Metals and Non-metals Class 10 Notes Science Chapter 3
Metals: Physical properties of metals, chemical properties of metals and non-metal oxide.
Metals are malleable and ductile elements that conduct heat and electricity.
Examples include lead (Pb), copper (Cu), gold (Au), platinum (Pt), iron (Fe), aluminium (Al), silver (Ag), calcium (Ca), magnesium (Mg), and others.
Metals are elements that lose electrons to form positive ions. Metals are thus known as Electropositive Elements.
Physical Properties of Metals
Hardness: Except for alkali metals such as sodium, potassium, lithium, and others, most metals are hard. These can be cut with a knife..Metals can be very hard or very soft depending on their composition.
Strength:
The majority of metals are strong and have a high tensile strength. As a result, large structures are made of metals such as copper (Cu) and iron (Fe). (With the exception of sodium (Na) and potassium (K), which are soft metals).
State: Except for mercury, metals are solid at room temperature (Hg).
Sound: Sonorous metals are those that produce a ringing sound. Metallic sound is another name for the sound of metals. This is why metal wires are used in the manufacture of musical instruments.
Conduction: Metals are excellent conductors of heat and electricity. This is why electric wires are made of metals such as copper and aluminium.
Malleability: Metals can be made malleable. Metals can thus be beaten into a thin sheet. Iron is used in the construction of large ships because of this property.
Ductility: Metallic ductility: Metals are ductile. Metals can thus be drawn into thin wires. A wire is made of metals because of this property.
Melting and Boiling Point: Melting and boiling points: Metals have high melting and boiling points in general. (With the exception of sodium and potassium metals, which have low melting points.)
Density:Most metals have a high density.
Colour: The majority of metals are grey in colour. However, gold and copper are exceptions.
Chemical Properties of Metals
1. Reaction with oxygen: When most metals react with oxygen, they form metal oxides.
example :- Metal Oxide = Metal + Oxygen
Potassium Metal Reaction with Oxygen: When potassium metal reacts with oxygen, potassium oxide is formed.
Sodium's reaction with oxygen results in the formation of sodium oxide.
Alkali-metals include lithium, potassium, sodium, and other substances. Oxygen and alkali metals react violently.
Reaction of Copper metal with Oxygen: Copper does not react with oxygen at ambient temperature, but when it is burned in the presence of air, it produces oxide.
Even at very high temperatures, the oxygen in the air does not combine with silver, gold, or platinum. The least reacting are them.
2. Reaction of metals with water: Metals form respective hydroxide and hydrogen gas when reacting with water. Metal + Water → Metal hydroxide + Hydrogen Most metals do not react with water. However, alkali metals react violently with water.
Reaction of Sodium metal with Water: When sodium metal reacts with water, it produces sodium hydroxide and releases hydrogen gas as well as a lot of heat.
Reaction of Calcium metal with Water: When calcium reacts with water, it produces calcium hydroxide, hydrogen gas, and heat.
Reaction of Magnesium metal with Water: Magnesium metal slowly reacts with water, forming magnesium hydroxide and hydrogen gas.
Reaction of Aluminium metal with Water: The reaction of aluminium metal with cold water is far too slow to be noticed. However, when steam is passed over aluminium metal, it produces aluminium oxide and hydrogen gas.
2Al + 3H2O → Al2O3 + 2H2
Reaction of Zinc metal with Water: When steam is passed over zinc metal, it produces zinc oxide and hydrogen gas. Cold water has no effect on zinc.
Reaction of Iron with Water: The reaction of iron with cold water is very slow and takes a long time to notice. When iron reacts with moisture in the atmosphere, it produces rust (iron oxide). By passing steam over iron metal, iron oxide and hydrogen gas are formed.
Both calcium (Ca) and magnesium (Mg) are heavier than water but still float over it: Calcium and magnesium both float on the water's surface because when these metals react with water, hydrogen gas is produced. It takes the form of bubbles that adhere to the metal surface. As a result, they float over it.
