Metals and Non-Metals Notes

Metals and Non-Metals Notes
Metals and Non-Metals Notes

Metals and Non-Metals:- Those elements, which conduct heat and electricity, are called metals, metals are those elements which give electrons to form cations. 3- It is solid at room temperature, but mercury (Hg) is in liquid state at room temperature.

Example- Aluminum (Al), Copper (Cu), Iron (Fe) etc.

Metals have the largest number of elements.

Physical Properties of Metals

1. Metallic luster Metallic lustre

The surface of pure metals is shiny due to the presence of free electrons. The surface of metals can also be made shiny by coating or polishing them. Example – copper (Cu), gold (Au), silver (Ag), etc.

Due to their metallic luster, they are used in making jewelery and decoration items.

(2) Thermal Conductivity

Metals are good conductors of heat. Silver and copper are the best conductors of heat but lead and mercury are poor conductors of heat.

 (3) Electrical conductivity

Metals are good conductors of electricity. They easily allow electricity to flow through them because free electrons are present in them.

Silver (Ag) and copper (Cu) are the best conductors of electricity.


Some metals can be beaten to form a thin sheet of metal. The property of turning metals into sheets (without breaking) is called malleability of the metal.

Example- Silver (Ag), gold (Au), aluminum (Al) and copper (Cu) are highly malleable metals.

5.Tensile Ductility

Some metals can be pulled into thin wires. This property of metals is called ductility.

Example – A fine wire 2 km long can be drawn from 1 gram of gold. Silver metal can be drawn into fine wire.

Almost invisible (very thin) wires can be made of tungsten (W). Tungsten wires are used to make filaments of electric bulbs.

6. Hardness Hardness Metals

Are generally hard. Example- aluminum (A), copper (Cu), iron (Fe), etc.

Sodium (Na) and potassium (K) are soft metals, they can be cut with a knife.

7. Melting Point Melting Point

The temperature at which a metal changes from its solid state to liquid is called the melting point of that metal. Metals have very high melting points.

Example – The melting point of copper (Cu) metal is 1058°C.

8. Brittleness Metals

are generally hard, due to which the property of brittleness is found in them. For example – Zic ( Zn ). They do not break by hitting.

9. Sound Sound

A sound is produced when metals are beaten with a hammer, which is also called metallic sound.

Example- The bell in schools is made of metal, which when struck with a wooden or metal hammer, emits a loud sound.

Chemical Properties of Metals

1.Reaction of Metals with Air Reaction of Metals with Oxygen

On combustion of metals in air, metals react with oxygen to form metal oxides.

Metal + Oxygen → Metal oxide

a) When copper is heated in air, copper reacts with oxygen to form black colored copper oxide.



Copper. Copper oxide

b) Some metal oxides dissolve in water to form alkali. Like

sodium oxide and potassium oxide Na2O+




sodium and potassium metals They react rapidly and catch fire when kept in the open, so they are kept immersed in kerosene oil to prevent them from fire. ,

c) amphoteric oxide;

Combustion of aluminum in the presence of air forms aluminum oxide (Al2O3). Which is amphoteric in nature. Because these metal oxides react with both base and acid to give salt and water


2.Reaction of Metals with Water

Metals react with water to form hydrogen gas and metal oxides or metal hydroxides. Metal oxides that dissolve in water, dissolve in water to form metal hydroxides.

Metal+Water—–>Metal Oxide+Hydrogen

Metal Oxide+Water—->Metal Oxide


(a) Sodium (Na) and Potassium (K) metals react rapidly with cold water (H2O) This process is

so fast that the hydrogen gas released by the released thermal energy starts burning.

 2Na + 2H2O——> 2NaOH + H2+heat energy

(b) The reaction of calcium (Ca) metal with cold water is slightly slower. The heat released in this reaction is not enough to burn the hydrogen gas.

 Ca + 2H20 —–>Ca ( OH ) 2 + H2

In the reaction, bubbles of hydrogen gas stick to the surface of calcium (Ca) metal, due to which it floats on the surface of water.

