Metals and Non-Metals Class 10 Notes: CBSE Science Chapter 3

The periodic table contains $118$  elements, $92$  of which are naturally occurring. The physical and chemical properties of metals and non-metals are vastly different. Currently, we are aware of roughly $80$  metals.

Except for bromine, which is a liquid, over half of the non-metals are gases at ambient temperature. The most abundant non-metal in the earth's crust is oxygen, which makes up about half of the crust and is one of the main elements of air, along with nitrogen.

Silicon is the next most prevalent non – metal, accounting for around $26\% $  of the earth's crust. Earth's two major elements are oxygen and silicon. The two principal ingredients of the oceans are hydrogen and oxygen.

Positions of Metals and Non – Metals in the Periodic Table

(Image will be uploaded soon)

Metals are found in the periodic table's leftmost groupings. Group \[{\text{I A}}\]  elements are alkali metals, which are highly reactive metals, and group ${\text{II A}}$  elements are alkaline earth metals. Transition metals are elements that belong to groups ${\text{II A}}$ and ${\text{III A}}$ .

Non-metals are elements that are found to the right of the Periodic Table, in groups ${\text{IV A}}$ , ${\text{V A}}$ , ${\text{VI A}}$ , and ${\text{VII A}}$ , with the exception of hydrogen. From the top to the bottom of the group, these elements get more non-metallic. The first and second members of group ${\text{V A}}$ , for example, are non-metals, the third and fourth members are metalloids, and the last member is metal. Metalloids are a class of elements that exhibit properties that are akin to both metals and nonmetals. Boron, silicon, germanium, arsenic, antimony, tellurium, and astatine are the metalloids in question. The non-metals, which include the element hydrogen, are found to the right of these metalloids.

Group ${\text{V A}}$:

Physical Properties of Metals

Physical State - With the exception of mercury and gallium, which are liquids at room temperature, metals are solids at room temperature.

Lustre – The property of metals which makes the light reflect from their surfaces is called lustre. This property of the metals can be due to the polished metal surfaces. Eg., gold and silver.

Malleability - Metals may be formed into thin sheets known as foils and can withstand hammering. With the exception of Zinc, which is fragile. 

Ductility - Wires can be made out of metals. With the exception of Zinc, which is fragile.

Hardness - Except for sodium and potassium, which are soft and can be cut with a knife, all metals are hard. 

Conduction - Because metals have free electrons, they are good conductors. Silver and copper are the best heat and electricity conductors. Lead is the least efficient heat conductor. Iron, bismuth, and mercury are likewise poor conductors.

Density - Metals have a high density and weigh a lot. The densities of iridium and osmium are the greatest, whereas lithium has the lowest density. 

Melting and Boiling Point - Metals are known for their high melting and boiling points. The melting point of tungsten is the highest, while the boiling point of silver is the lowest. The melting values of sodium and potassium are both low. 

Alloy Formation - Metals combine to create an alloy, which is a homogeneous combination of metals. Brass is a copper and zinc alloy.

Sonorous – Metals, when hit by a solid object, produce sound. This property of a metal is known as sonorous.

Physical Properties of Non – Metals 

Physical State - At ambient temperature, the majority of non-metals exist in two of the three states of matter: gases (oxygen) and solids (iodine, carbon, sulphur). There is no metallic sheen to them (save iodine) and they do not reflect light. (With the exception of carbon in the form of a diamond.)

Nature - Non – metals are extremely fragile, and they can't be coiled into wires or hammered into sheets. Except for diamond, which is the world's hardest substance. 

Conduction - Non – metals are poor heat and electrical conductors. (Except graphite conducts heat, both graphite and gas carbon conduct electricity.)

Electronegative Character - Non – metals have a proclivity for gaining or sharing electrons with neighbouring atoms. Hence, non – metals are known for their electronegative nature.

Reactivity - When they come into contact with oxygen, they produce acidic or neutral oxides. Hence, non – metals are reactive.