Other metals typically do not or only slowly react with water. Steam has no effect on lead, copper, silver, or gold. Thus, the reactivity of different metals towards water can be written in the following order: K > Na > Ca > Mg > Ae > Zn > Fe > Pb > Cu > Ag > Au
3. Reaction of metals with dilute acid: When metals react with dilute acid, they form salts. Metal salt + Hydrogen Metal salt + dil. acid
Reaction of Sodium metal with dilute hydrochloric acid: When sodium metal reacts with dilute hydrochloric acid, it produces sodium chloride and hydrogen gas.
Reaction of Magnesium metal with dilute hydrochloric acid: When magnesium reacts with dilute hydrochloric acid, magnesium chloride and hydrogen gas are formed.
Reaction of Zinc with dilute sulphuric acid: When zinc reacts with dilute sulfuric acid, zinc sulphate and hydrogen gas are formed. In the laboratory, this method is used to generate hydrogen gas.
Hydrogen (H2) gas is not evolved when metal is treated with nitric acid (HNO3):
Nitric acid is a strong oxidising agent that oxidises the hydrogen gas (H2) liberated into water (H2O) and reduces itself to some nitrogen oxides such as nitrous oxide (N2O)3, nitric oxide (NO), and nitrogen dioxide (NO2).
Noble metals include copper, gold, and silver. These are not affected by water or dilute acids.
The order of metal reactivity towards dilute hydrochloric acid or sulphuric acid is as follows;
K > Na > Ca > Mg > Al > Zn > Fe > Cu > Hg > Ag
Metal Oxides
Chemical Properties: Metal oxides occur naturally. Metal oxide aqueous solution turns red litmus blue.
Reaction of Metal oxides with Water: The majority of metal oxides are insoluble in water. Water dissolves alkali metal oxides. When alkali metal oxides are dissolved in water, they produce a strong base.
Reaction of Sodium oxide with Water: When sodium oxide reacts with water, it produces sodium hydroxide.
Reaction of Potassium oxide with Water: When potassium oxide reacts with water, it produces potassium hydroxide.
Reaction of Zinc oxide and Aluminium oxide: Water does not dissolve aluminium oxide or zinc oxide. Aluminium oxide and zinc oxide are both amphoteric. Amphoteric substances exhibit both acidic and basic properties. It reacts with bases like acids and with acids like bases.
When zinc oxide reacts with sodium hydroxide, it takes on the properties of an acid. This reaction produces sodium zincate and water.
When zinc oxide reacts with acid, it takes on the properties of a base. When zinc oxide reacts with hydrochloric acid, it produces zinc chloride and water.
Similarly, aluminium oxide behaves as a base when reacting with acid and as an acid when reacting with a base.
When aluminium oxide reacts with sodium hydroxide, it produces sodium aluminate as well as water.
When aluminium oxide reacts with hydrochloric acid, it produces aluminium chloride as well as water.
Reactivity Series of Metals: The Reactivity Series refers to the order of metal intensity or reactivity. Moving from top to bottom in the given reactivity series, the reactivity of elements decreases.
Copper, gold, and silver are at the bottom of the reactivity series and thus the least reactive. These metals are referred to as Noble metals. Potassium is at the top of the series and thus the most reactive element.
Some metals' reactivity is listed in descending order:
K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu
4. Reaction of metals with solution of other metal salts: The displacement reaction occurs when metals react with the solution of another metal salt. In this reaction, more reactive metal displaces the less reactive metal from its salt.
Metal A + Metal B salt Metal A + Metal B salt
Iron, for example, displaces copper from copper sulphate solution.
Similarly, aluminium and zinc remove copper from copper sulphate solutions.
In all the above examples, iron, aluminium and zinc are more reactive than copper. This is why they displace copper from its salt solution.
When copper is dipped in silver nitrate solution, it displaces silver and forms copper nitrate.
Copper is more reactive than silver in the reaction and thus displaces silver from the silver nitrate solution.
Because silver is less reactive than copper and cannot displace copper from its salt solution, it does not react with copper sulphate solution.
Similarly, no reaction occurs when gold is dipped in a solution of copper nitrate because copper is more reactive than gold.
