Chapter 3: Metals and Non-Metals Notes

What are Metals?

Metals are substances that are formed naturally below the surface of the Earth. Most of the metals are lustrous, i.e. they are shiny. Metals are made of substances that were never alive.
Physical Properties of Metals
The various physical properties of the metal are,
Conductivity: Metals are good conductors of heat and electricity and thus they find applications in day-to-day life like cooking utensils which are made of iron or aluminium as they are good conductors of heat.
Malleability: Metals are malleable in nature. Malleability is the property of metals that allows them to be beaten into flat sheets.
Ductility: The capacity of a substance to be drawn into a wire is known as ductility, and it is this property that permits metals to be used as cable wires and for soldering.
Sonorous: Metals are sonorous in nature they produce a deep or ringing sound when struck with another hard object.
Lustrous: Most metals are lustrous in nature, i.e. they have a shiny appearance.
Solid: Metals are generally solid in nature except for Mercury which is a liquid metal.
Chemical Properties of Metals
Metals have various chemical properties some of the properties of metals are listed below.
Metals produce metal oxide when reacting with the oxygen in the air.
Some metals like sodium and potassium are highly reactive and they can react vigorously with moisture in the air and are thus, stored in an oil bottle.
Metals are highly corrosive and they react with oxygen and water in the air to form rust.
Meals react with bases to form salt and liberate hydrogen.
Metals are good reducing agents
Uses and Applications of Metals
Metals are usually very strong, durable and highly immune to everyday wear and tear. As such, they need been used in past for tons of things. Even now, with developments in technology and a slew of other factors, metals’ applications have expanded significantly. Metals are even important in the economy.
Construction Industry: Metals are the most component within the housing industry. Iron and steel are amongst the most utilized metals in the construction of buildings and even homes.
Electronics: Metals are utilised to make cables and parts for electrically powered devices and gadgets because they are good conductors of electricity. TVs, cell phones, refrigerators, irons, and computers are just a few examples.
Medicine: Metal elements are required for a variety of functions, including nerve impulse transmission, oxygen flow, enzyme reaction, and so on. To treat particular deficiencies or illnesses, several medicines are combined with metal compounds. Antacids contain metals such as iron, calcium, magnesium, potassium, titanium, and aluminium, which are commonly used in medicine.
Automobiles and Machinery: They are widely employed in the production of machines for industry, agriculture, and farming, as well as autos such as road vehicles, railways, aeroplanes, and rockets. Iron, aluminium, and steel are the most often used metals in this area. The majority of cooking utensils are constructed of metals such as steel, aluminium, and copper. Metals are preferred because of their excellent thermal resistance.
Other Uses: These days, most furniture is made of metal. Metals are also employed in the military, where they are used in the production of weapons and ammunition. Galvanizing protects metals from rusting by using certain metals.
What are Non-Metals?
Elements that lack the attributes of metals are called non-metals.
Non-metals are good insulators of heat and electricity. They are mostly gases and liquids. Some non-metals are solid at room temperature. E.g. Carbon, sulphur and phosphorus.
Some common non-metals are shown in the image below.

Physical Properties of Non-Metals
The various physical properties of the non-metal are,
Non-metals are poor conductors of heat and electricity. A notable exception is “Graphite” which is a good conductor of heat and electricity.
Non-metal are not ductile in nature except for the carbon nanotube which is ductile in nature.
Non-metal are not malleable and they are brittle i.e. they break when hit by a hammer or when force is applied.
Non-metal are not lustrous as they do not have any shiny appearance.
Non-metals don’t produce a deep ringing sound when they are hit with another material. Thus, they are not sonorous.
Chemical Properties of Non-metals
Non-metals have various chemical properties some of the properties of non-metals are listed below.
Non-Metals are more reactive than metals.
Non-Metals react with oxygen to form acidic oxide.
Non-metals react with other non-metals at high temperatures
Uses and Applications of Non-Metals
Various uses and applications of non-metals are,
Daily Life: The respiration process is aided by oxygen, which is 21 % by volume. It’s also utilised to make steel and maintain a high temperature during the metal fabrication process. In the hospital, oxygen cylinders are used. As a bleaching chemical, chlorine is effective in eliminating stains and colour patches. Chlorine is used to make a variety of polymers and pesticides. It aids with water filtration. How? Bacteria are killed when chlorine is added to drinking water. For scientific experiments, helium is employed as an inert gas. Weather balloons use it as well. Iodine is used as an antiseptic in the treatment of wounds and cuts, as well as in the treatment of throat infections.
Fertilizers: Nitrogen is found in fertilizers. It aids in the growth of plants. It boosts the plant’s growth rate. Plants can also benefit from non-metallic phosphorus. These two nonmetals are essential for plant growth.
Crackers: Sulphur and phosphorus are used in fireworks.
Differences Between Metals And Non-Metals

Various differences between Metals And Non-Metals are discussed below in the table,

Metals	Non-metals
Metals are solids at room temperature except for mercury which is liquid at room temperature.	Non-metal exists in all three states i.e. solid, liquid and gases
Metals are generally very hard except for Sodium, Potassium and Mercury.	Non-metals are generally soft in nature except for Diamond and Graphite which are solid in nature.
Metals are malleable and ductile in nature	Non-metals are brittle in nature and can break into pieces if struck by a hammer.
Metals are electropositive in nature.	Non-metals are electronegative in nature.
Metals have high densities.	Non-metals have very low densities.
Metals are lustrous in nature.	Non-metals are non-lustrous in nature.

Reactivity Series of Metals

To determine if a metal can displace another in a single displacement reaction, the data provided by the reactivity series can be utilized. It can also be used to find out how reactive various metals are to acids and water.

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Below is a chart showing the reactivity series of common metals:

Salient Features of Reactivity Series

The reducing tendency of metals at the top of the table has high, which is why they are easily oxidized. These metals can get tarnished or corrode very easily. For example, sodium due to its high reactivity explodes as came in contact with air or water, thus needs to be kept away from moisture and air, and dipped in kerosene.
The electro-positivity (tendency to lose electrons) of the elements gets reduced while moving down the reactivity series of metals, as we go down in the series it’s hard for elements to lose electrons, hence the less reactivity.
The tendency to reduce the metals becomes weaker while traversing down the series.
On reaction with dilute HCl or dilute H₂SO₄, all metals that are found above hydrogen in the activity series liberate H₂ gas and form salt.
Higher-ranking metals require greater amounts of energy for their isolation from ores and other compounds.

Long Tabular Form of the Reactivity Series

The Reactivity series of some of the most common metals, arranged in descending order of reactivity which can be given as follows:

Metal	Ion	Reactivity
Potassium	K⁺	React with Cold Water
Sodium	Na⁺
Lithium	Li⁺
Barium	Ba⁺
Strontium	Sr⁺
Calcium	Ca⁺
Magnesium	Mg⁺	Reacts very slowly with cold water, whereas quickly in boiling water, and very strongly with acids.
Beryllium	Be²⁺	Reacts with steam and acids
Aluminium	Al³⁺	Reacts with steam and acids
Titanium	Ti⁴⁺	Reacts with concentrated mineral acids
Manganese	Mn²⁺	Reacts with acids
Zinc	Zn²⁺
Chromium	Cr³⁺
Iron	Fe²⁺
Cadmium	Cd²⁺
Cobalt	Co²⁺
Nickel	Ni²⁺
Tin	Sn²⁺
Lead	Pb²⁺
Antimony	Sb³⁺	May react with some strong oxidizing acids
Bismuth	Bi³⁺	May react with some strong oxidizing acids
Copper	Cu²⁺	React slowly with water
Tungsten	W³⁺	Highly unreactive May react with some of the strong oxidizing acids
Mercury	Hg²⁺
Silver	A^g+
Gold	Au³⁺

Important Uses of Reactivity Series

The reactivity series provides the study of properties and reactivities of the metals, Apart from this reactivity series also provides several other important applications. For example, the result we get out of the reactions between metals and acids, metals and water, and single displacement reactions between metals can be predicted.

Reaction Between Metals and Water

Calcium and the metals that are more reactive than calcium in the reactivity series can react with cold water to form the corresponding hydroxide while liberating hydrogen gas. For example, the reaction between potassium and water yields potassium hydroxide and H₂ gas, as described by the chemical equation provided below.

2K + 2H₂O → 2KOH + H₂

Single Displacement Reactions between Metals

Highly reactive metals can reduce the less reactive metals from their salts in an aqueous solution and this happens through a single displacement reaction.

For example, copper can easily displace silver from the silver nitrate solution, as copper is much more reactive than silver and the balanced chemical reaction of the same is given as follows:

Cu(s) + 2AgNO₃ (aq) → Cu(NO₃)₂(aq) + 2Ag(s)

Reaction with Acid and Base

Metals that are more reactive than hydrogen can displace hydrogen from dilute acids, and form respective metal salts and hydrogen gas.

