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Metals:

Physical properties of metals:

(i)  Metals are solid and shiny.

(ii) They conduct heat and electricity.

(iii) Metals are ductile. 

(iv) Metals are meliable.

(v) Metals are sonorous. 

Meliability: some metals can be beaten into thin sheets. This property is called malleability.

Ductility: The ability of metals to be drawn into thin wires is called ductility.

All metals have hight melting point. The best conductors of heat are silver and
copper. Lead and mercury are comparatively poor conductors of heat.

The full name of PVC : polyvinylchloride (PVC)

PVC and rubber like materials are bad conductor of heat and electricity.

Sonorous: Sonorous is a physical property of metals. By this property they produce sound on striking. Using this property of metals, school bells is made. 

Non-metals:

carbon, sulpher, oxygen, iodine and hydrogen are non-metals. 

Physical property of Non-metals:

(i)  Non-metals are not solid and shiny.

(ii) They are not good conductor of heat and electricity. 

(iii)  Non-metals are not ductile.

(iv) Non-metals are not melleable. 

(v) Non-metals do not produce sound on striking. 

 

The non-metals are either solids or gases except bromine which is a liquid.

Some other properties of metals nd non-metals:

(i) All metals except mercury exist as solids at room temperature.

(ii) Mercury is found in liquid form at room temperature.

(iii)  Gallium and caesium have very low melting points. These two metals will melt if you keep them on your palm.

(iv)  Iodine is a non-metal but it is lustrous (shiny).

(v)  Alkali metals (lithium, sodium, potassium) are so soft that they can be cut with a knife. They have low densities and low melting points.

Carbon And its Allotrope:

Carbon is a non-metal that can exist in different forms. Each form is called an allotrope.

Allotropes of carbon:

(i) Daimond 

(ii) Graphite

(iii) Buckminsterfullerene

(i) Diamond, an allotrope of carbon, is the hardest natural substance known and has a very high melting and boiling point.

(ii) Graphite, another allotrope of carbon, is a conductor of electricity.

Note: Most non-metals produce acidic oxides when dissolve in water. While metals produce basic oxides when dissolve in water. 

 

 

 

CHEMICAL PROPERTIES OF METALS

(i)  Almost all metals combine with oxygen to form metal oxides.

Metal + Oxygen → Metal oxide

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

2Cu     +    O₂    →    2CuO
(Copper)                (Copper(II) oxide)

Similarly, aluminium forms aluminium oxide.

4Al     +    3O₂      →        2Al₂O₃

(Aluminium)              (Aluminium oxide)

Amphoteric Oxides : some metal oxides, such as aluminium oxide, zinc oxide, etc., show both acidic as well as basic behaviour. Such metal oxides which react with both acids as well as bases to produce salts and water are known as amphoteric oxides.

Example: aluminium oxide and zinc oxide are amphoteric oxides. 

Reaction of Metal oxides with acids

Aluminium oxide reacts with hydrochloric acid and produces a salt aluminium chloride and water. 

The chemical equation is as;

Al₂ O₃ + 6HCl → 2AlCl₃ + 3H₂O

Reaction of Metal oxides with bases:

Aluminium oxide reacts with sodium hydroixe produces Sodium Aluminate and water. 

Al₂O₃    +    2NaOH      →       2NaAlO₂    +     H₂O
                                      (Sodium aluminate)

Solubility of metal oxides in water:

Most metal oxides are insoluble in water but some of these dissolve in water to form alkalis. Sodium oxide and potassium oxide dissolve in water to produce alkalis.

The disolving of sodium oxide and potassium oxide in water gives sodium hydroxide alkalis and potassium hydroxide alkalis respectively. 

Na₂O(s) +  H₂O(l)  →  2NaOH(aq)
K₂O(s)   +  H₂O(l)  →  2KOH(aq)

Reactivity of metals with oxygen: 

Different metals show different reactivities towards oxygen.

Reaction of sodium and potassium with oxygen: 

Metals such as potassium and sodium react so vigorously that they catch fire if kept in the open. Hence, to protect them and to prevent accidental fires, they are kept immersed in kerosene oil. 

Some metal oxides form protective layer:

At ordinary temperature, the surfaces of metals such as magnesium, aluminium, zinc, lead, etc., are covered with a thin layer of oxide. The protective oxide layer prevents the metal from further oxidation.

Some metal does not react with oxygen:

Iron does not burn on heating but iron filings burn vigorously when sprinkled in the flame of the burner. Copper does not burn, but the hot metal is coated with a black coloured layer of copper(II) oxide. Silver and gold do not react with oxygen even at high temperatures.