Similarly, no reaction occurs when copper is dipped in an aluminium nitrate solution because copper is less reactive than aluminium.
Non-Metals: Nonmetal physical properties, nonmetal chemical properties, nonmetal oxides Metal and nonmetal reactions, Ionic bonds and ionic bond formation. Non-metals are elements that are neither malleable nor ductile and do not conduct electricity. Examples: Carbon (C), Sulphur (S), Phosphorous (P), Silicon (Si), Hydrogen (H), Oxygen (O), Nitrogen (N), Chlorine (Cl), Bromine (Br), Neon (Ne) and Argon (Ar) etc. Non-metals are elements that gain an electron to form negative ions. As a result, nonmetals are also referred to as Electronegative Elements.
Physical properties of non-metals
Hardness: Non-metals are not hard rather they are generally soft. The diamond, on the other hand, is an exception; it is the hardest naturally occurring substance.
State: Nonmetals can exist as solids, liquids, or gases.
Lustre: Nonmetals are dull in appearance. The exceptions are diamond and iodine.
Sonority: Non-metals are not sonorous, which means they do not produce a typical sound when struck.
Conduction:
Malleability and ductility: Nonmetals conduct heat and electricity poorly. Graphite, a carbon allotrope, is an exception and a good conductor of electricity.Brittleness is a property of nonmetals.
Melting and boiling point: Nonmetals have low melting and boiling points in general.
Density: The majority of nonmetals have a low density.
Colour: Nonmetals come in a variety of colours.
- Carbon, as graphite, is a nonmetal that conducts electricity.
- Iodine is a nonmetal with a lustrous, gleaming surface.
- Diamond is a nonmetal that is extremely hard due to the presence of carbon.
- Diamond is a nonmetal with an extremely high melting and boiling point.
Chemical properties of Non-metals
1. Reaction of Non-metals with Oxygen: Non-metals form respective oxide when reacting with oxygen.
Non-metal + Oxygen → Non-metallic oxide
When carbon reacts with oxygen, carbon dioxide is formed along with the production of heat.
When carbon is burnt in an insufficient supply of air, it forms carbon monoxide. Carbon monoxide is a toxic substance. Inhaling of carbon monoxide may prove fatal.
Sulphur gives sulphur dioxide when reacting with oxygen. Sulphur catches fire when exposed to air.
When hydrogen reacts with oxygen it gives water.
Non-metallic Oxide: Non-metallic oxides are acidic in nature. The solution of non-metal oxides turns blue litmus red.
Carbon dioxide gives carbonic acid when dissolved in water.
Sulphur dioxide gives sulphurous acid when dissolved in water.
Sulphur dioxide gives sulphuric acid when reacts with oxygen.
2. Reaction of Non-metal with Chlorine: Non-metal gives respective chloride when they react with chlorine gas.
Non-metal + Chlorine → Non-metal chloride
Hydrogen gives hydrogen chloride and phosphorous gives phosphorous trichloride when reacting with chlorine.
3. Reaction of Non-metals with Hydrogen: Non-metals reactive with hydrogen to form covalent hydrides.
Non-metal + Hydrogen → Covalent Hydride
Sulphur combines with hydrogen to form a covalent hydride is called Hydrogen sulphide.
Nitrogen combines with hydrogen in presence of an iron catalyst to form covalent hydride ammonia.
Non-metals do not react with water (or steam) to evolve Hydrogen gas.
Non-metals do not react with dilute acids.
4. Reaction of Metal and Non-metal: Many metals form ionic bonds when they react with non-metals. Compounds so formed are known as Ionic Compounds.
Ions: Positive or negative charged atoms are known as ions. Ions are formed because of loss or gain of electrons. Atoms form ions obtain by the electronic configuration of the nearest noble gas.
Positive ion: A positive ion is formed because of the loss of electrons by an atom.
Following are some examples of positive ions:
Sodium forms sodium ion because of the loss of one electron. Because of the loss of one electron, one positive charge comes over sodium.