Melting and Boiling Points - Non – metals are known for their low melting and boiling points.

Comparative Properties of Metals and Non – Metals 

Chemical Properties of Metals

Metals are Electropositive Elements

Metals have a high reactivity. Metals are known as electropositive elements because they readily lose electrons and create positively charged ions. The metal sodium produces sodium ions ${\text{(N}}{{\text{a}}^{\text{ + }}}{\text{)}}$. Metals' electropositive nature makes it easy for them to form compounds with other elements.

Reaction of Metals with Oxygen

Some of the most reactive metals include sodium ${\text{(Na)}}$  and potassium ${\text{(K)}}$. In the presence of oxygen, potassium, sodium, lithium, calcium, and magnesium react and burn. Metals in the activity series of metals, from aluminium to copper, react slowly in air to generate metal oxides. Aluminum is the quickest, whereas copper is the slowest.

• At normal temperature, sodium metal interacts with oxygen in the air to generate sodium oxide. As a result, sodium is kept in kerosene to avoid reacting with oxygen, moisture, or carbon dioxide.

$4{\text{Na}} + {{\text{O}}_{\text{2}}} \to 2{\text{N}}{{\text{a}}_{\text{2}}}{\text{O}}$ 

• Sodium oxide is a basic oxide that forms sodium hydroxide when it combines with water.

${\text{N}}{{\text{a}}_{\text{2}}}{\text{O}} + {{\text{H}}_{\text{2}}}{\text{O}} \to 2{\text{NaOH}}$ 

• At room temperature, magnesium does not react with oxygen. Magnesium burns in air with intense light and heat to create magnesium oxide ${\text{(MgO)}}$  when heated.

${\text{2Mg}} + {{\text{O}}_{\text{2}}}\xrightarrow{\Delta }2{\text{MgO}}$ 

• Zinc when burned in air results in the formation of zinc oxide.

$2{\text{Zn}} + {{\text{O}}_{\text{2}}}\xrightarrow{\Delta }{\text{2ZnO}}$ 

• Iron gets oxidized and forms rust in the presence of moist air.

${\text{3Fe}} + {\text{2}}{{\text{O}}_{\text{2}}} + {\text{x}}{{\text{H}}_{\text{2}}}{\text{O}} \to {\text{F}}{{\text{e}}_{\text{3}}}{{\text{O}}_{\text{4}}}{\text{.x}}{{\text{H}}_{\text{2}}}{\text{O}}$ 

• When iron is heated in the presence of air (oxygen) the formation of triferric tetroxide takes place.

$3{\text{Fe}} + {\text{2}}{{\text{O}}_{\text{2}}}\xrightarrow{\Delta }{\text{F}}{{\text{e}}_{\text{3}}}{{\text{O}}_{\text{4}}}$ 

• Copper is the least reactive metal and, even when heated, does not burn in the air. Copper, on the other hand, reacts with oxygen and generates copper (II) oxide ${\text{(CuO)}}$ on the surface and copper (I) oxide ${\text{(C}}{{\text{u}}_{\text{2}}}{\text{O)}}$  on the inside when heated for a long time.

$2{\text{Cu}} + {{\text{O}}_{\text{2}}}\xrightarrow{\Delta }2{\text{CuO}}$ 

• When mixed with oxygen present in the air, gold and platinum do not react with it.

Reaction of Metal with Water

The metals that react with cold water are: potassium, sodium, lithium and calcium.

• Sodium when mixed with water is a highly vigorous reaction and forms sodium hydroxide and water.

$2{\text{Na}} + 2{{\text{H}}_{\text{2}}}{\text{O}} \to 2{\text{NaOH}} + {{\text{H}}_{\text{2}}}$ 

• Metals in the activity series, from magnesium to iron, react with steam (but not cold water) to generate metal oxide and hydrogen gas.