Metal + Dilute Acid or Base → Salt + Hydrogen gas

For example, sodium metal has a tendency to displace hydrogen from hydrogen chloride (HCl) to form sodium chloride and liberate hydrogen gas. The reaction is given below:

2Na(s) + 2HCl (dilute) → 2NaCl (aq) + H₂(g)

For another example, Iron can displace hydrogen from the sulfuric acid to make Iron Sulfate and Hydrogen gas, and the balanced chemical reaction of the same is as follows:

Fe(s) + H₂SO₄ (dilute) → FeSO₄(aq) + H₂(g)

For another example for Base, When zinc reacts with sodium hydroxide, it gives an aqueous solution of sodium zincate (as zinc can’t replace sodium from the base due to less reactivity) and hydrogen gas. The balance chemical reaction of the same is as follows:

Zn(s) + 2NaOH (aq) → Na₂ZnO₂(aq) + H₂(g)

Reactivity Series Trick

The basic strategy for learning the reactivity series is to group these elements into meaningful groups and then use a mnemonic for each group.

Here we have used the sentence “Please send charlie’s monkeys and zebras in lead & hydrogen cages in mountains securely guarded by Plato.” The first letter of each word represents a metal of the reactivity in the order from highest to lowest as:

Symbol	Metal of the Reactivity Series	Words/Mnemonic used to Remember the Series
K	Potassium	Please
Na	Sodium	Send
Ca	Calcium	Charlie’s
Mg	Magnesium	Monkey
Al	Aluminium	And
Zn	Zinc	Zebra
Fe	Iron	In
Pb	Lead	Lead &
H	Hydrogen	Hydrogen
Cu	Copper	Cages in
Hg	Mercury	Mountains
Ag	Silver	Securely
Au	Gold	Guarded by
Pt	Platinum	Plato

What is an Ionic Bond?

The ionic bond is the electrostatic force of attraction that holds two oppositely charged ions together.

The complete transfer of one or more electrons from one atom to the other forms a chemical link between two atoms, causing the atoms to acquire their closest inert gas configuration. There are three main ways for two atoms to unite in order to shed energy and become stable. One option is to complete their octet arrangement by donating or accepting electrons. An ionic bond, also known as an electrovalent bond, is created by this type of combination.

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Properties of Ionic Bond:

The following qualities are found in ionic bonded molecules due to the presence of a strong force of attraction between cations and anions:

Ionic bonds are the most powerful of all binds.
The ionic bond is the most reactive of all the bonds in the correct medium since it contains charge separation.
The melting and boiling points of ionic bound compounds are quite high.
Ionic bound molecules are strong conductors of electricity in their aqueous solutions or molten state. This is because ions, which act as charge carriers, are present. The following qualities are found in ionic bonded molecules due to the presence of a strong force of attraction between cations and anions.

Formation of ionic bonds

The ionic bond is formed by the transfer of electrons from one atom to another. In this case, one atom can signify electrons for the inert gas configuration, whereas another atom requires electrons for the inert gas configuration. The outer shell of metals has 1,2,3 electron, which are donated. Metals accept electrons with 4,5,6,7 electrons in their outer shell.

e.g. The formation of sodium chloride takes place as shown below:

What are Ionic Compounds?

Ions with opposing charges are closely packed to create crystalline solids. When metals react with non-metals, ionic compounds are created.

Ionic compounds are ionic compounds that have ionic bonds that bind them together. Elements can gain or lose electrons in order to obtain their closest noble gas configuration. Their stability is aided by the creation of ions for the completion of the octet (either by receiving or losing electrons).

In a reaction between metals and non-metals, metals lose electrons to complete their octet, whereas non-metals receive electrons to complete their octet. When metals and nonmetals react, ionic compounds are created. Crystalline solids are made up of ions with opposite charges packed close together. Ionic compounds are formed when metals react with nonmetals.

Structure of Ionic Compounds

Ionic molecule’s structure is determined by the relative sizes of its cations and anions. The majority of inorganic compounds, including salts, oxides, hydroxides, and sulphides, are ionic compounds. Ionic solids are held together by the electrostatic interaction of positive and negative ions.

Chloride ions attract sodium ions, and sodium ions attract chloride ions. As a result, Na⁺ and Cl^– ions alternate in a three-dimensional framework. The crystal in this picture is sodium chloride. The crystal is uncharged because the amount of sodium ions is equal to the number of chloride ions. The forces of attraction that exist between the ions keep them in place. A large ion structure is formed by the ionic interactions between the charged particles. Breaking all of the links in these huge complexes takes a lot of energy because the ions are tightly bonded. As a result, ionic compounds have high melting and boiling points.

Properties of Ionic Compounds

Physical properties of ionic compounds: Ionic compounds are solids that are difficult to break due to the strong attraction between the positive and negative ions. When pressure is applied to them, they usually split into fragments, making them brittle.
Melting and boiling points of ionic compounds: The presence of electrostatic forces of attraction between ions necessitates the use of a significant amount of energy to break the ionic bonds between atoms. Ionic compounds have high melting and boiling points as a result of this.
The solubility of ionic compounds: Ionic chemicals are generally soluble in polar solvents such as water, but non-polar solvents such as petrol, gasoline, and other hydrocarbons have a lower solubility.
Conduction of Electricity: In the solid-state, ionic compounds do not conduct electricity, but in the molten form, they are excellent conductors. The transfer of charge from one point to another is what electricity conducts. Ionic compounds do not conduct electricity in the solid state because ion mobility is not feasible. Ionic compounds conduct electricity in the molten state because the electrostatic forces of attraction between the ions are overcome by the heat generated.

Example of an Ionic Compound

In the solid-state, ionic compounds do not conduct electricity, but in the molten form, they are excellent conductors. The transfer of charge from one point to another is what electricity conducts. Ionic compounds do not conduct electricity in the solid state because ion mobility is not feasible. Ionic compounds conduct electricity in the molten state because the electrostatic forces of attraction between the ions are overcome by the heat generated.

The chlorine atom, on the other hand, contains seven electrons in its outermost shell. As a result, it only requires one electron to complete its octet. To become a magnesium ion, it can gain this one electron from the electrons shed by the magnesium atom. Because a magnesium atom loses two electrons whereas a chlorine atom can only get one, two chlorine atoms join with one atom of magnesium to create magnesium.

Sample Questions

Question 1: Is it feasible for two nonmetals to make ionic connections?

Answer:

Consider whether each element is a metal or a nonmetal to anticipate the type of bond that will form between them. Covalent bonds are formed by nonmetals, ionic bonds are formed by metals and nonmetals, and metallic bonds are formed by metals and metals in general.

Question 2: Why are the melting points of ionic compounds so high?

Answer:

Ionic compounds have both positive and negative charges, which is why they are called ionic compounds. As a result, there will be a strong attraction between them. Because breaking this force of attraction requires a lot of heat, ionic compounds have high melting temperatures.

Question 3: What are the two elements that make up an ionic compound?

Answer:

Ionic compounds are made up of ions, which are charged particles that occur when an atom (or group of atoms) acquires or loses electrons. An anion is a negatively charged ion, while a cation is a positively charged ion.

Question 4: How do you break ionic bonds?

Answer:

Because ionic compounds are polar, they dissolve in polar solvents like water. Polar solvents breakdown ionic bonds by disrupting them. The ionic bonds can be disrupted by dissolving the ionic material in water.

Question 5: What are some examples of common ionic compounds?

Answer:

With high melting and boiling temperatures, ionic compounds appear to be both robust and fragile. Ions can be simple groupings of atoms, such as sodium and chlorine in table salt (sodium chloride), or more complex groupings, such as calcium carbonate.
Occurrence of Metals
The earth’s crust is a key source of metals. Metal salts such as sodium chloride and magnesium chloride, among others, are found in seawater. Metals are not found in a free state in nature. As a result, the majority of metals are found in compounds with other elements, referred to as the combined state. The compounds of these metals found in nature are sulphides, oxides, carbonates, chloride, and some other metal compounds. Metals are present in these compounds in the form of positive ions or cations. Only a few metals, such as copper, silver, gold, and platinum, exist in a free state as metals. When a metal is discovered as a free state, it is referred to as being in its “native state.” Copper, silver, gold, and platinum are all found in nature in their native state. Copper and silver metals occur in the free state as well as in the combined state in nature.
Why do some metals occur in a combined state and others in a free state?
Potassium, sodium, calcium, magnesium, and aluminium, which are at the top of the reactivity series, are so reactive that they are never found as free elements in nature. They are always found in a combined state. Metals like zinc, iron, and lead, which are in the middle of the reactivity series, are moderately reactive metals that can also be found in the combined state.