Anodising: 

Anodising is a process of forming a thick oxide layer of aluminium. Aluminium
develops a thin oxide layer when exposed to air. This aluminium oxide coat makes it resistant to further corrosion. The resistance can be improved further by making the oxide layer thicker.

Anodising of Aluminium: 

During anodising, a clean aluminium article is made the anode and is electrolysed with dilute sulphuric acid. The oxygen gas evolved at the anode reacts with aluminium to make a thicker protective oxide layer. This oxide layer can
be dyed easily to give aluminium articles an attractive finish.

Reaction of metals with water:

Metals react with water and produce a metal oxide and hydrogen gas. Metal oxides that are soluble in water dissolve in it to further form metal hydroxide.

General equations:

Metal + Water → Metal oxide + Hydrogen
Metal oxide + Water → Metal hydroxide

Reaction of sodium and potassium with cold water:

Metals like potassium and sodium react violently with cold water. In case of sodium and potassium, the reaction is so violent and exothermic that the evolved hydrogen immediately catches fire.
2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g) + heat energy
2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + heat energy

Reaction of calcium with water:

The reaction of calcium with water is less violent. The heat evolved is not sufficient for the hydrogen to catch fire.
Ca(s) +  2H₂O(l) → Ca(OH)₂(aq) + H₂(g) 

Calcium starts floating because the bubbles of hydrogen gas formed stick to the surface of the metal.

Reaction of metals with hot water:

Magnesium does not react with cold water. It reacts with hot water to form magnesium hydroxide and hydrogen. It also starts floating due to the bubbles of hydrogen gas sticking to its surface.

Reaction of metals with steam:

Metals like aluminium, iron and zinc do not react either with cold or hot water. But they react with steam to form the metal oxide and hydrogen.
2Al(s) + 3H₂O(g) → Al₂O₃(s) + 3H₂(g)
3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

Some Metals do not react with water:

Metals such as lead, copper, silver and gold do not react with water at all.

Reaction of metals with Acids:

Metals react with acids give corresponding salt and hydrogen gas. 

Metal + Dilute acid → Salt + Hydrogen

Hydrogen gas is not evolved when a metal reacts with nitric acid. It is
because HNO₃ is a strong oxidising agent. It oxidises the H₂ produced to
water and itself gets reduced to any of the nitrogen oxides (N₂O, NO,
NO₂). But magnesium (Mg) and manganese (Mn) react with very dilute
HNO₃ to evolve H₂ gas.

Aqua regia: is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3:1.

It can dissolve gold, even though neither of these acids can do so alone. Aqua regia is a highly corrosive, fuming liquid. It is one of the few reagents that is able to dissolve gold and platinum.

Reaction of metals with other metal salt: 

Highly reactive metals can displace less reactive metals from their compounds in solution or molten form. this is called displacement reaction. 

Metal A + Salt solution of B → Salt solution of A + Metal B

The Reactivity Series: 

K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au 

Reaction with metals and non-metals: 

Mostly Metals form cation (postive charge) nad non-metals form anaion (negative charge).  

Cation And Anaion : To understand these both cation and anaion, we have to understand electronic configuration of elements and their valencies. 

Valency : The number of valence electrons present in the outer most shells of an atom is known as valency. Ex. Electronic configuration of Sodium (Na) is 

2    8     1 

There are three shells in sodium atom and the outer most shell has 1 electron can be shared, so valence electron of sodium is 1. 

• If outer most shell has 1, 2, 3 or 4 electrons these can be given in sharing of electrons. so 1, 2, 3, and for will be valance electrons.
• If outer most shell has 5, 6 or 7 electrons these can not be given in sharing of electrons as These need electrons to complete their octet. 

  Required valance electrons for outer most shell having 5 electrons = 8 - 5 = 3 

  Required valance electrons for outer most shell having 6 electrons = 8 - 6 = 2

  Required valance electron for outer most shell having 7 electrons = 8 - 7 = 1

        

Type of element	Element	Atomic Number	Number of electron in shells    K      L        M         N
Noble gases	Helium (He)  Neon (Ne)   Argon (Ar)	2     10     18	2          2       8     2       8         8
Metals	Sodium (Na) Magnesium (Mg) Aluminium (Al) Potassium (K) Calcium (Ca)	11       12       13       19       20	2       8         1   2       8         2   2       8         3   2       8         8          1   2       8         8          2
Non- metals	Nitrogen (N) Ogygen (O) Flurine (F) Phosphorus (P) Sulpher (S) Chlorine (Cl)	7       8       9       15       16       17	2       5   2       6   2       7   2       8        5   2       8        6   2       8        7