(c) Metals like iron (Fe), zinc (Zn), aluminum (AI), magnesium (Mg), etc., react with steam to form their oxides and hydrogen gas (H2).

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

Reaction of Metals with Acids Reaction of Metals with Dilute Acids

All metals react with dilute acids to form hydrogen gas and metal salts.

Metal + Dilute Acid → Metal Salt + Hydrogen


(a) Reaction of Metals with Dilute Hydrochloric Acid (HCI)

*When metals react with dilute hydrochloric acid, then metal chloride and hydrogen gas are made.

  Metal + HCl——-> Metal chloride + H2


  2Na + 2HCl —–> 2NaCl + H2

 Due to the low reactivity of copper (Cu), silver (Ag) and gold (Au), these metals do not react with dilute acids.

(b) Reaction of Metals with Very Dilute Nitric Acid (HNO3)*

When the metals magnesium(Mg) and manganese(Mn) react with nitric acid, the metal nitrates and hydrogen gas are formed.


Mg + 2HNO3——->Mg( NO3) 2+ H2


Amalraj or aqua-regia is a mixture of concentrated hydrochloric acid (HCI) and concentrated nitric acid (HNO3), amarose (Aqua- regia), it is strongly corrosive along with being a fiery liquid. Has the ability to smelt gold (Au) and platinum (Pt).

4.Reaction of metals with salt solutions A

more reactive metal from a solution of its compound with less reactive metal. displaces.


When a Zn rod is added to a solution of iron sulphate FeSO4, the zinc metal displaces iron from the iron sulphate solution to form zinc sulphate (ZnSO4).


Series of Metals Reactivity series of Metals. The activity series is a list in which the reactivity of metals is arranged in descending order. Those metals, which are located above hydrogen, are called active metals (K, Na, Ca) and those metals which are located below hydrogen are called inert metals.

Most reactive

K Potassium

Na Sodium

Ca Calcium

Mg Magnesium

Al Aluminum

Zn Zinc

Fe Iron

Pb Lead

H Hydrogen

Cu Copper

Hg Mercury

Ag Silver

Au Gold


Hydrogen also has non-metallic properties, but due to its electropositive, it has been kept in the activity category.


Elements that do not conduct heat and electricity are called non-metals.

Example:-  Carbon(C), Nitrogen(N), Sulfur(S)

Physical properties of non-metals Physical properties of Non – Metals

Following are the physical properties of non-metals

 (1) Metallic

 luster (Dull) is there. Example – No luster is found in non-metals like Sulfur (S).

Note – Iodine and graphite would have been shiny.

(2) Electrical and Thermal Insulators

Conductivity Non- metals are poor conductors of electricity and heat.

Example:- Phosphorus (P) and Sulfur (S) are both non-metals. Electricity and heat conduction is not done by both.

Note – Carbon is a non-metal. Iron is found in many forms. Which is called allotrope. Graphite is a conductor of electricity.

(3) Brittleness

Solid non-metals do not change into sheets when they are beaten, because they are brittle, that is, they break into small pieces when they are beaten with a hammer.

Example – Sulfur (S) and Phosphorous (P) non-metals are brittle.

(4) Elasticity Non-Ductility Non- metal

metals cannot be converted into wires by stretching, they lack ductility. Due to their brittleness, they break when pulled.

(5) Hardness

Most of the solid non-metals such as Sulfur (S) and Phosphorus (P) are soft while some are hard

(6) Melting point

Non-metals have low melting points.

For example, the melting points of gallium and cesium metals are so low that they start melting as soon as they are placed on the palm.

(7) Sound Sound

Non-metals are non-acoustic, that is, no sound is produced when they are beaten with a hammer. Reaction

between metals and non-metals Excellent gases are inert. Because his ashtak is fulfilled. Other elements also try to complete 8 electrons in their outermost shell, for this they gain or give up electrons.

Metals tend to give electrons while non-metals tend to take electrons.

When metals and non-metals react with each other and both try to complete their octet by exchange of electrons.

Example – Sodium – Let us try the electronic configuration of sodium (which is a metal).