All of the metals in the reactivity series above copper are only found in nature in form of their compounds. Copper, silver, gold, and platinum are among the least reactive or unreactive metals and are thus found in their free state as metals. Only a small amount of copper and silver are found in a free state. They are generally found as sulphides or oxides in their combined state.
What are Minerals?
A mineral is distinct from synthetic equivalents made in the laboratory by the fact that it must be developed through natural processes. Artificial minerals, such as emeralds, sapphires, diamonds, and other valuable gemstones, are regularly produced in industrial and research facilities and are almost identical to their natural counterparts. A mineral is a single solid substance of uniform composition that cannot be physically divided into simpler chemical compounds, according to its definition as a homogeneous solid.
Minerals have a regular geometric form and a highly ordered interior atomic structure. Minerals are classified as crystalline solids because of this property. Crystalline solids can express their structured internal framework through a well-developed external form, known as crystal form or morphology, under the right conditions. Solids that do not show such ordered internal arrangement are known as amorphous. The mineral formation can be divided into four groups as given below.
Minerals crystallise from a melt in igneous or magmatic rocks.
Minerals are formed as a result of sedimentation, a process in which the raw materials are particles from other rocks that have been weathered or eroded.
Metamorphic refers to the formation of new minerals at the expense of older ones as a result of changes in temperature, pressure, or both on an existing rock type.
Minerals are chemically precipitated from hot solutions within the Earth’s crust in a hydrothermal process.
The first three steps produce a variety of rocks with distinct mineral grains tightly intergrown in an interlocking weave. Hydrothermal and very low-temperature liquids tend to follow fracture zones in rocks, creating open spaces for mineral chemical precipitation. The majority of the exceptional mineral specimens have been obtained from such open spaces, which are partially filled by minerals deposited from solutions. Allowing a mineral in the process of growth (as a result of precipitation) to develop in free space generally results in a well-developed crystal shape, which contributes to the visual appeal of a specimen.
Thus, minerals are natural materials in which metals or their compounds can be found in the earth. Some minerals may contain a considerable amount of metal, whereas others may only contain a small percentage. Some minerals may be free of unwanted impurities, while others may include some unwanted impurities that restrict metal extraction. As a result, all minerals can not be used to extract metals.
What are Ores?
Ores are the minerals from which metals can be extracted efficiently and profitably. A previous meaning limited the term ore to metallic mineral deposits, although it has since been expanded to cover non-metallics in some cases. Only roughly 100 of the more than 2,800 mineral species have been recognised as ore minerals. Hematite, magnetite, limonite, and siderite are the main sources of iron; chalcopyrite, bornite, and chalcocite are the main sources of copper; and sphalerite and galena are the main sources of zinc and lead, respectively. No ore deposit is made up of a single ore mineral.
The ore is invariably mixed with gangue, which is a term for undesired or non-valuable rocks and minerals. In most cases, the ore and gangue are mined together—that is, a mass of ore and gangue is extracted from the host rock by either mechanical or human means. The ore is then separated from the gangue via a series of activities known as mineral processing, or ore dressing. Smelting, roasting, or leaching processes are used to extract the necessary metallic element from the ore. Some metals, such as copper, uranium, and gold, can now be extracted from host rock without drilling or blasting because of advances in hydrometallurgy.
In some cases, special bacteria are used as part of the process. Metals can be refined or alloyed with other metals once they’ve been recovered, as in a copper refinery or steel mill. Mining, processing, and refining are all processes in the process of extracting metal from an ore deposit. There are no unwanted impurities in ore, and it includes a high percentage of metal. As a result, all ores are minerals, but not all minerals are ores.
The following table lists some of the most common ores:

Metal (to be extracted)
Name of ore
Name of the compound in ore
Formula of ore
1.
Sodium
Rock salt
Sodium Chloride
NaCl
2.
Aluminium
Bauxite
Aluminium oxide
Al₂O₃.2H₂O
3.
Iron
Haematite
Iron (III) oxide
Fe₂O₃
4.
Manganese
Pyrolusite
Manganese dioxide
MnO₂
5.
Mercury
Cinnabar
Mercury (II) sulphide
HgS
6.
Zinc
Zinc blende
Calamine
Zinc sulphide
Zinc carbonate
ZnS
ZnCO₃
7.
Copper
Cuprite
Copper glance
Copper (I) oxide
Copper (I) Sulphide
Cu₂O
Cu₂S
Solved Questions
Question 1: Why do most of the metals do not occur in their free state?
Answer:
The most reactive metals are those at the top of the reactivity series, and are so reactive that they are never found as free elements in nature.
Question 2: Define the term gangue.
Answer:
The unwanted impurities or non-valuable rocks, earthy particles present in an ore are called gangue.
Question 3: Among the metals magnesium and mercury, which occur in a combined state and which occur in a free state?
Answer:
The metals above hydrogen in the reactivity series are so reactive that they are never found as free elements in nature, whereas the metals which are at the bottom of the reactivity series below hydrogen, are least reactive so they occur in a free state. Since magnesium exists above hydrogen in the reactivity series, so it occurs in a combined state and mercury exists below hydrogen in the reactivity series, so it occurs in a free state.
Question 4: Name a metal that occurs in a free state as well as in a combined state?
Answer:
Copper is a metal that exists both in its free and mixed states. Only a small quantity of copper exists in its free state. In its combined state, it usually exists as its sulphides or oxides.
Question 5: Name the metal which is extracted from the cinnabar ore.
Answer:
The metal that is extracted from the cinnabar ore is mercury.
Question 6: What is hydrometallurgy?
Answer:
Hydrometallurgy is the process of extracting metal from ore by making an aqueous solution of a metal salt and recovering the metal from it.

Extraction of Metals

Metals are extracted from their ores through a series of processes. The stages vary depending on the kind of ore, the metal’s reactivity, and the nature of impurities in the ore. Metallurgy refers to the processes involved in metal extraction and refinement. Most metal ores need to be transported to the Earth’s surface so that metal may be extracted. Mining is the term for this procedure. In general, the process of extracting metals consists of three basic phases. These are the following:

Enrichment or concentration of an ore
Extraction of metal from concentrated ore
Refining of the impure metal

Enrichment or Concentration of Ore

After being mined from the earth, and ore contains numerous undesirable impurities such as sand, rough minerals, and so on. Gangue is a term for undesirable impurities such as earthy materials, rocks, sandy materials, limestone, and so on.

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To get a concentrated ore having a considerably greater proportion of the metal, the first step in metallurgy is to eliminate these undesirable impurities from the ore. The physical or chemical characteristics of the gangue and ores determine the method employed to extract it from the ore. The following are some of the processes used to concentrate the ore:

Hand-picking: The ore is broken into little pieces, and the sand and mud that adheres to it are washed away by a stream of water.
Hydraulic washing: This procedure is also known as levigation or gravity separation. It is based on the specific gravities of the ore and gangue particles being different.
Electromagnetic separation: Magnetic ore is separated from impurities using this process. In this process, the powdered ore is placed on a leather belt that passes over two rollers, one of which is magnetic. When crushed ore travels over the magnetic roller, magnetic ore particles are attracted to it and sink below it, while impurities fall away from it. This method is used to extract chromite, rutile, and wolframite from siliceous gangue, chlorapatite, and cassiterite, respectively.
Froth floatation process: This method is widely used for sulphide ores and is based on the ore and gangue particles’ differing wetting properties. A large tank is filled with finely powdered ore, water, pine oil (frother), metal sulphide, and ethyl xanthate or potassium ethyl xanthate (collector). The entire mixture is then stirred with air. Oil-soaked ore particles enter the froth and are removed, whilst water-soaked impurities sink to the bottom. Pine oil is used as a foaming agent, and cresol and anisole are used as froth stabilisers. Ethyl xanthate and potassium ethyl xanthate are utilised as collectors. Activator in CuSO₄ and a wild depressant in KCN.
Liquation: This technique is appropriate for ore with readily fusible mineral particles and a high melting gangue.
Chemical separation (leaching): In this process, a suitable chemical reagent is used to dissolve the powdered ore while impurities remain insoluble in the reagent. With the assistance of NaOH, bauxite is separated from Fe₂O₃, SiO₂, and TiO₂, with Al₂O₃ soluble and the rest insoluble.

Al₂O₃+ 2NaOH → 2NaAlO₂+ H₂O

NaAlO₂+ 2H₂O → Al(OH)₃+ NaOH

2Al(OH)₃→ Al₂O₃+ 3H₂O

Ag₂S + 4NaCN → 2Na[Ag(CN)₂] + Na₂S

Extraction of Metal from Concentrated Ore

Metals are extracted from concentrated ore using a variety of techniques. The metals are divided into three categories based on their reactivity:

Less reactive metals or metals with low reactivity

Moderately reactive metals or metals with a medium reactivity
Moderately reactive metals or metals with a high reactivity
Metals with a high reactivity or highly reactive metals

Extraction of Less Reactive Metals

Mercury (Hg), gold (Au), and platinum (Pt) are the least reactive metals in nature and are found in a free state at the bottom of the activity series. As a result, these metals may be removed only by reducing their oxides with heat.

Extraction of Mercury: Cinnabar (HgS), a sulphide ore, is the most common mercury ore. The following steps can be used to extract the metal (Hg) from the ore.

Step 1: In the air, roast concentrated mercury (II) sulphide ore.

Roasting is the process of turning a sulphide ore into its equivalent metal oxides by rapidly heating it in the presence of air. As a result, concentrated mercury (II) sulphide is roasted in the air to produce mercury (II) oxide.