 

A sodium atom has one electron in its outermost shell. If it loses the electron from its M shell then its L shell now becomes the outermost shell and that has a stable octet. The nucleus of this atom still has 11 protons but the number of electrons has become 10, so there is a net positive charge giving us a sodium cation Na⁺ . On the other hand chlorine has seven electrons in its outermost shell and it requires one more electron to complete its octet. If sodium and chlorine were to react, the electron lost by sodium could be taken up by chlorine.
After gaining an electron, the chlorine atom gets a unit negative charge, because its nucleus has 17 protons and there are 18 electrons in its K, L and M shells. This gives us a chloride anion C1^-. So both these elements can have a give-and-take relation between them.

E.g : 

Na   →  Na⁺ + e^-
2,8,1     2,8
               (Sodium cation)

Cl      + e^- → Cl^-
2,8,7             2,8,8
                   (Chloride anion)

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

Properties Of Ionic Compound :

(i) Physical nature: Ionic compounds are solids and are somewhat hard because of the strong force of attraction between the positive and negative ions. These compounds are generally brittle and break into pieces when pressure is applied.
(ii) Melting and Boiling points: Ionic compounds have high melting and boiling points (see Table 3.4). This is because a considerable amount of energy is required to break the strong inter-ionic attraction.
(iii) Solubility: Electrovalent compounds are generally soluble in water and insoluble in solvents such as kerosene, petrol, etc.
(iv) Conduction of Electricity: The conduction of electricity through a solution involves the movement of charged particles. A solution of an ionic compound in water contains ions, which move to the opposite electrodes when electricity is passed through the solution. 

Ionic compounds in the solid state do not conduct electricity because movement of ions in the solid is not possible due to their rigid structure. But ionic compounds conduct in the molten state. This is possible in the molten state since the elecrostatic forces of attraction between the oppositely charged ions are overcome due to the heat. Thus, the ions move freely and conduct electricity.

OCCURRENCE OF METALS

The earth’s crust is the major source of metals. Seawater also contains
some soluble salts such as sodium chloride, magnesium chloride, etc.

Minerals:​ The elements or compounds, which occur naturally in the earth’s crust, are known as minerals.

Ores: The minerals contain a very high percentage of a particular metal and the metal can be profitably extracted from it. These minerals are called ores.

Gangue: Ores mined from the earth are usually contaminated with large amounts of impurities such as soil, sand, etc., are called gangue. 

Extraction of Metals:

All metals are extracted from its ore. Some metals are found in the earth’s crust in the free state. Some are found in the form of their compounds. 

Metals                                Ores                                    Formulla​

Iron                                    Haematite/magnetite           Fe₂O₃/Fe₃O₄

Mercury                             Cinnabar                             HgS 

Zinc                                   Zincite 

Lead                                  Galena

Aluminium                          Bauxite  

 

Extraction of Metals acording to their reactivity.

(i)   Highly reactive: Such as K, Na, Ca, Mg, Al like metals are extracted using Electrolysis. 

(ii)  Middle reactive :Such as Zn, Fe, Pb and Cu are extracted by Reduction using carbon.

(iii) Less reactive : Ag and Au are found in native state.

Contains of ores:

Ores mined from the earth are usually contaminated with large amounts
of impurities such as soil, sand, etc., called gangue. The impurities must
be removed from the ore prior to the extraction of the metal.

The processes used for removing the gangue from the ore are based on the differences between the physical or chemical properties of the gangue and the ore.

(i) Extracting Metals Low in the Activity Series:

Metals low in the activity series are very unreactive. Copper and silver are also found in the combined state as their sulphide or oxide ores. The oxides of these metals can be reduced to metals by heating alone. When ores are heated in the air, They are first converted into corresponding metals oxide. Further heating gives pure metals from its oxide.  

2HgS(s) + 3O (g) →2HgO(s) + 2SO (g)

(ii)  Extracting Metals in the Middle of the Activity Series: 

The metals in the middle of the activity series such as iron, zinc, lead, copper, etc., are moderately reactive. These are usually present as sulphides or carbonates in nature. The metal sulphides and carbonates are converted into metal oxides. As It is easier to obtain a metal from its oxide, as compared to its sulphides and carbonates.

Converting into oxides: 

(a)  Roasting: The sulphide ores are converted into oxides by heating
strongly in the presence of excess air. This process is known as roasting.

