Monday 4 April 2016

Notes on Metals and Non Metal

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
Lithium, potassium, sodium, etc. are known as alkali metals. Alkali metals react vigorously with oxygen.
Reaction of magnesium metal with oxygen: Magnesium metal gives magnesium oxide when reacts with oxygen. Magnesium burnt with dazzling light in air and produces lot of heat.
2Mg + O2 ⇨ 2MgO
Reaction of aluminium metal with oxygen: Aliminium metal does not react with oxygen at room temperature but it gives aluminium oxide when burnt in air.
4Al + 3O2 ⇨ 2Al2O3
Reaction of zinc metal with oxygen: Zinc does not react with oxygen at room temperature. But it gives zinc oxide when heated strongly in air.
2Zn + O2 ⇨ 2ZnO
Reaction of Iron metal with oxygen: Iron does not react with oxygen at room temperature. But when iron is heated strongly in air, it gives iron oxide.
3Fe + 2O2 ⇨ Fe3O4
Iron fillings give sparkle in flame when burnt.
Reaction of copper metal with oxygen: Copper does not react with oxygen at room temperature but when burnt in air, it gives copper oxide.
2Cu + O2 ⇨ 2CuO

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 magnesium metal with water: Magnesium metal reacts with water slowly and forms magnesium hydroxide and hydrogen gas.
Mg + 2H2O ⇨ Mg(OH)2 + H2
When steam is passed over magnesium metal, magnesium oxide and hydrogen gas are formed.
Mg + H2O ⇨ MgO + H2
Reaction of aluminium metal with water: Reaction of aluminium metal with cold water is too slow to come into notice. But when steam is passed over aluminium metal; aluminium oxide and hydrogen gas are produced.
2Al + 3H2O ⇨ Al2O3 + 2H2
Reaction of zinc metal with water: Zinc metal produces zinc oxide and hydrogen gas when steam is passed over it. Zinc does not react with cold water.
Zn + H2O ⇨ ZnO + H2
Reaction of Iron with water: Reaction of iron with cold water is very slow and come into notice after a long time. Iron forms rust (iron oxide) when reacts with moisture present in atmosphere.
Iron oxide and hydrogen gas are formed by passing of steam over iron metal.
3Fe + 4H2O ⇨ Fe3O4 + 4H2
Other metals usually do not react with water or react very slowly.

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
Reaction of magnesium metal with dilute hydrochloric acid: Magnesium chloride and hydrogen gas are formed when magnesium reacts with dilute hydrochloric acid.
Mg + 2HCl ⇨ MgCl2 + H2
Reaction of aluminium with dilute hydrochloric acid: Aluminium chloride and hydrogen gas are formed.
2Al + 6HCl ⇨ 2AlCl3 + 3H2
Reaction of zinc with dilute sulphuric acid: Zinc sulphate and hydrogen gas are formed when zinc reacts with dilute sulphuric acid. This method is used in laboratory to produce hydrogen gas.
Zn + H2SO4 ⇨ ZnSO4 + H2
CopperCopper, gold and silver are known as noble metals. These do not react with water or dilute acids.

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
Reaction of potassium oxide with water: Potassium oxide gives potassium hydroxide when reacts with water.
K2O + H2O ⇨ 2KOH
Reaction of zinc oxide and aluminium oxide: Aluminium oxide and zinc oxide are insoluble in water. Aluminium oxide and zinc oxide are amphoteric in nature. An amphoteric substance shows both acidic and basic character. It reacts with base like acid and reacts with acid like a base.
When zinc oxide reacts with sodium hydroxide, it behaves like an acid. In this reaction, sodium zicate and water are formed.
ZnO + 2NaOH ⇨ Na2ZnO2 + H2O
Zinc oxide behaves like a base when reacts with acid. Zinc oxide gives zinc chloride and water on reaction with hydrochloric acid.
ZnO + 2HCl ⇨ ZnCl2 + H2O
In similar way aluminium oxide behaves like a base when reacts with an acid and behaves like an acid when reacts with a base.
Aluminium oxide gives sodium aluminate along with water when reacts with sodium hydroxide.
Al2O3 + 2NaOH ⇨ 2NaAlO2 + H2O
Aluminium oxide gives aluminium chloride along with water when it reacts with hydrochloric acid.
Al2O3 + 6HCl ⇨ 2AlCl3 + 3H2O

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
reactivity series of metals

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
Similarly, aluminium and zinc displace copper from the solution of copper sulphate.
2Al + 3CuSO4 ⇨ Al2(SO4 )3 + 3Cu
Zn + CuSO4 ⇨ ZnSO4 + Cu
In all the above examples, iron, aluminium and zinc are more reactive than copper. That’s why they displace copper from its salt solution.
When copper is dipped in the solution of silver nitrate, it displaces silver and forms copper nitrate.
Cu + 2AgNO3 + Cu(NO3 )2 + 2Ag
In this reaction copper is more reactive than silver and hence displace silver from silver nitrate solution forming copper nitrate.
Silver metal does not react with copper sulphate solution. Because silver is less reactive than copper and not able to displace copper from its salt solution.
Ag + CuSO4 ⇨ No reaction
Similarly, when gold is dipped in the solution of copper nitrate, no reaction takes place. Because copper is more reactive than gold.
Au + CuSO4 ⇨ No reaction
In similar way no reaction takes place when copper is dipped in the solution of aluminium nitrate. Because copper is less reactive than aluminium.
Al(NO3 )3 + Cu ⇨ No reaction

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
Sulphur gives sulphur dioxide when react with oxygen. Sulphur caught fire when exposed to air.
S + O2 ⇨ SO2
When hydrogen reacts with oxygen it gives water.
2H2 + O2 ⇨ 2H2O

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
Sulphur dioxide gives sulphur trioxide when it reacts with oxygen.
2SO2 + O2 ⇨ 2SO3
Sulphur trioxide gives sulphuric acid when dissolved in water.
SO3 + H2O ⇨ H2SO4

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−
Magnesium forms positive ion because of loss of two electrons. Two positive charges come over magnesium because of loss of two electrons.
Mg ⇨ Mg+ + + 2e−
Similarly calcium gets two positive charges over it by loss of two electrons.
Ca ⇨ Ca+ + + 2e−
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−
Oxygen gets two negative charge over it by gain of two electrons forming oxide ion; in order to obtain stable configuration.
O + 2e− ⇨ O− −

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
lewis dot structure sodium chloride
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
lewis dot structure magnesium chloride
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.
lewis dot structure calcium chloride
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.
lewis dot structure calcium oxide
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

    flow chart metal extraction
    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.

    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 aluminium: Metal oxides are heated with aluminium (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 aluminium; 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
    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 the 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.



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