To complete its octet (ie 8 electrons in the outermost shell) sodium gives 1 electron instead of taking 7 electrons in its M shell.

Thus it acquires a positive charge (+) by giving up one electron from the M shell. Now it has 8 electrons in its L shell and it becomes the outermost shell. Positively charged sodium is called cation or sodium ion.

It is displayed as follows.

 Na +( configuration 2,8 )

 Na―>[Na + ]+ [e-]

 2,8,1 2,8 Similarly

if the electronic configuration of chlorine is 2,8 , 7, which is a nonmetal. It is easier for chlorine to gain one electron in its M shell than it is to give 7 electrons.

Thus it acquires an electron in its M shell and becomes negatively charged, now it is called chloride ion or anion, it has Cl configuration from KLM 2,8,8.

Sodium and chlorine are attracted to each other and by forming strong bonds, sodium chloride remains in the form of NaCl.


Bond or electrovalent bond

Compounds formed by the exchange of electrons from metal to non-metal are called ionic compounds or electrovalent compounds.

Properties of ionic compounds

(characteristics of Electrovalent or ionic compounds) Have you ever noticed the following general properties of ionic compounds :

1. Physical nature :(crystalline nature) Ionic compounds are solid and somewhat hard because of the strong force of attraction between positive and negative ions. Ionic compounds are often brittle and break into fragments under pressure.

2. Melting and Boiling Points: Ionic compounds have very high melting points and boiling points because a large amount of energy is required to break the strong intermolecular attraction.

3. Solubility: Electrovalent compounds are soluble in water and insoluble in solvents like kerosene, petrol etc.

4. Electrical conductivity: Ions are present in aqueous solutions of ionic compounds. When an electric current is passed through the solution, these ions start moving towards the opposite electrode. Ionic compounds cannot conduct electricity in the solid state because the movement of ions is not possible in the solid state due to the rigid structure. But ionic compounds conduct electricity in the molten state because the electrostatic force of attraction between oppositely charged ions in the molten state is weakened by the effect of heat. Hence ions can move freely and hence conduct electricity.

Occurrence of metals

Earth’s crust is the main source of metals.

Some soluble salts like sodium chloride, magnesium chloride etc. are also present in sea water.

Minerals:- The elements or compounds that occur naturally in the earth’s crust are called minerals.

Ores:- At some places, minerals contain a significant amount of a particular metal, which is beneficial to extract. These minerals are called ores.

Extraction of Metals

The extraction of pure metals from metal compounds is called extraction of metals.

Some metals are found in free state in the earth’s crust and some metals are found in the form of their compounds.

The metals that fall below the reactivity series are the least reactive. Hence, they are found in free state. Example – Gold (Gold), Silver (Silver), Platinum and Copper etc.

Copper and silver are also found in the combined state as ores of their sulfides or oxides.

The metals at the top of the activity series (K, Na, Ca, Mg and AI) are not found in the free state due to being highly

reactive, the reactivity of the metals in the middle of the activity series (Zn, Fe, Pb etc.) is moderate. They are found in the form of oxides, sulfides or carbonates.

The ores of many metals are oxides. The reason for this is the high reactivity of oxygen and its abundance on Earth. The process of extraction of elements from their ores is called metallurgy. Thus, on the basis of reactivity we can classify metals into following three classes

(1) Low reactive metals

(2) Medium reactive metals

(3) High reactive metals

Steps Involved in the Extraction of Metal

Extraction of pure metal from ore takes place in several stages.

1.Enrichment of Ores

The ores derived from the earth contain many impurities like clay, sand, which are called gangue. It is very important to remove the impurities in the ore before the extraction of metals. The chemical processes used to remove gangue from ores are based on the physical or chemical properties of the ore and gangue. Several techniques are used for this separation.

A) Extraction of metals falling in the lower activity series

Being less reactive, metals can be obtained only on heating.

For example,

1. Cinnabar (Hgs) is an ore of mercury. It first forms mercuric oxide at hot spring in the presence of air and on further heating, mercuricoxide turns into mercury.

3HgS+3O2——->2HgO+2SO2 2HgO——->


Similarly, copper present in Cu2S can be obtained by heating its ore in air.



