have high density.

Color: Most of the metals are grey in color. But gold and copper are exceptions.

 

 

Chemical Properties of Metals

Reaction with oxygen:

Most of the metals form respective metal oxides when react with oxygen.

Metal + Oxygen Metal oxide

Examples:

Reaction of potassium with oxygen: Potassium metal forms potassium oxide when reacts with oxygen.

4K + O2  2K2O

Reaction of sodium with oxygen: Sodium metal forms sodium oxide when reacts with oxygen.

4Na + O2  2Na2O

 

Reaction of metals with water:

Metals form respective metal hydroxide and hydrogen gas when react with water.

Metal + Water Metal hydroxide + Hydrogen

Most of the metals do not react with water. However, alkali metals react vigorously with water.

Reaction of sodium metal with water: Sodium metal forms sodium hydroxide and liberates hydrogen gas along with lot of heat when reacts with water.

Na + H2O NaOH + H2

Reaction of potassium metal with water: Potassium metal forms potassium hydroxide and liberates hydrogen gas along with lot of heat when reacts with water.

K + H2O KOH + H2

Reaction of calcium metal with water: Calcium forms calcium hydroxide along with hydrogen gas and heat when reacts with water.

Ca + 2H2O Ca(OH)2 + H2

 

Reaction of metals with dilute acid:

Metals form respective salts when react with dilute acid.

Metal + dil. acid Metal salt + Hydrogen

Reaction of sodium metal with dilute acid: Sodium metal gives sodium chloride and hydrogen gas when react with dilute hydrochloric acid.

2Na + 2HCl 2NaCl + H2

Reaction of potassium with dilute sulphuric acid: Potassium sulphate and hydrogen gas are formed when potassium reacts with dilute sulphuric acid.

2K + H2SO4  K2SO4 + H2

Metal Oxides: Chemical Properties

Metal oxides are basic in nature. Aqueous solution of metal oxides turns red litmus blue.

Reaction of metal oxides with water:
Most of the metal oxides are insoluble in water. Alkali metal oxides are soluble in water. Alkali metal oxides give strong base when dissolved in water.

Reaction of sodium oxide with water:  Sodium oxide gives sodium hydroxide when reacts with water.

Na2O + H2O 2NaOH

Reaction of magnesium oxide with water: Magnesium oxide gives magnesium hydroxide with water.

MgO + H2O Mg(OH)2

Reactivity Series of Metals

The order of intensity of reactivity is known as reactivity series. Reactivity of element decreases on moving from top to bottom in the given reactivity series.

In the reactivity series, copper, gold, and silver are at the bottom and hence least reactive. These metals are known as noble metals. Potassium is at the top of the series and hence most reactive.

Reactivity of some metals are given in descending order

K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu

Reaction of metals with solution of other metal salts:

Reaction of metals with solution of other metal salt is displacement reaction. In this reaction more reactive metal displace the less reactive metal from its salt.

Metal A + Salt of metal B Salt of metal A + Metal B

Examples:

Iron displaces copper from copper sulphate solution.

Fe + CuSO4  FeSO4 + Cu

Physical properties of non-metals

Hardness: Non-metals are not hard rather they are generally soft. But diamond is exception; it is most hard naturally occurring substance.

State: Non-metals may be solid, liquid or gas.

Lustre: Non-metals have dull appearance. Diamond and iodine are exceptions.

Sonority: Non-metals are not sonorous, i.e. they do not produce a typical sound no being hit.

Conduction: Non-metals are bad conductor of heat and electricity. Graphite which is allotrope of carbon is good conductor of electricity and is an exception.

Malleability and ductility: Non-metals are brittle.

Melting and boiling point: Non-metals have generally low melting and boiling points.

Density: Most of the non-metals have low density.

Color: Non-metals are of many colors.

 

 

Chemical properties of Non-metals

Reaction of non-metals with oxygen: Non-metals form respective oxide when react with oxygen.

Non-metal + Oxygen Non-metal oxide

When carbon reacts with oxygen, carbon dioxide is formed along with production of heat.

C + O2  CO2 + Heat

When carbon is burnt in insufficient supply of air, it forms carbon monoxide. Carbon monoxide is a toxic substance. Inhaling of carbon monoxide may prove fatal.