${\text{Mg}} + {{\text{H}}_{\text{2}}}{\text{O}} \to {\text{MgO}} + {{\text{H}}_{\text{2}}}$ 

${\text{2Al}} + 3{{\text{H}}_{\text{2}}}{\text{O}} \to {\text{A}}{{\text{l}}_{\text{2}}}{{\text{O}}_{\text{3}}} + 3{{\text{H}}_{\text{2}}}$ 

${\text{Zn}} + {{\text{H}}_{\text{2}}}{\text{O}} \to {\text{ZnO}} + {{\text{H}}_{\text{2}}}$ 

• When iron reacts with steam, it forms iron (I, II) oxide.

${\text{3Fe}} + 4{{\text{H}}_{\text{2}}}{\text{O}} \rightleftharpoons {\text{F}}{{\text{e}}_{\text{3}}}{{\text{O}}_{\text{4}}} + 4{{\text{H}}_2}$ 

${\text{2Fe}} + 3{{\text{H}}_{\text{2}}}{\text{O}} \rightleftharpoons {\text{F}}{{\text{e}}_{\text{2}}}{{\text{O}}_3} + {\text{3}}{{\text{H}}_{\text{2}}}$ 

Note: The reaction between iron and steam is irreversible. Tin, lead, copper, silver, gold and platinum do not react with water or steam.

Reaction of Metals with Acids

• The metal salt (either sulphate or chloride) and hydrogen gas are formed when potassium, sodium, lithium, and calcium react forcefully with dilute ${{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}$  and dilute ${\text{HCl}}$. The reaction is identical to the one that occurs when water is used.

${\text{2Na}} + 2{\text{HC}}{{\text{l}}_{{\text{(dil)}}}} \to 2{\text{NaCl}} + {{\text{H}}_{\text{2}}}$ 

• With dilute acid, magnesium, aluminium, zinc, iron, tin, and lead react safely. Magnesium is the fastest of the six elements, while lead is the slowest.

${\text{Mg}} + {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{MgS}}{{\text{O}}_{\text{4}}} + {{\text{H}}_{\text{2}}}$ 

${\text{Mg}} + 2{\text{HCl}} \to {\text{MgC}}{{\text{l}}_2} + {{\text{H}}_{\text{2}}}$ 

${\text{Zn}} + {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_4} \to {\text{ZnS}}{{\text{O}}_4} + {{\text{H}}_{\text{2}}}$ 

${\text{Fe}} + 2{\text{HCl}} \to {\text{FeC}}{{\text{l}}_2} + {{\text{H}}_{\text{2}}}$ 

For the laboratory synthesis of hydrogen, zinc with dilute sulphuric acid is commonly utilised. At room temperature, the reaction is sluggish, but it can be sped up by adding a little copper (II) sulphate. Copper metal, which works as a catalyst, is displaced by zinc.

Copper, silver, gold, and platinum are metals that do not react with dilute acid to liberate hydrogen. 

• Hydrochloric acid produces metal chlorides in general. 

• Sulphuric acid reacts with metals to form sulphates. 

• Nitric acid reactions are more complicated; nitrate is created, but the gas produced is rarely hydrogen and more frequently an oxide of nitrogen.

Reactions of Metals with Salt Solutions

Reactive metals can displace any metal less reactive than themselves in solution or in their molten state, such as the less reactive metal's oxide, chloride, or sulphate. Metal A is more reactive than B if it displaces metal B from its solution.

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

${\text{Fe}} + {\text{CuS}}{{\text{O}}_{\text{4}}} \to {\text{FeS}}{{\text{O}}_{\text{4}}} + {\text{Cu}}$ 

When dilute, copper (II) sulphate solution appears blue, while iron sulphate solution is practically colourless. The blue solution loses its colour during displacement, and the iron metal turns pink-brown as the displaced copper is deposited on it.