2ZnS(s) + 3O (g) Heat→ 2ZnO(s) + 2SO (g)

(b)  Calcination:​ The carbonate ores are changed into oxides by heating strongly in limited air. This process is known as calcination.

ZnCO (s) → ZnO(s) + CO (g)

The metal oxides are then reduced to the corresponding metals by using suitable reducing agents such as carbon. For example, when zinc oxide is heated with carbon, it is reduced to metallic zinc.

ZnO(s) + C(s) →  Zn(s)     +   CO(g) 

                           (reduced)  (oxidised)

There is used reduction reaction to reduce Metals from its compound. 

Using of displacement reaction: 

Besides using carbon (coke) to reduce metal oxides to metals, sometimes displacement reactions can also be used. The highly reactive metals such as sodium, calcium, aluminium, etc., are used as reducing agents because they can displace metals of lower reactivity from their compounds. For example, when manganese dioxide is heated with aluminium powder, the following reaction takes place –

3MnO₂(s) + 4Al(s) → 3Mn(l) + 2Al₂O₃(s) + Heat

Thermit Reaction: Some displacement reactions are highly exothermic. the amount of heat evolved is so large that the metals are produced in the molten state. such a reaction is known as thermit reaction. 

use of thermit reaction:

The reaction of iron(III) oxide (Fe₂O₃) with aluminium is used to join railway tracks or cracked machine parts. 

Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s) + Heat

(iii) Extracting Metals towards the Top of the Activity Series:

The metals high up in the reactivity series are very reactive. These metals are obtained by electrolytic reduction. For example, sodium, magnesium and calcium
are obtained by the electrolysis of their molten chlorides. The metals are deposited at the cathode (the negatively charged electrode), whereas, chlorine is liberated at the anode (the positively charged electrode).

Highly reactive metals have highly affinity for oxygen:

Highly reactive metals have more affinity for oxygen than carbon. So they can not be obtained from their compound by heating with carbon. For example, carbon cannot reduce the oxides of sodium, magnesium, calcium, aluminium, etc., to the respective metals.

Refining of Metals

The metals produced by various reduction processes described above are not very pure. They contain impurities, which must be removed to obtain pure metals. The most widely used method for refining impure metals is electrolytic refining.

Electrolytic Refining:

Many metals, such as copper, zinc, tin, nickel, silver, gold, etc., are refined electrolytically. In this process, the impure metal is made the anode and a thin strip of pure metal is made the cathode.

A solution of the metal salt is used as an electrolyte. On passing the current through the electrolyte, the pure metal from the anode dissolves into the electrolyte. An equivalent amount of pure metal from the electrolyte is deposited on the cathode. The soluble impurities go into the solution, whereas, the insoluble impurities settle down at the bottom of the anode and are known as anode mud.

 

 

Corrosion: 

When some metals articles come in the contact of oxygen or moistures, this reacts with oxygen or water vapour and forms metal oxides which get rusted, resulting in fatal surface loss of metals. This surface loss of metal is known as corrsion. 

Prevention of Corrosion:

Corrosion can be prevented by 

(i)  By painting, Oiling and by greasing.

(ii)  By Galvanisation, chrome plating or by anodising.

(iii)  By making alloys. 

Galvanisation: A method of protecting steel and iron from rusting by coating them with a thin layer of zinc. This process is known as Galvanisation.

Alloying: Pure metals are not used to make articles. So there are mixed some other substances to make them hard and strong and causing changing in metal's properties. This process is called alloying.  

• Alloying is a very good method of improving the properties of a metal.
• An alloy is a homogeneous mixture of two or more metals, or a metal and a nonmetal.
• It is prepared by first melting the primary metal, and then, dissolving the other elements in it in definite proportions. It is then cooled to room temperature.

Some examples of alloys:

(I) Iron is the most widely used metal. But it is never used in its pure state. This is because pure iron is very soft and stretches easily when hot. But, if it is mixed with a small amount of carbon (about 0.05 %), it becomes hard and strong steel.

(ii) Iron is mixed with nickel and chromium to make stainless steel. 

(iii) Pure gold is of 24 carat gold, which is very shoft. so there can not be made jewellery by it. It is alloyed with either silver or copper to make it hard. 22 parts of pure gold is alloyed with 2 parts of either copper or silver such a gold is called 22 carat of gold.                                                                                                  

Alloy	Mixture	Symbols
Brass (पीतल)	Copper + zinc	Cu + Zn
Bronze (कांसा)	Copper + tin	Cu + Sn
Solder (सोल्डर)	Lead + tin	Pb + Sn

 

Important Questions Based on Chapter

 

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