2C + O2  2CO + Heat

Non-metal oxide:

Non-metal oxides are acidic in nature. Solution of non-metal oxides turns blue litmus red.

Carbon dioxide gives carbonic acid when dissolved in water.

CO2 + H2O H2CO3

Sulphur dioxide gives sulphurous acid when dissolved in water.

SO2 + H2O H2SO3

 

 

Reaction of non-metal with chlorine:

Non metals give respective chloride when they react with chlorine gas.

Non-metal + Chlorine Non-metal chloride

Hydrogen gives hydrogen chloride and phosphorous gives phosphorous trichloride when react with chlorine.

H2 + Cl2  2HCl

P4 + 6Cl2  4PCl3

Reaction of Metal and Non-metal

Many metals form ionic bonds when they react with non-metals. Compounds so formed are known as ionic compounds.

Ions: Positive or negative charged atoms are known as ions. Ions are formed because of loss or gain of electrons. Atoms form ion to obtain electronic configuration of nearest noble gas, this means to obtain stable configuration.

Positive ion: A positive ion is formed because of loss of electrons by an atom. Following are some examples of positive ions.

Sodium forms sodium ion because of loss of one electron. Because of loss of one electron; one positive charge comes over sodium.

Na Na+ + e

Similarly; potassium gets one positive charge by loss of one electron.

K K+ + e

Negative ion: A negative ion is formed because of gain of electron. Some examples are given below.

Chlorine gains one electron in order to achieve stable configuration. After loss of one electron chlorine gets one negative charge over it forming chlorine ion.

Cl + e  Cl

Similarly, fluorine gets one negative charge over it by gain of one electron forming chloride ion; in order to achieve stable configuration.

F + e  F

 

 

Ionic Bonds

Ionic bonds are formed because of transfer of electrons from metal to non-metal. In this course, metals get positive charge because of transfer of electrons and non-metal gets negative charge because of acceptance of electrons. In other words bond formed between positive and negative ion is called ionic bond.

Since, a compound is electrically neutral, so to form an ionic compound negative and positive both ions must be combined. Some examples are given below.

Formation of sodium chloride (NaCl):
In sodium chloride; sodium is a metal (alkali metal) and chlorine is non-metal.

Atomic number of sodium = 11
Electronic configuration of sodium: 2, 8, 1
Number of electrons in outermost orbit = 1
Valence electrons = Electrons in outermost orbit = 1
Atomic number of chlorine = 17
Electronic configuration of chlorine: 2, 8, 7
Electrons in outermost orbit = 7
Therefore, valence electrons = 7

Sodium has one valence electron and chlorine has seven valence electrons. Sodium requires losing one electron to obtain stable configuration and chlorine requires gaining one electron in order to obtain stable electronic configuration. Thus, in order to obtain stable configuration sodium transfers one electron to chlorine.
After loss of one electron sodium gets one positive charge (+) and chlorine gets one negative charge after gain of one electron. Sodium chloride is formed because of transfer of electrons. Thus, ionic bond is formed between sodium and chlorine. Since, sodium chloride is formed because of ionic bond, thus it is called ionic compound. In similar way; potassium chloride (KCl) is formed.

Formation of Magnesium Chloride (MgCl2):

The atomic number of magnesium is 12
Electronic configuration of magnesium: 2, 8, 2
Number of electrons in outermost orbit = 2
Valence electron = 2
Atomic number of chlorine = 17
Electronic configuration of chlorine: 2, 8, 7
Electrons in outermost orbit = 7
Therefore, valence electrons = 7

Magnesium loses two electrons in order to obtain stable electronic configuration. Each of the two chlorine atoms gains one electron lost by magnesium to obtain stable electronic configuration. The bonds so formed between magnesium and chlorine are ionic bonds and compound (magnesium chloride) is an ionic compound.

Formation of calcium chloride: (CaCl2):
Atomic number of calcium is 20.
Electronic configuration of calcium: 2, 8, 8, 2
Number of electrons in outermost orbit = 2
Valence electron = 2
Valence electrons of chlorine = 7

Calcium loses two electrons in order to achieve stable electronic configuration. Each of the two chlorine atoms on the other hand gains one electron losing from calcium to get stability. By losing of two electrons calcium gets two positive charges over it. Each of the chlorine atoms gets one positive charge over it.