When magnesium powder and black copper (II) oxide are heated together, white magnesium oxide is created with brown copper bits:

${\text{Mg}} + {\text{CuO}}\xrightarrow{\Delta }{\text{MgO}} + {\text{Cu}}$ 

The blue colour disappears as colourless magnesium sulphate forms and brown particles of copper metal form a precipitate when magnesium is added to a blue copper (II) sulphate solution:

${\text{Mg}} + {\text{CuS}}{{\text{O}}_4} \to {\text{MgS}}{{\text{O}}_4} + {\text{Cu}}$ 

(Image will be uploaded soon)

Electronic Nature of Metals and Non – Metals 

With the exception of noble gases, all atoms have an incomplete outermost shell. Noble gases have a full outer shell and are hence non-reactive or "inert." 

The stability of noble or inert gases is achieved via electron transfer or electron sharing in most reactive components. Metals are elements that can donate electrons. They lose electrons and create positive ions as a result.

Non-metals are elements that can accept electrons. They gain electrons and generate negative ions as a result. The outermost shell of a metal has $1$  to $3$  electrons, while the outermost shell of a non-metal has $4$  to $8$  electrons.

Hydrogen and helium are the only two exceptions to this rule. Hydrogen is a non-metal with one electron in its valence shell, while helium has two electrons in its valence shell.

Important Points

Donor metals get a positive charge equal to the number of electrons they donate. Aluminum, for example, has an atomic number of 13, which means its electrical configuration is 2, 8, and 3. The valence shell of aluminium has three electrons; it loses three electrons to create ${\text{A}}{{\text{l}}^{{\text{3 + }}}}$.

${\text{A}}{{\text{l}}^0} - 3{{\text{e}}^ - } \to {\text{A}}{{\text{l}}^{{\text{3 + }}}}$ 

Other examples are given below:

${\text{N}}{{\text{a}}^0} - 1{{\text{e}}^ - } \to {\text{N}}{{\text{a}}^ + }$ 

${\text{M}}{{\text{g}}^0} - 2{{\text{e}}^ - } \to {\text{M}}{{\text{g}}^{2 + }}$ 

Non – metals gain electrons and, as a result, gain a negative charge equal to the number of electrons accepted.

The examples are as follows:

${\text{C}}{{\text{l}}^0} + {\text{1}}{{\text{e}}^ - } \to {\text{C}}{{\text{l}}^ - }$ 

${{\text{S}}^0} + 2{{\text{e}}^ - } \to {{\text{S}}^{2 + }}$ 

${{\text{P}}^0} + 3{{\text{e}}^ - } \to {{\text{P}}^{3 + }}$ 

The Reactivity Series of Metals

Despite the fact that most metals are electropositive in nature and lose electrons in chemical reactions, they do not respond with the same vigour or speed. Different metals react differently to different chemicals. The stronger an element's reactivity, the easier it is for it to lose its electrons and get a positive charge. Furthermore, the higher the number of shells and the lower the number of valence electrons, the higher the metal's reactivity. The metal activity series puts all metals in decreasing chemical activity order. The ease with which a metal loses electrons and generates positive ions in solutions reduces as we progress down the activity series from potassium to gold.

Potassium, the most active metal, is at the top of the list, while gold, the least reactive metal, is at the bottom. Despite the fact that hydrogen is a non-metal, it is included in the activity series because it behaves like a metal in most chemical reactions, i.e., the hydrogen ion has a positive charge [H+] similar to other metals.

(Image will be uploaded soon)

• The more reactive a metal is in the series, the faster and more exothermic its reaction will be.

• This also means that the reverse reaction becomes more complex, implying that the more reactive a metal is, the more difficult it is to remove it from its ore.  With oxygen and water, the metal is also more prone to corrosion.

• The reactivity series can be determined by watching metals react with water, oxygen, or acids. 

• There are certain periodic table trends within the overall reactivity or activity series:

Formation of Sodium Chloride

Because sodium has only one electron in its valence shell, it loses one electron to complete its octet during the creation of an ionic bond between the metal sodium and the non-metal chlorine. It takes on the form of a noble gas, neon ${\text{(2, 8)}}$ . The chlorine atom contains seven electrons in its valence shell and gains one to complete its octet, as well as acquiring stable electronic configurations.

