CBSE NOTES Class 10 Science Chapter-3. Metals and Non-metals Chemical Properties of Metals

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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

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.