The bonds formed in the calcium chloride are ionic bonds and compound (calcium chloride) is an ionic compound. In similar way; Barium chloride is formed.

Formation of Calcium oxide (CaO):

Valence electron = 2
Atomic number of oxygen is 8
Electronic configuration of oxygen is: 2, 6
Number of electrons in outermost orbit = 6
Valence electron = 6
Calcium loses two electrons and gets two positive charges over it in order to get stability. Oxygen gains two electrons; lost by calcium and thus gets two negative charges over it.

Bond formed between calcium oxide is ionic bond. Calcium oxide is an ionic compound. In similar way; magnesium oxide is formed.

Properties of Ionic compound:

  • Ionic compounds are solid. Ionic bond has greater force of attraction because of which ions attract each other strongly. This makes ionic compounds solid.
  • Ionic compounds are brittle.
  • Ionic compounds have high melting and boiling points because force of attraction between ions of ionic compounds is very strong.
  • Ionic compounds generally dissolve in water.
  • Ionic compounds are generally insoluble in organic solvents; like kerosene, petrol, etc.
  • Ionic compounds do not conduct electricity in solid state.
  • Solution of ionic compounds in water conduct electricity. This happens because ions present in the solution of ionic compound facilitate the passage of electricity by moving towards opposite electrodes.
  • Ionic compounds conduct electricity in molten state.

 

Occurance and Extraction of Metals

Source of metal: Metals occur in earth’s crust and in sea water; in the form of ores. Earth’s crust is the major source of metal. Sea water contains many salts; such as sodium chloride, magnesium chloride, etc.

Mineral: Minerals are naturally occurring substances which have uniform composition.

Ores: The minerals from which a metal can be profitably extracted are called ores.

Metals found at the bottom of reactivity series are least reactive and they are often found in nature in free-state; such as gold, silver, copper, etc. Copper and silver are also found in the form of sulphide and oxide ores.

Metals found in the middle of reactivity series, such as Zn, Fe, Pb, etc. are usually found in the form of oxides, sulphides or carbonates.

Metals found at the top of the reactivity series are never found in free-state as they are very reactive, e.g. K, Na, Ca, Mg and Al, etc.

Many metals are found in the form of oxides because oxygen is abundant in nature and is very reactive.

 

 

Extraction of Metals

Metals can be categorized into three parts on the basis of their reactivity: most reactive, medium reactive and least reactive.

Steps of Extraction of Metals

Concentration of ores: Removal of impurities, such as soil, sand, stone, silicates, etc. from mined ore is known as Concentration of Ores.

Ores which are mined often contain many impurities. These impurities are called gangue. First of all, concentration is done to remove impurities from ores. Concentration of ores is also known as enrichment of ores. Process of concentration depends upon physical and chemical properties of ores. Gravity separation, electromagnetic separation, froth flotation process, etc. are some examples of the processes which are applied for concentration of ores.

 

 

Conversion of metals ores into oxides:

It is easy to obtain metals from their oxides. So, ores found in the form of sulphide and carbonates are first converted to their oxides by he process of roasting and calcination. Oxides of metals so obtained are converted into metals by the process of reduction.

Roasting: Heating of sulphide ores in the presence of excess air to convert them into oxides is known as ROASTING.

Calcination: Heating of carbonate ores in the limited supply of air to convert them into oxides is known as CALCINATION.

Reduction: Heating of oxides of metals to turn them into metal is known as REDUCTION.

Purification: Metal; so obtained is refined using various methods.

 

 

Extraction of Metals of Least Reactivity

Mercury and copper, which belong to the least reactivity series, are often found in the form of their sulphide ores. Cinnabar (HgS) is the ore of mercury. Copper glance (Cu2S) is the ore of copper.

Extraction of mercury metal: Cinnabar (HgS) is first heated in air. This turns HgS [mercury sulphide or cinnabar] into HgO (mercury oxide); by liberation of sulphur dioxide.