${\text{Na}}\xrightarrow{{{\text{loss of electrons}}}}{\text{N}}{{\text{a}}^{\text{ + }}} + {{\text{e}}^ - }$ 

${\text{Cl}} + {{\text{e}}^ - }\xrightarrow{{{\text{gain of electron}}}}{\text{C}}{{\text{l}}^ - }$ 

Formation of Magnesium Chloride

Magnesium, with an atomic number of 12, has electronic configuration as ${\text{2, 8, 2}}$ . It has two electrons in its valence shell. Chlorine's electronic configuration (atomic number: 17) is ${\text{2, 8, 7}}$ . There are seven valence electrons in it. Because magnesium has two electrons more than the neon configuration ${\text{(2, 8)}}$  and chlorine has one electron less than the argon configuration ${\text{(2, 8, 8)}}$ , one atom of magnesium will seek out two atoms of chlorine to transfer its two electrons to (one to each), as seen below:

${\text{Mg}} \to {\text{M}}{{\text{g}}^{{\text{2 + }}}}{\text{ +  2}}{{\text{e}}^ - }$ 

${\text{2Cl  +  2}}{{\text{e}}^ - } \to {\text{2C}}{{\text{l}}^ - }$ 

The ion of magnesium ${\text{(M}}{{\text{g}}^{{\text{2 + }}}}{\text{)}}$ and the two ions of chlorine ${\text{(C}}{{\text{l}}^{\text{ - }}}{\text{)}}$form magnesium chloride which has an ionic bond between them.

 ${\text{M}}{{\text{g}}^{2 + }} + {\text{2C}}{{\text{l}}^ - } \to {\text{[C}}{{\text{l}}^ - } - {\text{M}}{{\text{g}}^{{\text{2 + }}}} - {\text{C}}{{\text{l}}^ - }] \to {\text{MgC}}{{\text{l}}_{\text{2}}}$ 

In the terms of Lewis dot structure

Properties of Electrovalent Compounds

Occurrence of Metals 

Minerals and Ores

Minerals are metals and their compounds that are present in the earth's crust. Ores are minerals from which metals can be profitably and easily mined. Metal complexes with a lower amount of impurities are found in ore. Therefore, it is clear that all ores are minerals, but not all minerals are ores.

In the Free State

Only a few metals exist in their natural condition. Only pure metals such as gold, platinum, and mercury are occasionally found in the Free State. Copper and silver are occasionally discovered in the Free State. Air and water have no effect on such metals.

In Combined State

The remaining metals are found as oxides, carbonates, sulphides, sulphates, silicates, chlorides, nitrates, phosphates, and other compounds. Copper and silver are two metals that may be found in both free and mixed forms in the form of sulphide, oxide, or halide ores. Because the metals at the top of the activity series \[{\text{(K, Na, Ca, Mg, Al)}}\]  are so reactive, they are never found as free elements in nature. The intermediate metals in the activity series \[\left( {{\text{Zn, Fe, Pb}}} \right)\] are moderately reactive. They are mainly found as oxides, sulphides, or carbonates in the earth's crust.

Extraction of Metals – Metallurgy

Metallurgy is the study of the different processes involved in extracting metals from their ores and purifying them.

(Image will be uploaded soon)

Concentration – Enrichment of Ores

Ore is an impure metal with a lot of sand and stony particles in it. The impurities in the ore, such as sand, stony minerals, limestone, mica, and so on, are referred to as gangue or matrix. Before the metal can be extracted, these contaminants must be eliminated from the ore. A fusible compound called slag is formed when a chemical called flux is introduced to the ore to remove the matrix.