Step 2: Mercury (II) oxide ore is converted to mercury metal

Sulphide ore roasting: To transform cinnabar ore (HgS) into metal oxide, it is heated in the presence of air (HgO).

2HgS + 3O₂→ 2HgO + 2SO₂

Reduction of metal oxide to metal: As the metal oxide (HgO) is heated further, it is reduced to metal.

2HgO → 2Hg + O₂

Thus, HgS can be converted into Hg by heating alone.

Extraction of Moderately Reactive Metals

The metals lying in the reactivity series are respectably receptive and typically happen as sulfides or carbonates in nature. Accordingly, extraction of these metals is additionally done in two stages:

The respectably receptive metals can be removed by the decrease of their oxides with carbon (C), aluminium (Al), sodium (Na), or calcium (Ca). Some respectably receptive metals likewise happen in nature as their carbonates or sulfides. In any case, we can say that metals can be more effortlessly removed from their oxide minerals than carbonates or sulfide metals. The oxide minerals can be straightforwardly changed over into metals by warming, while the carbonate or sulfide metals should initially be changed over into metal oxide.
The concentrated minerals can be changed over into metal oxide by utilizing calcination or broiling dependent on the idea of the metal. The method involved with warming carbonate metal firmly without air is called calcination.

Metals like zinc (Zn), tin (Sn), lead (Pb), and iron (Fe) can be extricated by the course of calcination. The metal oxide framed is then changed over into metal by warming it within the sight of decreasing specialists like carbon (C), aluminium (Al), sodium (Na), or calcium (Ca). The utilization of diminishing specialists relies upon the compound reactivity of the metal to be extricated.

Extraction of Zinc:

Conversion of ore into metal oxide: Zinc is found in nature as both sulphide and a carbonate. As a result, they must be transformed into zinc oxide before being reduced.

Roasting of zinc sulphide:

2ZnS+3O₂→ 2ZnO + 2SO₂

Calcination of zinc carbonate:

ZnCO₃→ ZnO + CO₂

Reduction of metal oxide to metal: Carbon is used as a reducing agent to convert zinc oxide to zinc metal.

ZnO + C → Zn + CO

The type of reducing agent employed is determined by the metal oxide to be reduced.

Extraction of Highly Reactive Metals

The metals lying high in the reactivity series are extremely responsive and can’t be acquired by diminishing their oxides and different mixtures utilizing normal decreasing specialists like carbon. The metal oxides of these metals are hard to decrease as these metals have a high proclivity for oxygen. Such profoundly responsive metals are separated by the electrolytic decrease of their liquid chlorides or oxides.

Electrolytic decrease: When metals are separated from their liquid chlorides or oxides by passing an electric flow through them. This course of electrolytic decrease is likewise called electrolysis. In an electrolytic decrease technique, metal particles on electrolysis move towards the cathode, acquiring an electron to become metal atoms.

Electrolysis of Molten Sodium Chloride

Sodium metal is removed from the liquid sodium chloride by the course of electrolytic decrease. At the point when liquid sodium chloride is electrolyzed by passing electric flow, it deteriorates into sodium (Na⁺) particles and chloride (Cl^–) particles. The sodium Na⁺ particles move towards the cathode (negative terminal) while chloride Cl^– particles move towards the anode (positive cathode). These sodium Na⁺ particles acquire electrons at the cathode and get diminished to sodium molecules, and chloride Cl^–particles lose electrons at the anode and get oxidized to chlorine iotas. The response included is as per the following:

At cathode: 2Na⁺+ 2e^–→ 2Na
At anode: 2Cl^–→ Cl₂ + 2e^–
Overall reaction: 2Na → 2Na + Cl₂

At the cathode, sodium metal is obtained, while at the anode, chlorine gas is released.

Refining of Impure Metals

The metal acquired by any strategy for decrease measure normally contains a few pollutants, so they are tainted. The metal got alongside the pollutants is called rough metal. Presently, we need to eliminate these contaminations to get 99.9% unadulterated metal. The most common way of cleaning debased metals (unrefined metals) is called refining of the metal.

Diverse refining strategies are utilized for various metals. The strategy to be utilized for refining a tainted metal relies upon the idea of the metal and the idea of pollutants present in it. The most significant and broadly utilized strategy for purging debased metals is electrolytic refining. Since the refining of the metal is finished by electrolysis, this technique is called electrolytic refining. Numerous ways are utilized to clean metals, out of which electrolytic refining is most generally utilized.

Sample Questions

Question 1: What are the methods of extraction of metals?

Answer:

Metals may be extracted from ore in three different ways. Electrolysis, reducing an ore with a more reactive metal, and reducing the ore with carbon are the methods used.

Question 2: What type of chemical reaction is used to extract metals from ores?

Answer:

The decomposition reaction, which is powered by electricity, removes metals from naturally occurring compounds such as oxides and chlorides.

Question 3: How is a metal extracted and processed?

Answer:

Although other non-ferrous metals have lower melting points than aluminium and may thus be treated at lower temperatures, the same process stages are often used: crushing, grinding, flotation or other methods of concentration, smelting, refining, and electrolytic purification.

Question 4: What is the major source of metals?

Answer:

The primary natural sources of heavy metals in the environment are rocks and soils. When magma cools, the main rocks, also known as magmatic or igneous rocks, crystallise.

Question 5: What are the steps involved in the extraction of metals from ores?

Answer:

The process of extracting a metal from its ore is divided into three steps. They are ore enrichment or concentration, metal extraction from concentrated ore, and impure metal refining.

Question 6: What is the froth floatation process?

Answer:

This method is widely used for sulphide ores and is based on the ore and gangue particles’ differing wetting properties. A large tank is filled with finely powdered ore, water, pine oil (frother), metal sulphide, and ethyl xanthate or potassium ethyl xanthate (collector). The entire mixture is then stirred with air. Oil-soaked ore particles enter the froth and are removed, whilst water-soaked impurities sink to the bottom. Pine oil is used as a foaming agent, and cresol and anisole are used as froth stabilisers. Ethyl xanthate and potassium ethyl xanthate are utilised as collectors. Activator in CuSO4 and a wild depressant in KCN.

Extraction of Highly Reactive Metals

Electrolytic reduction is used to extract metals with high reactivity from their ores, such as sodium, calcium, magnesium, and aluminium. As carbon is less reactive than these metals, it cannot be used to reduce them.

Electrolytic Reduction

Electrolytic reduction is a type of electrolysis in which an electric current is passed through a molten or dissolved ionic substance, causing the electrodes to chemically react and the materials to decompose. The hydroxides, oxides, and chlorides of metals in their combined state are electrically reduced using this process.

Metals are collected at the cathode. Some metals, such as K, Na, and Al, are obtained through the electrolytic reduction process. When an electric current is passed through molten sodium chloride or a solution of sodium chloride, sodium metal is deposited over the cathode.

Na⁺ + e⁻ ⇢ Na

2Cl⁻ − e⁻ ⇢ Cl2

2NaCl ⇢ 2Na + Cl₂

Electrolysis of Aqueous Sodium Chloride: In an aqueous solution, sodium chloride is dissociated and exists as sodium and chloride ions. In an aqueous solution, the electrolysis of sodium chloride is simpler. Water, on the other hand, can undergo reduction and oxidation reactions at various potentials. As a result, the substance that is oxidized or reduced is not only sodium and chloride ions, but also the water molecule. At both the cathode and the anode, two competing reactions are possible.

At Cathode: A reduction reaction occurs when the pH is 7. Water can be converted to hydrogen gas, and sodium ions can be converted to sodium metal.

2H₂O (l) + 2e^– → H₂(g) + 2OH^–

Na⁺(l) + e^– → Na (l)

At anode: With a pH of 7, the oxidation reaction occurs. Water can be oxidised to produce oxygen, or a chloride ion can be oxidised to produce a chlorine molecule.

2H₂O → O₂(g) + 4H⁺

2Cl^– → Cl₂ + 2e^–

As a result, the product of aqueous sodium chloride electrolysis can be anything between sodium metal or hydrogen gas at the cathode and chlorine or oxygen gas at the anode, with a sodium hydroxide byproduct resulting from the reaction of sodium and water. The electrolysis product is determined by the concentration of sodium chloride aqueous solution.

Electrolytic Reduction of Alumina

The negative electrode is a steel container that has been coated with carbon. Aluminium oxide is a type of ionic compound. The Al⁺³ and O^-2 ions are free to move and conduct electricity when it is melted. The cathode of the electrolysis of the alumina/cryolite solution yields aluminium and the anode yields oxygen.

Cathode Reaction: 4Al⁺³ + 12e^– → 4Al

Anode Reaction: 6O^-2 – 12e^– → 3O₂

Because aluminium is denser than alumina, it settles to the bottom of the cell and can be tapped off as pure liquid metal. At the positive carbon anode, oxygen is released. Carbon dioxide is also produced at the carbon anode as a result of hot oxygen reacting with the carbon anode to produce carbon dioxide gas.

Carbon + Oxygen → Carbon dioxide.

C (s) + O₂(g) → CO₂(g)

Carbon anodes gradually degrade because each molecule of carbon dioxide emitted carries a small amount of carbon with it. When the carbon anodes become too small, they must be replaced.

