Reactions with Metals - JEE Important Topic

VSAT 2022

Introduction to Metals

Metals are compounds that are mainly characterised by distinctive properties like high thermal and electrical conductivity and high reactivity.  Some metals like copper, gold, silver, etc are found in a free state. Most metals are found in the earth's crust as metal ores like oxides, sulphides, silicates, etc. Most of the metals have high malleability, ductility and strength. Sodium and potassium are soft metals that do not show these properties. Except for mercury, which is found in a liquid state, almost all other metals are found in a solid state.w


Around one-third of elements are metals. They are characterised by metallic bonding, where the valence electrons are delocalised or free to move. Due to this, most metals show high conductivity. Metals find extensive applications in automobiles, satellites, electronics, manufacturing industries, utensils, etc. because of their distinctive properties and their abundance in nature. Generally, metal atoms contain less than half the full amount of electrons in their outermost shell. Some metals like lithium and sodium are highly reactive, but noble metals like gold and silver are non-reactive. The most reactive metal in the periodic table is Cesium. The reactivity series of metals is an arrangement of metals in descending order of reactivity. 


Chemical Reaction of Metals

The chemical reaction of metal is based on the reactivity of metals. Most of the metals react with oxygen, water and acids. Generally, alkali metals and alkaline metals show high reactivity and react with dilute acids, oxygen and water. Let us see some examples of reactions with metals.


Reaction of Metals with Oxygen

Oxygen is a strong reactive non-metal and a powerful oxidising agent. Most of the metals react with oxygen and form metal oxide. These metal oxides are basic, with some exceptions like aluminium and zinc oxide. Some reactions are as follows:

  • $2 \mathrm{Zn}+\mathrm{O}_{2} \rightarrow 2 \mathrm{ZnO}$

  • When copper is heated in the air, it forms copper(II) oxide, a black oxide. 

$2 \mathrm{Cu}+\mathrm{O}_{2} \rightarrow 2 \mathrm{CuO}$

  • Alkali metals burn vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide and other alkali metals form superoxides (like KO2). 

$(1) \ \mathrm{Li}+\mathrm{O}_{2} \rightarrow 2 \mathrm{Li}_{2} \mathrm{O} \\$ $(2) \ \mathrm{Na}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{~g}) \rightarrow \mathrm{Na}_{2} \mathrm{O}_{2}(\mathrm{~s}) $

Compounds like Na2O2, which are rich in oxygen are called ‘peroxides’. 

  • Aluminium forms a very thin oxide (Al2O3) layer on the surface which protects metals from further chemical attack. 

  • As mentioned, metal oxides such as aluminium oxide and zinc oxide are amphoteric oxides, as they show both acidic and basic behaviour. They react with both acids as well as bases to produce corresponding salt and water.

$\mathrm{Al}_{2} \mathrm{O}_{3}+6 \mathrm{HCl} \rightarrow 2 \mathrm{AlCl}_{3}+3 \mathrm{H}_{2} \mathrm{O} \\$ $\mathrm{Al}_{2} \mathrm{O}_{3}+2 \mathrm{NaOH} \rightarrow \mathrm{NaAlO}_{2}+\mathrm{H}_{2} \mathrm{O}$

Thus, oxygen’s reaction with metals shows that metals have different reactivity.


Reaction of Metals with Water

Metals react with water and produce metal hydroxides and hydrogen gas.

Metal + Water $\rightarrow$ Metal oxide + Hydrogen gas

Metal oxide + Water $\rightarrow$ Metal hydroxide

  • Metals like potassium and sodium react violently with cold water because the reaction is exothermic and the evolved hydrogen immediately catches fire. 

$2 \mathrm{Na}(\mathrm{s})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{I}) \rightarrow 2 \mathrm{NaOH}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{~g})+\text { heat energy }$

The reaction of calcium is not that violent.

$\mathrm{Ca}(\mathrm{s})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{I}) \rightarrow \mathrm{Ca}(\mathrm{OH})_{2}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{~g})$

  • Magnesium reacts with hot water and produces magnesium hydroxide and hydrogen. 

  • Metals like aluminium, iron, zinc, etc. react with steam to produce corresponding metal oxide and hydrogen.

$(2) \ \mathrm{Al}(\mathrm{s})+3 \mathrm{H}_{2} \mathrm{O}(\mathrm{g}) \rightarrow \mathrm{Al}_{2} \mathrm{O}_{3}(\mathrm{~s})+3 \mathrm{H}_{2}(\mathrm{~g}) \\$

$(3) \ \mathrm{Fe}(\mathrm{s})+4 \mathrm{H}_{2} \mathrm{O}(\mathrm{g}) \rightarrow \mathrm{Fe}_{3} \mathrm{O}_{4}(\mathrm{~s})+4 \mathrm{H}_{2}(\mathrm{~g})$

  •  Metals like lead, copper, silver and gold do not react with water.


Reaction of Metals with Acid

The reaction of acid with metals is discussed below:

  • Most metals react with acids and form metal salt and hydrogen gas.

Metal + Dilute acid $\rightarrow$ Salt + Hydrogen

  • HCl reaction with metals is discussed here:

For example, magnesium, aluminium, zinc and iron react with dilute hydrochloric acid.

$2 \mathrm{Mg}+2 \mathrm{HCl} \rightarrow 2 \mathrm{MgCl}+\mathrm{H}_{2}$

Iron reacts with hydrochloric acid and sulphuric acid when they are heated. But copper does not react with HCl even on heating. Copper reacts with sulphuric acid on heating. 