Gangue + Flux = Slag

The procedures for separating the gangue from the ore are based on the physical or chemical differences between the gangue and the ore. The ore is first crushed to a powder. Physical techniques such as hydraulic washing, froth flotation, and magnetic separation, as well as chemical processes, are used to separate the crushed ore, depending on the type of the ore and its impurities. Ore concentration is sometimes known as ore 'dressing' or 'enrichment.'

Extracting Metals Low in the Activity Series

Metals with a low activity series are extremely non-reactive. By simply heating these metals' oxides, they may be reduced to metals. Mercury, for example, is extracted from its ore, cinnabar ${\text{(HgS)}}$ , by a heating process.

Reduction using heat

${\text{2HgS  +  3}}{{\text{O}}_2}\xrightarrow{{{\text{Heat}}}}{\text{2HgO  +  2S}}{{\text{O}}_{\text{2}}}$ 

${\text{2HgO}}\xrightarrow{{{\text{Heat}}}}2{\text{Hg  +  }}{{\text{O}}_{\text{2}}}$ 

By treating copper sulphide ${\text{(C}}{{\text{u}}_{\text{2}}}{\text{S)}}$ ore in the same way, pure copper can be obtained by the similar method as the above method.

Note: Decomposition of mercury and copper oxide formed takes place after the reaction is completed.

Extracting Metals in the Middle of the Activity Series

Iron, zinc, lead, copper, and other metals are in the centre of the reactivity series. Sulphides and carbonates are the most common forms of these moderately reactive metals.

Reduction techniques are used to extract these metals from their ores. The oxide ore is reduced during the reduction process.

When compared to sulphides and carbonates, an oxide ore is easier to reduce. If the ore is not an oxide ore, it must first be transformed to an oxide by calcination or roasting.

Roasting

Sulphide ores are transformed to oxides by heating them rapidly in the presence of abundant air, causing oxygen to be added, resulting in the formation of the appropriate oxides. Impurities of sulphur escape as gas. Roasting is the term for this procedure.

\[2{\text{ZnS (s)  +  3}}{{\text{O}}_{\text{2}}}{\text{(g)}} \to {\text{2ZnO (s)  +  2S}}{{\text{O}}_{\text{2}}}{\text{(g)}}\] 

${\text{2PbS (s)  +  3}}{{\text{O}}_{\text{2}}} \to {\text{2PbO (s)  +  2S}}{{\text{O}}_{\text{2}}}{\text{(g)}}$ 

${\text{4FeS (s)  +  7}}{{\text{O}}_{\text{2}}}{\text{(g)}} \to {\text{2F}}{{\text{e}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{(s)  +  4S}}{{\text{O}}_{\text{2}}}\left( {\text{g}} \right)$ 

Calcination

The ore is heated to a high temperature in the absence of air, or when air is not present throughout the reaction. Carbonate ores, as well as ores containing water, are usually calcined to remove carbonate and moisture impurities.

${\text{ZnC}}{{\text{O}}_{\text{3}}}{\text{(s)}} \to {\text{ZnO (s)  +  C}}{{\text{O}}_{\text{2}}}{\text{(g)}}$

${\text{A}}{{\text{l}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{.2}}{{\text{H}}_{\text{2}}}{\text{O (s)}} \to {\text{A}}{{\text{l}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{(s)  +  2}}{{\text{H}}_{\text{2}}}{\text{O (l)}}$ 

Differences between Roasting and Calcination

Reduction of Ores

Carbon or hydrogen is used to decrease the oxide produced by calcination or roasting. Coke, or carbon monoxide, is the most common form of carbon. Carbon or carbon monoxide cannot reduce all metallic oxides. Thus, the technique used to extract metal from its ore is determined by the metal's reactivity, or its location in the metal activity series.

Reduction Using Carbon (Coke)

For oxides of moderately reactive metals, this technique is utilised. Coke is the most commonly used reducing agent since it is affordable.