Sample Questions

Question 1: What does a reactivity series show?

Answer:

Metals are arranged in descending order of reactivity in the reactivity series. The reactivity of a metal can be determined by studying its reactions to competition and displacement.

Question 2: What metal is the least reactive?

Answer:

Transition metals are elements in the periodic table that are much less reactive, and metals like gold and platinum are at the bottom of the list, exhibiting little chemical reaction with any common reagents.

Question 3: Which is the most reactive element?

Answer:

Alkali metals are the most reactive elementary group (situated far apart from intermediate metals and noble gases). Cesium is second from the bottom of this group, has six electron shells, and has the characteristics of a reactive atom, making it the most reactive element.

Question 4: Why electrolytic reduction is done to obtain aluminium?

Answer:

Since aluminium has a high affinity for oxygen, it is a stable compound. It is resistant to common reducing agents such as carbon. As a result, electrolytic reduction is used to obtain aluminium.

Question 5: Define electrolytic reduction.

Answer:

Electrolytic reduction is the process by which highly reactive metals are removed from their compound state. It is usually done to obtain the purest form of the desired metal. Magnesium and aluminium are two examples.

Extraction of Moderately Reactive Metals

In the middle of the reactive series are moderately reactive metals such as zinc, iron, tin, and lead, among others. As a result, extracting moderately reactive metals involves extracting metals in the middle of the reactivity series. The middle of the reactivity series metals is extracted by reducing their oxides with carbon, aluminium, sodium, or calcium.

Some moderately reactive metals are found as oxides in nature, while others are found as carbonates or sulphide ores. Metals are now easier to be obtained by reduction from their oxides than from carbonates and sulphides. As a result, the ore must first be turned into a metal oxide, which may then be reduced. Calcination or roasting can be used to transform the concentrated ores into metal oxide. The process to be utilised is determined by the ore’s nature. Calcination converts a carbonate ore to oxide, whereas roasting converts a sulphide ore to oxide.

Calcination- Calcination is the process of converting a carbonate ore into metal oxide by rapidly heating it in the absence of air. Zinc, for example, is found in calamine ore as zinc carbonate (ZnCO₃). In order to extract zinc metal from zinc carbonate, the latter must first be transformed to zinc oxide. Calcination is used to accomplish this. Calamine ore or zinc carbonate decomposes into zinc oxide and carbon dioxide when heated strongly in the absence of air. As a result of calcination, zinc carbonate is converted to zinc oxide.

ZnCO₃ → ZnO + CO₂
(Zinc carbonate) (Zinc oxide) (Carbon dioxide)

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Roasting- Roasting is the process of converting a sulphide ore into metal oxide by rapidly heating it in the presence of air. Zinc blende ore, for example, contains zinc as a sulphide, ZnS. So, before extracting zinc metal from zinc sulphide, zinc sulphide must first be converted to zinc oxide. Roasting is used to accomplish this. Zinc oxide and sulphur dioxide are formed when zinc blende ore or zinc sulphide is vigorously heated in the air or roasted. As a result, roasting turns zinc sulphide into zinc oxide.

2ZnS + 3O₂ → 2ZnO + 2SO₂
(Zinc sulphide) (Oxygen) (Zinc oxide) (Sulphur dioxide)

Reduction of Metal Oxides

Using reducing agents such as carbon, aluminium, sodium, or calcium, metal oxides formed by calcination and roasting of ores are converted to free metal. The chemical reactivity of the metal to be extracted determines the reducing agent used.

Reduction of Metal Oxide with Carbon- Carbon is commonly used to reduce the oxides of comparatively less reactive metals such as zinc, iron, nickel, tin, lead, and copper. The metal oxide is combined with carbon, in the form of coke and heated in a furnace in the carbon reduction process. The metal oxide is reduced to free metal by carbon.

For example, zinc is recovered from its oxide by reducing it with carbon or coke. When zinc oxide is heated with carbon, zinc metal is produced.

ZnO + C → Zn + CO
(Zinc oxide) (Carbon) (Zinc) (Carbon monoxide)

Iron metal is also recovered from its oxide ore, ‘haematite’ (Fe₂O₃), by carbon reduction in the form of coke. The reduction of oxides with carbon is also used to recover tin and lead metals. Copper, which is a less reactive metal, is extracted by reducing its oxide with carbon.

Reduction of Metal Oxide with Aluminium- In the extraction of metals from their oxides, a more reactive metal like aluminium can be utilised as a reducing agent. When the metal oxide is of a more reactive metal than zinc, for instance, and cannot be reduced effectively by carbon, aluminium is utilised as a reducing agent. This is due to the fact that a more reactive metal, such as aluminium, can displace a less reactive metal from its metal oxides, resulting in free metal. As a result, displacement reactions can be employed to convert some metal oxides to free metals. The aluminium powder oxidises to aluminium oxide after reducing the metal oxide to metal.

For example, Manganese is removed from its oxide by reducing it with aluminium powder as the reducing agent. Manganese metal is generated when manganese dioxide is heated with aluminium powder.

3MnO₂ + 4Al → 3Mn + 2Al₂O₃ + Heat
(Manganese dioxide) (Aluminium powder) (Manganese metal) (Aluminium oxide)

This oxidation and reduction reaction between MnO₂ and Al is a displacement reaction.

Extraction of Less Reactive Metals

Mercury, copper, and other less reactive metals are placed near the bottom of the reactivity series. As a result, the extraction of less reactive metals involves the extraction of metals with low reactivity. The less reactive metals, which are near the bottom of the reactivity scale, are removed solely by reducing their oxides with heat. Mercury and copper, for example, are less reactive metals that are near the bottom of the reactivity series. As a result, we’ll look at how to extract mercury and copper by reducing their oxides with heat alone.

Extraction of Mercury- Mercury is a relatively less reactive metal with a low reactivity level. Mercury may be extracted from its sulphide ore by simply heating it in the air. This occurs in the following manner. The sulphide mineral cinnabar, HgS, which is actually mercury (II) sulphide, is used to make mercury metal. The following are the two steps involved in extracting mercury from cinnabar.

When mercury (II) oxide is formed, the concentrated mercury (II) sulphide ore (cinnabar ore) is roasted in the air.

2HgS + 3O₂ → 2HgO + 2SO₂
(Mercury (II) sulphide) (Oxygen) (Mercury (II) oxide) (Sulphur dioxide)

When mercury (II) oxide is heated to around 300⁰C, it decomposes or reduces to generate mercury metal.

2HgO → 2Hg + O₂
(Mercury (II) oxide) (Mercury metal) (Oxygen)

As a result, mercury metal was produced by reducing mercury (II) oxide with heat alone.

Extraction of Copper- Copper is a less reactive metal that ranks near the bottom of the reactivity scale. Copper can be extracted from its sulphide ore by simply heating it in the air. This occurs in the following manner. A copper glance, CuS, is actually copper (I) sulphide, and is one among the ores from which copper metal is produced. The following two steps are involved in extracting copper from copper glance ore.

When a part of copper (I) sulphide is oxidised to copper (I) oxide, the concentrated copper (I) sulphide ore, also known as a copper glance, is roasted in the air.

2Cu₂S + 3O₂ → 2Cu₂O + 2SO₂
(Copper (I) sulphide) (Oxygen) (Copper (I) oxide) (Sulphur dioxide)

The supply of air for roasting is turned off once a significant amount of copper (I) sulphide has been converted to copper (I) oxide. Copper (I) oxide generated above interacts with the residual copper (I) sulphide to form copper metal and sulphur dioxide in the absence of air.

2Cu₂O + Cu₂S → 6Cu + SO₂
(Copper (I) oxide) (Copper (I) sulphide) (Copper metal) (Sulphur dioixde)

Sample Questions

Question 1: Name two reducing agents that are utilised in metal extraction.

Answer:

Carbon, which comes in the form of coke, and aluminium powder are the two reducing agents utilised in metal extraction.

Question 2: What is a thermite reaction, and how does it work?

Answer:

A thermite reaction is the reduction of a metal oxide to metal utilising aluminium powder as a reducing agent.

Question 3: When is a mineral called an ore?

Answer:

A mineral is called an ore if the mineral can be extracted profitably from it.

Question 4: Name a metal oxide that can be extracted from its ore by reduction with carbon?

Answer:

Zinc can be extracted from its oxides by reduction with carbon.

Question 5: How do you extract iron out of its ore haematite?

Answer:

Iron metal is recovered from its oxide ore, ‘haematite’ (Fe₂O₃), by carbon reduction in the form of coke.

Question 6: How can mercury be extracted from its sulphide ore?

Answer:

Mercury may be extracted from its sulphide ore by simply heating it in the air.

What is Electrolytic Refining?

Electrolytic refining is the process of purifying using electrolysis. Copper, zinc, tin, lead, chromium, nickel, silver, and gold are just a few of the metals that can be processed electrolytically. A thick block of impure metals is used as an anode for electrolysis refining of impure metals, and it is connected to the positive terminal of the battery.