Magnesium forms positive ion because of the loss of two electrons. Two positive charges come over magnesium because of loss of two electrons.
Negative ion: A negative ion is formed because of the gain of an electron.
Some examples are given below :
Chlorine gains one electron in order to achieve a stable configuration. After the loss of one electron, chlorine gets one negative charge over it forming chlorine ion.
Ionic Bonds: Ionic bonds are formed because of transfer of electrons from metal to non¬metal. In this course, metals get positive charge because of transfer of electrons and non-metal gets negative charge because of acceptance of electrons. In other words, bond formed between positive and negative ion is called Ionic Bond.
Since, a compound is electrically neutral, so to form an ionic compound, negative and positive both ions must be combined.
Some examples are given below:
Formation of Sodium Chloride (NaCl): In sodium chloride, sodium is a metal (alkali metal) and chlorine is a non-metal.
Atomic number of sodium = 11
Electronic configuration of sodium : 2, 8, 1
Number of electrons in outermost orbit = 1
Valence electrons = Electrons in outermost orbit = 1
Atomic number of chlorine = 17
Electronic configuration of chlorine : 2, 8, 7
Electrons in outermost orbit = 7
Therefore, valence electrons = ?
Sodium has one valence electron and chlorine has seven valence electrons. Sodium requires losing one electron to obtain stable configuration and chlorine requires gaining one electron in order to obtain stable electronic configuration. Thus, in order to obtain stable configuration, sodium transfers one electron to chlorine. After loss of one electron, sodium gets one positive charge (+) and chlorine gets one negative charge after gain of one electron. Sodium chloride is formed because of transfer of electrons. Thus, ionic bond is formed between sodium and chlorine. Since, sodium chloride is formed because of ionic bond, thus, it is called Ionic compound. In similar way, potassium chloride (KCl) is formed.
Properties of Ionic compound
Ionic compounds are solid. Ionic bond has a greater force of attraction because of which ions attract each other strongly. This makes ionic compounds solid.
Ionic compounds are brittle.
Ionic compounds have high melting and boiling points because force of attraction between ions of ionic compounds is very strong.
Ionic compounds generally dissolve in water.
Ionic compounds are generally insoluble in organic solvents; like kerosene, petrol, etc.
Ionic compounds do not conduct electricity in the solid state.
The solution of ionic compounds in water conduct electricity. This happens because ions present in the solution of ionic compound facilitate the passage of electricity by moving towards opposite electrodes.
Ionic compounds conduct electricity in the molten state.
Occurrence and Extraction of Metals: Minerals, ores, extraction of metals of least reactivity, extraction of metals of middle reactivity, extraction of metals of high reactivity, refining or purification of metals and corrosion.
Occurrence and Extraction of Metals:
Source of metal: Metals occur in Earth’s crust and in seawater; in the form of ores. Earth’s crust is the major source of metal. Seawater contains many salts such as sodium chloride, magnesium chloride, etc.
Mineral: Minerals are naturally occurring substances which have a uniform composition.
Ores: The minerals from which a metal can be profitably extracted are called Ores.
Metals found at the bottom of reactivity series are least reactive and they are often found in nature in free-state; such as gold, silver, copper, etc. Copper and silver are also found in the form of sulphide and oxide ores.
Metals found in the middle of reactivity series, such as Zn, Fe, Pb, etc. are usually found in the form of oxides, sulphides or carbonates.
Metals found at the top of the reactivity series are never found in free-state as they are very reactive, example; K, Na, Ca, Mg and Al, etc.
Many metals are found in the form of oxides because oxygen is abundant in nature and is very reactive.
Extraction of Metals: Metals can be categorised into three parts on the basis of their reactivity: Most reactive, medium reactive and least reactive.
The three major steps involved in the extraction of a metal from its ore are
Concentration or enrichment of ores.
Conversion of concentrated ore into crude metal and,
Refining of impure or crude metal.
1. Concentration of Ores: Removal of impurities, such as soil, sand, stone, silicates, etc. from mines ore is known as Concentration of Ores.