${\text{2ZnO  +  C}}\xrightarrow{\Delta }{\text{2Zn  +  C}}{{\text{O}}_2}$ 

${\text{2PbO  +  C}}\xrightarrow{\Delta }{\text{2Pb  +  C}}{{\text{O}}_{\text{2}}}$ 

Reduction Using Carbon Monoxide

Carbon monoxide is an extremely effective reducing agent. It is used in the blast furnace for hematite reduction.

${\text{F}}{{\text{e}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{ +  3CO}}\xrightarrow{\Delta }{\text{2Fe  +  3C}}{{\text{O}}_{\text{2}}}$ 

Extracting Metals towards the Top of the Activity Series

This method is used to make oxides of highly reactive metals that are higher in the reactivity series than aluminium. It can also be used to reduce aluminium oxide. Example:

${\text{2A}}{{\text{l}}_{\text{2}}}{{\text{O}}_{\text{3}}}\xrightarrow{{{\text{electricity}}}}{\text{4Al  +  3}}{{\text{O}}_{\text{2}}}$ 

The expense of keeping the electrolyte molten is quite expensive since alumina has a very high melting point. When alumina is combined with cryolite ${\text{(N}}{{\text{a}}_{\text{3}}}{\text{Al}}{{\text{F}}_{\text{6}}}{\text{)}}$  and fluorspar ${\text{(Ca}}{{\text{F}}_{\text{2}}}{\text{)}}$, the melting point is significantly reduced, and the cost is also reduced. The conductivity of the electrolyte is increased by these compounds.

Refining of Metals

The majority of metals acquired during the reduction process are not very pure. These must be processed or purified further. The final phase in metallurgy is metal purification. The difference between the characteristics of metals and their impurities is used in refining. For refining, we utilise the following procedure.

Electrorefining

Electrolysis can be used to recover metals that cannot be separated via a chemical reduction technique, as well as to purify metals acquired through other means. The anode in the electrorefining process is a block of impure metal, while the cathode is a thin sheet of pure metal. An aqueous solution of the metal salt is included in the electrolytic cell. When an electric current of a sufficient voltage is passed via the anode, impure metal is dissolved and pure metal is deposited at the cathode. The following is how metal ions from the anode enter the electrolyte:

${\text{M}} \to {{\text{M}}^{{\text{ + n}}}}{\text{ + n}}{{\text{e}}^ - }$ 

These ions get deposited on the cathode in the following manner

${{\text{M}}^{{\text{ + n}}}}{\text{ +  n}}{{\text{e}}^ - } \to {\text{M}}$ 

Near the anode, the contaminants are left as anode mud. Due to the accumulation of pure metal, the anode eventually disintegrates, while the cathode accumulates weight.

(Image will be uploaded soon)

This technique is used to refine volatile metals with lower boiling points than their impurities, such as copper, silver, tin, and nickel. For example: Mercury and Zinc.

• An electrolyte is a substance (salt, acid, or base) that transmits an electric current in solution or in a molten form while also being decomposed by it. The current is carried by ionised electrolytes, which are electrically charged ions.

• Charged ions migrate towards oppositely charged electrodes in order to lose their electric charge and form atoms, which are then either released or deposited at the electrodes.

Corrosion of Metals

We've learned that in the presence of a wet environment, chemically active metals corrode. Corrosion is an oxidation process involving ambient oxygen and water on a metal's surface. Iron corrodes faster than most other transition metals, resulting in the formation of iron oxide. In the presence of carbon dioxide, as well as salt solution, the corrosion or rusting of iron is increased.

Rusting is: ${\text{2Fe (s)  +  }}{{\text{O}}_{\text{2}}}{\text{(g)  +  }}{{\text{H}}_{\text{2}}}{\text{O (l}}) \to {\text{F}}{{\text{e}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{.x}}{{\text{H}}_{\text{2}}}{\text{O (s)}}$ 

Rust is also known as hydrated iron oxide.

Rusting is an oxidation reaction because it includes iron atoms losing electrons or iron atoms obtaining oxygen. The amount of water 'x' is variable, ranging from dry to wet, and the equation is not supposed to be balanced.