  • Nitric acid reaction with metals: Most of the time, hydrogen gas is not evolved when a metal reacts with nitric acid. It is because HNO3 is a strong oxidising agent which oxidises H2 produced in water and itself and reduces to any of the nitrogen oxides (N2O, NO, NO2). 

$\mathrm{Zn}+4 \mathrm{HNO}_{3} \rightarrow \mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2}+2 \mathrm{NO}_{2}+2 \mathrm{H}_{2} \mathrm{O}$

  • Dilute nitric acid reactions with metals like magnesium (Mg) and manganese (Mn)  evolve hydrogen gas.

$\mathrm{Mg}+2 \mathrm{HNO}_{3} \rightarrow \mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}+2 \mathrm{H}_{2}$

 

Reaction of Metals with Base

Some metals react with base to give corresponding salt and hydrogen gas.

Metal + Base $\rightarrow$ Salt + Hydrogen gas

  • Zinc reacts with NaOH and produces hydrogen gas and sodium zincate. 

$2 \mathrm{NaOH}+\mathrm{Zn} \rightarrow \mathrm{Na}_{2} \mathrm{ZnO}_{2}+\mathrm{H}_{2}$

  • Aluminium reacts with  NaOH and produces sodium aluminate and hydrogen gas.

$\mathrm{Al}+2 \mathrm{NaOH} \rightarrow \mathrm{Na}_{2} \mathrm{AlO}_{2}+\mathrm{H}_{2}$


Reaction of Metals with Non-metals

When metals react with non-metals, metal atoms lose electrons to the non-metal atoms, forming ions. The resulting product is called an ‘ionic compound’. Metal atoms lose electrons to form positive ions called ‘cations’ and non-metal atoms gain electrons to form negative ions called ‘anions’.

Some reactions between metals and nonmetals are:

  • $2 Na+Cl_{2} \rightarrow 2 NaCl$

Here,

$2Na\rightarrow 2Na^+ + 2e^-$

$Cl_2 + 2e^-\rightarrow 2Cl^-$ 

And finally, anions and cations react to form NaCl.

$2Na^+ + 2Cl^-\rightarrow 2NaCl$

  • $Zn + S \rightarrow Zn^{2+} + S^{2-} \rightarrow ZnS$

  • $2Mg + Cl_2 \rightarrow 2Mg^{2+} + 2Cl^- \rightarrow MgCl_2$

Ionic compounds have strong electrostatic forces between oppositely charged ions. Hence, they have high melting points. Ionic compounds have free electrons in their crystal lattice which allow them to conduct electricity when molten or in an aqueous solution.

 

Reaction of Metals with Solution of Other Metal Salts

Based on the position in the reactivity series, metals undergo a reaction with a metal salt solution. Metals that have high reactivity, which means it is in a higher position in the reactivity series can displace less reactive metal from its salt in its solution form or molten form.

  • For example, zinc reacts with copper sulphate solution and forms zinc sulphate and copper metal. Here, Zn is more reactive than Copper as it is higher in the reactivity series.

$\mathrm{Zn}(\mathrm{s})+\mathrm{CuSO}_{4}(\mathrm{aq}) \rightarrow \mathrm{ZnSO}_{4}(\mathrm{aq})+\mathrm{Cu}(\mathrm{s})$

  • Copper reacts with silver nitrate solution and forms copper nitrate and silver. Here, the more reactive copper displaces silver from the solution of silver nitrate.

$\mathrm{Cu}(\mathrm{s})+2 \mathrm{AgNO}_{3}(\mathrm{aq}) \rightarrow \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})+2 \mathrm{Ag}(\mathrm{s})$

  • When silver metal is kept in the copper sulphate solution, no reaction takes place because silver is less reactive than copper.

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Reactivity Series of Some Common Metals

 

Conclusion

Elements are mainly divided into two groups: metals and non-metals. Most of the metals are lustrous, malleable, ductile and have high conductivity for heat and electricity. Some metals can form positive ions by losing electrons to nonmetals and they form ionic compounds. Most of the metals combine with oxygen to form basic oxides but aluminium oxide and zinc oxide are amphoteric. Based on reactivity, different metals react with water and dilute acids and form salt and hydrogen. In the reactivity series based on the reactivity, metals are arranged in order of their decreasing reactivity. A more reactive metal that is higher in the reactivity series displaces a less reactive metal from its salt solution.

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FAQs on Reactions with Metals - JEE Important Topic

1. What is an Electrochemical Cell?

An electrochemical cell is a device that generates electrical energy from chemical reactions or uses electrical energy to speed chemical reactions. The voltaic or galvanic cells are electrochemical cells that generate an electric current. Electrolytic cells generate chemical reactions via electrolysis. In an electrolytic cell, a non-spontaneous redox reaction happens by the application of electrical energy. For example, electroplating. In a galvanic cell, electrical energy is produced from spontaneous redox reactions taking place within the cell. It generally consists of two different metals connected by a salt bridge or individual half-cells separated by a porous membrane in an electrolytic solution. When electrons shift from species to species through a spontaneous redox reaction, energy is released.

2. What is corrosion?

Corrosion is a natural process that happens due to electrochemical reactions between materials and substances in their environment. Corrosion is the gradual destruction of pure metals by the action of air, moisture or a chemical on their surface. For example, rusting of iron. Most metals can be easily oxidised as they tend to lose electrons to oxygen in air or water. The oxygen gets reduced and forms an oxide layer on the metal surface. When reduction and oxidation take place on different kinds of metal in contact with one another, the process is called ‘galvanic corrosion’. Some materials gain natural passivity or resistance to corrosion. 

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