As it is connected to the negative terminal of the battery, a thin strip of pure metal is used as the cathode. As an electrolyte, a water-soluble salt of the metal to be refined is used. When an electric current passes through the anode, the impure metal dissolves and enters the electrolyte solution, while pure metal from the electrolyte deposits on the cathode. The soluble impurities in impure metal dissolve in solution, whereas insoluble impurities settle at the bottom of the anode as anode mud. Consider an example to better understand electrolytic metal refining. Let’s look at how this method works for refining copper metal.

Electrolytic refining of Copper

An electrolytic tank containing an acidified copper sulphate solution as an electrolyte is used in the electrolytic refinement of copper. As the anode, a thick block of impure copper metal is attached to the positive terminal of the battery. Since it is connected to the negative terminal of the battery, a thin strip of pure copper metal is used as the cathode.

Impure copper from the anode dissolves and enters the copper sulphate solution when an electric current passes through it, whereas pure copper from the copper sulphate solution deposits on the cathode. Thus, on the cathode, pure copper metal is produced. The soluble impurities dissolve in the solution, while the insoluble impurities gather as anode mud below the anode.

Explanation of Electrolytic refining of Copper: Copper ions and sulphate ions are present in a copper sulphate solution. The following reaction occurs at the two electrodes when an electric current is passed through the copper sulphate solution.

The positively charged copper ions, Cu²⁺, from the copper sulphate solution go to the negative electrode (cathode) and are reduced to copper atoms by taking electrons from the cathode. Pure copper metals are formed when these copper atoms are placed on the cathode.

At cathode: Cu²⁺ + 2e^– → Cu

(Copper ions) (Electrons) (Copper atom)

The impure anode’s copper atoms each lose two electrons to the anode, forming copper ions Cu²⁺ in the electrolytic solution.

At anode: Cu – 2e^– → Cu²⁺

(Copper atom) (Electrons) (Copper ion)

The copper ions are removed from the copper sulphate solution at the cathode and added to the solution at the anode in this way. The impure anode gets thinner and thinner as the process progresses, whereas the pure cathode gets thicker and thicker. At the cathode, pure copper is thus obtained.

What happens to the impurities in impure copper that are metallic?

Impure copper contains metallic impurities that are either more reactive or less reactive. The more reactive metals in impure copper, such as iron, now flow into the electrolyte solution and stay there. The less reactive metals in impure copper, such as gold and silver, settle in the bottom of the electrolytic cell below the anode in the form of anode mud. The anode mud can be recovered for gold and silver metals. As a result, electrolytic metal refining serves two objectives.

It allows other valuable metals, such as gold and silver, to be recovered from impurities in the metal being refined.

It purifies the metal in consideration.

Solved Questions

Question 1: Define anode mud.

Answer:

During electrolytic refining, the soluble impurities in impure metal dissolve in solution, whereas insoluble impurities settle at the bottom of the anode and are known as anode mud.

Question 2: What is the cathode made up of in the electrorefining of copper?

Answer:

The cathode is the negative electrode in the electrorefining of copper. It is made entirely of copper. As a result, the cathode in copper electrorefining is made up of pure copper.

Question 3: What is the range of purity for copper obtained by electrolytic refining?

Answer:

The transfer of pure metal from the anode to the cathode is the result of electrolytic refining. More basic metal impurities remain in the electrolytic solution as ions, while less basic metal impurities settle down as anode mud. The copper that has been purified by electrolytic refining is 99.95-99.99% pure.

Question 4: What is the anode made up of in the electrorefining of copper?

Answer:

The anode is the positive electrode in the electrorefining of copper and it is made from impure copper blocks. As a result, the anode in copper electrorefining is made up of impure copper blocks.

Question 5: During the electrolytic refining of copper, which electrolyte is used?

Answer:

The electrolyte in the electrolytic refining of copper is a solution of copper sulphate acidified by sulphuric acid.
What is Corrosion?
One of the most typical phenomena we see in our daily lives is corrosion. You’ve probably seen that over time, some iron objects become covered in an orange or reddish-brown coloured layer. This layer is formed as a result of a chemical reaction known as rusting, which is a type of corrosion.
Corrosion is the process by which refined metals are transformed into more stable compounds such as metal oxides, metal sulphides, and metal hydroxides. The development of iron oxides occurs as a result of the action of air moisture and oxygen on iron. Corrosion is commonly regarded as a bad phenomenon since it compromises the metal’s good characteristics.