B) Extraction of metals located in the middle of the activity series

In nature, these metals are often found in the form of sulfides or carbonates. It is easier to obtain metals from their oxides than from sulfides or carbonates.

Sulfides are converted into oxides by roasting and carbonates to oxides by calcination. After this the metal is obtained from the metal oxide using a suitable reducing agent such as carbon.

The following chemical reactions take place during the extraction of zinc

(a) Roasting:- The sulphide ore is converted into oxide on heating in the presence of air at a temperature below its melting point. This process is called roasting. Roasting Example –

2ZnS + 3O2 → 2ZnO+ 2SO2

(b) Calcination  The process of converting a concentrated ore into an oxide by heating it below its melting point but at a higher temperature in the absence of air is called calcination. Example-

ZnCO3——>ZnO + CO2

(c) Reduction of Oxide Ore  It is the process of converting metal oxide ore into metal. This can be done by heating the oxide.

*ZnO + C → Zn + CO

Sometimes displacement reactions are used in the reduction of metal oxides. Highly reactive metals such as sodium, calcium, aluminum are used as reducing agents, because they displace metals with low reactivity from their compounds.

Example – Reaction of manganese dioxide with aluminum powder

3MnO2+ 4Al —> 3Mn + 2Al2O3 + Heat

These displacement reactions are highly exothermic, so the amount of heat generated is high, which is produced in the liquefied state of metals. The iron produced by the reaction of aluminum with iron (III) oxide (FeO) is used to join cracks in railway tracks or machine parts. The formation of a metal by the reaction of metal oxide with the use of heat as a reducing agent of aluminum powder is called thermit reaction.

Fe2O3 + 2AI→ 2Fe + AI2O3+ heat

(iii) Extraction of metals at the top of the reactivity series (highly reactive)

These metallic compounds cannot be reduced due to high affinity with carbon or other reducing agents, so electrolytic reduction experiments are done for these metals, such as Na. , Mg, Ca etc.

Refining of Metals

It is the process of purification of metals obtained after reduction.

Many methods are used for refining, but the most common is electrolytic refining.

Many metals like copper, tin, nickel, silver, gold etc. are refined by electrolysis.

Electrolytic Refining: 

In this process, the impure metal is made as the anode and a thin layer of pure metal is made as the cathode. The salt solution of the metal is used as an electrolyte.

When a current is passed through an electrolyte, the impure metal at the anode dissolves in the electrolyte. The soluble impurities go into the solution and the insoluble impurities get deposited at the bottom of the anode which is called anode mud.

Corrosion of Metals

The corrosion of a surface metal when it comes in contact with moist air is called corrosion. It is an oxidising process.


1. Rusting of iron

2. Discoloration of silver

3. Formation of

green coating on copper and bronze are examples of corrosion.

Prevention of Corrosion

Iron can be prevented from rusting by painting, oiling, greasing, lacquering, chromium coating, anodizing or alloying. Some methods are described below

(1) Galvanization

The process of plating a thin layer of zinc on iron and steel objects is called slag.

This is done by immersing the object in liquefied zinc in the process.

(2) Turning metals into alloys is a method of improving the properties of metals. In which two or more metals are mixed. For example, brass is made from a mixture of copper and copper.

(3) By painting the surface of the metal with an acid-resistant paint, the metal is protected from the effects of air or any solution.

(4) When grease or oil is applied to the surface of an iron object, moisture cannot come in contact with it, thereby protecting it from corrosion.

For example, iron parts and machines are smeared with grease. Various alloys are made by mixing iron with many metals.


A homogeneous mixture of two or more metals is called an alloy.

To make an alloy, the parent metal is first brought to a molten state and after that other elements are mixed in a certain proportion. Then it is cooled to room temperature.

If any one metal in the alloy is mercury, then the alloy is called amalgam. The electrical conductivity and melting point of its alloy are lower than that of pure metals.

Example The name of an alloy of lead (Pb) and tin (Sn) is solder, which has a very low melting point. Solder is used for welding electrical wires together.

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