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 a 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 a 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 reacts with oxygen.
Metal + Oxygen ⇨ Metal oxide
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 metals with water:
Metals form respective metal hydroxide and hydrogen gas when reacts 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 a 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 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 reacts 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 aluminum with dilute hydrochloric acid: Aluminum 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
Copper, gold and silver are known as noble metals. These do not react with water or dilute acids.
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 the solution of other metal salt is called the 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
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, aluminum and zinc are more reactive than copper. That’s why they displace copper from its salt solution.
Physical properties of non-metals
Hardness: Non-metals are not hard rather they are generally soft. But diamond is an exception; it is most hard naturally occurring substance.
State: Non-metals may be solid, liquid or gas.
Lustre: Non-metals have a dull appearance. Diamond and iodine are exceptions.
Sonority: Non-metals are not sonorous, i.e. they do not produce a typical sound any being hit.
Conduction: Non-metals are the bad conductor of heat and electricity. Graphite which is an 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 the production of heat.
C + O2 ⇨ CO2 + Heat
When carbon is burnt in the 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 reacts with oxygen. Sulphur caught fire when exposed to air.
S + O2 ⇨ SO2
When hydrogen reacts with oxygen it gives water.
2H2 + O2 ⇨ 2H2ONon-metal oxide:
Non-metal oxides are acidic in nature. The solution of non-metal oxides turns blue litmus red.
Carbon dioxide gives carbonic acid when dissolved in water.
CO2 + H2O ⇨ H2CO3
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 reacts 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.
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.
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 the loss of two electrons.
Ca ⇨ Ca+ + + 2e−
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 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 the loss of one electron sodium gets one positive charge (+) and chlorine gets one negative charge after the gain of one electron. Sodium chloride is formed because of transfer of electrons. Thus, an ionic bond is formed between sodium and chlorine. Since sodium chloride is formed because of ionic bond, thus it is called the ionic compound. In a similar way, potassium chloride (KCl) 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 conducts 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.
Occurrence and Extraction of Metals
Source of metal: Metals occur in earth’s crust and in seawater; in the form of ores. Earth’s crust is the major source of metal. Seawater contains many salts; such as sodium chloride, magnesium chloride, etc.
Mineral: Minerals are naturally occurring substances which have the 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 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.
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
Reduction of metal oxide by heating with aluminum: 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
Extraction of Metals of high reactivity
Metals of high reactivity; such as sodium, calcium, magnesium, aluminium, etc. are extracted from their ores by electrolytic reduction. These metals cannot be reduced using carbon because carbon is less reactive than them.
Electrolytic Reduction: Electric current is passed through the molten state of metal ores. Metal; being positively charged; is deposited over the cathode.
Example: When electric current is passed through molten state or solution of sodium chloride, sodium metal deposited over cathode.
Na+ + e− ⇨ Na
2Cl− − e− ⇨ Cl2
2NaCl ⇨ 2Na + Cl2
Metals obtained from the process of electrolytic reduction are pure in form.
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 an electric current is passed through the solution, pure metal is deposited over a thin strip of pure metal; from the lump of impure metal. In this, impure metal is used as the anode and pure metal is used as the 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
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. If this phenomenon of 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.