Mercury oxide so obtained is again heated strongly. This reduces mercury oxide to mercury metal.

2HgS + 3O2  2HgO + 2SO2

2HgO 2Hg + O2

Extraction of copper metal: Copper glance (Cu2S) is roasted in the presence of air. Roasting turns copper glance (ore of copper) into copper (I) oxide. Copper oxide is then heated in the absence of air. This reduces copper (I) oxide into copper metal.

2Cu2S + 3O2  2Cu2O + 2SO2

2Cu2O + Cu2S 6Cu + SO2

Extraction of Metals of middle reactivity:

Iron, zinc, lead, etc. are found in the form of carbonate or sulphide ores. Carbonate or sulphide ores of metals are first converted into respective oxides and then oxides are reduced to respective metals.

Extraction of zinc: Zinc blende (ZnS: zinc sulphide) and smithsonite or zinc spar or calamine (ZnCO3: zinc carbonate) are ores of zinc. Zinc blende is roasted to be converted into zinc oxide. Zinc spar is put under calcination to be converted into zinc oxide.

2ZnS + 3O2  2ZnO + 2SO2

ZnCO3  ZnO + CO2

Zinc oxide so obtained is reduced to zinc metal by heating with carbon (a reducing agent).

ZnO + C Zn + CO

Extraction of iron from Hematite (Fe2O3): Hematite ore is heated with carbon to be reduced to iron metal.

Fe2O3 + 3C 4Fe + 3CO2

Extraction of lead from lead oxide: Lead oxide is heated with carbon to be reduced to lead metal.

2PbO + C 2Pb + CO2

Reduction of metal oxide by heating with aluminum: Metal oxides are heated with aluminum (a reducing agent) to be reduced to metal. Following is an example:

Manganese dioxide and copper oxide are reduced to respective metals when heated with aluminium.

3MnO2 + 4Al 3Mn + 2Al2O3

3CuO + 2Al 3Cu + Al2O3 + heat

Thermite Reaction: Ferric oxide; when heated with aluminum; is reduced to iron metal. In this reaction, lot of heat is produced. This reaction is also known as Thermite Reaction. Thermite reaction is used in welding of electric conductors, iron joints, etc. such as joints in railway tracks. This is also known as Thermite Welding (TW).

Fe2O3 + 2Al 2Fe + Al2O3 + heat

Refining or purification of metals:

Metals extracted from various methods contains some impurities, thus they are required to be refined. Most of the metals are refined using electrolytic refining.

Electrolytic Refining: In the process of electrolytic refining a lump of impure metal and a thin strip of pure metal are dipped in the salt solution of metal to be refined. When electric current is passed through the solution, pure metal is deposited over thin strip of pure metal; from lump of impure metal. In this, impure metal is used as anode and pure metal is used as cathode.

 

Electrolytic refining of copper:

A lump of impure copper metal and a thin strip of pure copper are dipped in the solution of copper sulphate. Impure lump of metal is connected with the positive pole and thin strip of pure metal is connected with the negative pole. When electric current is passed through the solution, pure metal from anode moves towards cathode and is deposited over it. Impurities; present in metal are settled near the bottom of anode in the solution. Settled impurities in the solution are called anode mud.

Cu − 2e  Cu+ +

Cu+ + + 2e  Cu

Corrosion:

Most of the metals keep on reacting with the atmospheric air. This leads to formation of a layer over the metal. In the long run, the underlying layers of the metal keep on getting lost due to conversion into oxides or sulphides or carbonate, etc. As a result, the metal gets eaten up. This process is called corrosion.

Rusting of Iron: Rusting of iron is the most common form of corrosion. When iron articles; like gate, grill, fencing, etc. come in contact with moisture present in air, the upper layer of iron turns into iron oxide. Iron oxide is brown-red in color and is known as rust. This phenomenon is called rusting of iron.

If rusting is not prevented in time, the whole iron article would turn into iron oxide. This is also known as corrosion of iron. Rusting of iron gives huge loss every year.

Prevention of Rusting: For rusting, iron must come in contact with oxygen and water. Rusting is prevented by preventing the reaction between atmospheric moisture and the iron article. This can be done by painting, greasing, galvanization, electroplating, etc.


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