Iron, for example, is recognised for its tensile strength and stiffness (especially alloyed with a few other elements). Rusting, on the other hand, causes iron items to become brittle, flaky, and structurally unsound. Corrosion is an electrochemical process because it usually involves redox interactions between the metal and certain atmospheric agents including water, oxygen, and Sulphur dioxide, among others.
Do All Metals Corrode?
Metals with a greater reactivity series, such as iron and zinc, corrode quickly, whereas metals with a lower reactivity series, such as gold, platinum, and palladium, do not corrode. The reason for this is because corrosion requires the oxidation of metals. The tendency to oxidise decreases as we progress down the reactivity series (oxidation potentials is very low). Interestingly, although being reactive, aluminium does not corrode like other metals. This is due to the fact that aluminium is already covered with an oxide layer. It is protected from further corrosion by this layer of aluminium oxide.
Factors Affecting Corrosion:
Metals are exposed to gases such as CO₂, SO₂, and SO₃ in the air.
Metals are exposed to moisture, particularly saltwater (which increases the rate of corrosion).
Impurities such as salt are present (e.g. NaCl).
Temperature: As the temperature rises, so does the rate of corrosion.
The nature of the first oxide layer that forms: some oxides, such as Al₂O₃, generate an insoluble protective coating that can prevent further corrosion. Rust, for example, crumbles readily and exposes the rest of the metal.
Presence of acid in the atmosphere: acids have the ability to speed up the corrosion process.
Types of Corrosion
The following are the types of corrosion types:
Crevice Corrosion: A limited kind of corrosion known as crevice corrosion can occur whenever there is a difference in ionic concentration between any two local locations of a metal. Gaskets, the underside of washers, and bolt heads are all places where crevice corrosion can occur. Crevice corrosion occurs in all grades of aluminium alloys and stainless steel, for example.
Stress Corrosion Cracking: Corrosion Due to Stress SCC refers to the breaking of metal as a result of the corrosive environment and the tensile stress exerted on it. It happens a lot when the weather is hot. In a chloride solution, stress corrosion cracking of austenitic stainless steel is an example.
Intergranular Corrosion: The presence of contaminants in the grain boundaries that separate the grain generated during the solidification of the metal alloy causes intergranular corrosion. Depletion or enrichment of the alloy at these grain boundaries can also cause it. IGC, for example, has an impact on aluminium-base alloys.
Galvanic Corrosion: Galvanic corrosion can occur when an electric contact develops between two metals that are electrochemically different and are in an electrolytic environment. It describes the breakdown of one of these metals at a joint or junction. The degradation that occurs when copper comes into contact with steel in a saltwater environment is a good illustration of this form of corrosion. When aluminium and carbon steel are linked and submerged in seawater, the aluminium corrodes faster while the steel is protected.
Pitting Corrosion: Pitting Corrosion is unpredictably unpredictable, making it difficult to detect. It is regarded as one of the most hazardous forms of corrosion. It starts at a single location and progresses to the production of a corrosion cell encircled by the regular metallic surface. Once established, the ‘Pit’ continues to develop and can take on a variety of shapes. The pit progressively eats away at metal from the surface in a vertical direction, eventually leading to structural failure if not addressed. Consider a droplet of water on a steel surface; pitting will begin near the water droplet’s centre (anodic site).
Uniform Corrosion: This is the most prevalent type of corrosion, in which the environment attacks the metal’s surface. The degree of the rusting can be seen clearly. This sort of corrosion has a minimal impact on the material’s performance. A piece of zinc or steel immersed in diluted sulphuric acid would normally dissolve at a constant rate throughout its whole surface.
Corrosion Examples and Reactions
Here are some common examples of corrosion, which are typically encountered in metals.
Copper Corrosion
When copper metal is exposed to the environment, it combines with oxygen in the air to produce copper (I) oxide, which is a reddish-brown substance.
2Cu + 1/2 O₂ → Cu₂O
Cu₂O is oxidised further to generate CuO, which is black in colour.
Cu₂O+ 1/2O₂ → 2CuO
CuO interacts with CO₂, SO₃, and H₂O in the environment to produce Cu₂(OH)₂ (Malachite), a blue mineral, and Cu₄SO₄(OH)₆ (Brochantite), a green mineral. The colour of the copper plating on the Statue of Liberty, which has turned blue-green, is a good example of this.
Silver Tarnishing
Silver combines with Sulphur in the air to form silver sulphide (Ag₂S), which is a dark substance. Exposed silver reacts with H₂S in the environment, which is present due to some industrial processes, to generate Ag₂S.
2Ag + H₂S → Ag₂S+ H⁺₂
Corrosion of Iron (Rusting)
When iron comes into touch with air or water, rusting occurs, which is the most typical occurrence. The reaction resembles that of a normal electrochemical cell. Metal iron loses electrons and is converted to Fe²⁺ in this process (this could be considered as the anode position). The electrons that are lost will travel to the opposite side and interact with H+ ions. H+ ions are emitted in the atmosphere by either H₂O or H₂CO₃ (this could be considered as the cathode position).
H₂O ⇌ H⁺ + OH^–
H₂CO₃⇌ 2H⁺ + CO₃²
Prevention from Corrosion
Corrosion can be prevented in a number of ways. We’ll go over a few of the more popular ones below.
Electroplating: It’s an electrolysis-based method that coats a metal (I) with a thin layer of another metal (II). The new metal covering protects the metal (I) from corrosion in this way. Metal (I) (metal to be plated) is used as the anode and metal (I) (metal to be plated) is used as the cathode in this procedure. Metal ‘I’ is connected to the negative terminal, while metal ‘II’ is connected to the positive terminal. When electricity is applied to these two electrodes, oxidation occurs in the anode, resulting in the dissolution of metal II ions in the electrolyte. At the cathode, these dissolved metal II ions are reduced, resulting in a coating on metal I. Copper, Nickel, Gold, Silver, Zinc, and other metals are often used as anodes.
Cathodic Protection: The base metal is connected to a sacrificial metal that corrodes instead of the base metal in this procedure. This sacrificial metal (which is more reactive than the base metal) will release electrons and become oxidised as a result. The ions produced as a result of this process participate in corrosion reactions, preserving the base metal.
Galvanization: This procedure includes applying a thin layer of zinc to iron. In most cases, this is accomplished by dipping iron in molten zinc. As a result, the zinc layer protects the iron against corrosion.
Painting and Greasing: Applying a layer of paint or grease to the metal can keep it from coming into contact with the outside world, preventing corrosion.
Using Corrosion Inhibitor: Corrosion inhibitors are substances that, when introduced into a corrosion environment, reduce the rate of corrosion.
Make use of the Right Material: Corrosion can also be avoided by selecting the proper material. Aluminium and stainless steel, for example, are extremely corrosion resistant.
Sample Problems
Question 1: What do you understand by Corrosion?
Answer:
Corrosion is the process by which refined metals are transformed into more stable compounds such metal oxides, metal sulphides, and metal hydroxides. The development of iron oxides occurs as a result of the action of air moisture and oxygen on iron. Corrosion is commonly regarded as a bad phenomenon since it compromises the metal’s good characteristics. Iron, for example, is recognised for its tensile strength and stiffness (especially alloyed with a few other elements).
Rusting, on the other hand, causes iron items to become brittle, flaky, and structurally unsound. Corrosion is an electrochemical process because it usually involves redox interactions between the metal and certain atmospheric agents including water, oxygen, and Sulphur dioxide, among others.
Question 2: What are the factors that affect corrosion?
Answer:
The factors that affect corrosion are as follow:
Metals are exposed to gases such as CO2, SO2, and SO3 in the air.
Metals exposed to moisture, particularly salt water (which increases the rate of corrosion).
Impurities such as salt are present (eg. NaCl).
As the temperature rises, so does the rate of corrosion.
The nature of the first oxide layer that forms: some oxides, such as Al2O3, generate an insoluble protective coating that can prevent further corrosion. Rust, for example, crumbles readily and exposes the rest of the metal.
Presence of acid in the atmosphere: acids have the ability to speed up the corrosion process.
Question 3: What will happen when copper metal is exposed to the environment?
Answer:
When copper metal is exposed to the environment, it combines with oxygen in the air to produce copper (I) oxide, which is a reddish-brown substance.
2Cu + 1/2 O₂ → Cu₂O
Cu₂O is oxidised further to generate CuO, which is black in colour.
Cu₂O+ 1/2O₂ → 2CuO
Question 4: What will happen when iron comes into touch with air or water?
Answer:
When iron comes into touch with air or water, rusting occurs, which is the most typical occurrence. The reaction resembles that of a normal electrochemical cell.
Metal iron loses electrons and is converted to Fe²⁺ in this process (this could be considered as the anode position). The electrons that are lost will travel to the opposite side and interact with H+ ions. H+ ions are emitted in the atmosphere by either H₂O or H₂CO₃ (this could be considered as the cathode position).
H₂O ⇌ H⁺ + OH^–
H₂CO³ ⇌ 2H⁺ + CO₃²
Question 5: What will happen when Silver combines with Sulphur in the air?
Answer
Silver combines with Sulphur in the air to form silver sulphide (Ag₂S), which is a dark substance. Exposed silver reacts with H₂S in the environment, which is present due to some industrial processes, to generate Ag₂S.
2Ag + H₂S → Ag₂S+ H⁺₂
Question 6: How do you prevent metals from getting corrode?
Answer:
Corrosion prevention is critical in order to avoid significant losses. Metals make up the majority of the structures we employ. Bridges, autos, machines, and home items such as window grills, doors, and railway lines are all examples. Electroplating, galvanization, painting and lubrication, and the use of corrosion inhibitors are just a few of the popular methods for preventing corrosion.

Prevention of Rusting of Iron

The loss of iron objects due to rusting has a huge economic impact on the country, and it must be avoided. To keep iron things from rusting, a variety of techniques are employed. To keep air and water out, the majority of the ways require covering the iron piece with something. The following are some of the most prevalent ways to keep iron from rusting:

Rusting of iron can be prevented by painting: Coating the surface of the iron with paint is the most popular way to keep it from rusting. When the paint is placed on the surface of an iron object, it prevents air and moisture from getting into touch with the object, preventing rusting. To prevent rusting, window grills, railings, iron bridges, steel furnishings, railway coaches, and the bodies of automobiles, buses, and trucks, among other things, are all painted on a regular basis.
Rusting of iron can be prevented by applying grease or oil: When grease or oil is placed on the surface of an iron object, air and moisture are kept from coming into touch with it, preventing corrosion. Iron and steel tools and machine parts, for example, are rubbed with grease or oil to prevent corrosion.
Rusting of iron can be prevented by galvanisation: Galvanizing protects articles exposed to excessive moisture, such as roof sheets and pipelines, against rusting. Galvanization is the technique of applying a thin layer of zinc to steel and iron to prevent rust. Galvanised iron is iron that has been zinc-coated. Zinc is more reactive than iron, therefore in the presence of moisture, it interacts with oxygen to generate an invisible layer of zinc oxide that protects it from further rusting. It’s worth noting that even if the zinc coating on galvanised iron products is broken, they remain rust-free. Because zinc is more reactive than iron, this is the case.
Rusting of iron can be prevented by electroplating: Electroplating is another method for keeping items from rusting. In this procedure, noncorroding metals including tin, nickel, and chromium are electroplated on iron. This technique not only keeps the goods from rusting but also improves their beauty. Bathroom fittings and vehicle elements such as bicycle handlebars, car bumpers, and so on are examples of chromium-plated items.
Rusting of iron can be prevented by alloying it to make stainless steel: Stainless steel is created when the iron is alloyed with chromium and nickel. Stainless steel is impervious to rust. Stainless steel cooking utensils, scissors, and medical equipment, for example, do not corrode. Stainless steel, on the other hand, is too expensive to be utilised in big quantities.
Rusting of iron can be prevented by tinning: Tin is non-toxic, and its reactivity is lower than that of iron. Food cans are tinned, which implies that they have a thin layer of tin on them. As a result, when an electroplated thin coating of tin metal is deposited on iron and steel items, the iron and steel objects are protected from rusting. Tin-plated tiffin boxes are utilised because they are non-toxic and do not contaminate the food within.
Rusting of iron can be prevented by Enameling: Enameling is a high-heat procedure that involves fusing powdered glass into a metal substrate. Enamels can be used on a variety of surfaces, including glass and ceramics.

Sample Problems

Question 1: What is the process of rusting iron?

Answer:

Iron rusting is an oxidation reaction. In the presence of water, the iron metal interacts with oxygen in the air to generate hydrated iron (III) oxide, Fe₂O₃.xH₂O. This hydrated iron (III) oxide is referred to as rust. Rust is largely hydrated iron (III) oxide, Fe₂O₃.xH₂O, as a result. Rust is a reddish-brown hue

Question 2: What is rusting of iron called?

Answer:

Rusting is the phenomena of a reddish-brown coating forming on the surface of iron due to the action of wet air, and the reddish-brown coating is referred to as rust.

Question 3: How rusting of iron can be prevented?

Answer:

Rusting of iron can be prevented by

Applying paint

Applying grease or oil

By galvanisation

By electroplating

Using alloying iron to make stainless steel

By tinning

Using Enameling

Question 4: What is rust? Give the equation for the formation of rust?

Answer:

When iron is exposed to air for an extended period of time, it oxidises and develops a reddish-brown iron oxide on the surface. Rust is the name for this reddish-brown material.

Rust is formed via the following equation:

4Fe + 3O₂ +2xH₂O → 2Fe₂O₃.xH₂O

Question 5: How does rust damage iron objects?

Answer:

Rust is permeable and soft, and as it slips off the surface of a rusty iron object, the iron beneath rusts. As a result, iron rust is a constant process that eats away at iron items over time, rendering them worthless. Rusting of iron causes significant damage over time since it is used to build a wide range of structures and commodities, including bridges, grills, railings, gates, and the bodies of cars, buses, trucks, and ships. It goes without saying that we should have a way to keep iron from rusting.

Question 6: What are the conditions necessary for rusting?

Answer:

Many factors contribute to the rusting of iron, including the amount of moisture in the air and the pH of the surrounding environment. The following are a few of these elements.

Moisture: The availability of water in the environment limits the corrosion of iron. The most prevalent cause of rusting is exposure to rain.

The rusting process is accelerated if the pH of the environment around the metal is low. When iron is exposed to acid rain, it rusts more quickly. Iron corrosion is slowed by a higher pH.

Due to the presence of various salts in the water, iron rusts more quickly. Many ions in saltwater speed up the rusting process through electrochemical processes.

Impurity: When compared to iron having a variety of metals, pure iron rusts more slowly.

Question 7: How does rust of iron be a chemical change?

Answer:

Rust is made up of iron oxide (Fe2O3). As a result, rust and iron are not synonymous. Rust isn’t the same thing as the iron it’s deposited on. Because a new component termed “iron oxide” is created during the rusting of iron, it represents a chemical change.

What are Alloys?

An alloy is a mixture of two or more metals or an alloy is a mixture of metal and small amounts of non-metals.

Pure metals are never used in industries for manufacturing purposes. A combination of metals is used to enhance the properties of a single metal and this combination of metals is known as an alloy. It may also contain metal and non-metal. In general, an alloy of metals is made by melting the various metals in the proper proportions and then cooling the mixture to room temperature. An alloy of a metal and a non-metal can be prepared by first melting the metal and then dissolving the non-metal in it, which is followed by cooling it to room temperature. In comparison to metals, alloys have more strength and last longer.

For example-

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Aluminium metal is light but not strong, but an alloy of aluminium with copper, magnesium and manganese is light as well as strong.
Aluminium metal is light but not hard, but an alloy of aluminium with magnesium is light as well as hard.
Iron is the most widely used metal. But it is never used in the pure form because pure iron is very soft and stretches very easily when hot. When a small amount of carbon is mixed with iron, an alloy called steel is obtained. Also when the iron is mixed with chromium and nickel, we get an alloy called stainless steel, which is strong, tough and does not rust at all.

Various Composition of Alloys

Some of the common alloys are Brass, Steel, Stainless steel, Bronze, Solder, amalgam etc. The compositions of various alloys are given below:

Bronze was the first alloy to be discovered. It’s made up of copper and tin. It has a copper content of 90% and a tin content of 10 %. To improve the overall characteristics, very small amounts of zinc, nickel, or manganese may be added.
Copper and zinc are combined to make brass. It contains approximately 80 % copper and 20 % zinc. Other components can be added in smaller amounts. Brass is used to improving copper’s electrical characteristics.
Steel is made by mixing 90 % iron and 1% carbon. It’s more corrosion-resistant and long-lasting.
Stainless steel is made by mixing iron with chromium and nickel. It contains approximately 18 % chromium and 5 % nickel.
Alnico is a metal alloy made up of iron, nickel, cobalt, and aluminium.
Tin and lead alloys are used to make solder. It’s made up of 50 % lead and 50 % tin.
Cast iron is formed by mixing iron with carbon. It contains 96-98% of iron and 2-4% of carbon. Silicon traces may also be discovered.
Sterling silver is made by combining 92.5 % silver and 7.5 % other metals, most commonly copper. If the air contains sulphur compounds, silver will corrode and turn black. Copper or other metals can be blended with silver to make this alloy that reduces tarnishing.
Nickel, chromium, and iron are used to make nichrome. It has a high resistance, melting point, ductility, and other properties. It has a high resistance to electron flow and is difficult to oxidize.
Amalgam is a mercury alloy including one or more additional metals. A solution of sodium metal in liquid mercury metal is known as sodium amalgam.
Gold with a purity of 24 carats is regarded as the purest. Pure gold is very soft due to which it is not suitable for making jewelry. To make gold harder, it is mixed with a small amount of silver or copper. In India gold ornaments are made of 22 carats of gold, which means that 22 parts of pure gold is alloyed with 2 parts of either silver or copper.

Properties of Alloys

Each alloy has certain useful properties. An alloy’s properties are distinct from those of the individual metals from which it is produced. Some properties of alloys are given below.

Alloys are harder than their constituent metals.
Alloys are more resistant to corrosion than pure metals.
Alloys are more durable than the metals they are made from.
The electrical conductivity of alloys is lower than that of pure metals.
Alloys have a lower melting point than the metals from which they are made.
Alloys have greater ductility than their constituent metals.

Uses of Alloys

Alloys are used in a number of ways in our daily lives. Some of the most common uses of alloys are given below.

Brass is used for making cooking utensils, screws, locks, doorknobs, electrical appliances, zippers, musical instruments, decoration and gift items
Bronze is used to make statues, coins, medals, cooking utensils, and musical instruments, among other things.
Alnico is a ferromagnetic substance and is used in permanent magnets.
Solder is used to repair or join two pieces of metals, i.e., it is used to permanently join electrical components.
Surgical devices, musical instruments, cutlery, and jewellery are all made of sterling silver.
Stainless steel is used for the construction of railways, bridges, roads, airports etc. It is also used for making cooking utensils and other products.
Alloys of aluminium are lightweight, so they are used for making bodies of aircraft and their parts.
Because of their high-temperature strength and superplastic behaviour, titanium alloys are widely used in the aerospace industry.
Amalgam is a mercury alloy that is utilized in medicinal procedures. Dentists also use it to fix cavities in teeth.
Certain alloys of gold such as rose gold, are used for jewellery making purposes.

The Iron Pillar at Delhi

The Iron Pillar near Qutub Minar in Delhi is made up of wrought iron which is low carbon steel. It is 8 meters high and 6000 kg in weight. Indian iron craftsmen constructed this pillar in 400 BC. Though wrought iron rusts slowly with time, the ironworkers have developed a process that prevented the wrought iron pillar from rusting even after thousands of years.

The formation of a thin film of magnetic oxide of iron $Fe_{3}O_{4}$ on the surface has prevented corrosion. This thin layer was formed on the surface of the pillar as a result of finishing treatment given to the pillar by painting it with a mixture of different salts, then heating and rapid cooling. This pillar stands in a good condition more than 2000 years after it was made. This pillar has not rusted at all. It suggests that ancient Indians were well-versed in metals and alloys.

Sample Questions

Question 1: What is meant by 22 carats of gold?

Answer:

22 carats of gold, means that 22 parts of pure gold is alloyed with 2 parts of either silver or copper.

Question 2: How is an alloy made?

Answer:

An alloy of metals is prepared by mixing the various metals in molten state in required proportions and then cooling their mixture to the room temperature.

Question 3: How are alloys used in the Aerospace industry?

Answer:

Aluminum is a lightweight metal and its alloys are used in the aerospace industry. These alloys are used for making bodies of aircrafts and to form high strength parts for jet engines. These parts deal with the extremities of temperature, pressure and vibration. They provide high strength and the ability to function at very high temperatures.

Question 4: Which alloy is used by the dentists?

Answer:

Amalgam is an alloy of mercury metal. An amalgam consisting of mercury, silver, tin and zinc is used by dentists for filling in teeth.

Question 5: Why is the Iron Pillar at Delhi still not rusted?

Answer:

The formation of a thin film of magnetic oxide of iron $Fe_{3}O_{4}$ on the surface of the pillar has prevented the rusting of Iron Pillar, as a result of finishing treatment given to the pillar by painting it with a mixture of different salts, then heating and rapidly cooling.

Question 6: What are the constituents of stainless steel?

Answer:

Stainless steel is a mixture of iron with chromium and nickel. It is very strong and do not rust. It is commonly used for making cooking utensils.

What are Metals?

Reactivity Series of Metals

Salient Features of Reactivity Series

Long Tabular Form of the Reactivity Series

Important Uses of Reactivity Series

Reaction Between Metals and Water

Single Displacement Reactions between Metals

Reaction with Acid and Base

Reactivity Series Trick

What is an Ionic Bond?

Formation of ionic bonds

What are Ionic Compounds?

Properties of Ionic Compounds

Sample Questions

Extraction of Metals

Enrichment or Concentration of Ore

Extraction of Metal from Concentrated Ore

Refining of Impure Metals

Sample Questions

Extraction of Highly Reactive Metals

Sample Questions

Extraction of Moderately Reactive Metals

Reduction of Metal Oxides

Extraction of Less Reactive Metals

Sample Questions

What is Electrolytic Refining?

Electrolytic refining of Copper

What happens to the impurities in impure copper that are metallic?

Solved Questions

Prevention of Rusting of Iron

Sample Problems

What are Alloys?

Various Composition of Alloys

Properties of Alloys

Uses of Alloys

The Iron Pillar at Delhi

Sample Questions

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