Ores which are mined often contain many impurities. These impurities are called gangue. First of all, concentration is done to remove impurities from ores. The concentration of ores is also known as enrichment of ores. Process of concentration depends upon physical and chemical properties of ores. Gravity separation, electromagnetic separation, froth flotation process, etc. are some examples of the processes which are applied for concentration of ores.
2. Conversion of Concentrated Ore into Crude Metal
Conversion of metals ores into oxides: It is easy to obtain metals from their oxides. So, ores found in the form of sulphide and carbonates are first converted to their oxides by the process of roasting and calcination. Oxides of metals so obtained are converted into metals by the process of reduction.
Roasting: Heating of sulphide ores in the presence of excess air to convert them into oxides is known as Roasting.
Calcination: Heating of carbonate ores in the limited supply of air to convert them into oxides is known as Calcination.
3. Reduction: Heating of oxides of metals to turn them into metal is known as Reduction.
(i) Extraction of Metals of Least Reactivity: Mercury and copper, which belong to the least reactivity series, are often found in the form of their sulphide ores. Cinnabar (HgS) is the ore of mercury. Copper glance (Cu2S) is the ore of copper.
Extraction of Mercury Metal: Cinnabar (HgS) is first heated in air. This turns HgS (mercury sulphide or cinnabar) into HgO (mercury oxide) by liberation of sulphur dioxide. Mercury oxide so obtained is again heated strongly. This reduces mercury oxide to mercury metal.
Extraction of Copper Metal: Copper glance (Cu2S) is roasted in the presence of air. Roasting turns copper glance (ore of copper) into copper (l) oxide. Copper oxide is then heated in the absence of air. This reduces copper (l) oxide into copper metal.
(ii) Extraction of Metals of Middle Reactivity: Iron, zinc, lead, etc. are found in the form of carbonate or sulphide ores. Carbonate or sulphide ores of metals are first converted into respective oxides and then oxides are reduced to respective metals.
Extraction of Zinc: Zinc blende (ZnS: zinc sulphide) and smithsonite or zinc spar or calamine (ZnCO3: zinc carbonate) are ores of zinc. Zinc blende is roasted to be converted into zinc oxide. Zinc spar is put under calcination to be converted into zinc oxide.
Zinc oxide so obtained is reduced to zinc metal by heating with carbon (a reducing agent).
Extraction of Iron from Haematite (Fe2O3): Haematite ore is heated with carbon to be reduced to iron metal.
Extraction of Lead from Lead oxide: Lead oxide is heated with carbon to be reduced to lead metal.
Reduction of Metal oxide by Heating with Aluminium: Metal oxides are heated with aluminium (a reducing agent) to be reduced to metal. Following is an example: Manganese dioxide and copper oxide are reduced to respective metals when heated with aluminium.
Thermite Reaction: Ferric oxide; when heated with aluminium; is reduced to iron metal. In this reaction, a lot of heat is produced. The thermite reaction is used in the welding of electric conductors, iron joints, etc. such as joints in railway tracks. This is also known as Thermite Welding (TW).
(iii) Extraction of Metals of High Reactivity: Metals of high reactivity; such as sodium, calcium, magnesium, aluminium, etc. are extracted from their ores by electrolytic reduction. These metals cannot be reduced using carbon because carbon is less reactive than them.
Electrolytic Reduction: Electric current is passed through the molten state of metal ores. Metal being positively charged is deposited over the cathode.
Example: When an electric current is passed through molten state or solution of sodium chloride, sodium metal gets deposited over the cathode.
Metals obtained from the process of electrolytic reduction are pure in form.
4. Refining or purification of metals: Metals extracted from various methods contains some impurities, thus, they are required to be refined. Most of the metals are refined using electrolytic refining.
Electrolytic Refining: In the process of electrolytic refining, a lump of impure metal and a thin strip of pure metal are dipped in the salt solution of metal to be refined. When an electric current is passed through the solution, pure metal is deposited over a thin strip of pure metal
from a lump of impure metal. In this, impure metal is used as anode and pure metal is used as a cathode.