Corrosion is a serious issue in sectors that utilise iron (or steel) as a structural material, such as construction, infrastructure, bridges, rail transport power transmission, shipbuilding, cars, and heavy industries.

Aluminum, another valuable structural metal, undergoes an oxidation process as well, but not as fast as its reactivity indicates. Once a small coating of ${\text{A}}{{\text{l}}_{\text{2}}}{{\text{O}}_{\text{3}}}$  has developed on the surface, it acts as a barrier to oxygen and water, preventing the aluminium from corroding further. As a result, aluminium is known as a self-protective metal. Aluminum alloys may be used to make it harder, stronger, and stiffer by combining it with tiny quantities of other metals (such as magnesium).

Alkali metals, such as sodium, which are used in chemistry laboratories and some chemical industries, corrode quickly in the air and must be kept under oil.

Copper and lead are both utilised in roofing because they are both non-reactive. The surface chemicals do not peel away as readily as rust does off iron. Copper corrodes to create a basic green carbonate (combination of the hydroxide and the carbonate), while lead corrodes to form a white lead oxide or carbonate. Both metals have been used for pipes in the past, but because lead is deemed too hazardous, copper is now the most common choice. Gold, platinum, and mercury are non-reactive metals that do not corrode.

Prevention of Corrosion

Paint, which creates a barrier between the metal and air/water, is the best way to preserve iron and steel (alloys of iron). A water resistant oil or grease coating can be used to protect moving components on machinery. Another option is to use enamel and lacquers to cover the surface.

Sacrificial Protection

By linking iron to a more reactive metal, 'rusting' can be avoided (e.g., zinc or magnesium). Because the more reactive shielding metal is preferentially oxidised away, leaving the protected metal intact, this is referred to as sacrificial protection or sacrificial corrosion.

Alloying

Iron and steel, as well as other metals, may be made non-rusting alloys by 'alloying,' or combining with other metals (e.g., chromium). Stainless steel is an iron and carbon alloy that does not rust. Brass, a copper-based alloy, is another less costly and non-reactive metal alloy.

5 Important Topics of Class 10 Chapter 3 You Shouldn’t Miss!

Importance of Science Chapter 3 Notes Metals and Non Metals Class 10 PDF

• Understanding the properties of metals and non-metals helps in knowing how they react with each other and their environment.

• Learning about the extraction processes from ores is crucial for understanding industrial applications and the source of metals.

• Knowing the methods to prevent corrosion is important for real-life applications, such as maintaining infrastructure and machinery.

• Recognising the differences in physical and chemical properties helps in choosing the right materials for various uses.

• Studying these concepts prepares students for exams by providing a clear understanding of key reactions and processes.

Tips For Learning the Class 10 Science Chapter 3 Metals and Non Metals

• Memorise key reactions between metals and nonmetals, such as how they react with oxygen, water, and acids.

• Use diagrams to understand processes like extracting metals from ores and the differences between metals and nonmetals.

• Work through examples of metal and non-metal reactions in various conditions to strengthen your understanding.

• Relate concepts to real-life examples, like how rust forms on iron or why aluminium is used in cooking utensils.

• Review past exam questions on metals and nonmetals to familiarise yourself with common question types and improve your exam skills.

Conclusion 

The chapter on Metals and Non-Metals explains the key properties and reactions of these materials. Metals are usually shiny, good conductors of heat and electricity, and can be shaped easily. Non-metals, on the other hand, are often not shiny, do not conduct heat or electricity well, and are usually brittle. The chapter also covers how metals are extracted from ores and how they react with acids and bases. Knowing these details helps in understanding the practical uses of metals and non-metals and is important for solving exam questions effectively.

Related Study Materials for Class 10 Science Chapter 3 Metals and Non-metals

Students can also download additional study materials provided by Vedantu for Science Class 10, Chapter 3–

Chapter-wise Links for Science Class 10 Notes

Important Study Materials for Class 10 Science