Electrolytic Refining of Copper: A lump of impure copper metal and a thin strip of pure copper are dipped in the solution of copper sulphate. Impure lump of metal is connected with the positive pole and thin strip of pure metal is connected with negative pole. When electric current is passed through the solution, pure metal from anode moves towards cathode and is deposited over it. Impurities present in metal are settled near the bottom of anode in the solution. Settled impurities in the solution are called Anode Mud.
5. Corrosion: Most of the metals keep on reacting with the atmospheric air. This leads to the formation of a layer over the metal. In the long run, the underlying layer of metal keeps on getting lost due to conversion into oxides or sulphides or carbonate, etc. As a result, the metal gets eaten up. The process is called Corrosion.
Rusting of Iron: Rusting of iron is the most common form of corrosion. When iron articles like the gate, grill, fencing, etc. come in contact with moisture present in the air, the upper layer of iron turns into iron oxide. Iron oxide is brown-red in colour and is known as Rust. The phenomenon is called Rusting of Iron.
If rusting is not prevented in time, the whole iron article would turn into iron oxide. This is also known as Corrosion of Iron. Rusting of iron gives a huge loss every year.
Prevention of Rusting: For rusting, iron must come in contact with oxygen and water. Rusting is prevented by preventing the reaction between atmospheric moisture and the iron article. This can be done by:
Painting
Greasing
Galvanization
Electroplating
Alloying
6. Alloys: The homogeneous mixture of two or more metals, or a metal and a non-metal is called Alloy.
Types of alloys :
Ferrous alloys: An alloy in which iron (Fe) is present. For example : manganese steel (Fe = 86% ; Mn = 13% ; C = 1%) and Nickle steel (Fe = 98% ; Ni = 2%).
Non-ferrous alloys: An alloy does not contain iron. For example : Brass (Cu = 80% ; Zn = 20%), and Bronze (Cu = 90% ; Sn = 10%).
Amalgams: An alloy in which mercury (Hg) is present. For example Sodium amalgams [Na(Hg)] and Zinc amalgams [Zn(Hg)].
Properties of an Alloy
Alloys are stronger than the metal from which they are obtained.
It is harder than the constituent metals.
More resistance to corrosion.
The melting point of alloys is lower than the constituent metals.
Example: Solder [Sn(80%) + Pb(50%)] has lower m. p. than Pb and Sn.The electrical conductivity of alloys is lower than the constituent metals.
Some examples of Alloys:
Brass: [80% Cu + 20% Zn ]
Bronze: [90% Cu + 20% Sn]
Solder: [50% Pb + 50% Sn]
Duralumin: [95% Al + 4% Cu + 0.5% Mg + 0.5 Mn]
Steel: [99.95% Fe + 0.05% C]
Stainless steel: [74% Fe + 18% Cr + 8% Ni]
Magnesium: [95% Al + 5% Mg]
German Silver: [60% Cu + 20% Zn + 20% Ni]
Alloys of Gold: Pure gold is said to be of 24 carats. Gold is alloyed with a small amount of silver or copper to make it hard.
Metals and Non-metals:
Chemical Properties of Metals and Non-metals.
The reaction of metals with oxygen. Metals form their oxides when reacting with oxygen.
Metal + Oxygen → Metal oxide
Metal oxides are basic in nature. Example, Reaction of Iron metal with oxygen When iron reacts with moist air, it forms rust.
Rust is iron oxide. Articles made of iron, such as grills, fencing, etc. are getting rusted because of reaction with moist air.
Iron (Fe) + Water (H2O) + Oxygen (O2) → Fe3O4n.H2O (Iron II, III) Oxide (Rust)
Rust is reddish brown in colour and is iron oxide. Iron oxide is basic in nature. It turns red litmus blue.
Rusting of iron can be prevented:
by galvanizing the iron articles with zinc coating.
by painting and applying grease on the articles.
The reaction of Magnesium metal with oxygen: When magnesium is burnt in air, it forms magnesium oxide. Burning in the air means reaction with oxygen.
Magnesium + Oxygen (O2) → MgO (Magnesium oxide)
Magnesium oxide forms magnesium hydroxide with water. The solution of Magnesium oxide turns red litmus paper blue. This means magnesium oxide is basic in nature.
MgO + H2O → Mg(OH)2 (Magnesium Hydroxide)
The reaction of Non-metals with oxygen: Non-metals form their oxides when they react with oxygen.
Non-metal + Oxygen → Non-metal oxide
Non-metal oxides are acidic in nature.
Example., Reaction of sulphur with oxygen.
When sulphur is burnt in air, it forms sulphur dioxide.
Sulphur + Oxygen (O2) → SO2 (Sulphur dioxide)
The solution of sulphur dioxide turns blue litmus paper red. Sulphur dioxide forms sulphurous acid when dissolved in water. Thus, sulphur dioxide is acidic in nature.
SO2 + H2O → Sulphurous acid (H2SO3)
The reaction of carbon with oxygen—When carbon is burnt in air, it forms carbon dioxide.
Carbon + Oxygen (O2) → CO2 (Carbon dioxide)
You can observe that when coal (carbon) is burnt it forms smoke, which contains carbon dioxide. Carbon dioxide is acidic in nature. The solution of carbon dioxide in water turns blue litmus paper red.
CO2 + H2O → Carbonic acid (H2CO3)
The reaction of Metals and Non-metals with water: Generally, metals form respective hydroxides when they react with water.
Metal + Water → Metal hydroxide
The reaction of sodium metal with water: Sodium metal vigorously reacts with water and forms sodium hydroxide along with a lot of heat.
Na + H2O → NaOH (Sodium hydroxide) + H2 (Hydrogen) + Heat
Non-metals generally do not react with water. Rather some non-metals which react with air vigorously are stored in water. The reaction of metals and non-metals with dilute acid. Metals give hydrogen gas when they react with dilute acid.
Metal + Acid → Hydrogen gas + Salt
The reaction of zinc with dilute acid. Zinc gives hydrogen gas along with zinc chloride when it reacts with hydrochloric acid. Similarly, zinc gives hydrogen gas along with zinc sulphate when it reacts with sulphuric acid. This method is used to produce hydrogen gas in the laboratory.
Zn + H2SO4 (Sulphuric acid) → ZnSO4 (Zinc sulphate) + H2 (Hydrogen)
The reaction of Aluminium with dilute acid. Aluminium gives hydrogen gas along with aluminium chloride when it reacts with dilute hydrochloric acid.
2Al + 6HCl (Hydrochloric acid) → 2AlCl3 (Aluminium Chloride) + 3H2 (Hydrogen)
Copper does not react with dilute sulphuric acid even on heating, but it reacts with concentrated sulphuric acid. Copper, silver and gold are considered as noble metals as do not react with dilute acid.
Generally, non-metals do not react with dilute acid.
The reaction of metals and non-metals with the base. Metals give hydrogen gas when they react with a base.
Metal + Base → Hydrogen gas + Salt
The reaction of aluminium metal with sodium hydroxide.
Al + NaOH (Sodium hydroxide) → NaAlO2 (Sodium aluminate) + H2 (Hydrogen)
Aluminium metal forms hydrogen gas and sodium aluminate when it reacts with sodium hydroxide. Similarly, zinc gives sodium zincate and hydrogen gas when it reacts with sodium hydroxide.
Displacement Reaction: When a more reactive metal reacts with the salt solution of less reactive metal, more reactive metal displaces the less reactive metal from its solution.
Metal A + Salt Solution of metal B → Salt Solution of metal A + metal B
In the above equation, metal A is more reactive than metal B.
Example., When aluminium metal is dipped in the solution of copper sulphate, it forms aluminium sulphate and copper.
Al + CuSO4 (Copper sulphate) → Al2(SO4)3 (Aluminium Sulphate) + Cu (Copper)
In the above reaction, aluminium is more reactive than copper, that is why it replaces copper from the solution of copper sulphate.
When copper metal is dipped in the solution of aluminium nitrate, no reaction takes place. Because copper is less reactive than aluminium.
Roasting and Calcination:
