Lesson 3 - Metals and Non - metals-Class 10 Science Notes
Metals and Non - metals
Metals
Metals are the elements that conduct heat and electricity and are malleable and ductile.
Example:- Iron (Fe), Aluminium (Al), Silver (Ag), Copper (Cu), Gold (Au), Platinum (Pt), Sodium (Na), Potassium (K), Calcium (Ca), Lead (Pb), Magnesium (Mg), Lithium (Li) etc.
Metals are the elements which form positive ions by losing electrons. Thus metals are known as Electropositive elements.
Physical Properties of Metals
1. Metals are hard and shiny.
Most metals are hard, except alkali metals, such as sodium, potassium, lithium etc. are very soft metals. These can be cut by using a knife.
2. Metals are strong and tensile.
Most of the metals are strong and have high tensile strength. Because of this, big structures are made using metals, such as copper (Cu) and Iron (Fe). Except sodium (Na) and potassium (K) which are soft metals.
3. Metals are solid.
Metals are solid at room temperature except for Mercury (Hg), which is liquid at room temperature.
4. Metals conduct heat and electricity.
Metals are a good conductor of heat and electricity. This is the cause that electric wires are made of metals like copper and aluminium. Lead and Mercury are poor conductors of heat.
5. Metals are Sonorous.
Metals are sonorous because they produce ringing sound. Sound of metals is also known as metallic sound. This is the cause that metal wires used in making instruments.
6. Metals are melliable.
This means metals can be beaten into a thin sheet. Because of this property, iron is used in making big ships.
7. Metals are ductile.
This means metals can be drawn into thin wire. Because of this property, a wire is made of metals.
8. Metals have high melting point and boiling point.
Metals have generally high melting and boiling point except Sodium (Na) and Potassium (K).
9. Metals have high density.
10. Most metals are grey in colour. But gold and copper are exceptions.
Chemical Properties of Metals
1. Reaction with Oxygen
Most of the metals form respective metal oxides when reacting with oxygen.
Metal + Oxygen → Metal Oxide
Example:-
1. Potassium and Sodium metal forms potassium oxide when reacts with oxygen.
4K + O2 → 2K2O
4Na + O2 → 2K2O
Sodium (Na), Potassium (K), Lithium (Li) etc are known as Alkali-metals. Alkali metals react vigorously with oxygen.
Why Sodium metal kept immersed in Kerosene oil?
Sodium and Potassium react so vigorously that they catch fire if kept in the open. Hence, to protect them and to prevent accidental fires, they are kept immersed in kerosene oil.
Some metal oxide form protective layer.
1. At ordinary temperature, the surfaces of metals such as Magnesium (Mg), Aluminum (Al), Zinc (Zn), Lead (Pb) etc., are covered with a thin layer of oxide. The protective oxide layer prevents the metal further oxidation.
2Mg + O2 → 2MgO
4Al + 3O2 → 2Al2O3
2Zn + O2 → 2ZnO
2Pb + O2 → 2PbO
2. Copper does not react with oxygen at room temperature but when burnt in air, it gives oxide.
2Cu + O2 → 2CuO
3.Iron (Fe) does not burn on heating but iron filings burn vigorously when sprinkled in the flame of the burner.
4Fe + 3O2 → 2Fe2O3
4. Silver (Ag), Gold (Au) and Platinum (Pt) do not combine with the oxygen of air even at high temperature. They are least reactive.
2. Reaction of metals with water
Reaction of metal with cold water
Metals react with water and produce Metal hydroxide and hydrogen gas.
Metal + Water → Metal hydroxide + Hydrogen
Example:-
1. Sodium and Potassium metals react violently with cold water and makes Sodium and Potassium hydroxide and hydrogen gas. The reaction is exothermic so the evolved hydrogen gas catches fire.
2K (s) + 2H2O (l) → 2KOH (aq) + H2 (g) + heat energy
2Na(s) + 2H2O (l) → 2NaOH (aq) + H2 (g) + heat energy
2. Calcium forms calcium hydroxide and hydrogen gas when react with water. The reaction of calcium with water is less violent. The heat is evolved is not sufficient for the hydrogen to catch fire.
Ca (s) + 2H2O (l) → Ca(OH)2 (aq) + H2 (g)
Calcium starts floating because the bubbles of hydrogen gas formed stick to the surface of the metal.
Reaction of metal with hot water
3. Magnesium does not react with cold water. It reacts with hot water to form magnesium hydroxide and hydrogen. It also starts floating due to the bubbles of hydrogen gas sticking to its surface.
Mg (s) + H2O (l) → Mg(OH)2 (aq) + H2 (g)
Reaction of metal with steam
4. When steam is over magnesium metal, magnesium oxide and hydrogen gas are formed.
Mg (s) + H2O (g) → MgO (s) + H2 (g)
5. Metals like aluminium, iron and zinc do not react with cold or hot water. But they react with steam to form the metal oxide and hydrogen.
2Al (s) + 3H2O (g) → Al2O3 (s) + 3H2 (g)
3Fe (s) + 4H2O (g) → Fe3O4 (s) + 4H2 (g)
Zn (s) + H2O (g) → ZnO (s) + H2 (g)
(**Metals such as lead, copper, silver and gold do not react with water at all.)
The order of reactivity of different metals towards water may be written as-
K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu > Ag > Au
3. Reaction of metal with dilute acid
Metals form respective metal salt when reacting with dilute acid.
Metal + dilute acid → metal salt + hydrogen
Example:- Reaction with hydrochloric acid
1. Sodium and Magnesium metal gives sodium and magnesium chloride and hydrogen gas when react with dilute hydrochloric acid.
2Na (s) + 2HCl (aq) → 2NaCl (aq) + H2 (g)
Mg (s) + 2HCl (aq) → MgCl2 (aq) + H2 (g)
Reaction with sulphuric acid
2. Zinc sulphate and hydrogen gas are formed when zinc reacts with dilute sulphuric acid. This method is used in laboratory to produce hydrogen gas.
Zn (s) + H2SO4 (aq) → ZnSO4 (aq) + H2 (g)
Hydrogen gas is not evolved when metal is treated with nitric acid (HNO3)
3. Nitric acid is strong oxidizing agent and it oxidizes the hydrogen gas (H2) liberated into water (H2O) and itself get reduced to some oxide of nitrogen like nitrus oxide (N2O), nitric oxide (NO) and nitrogen dioxide (NO2).
Zn (s) + 4HNO3 (aq) → Zn(NO3)2 (aq) + 2NO2 (g) + 2H2O (l)
3Cu (s) + 8HNO3 (aq) → 3 Cu(NO3)2 (aq) + 2NO (g) + 4H2O (l)
Exception:- Magnesium and Maganese react with very dilute nitric acid to liberate hydrogen gas.
Mg (s) + 2HNO3 (aq) → Mg(NO3)2 (aq) + H2 (g)
Mn (s) + 2HNO3 (aq) → Mn(NO3)2 (aq) + H2 (g)
(**Copper, gold and silver are known as noble metals. These do not react with water or dilute acids.)
The order of reactivity of metal towards dilute hydrochloric acid or sulphuric acid is as-
K > Na > Ca > Mg > Al > Zn > Fe > Cu > Hg > Ag
Aqua Regia
Aqua regia is a freshly prepared mixture of concentrated hydrochoric acid and nitric acid in the ratio of 3:1.
It can dissolve gold, even though neither of these acids can do so alone. Aqua regia is a highly corrosive, fuming liquid. It is one of the few reagents that is able to dissolve gold and platinum.
Reactivity Series of Metals
The order of intensity or reactivity of metal is known as Reactivity Series. Reactivity of elements decreases on moving from top to bottom in the given reactivity series.
In the reactivity series, copper, gold and silver are at the bottom and hence, least reactive. These metals are known as Noble metals. Potassium is at the top of the series and hence, most reactive.
Reactivity of some metals are given in descending order-
K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au
4. Reaction of metal with solution of other metal salt
Reaction of metals with the solution of other metal salt is displacement reaction. In this reaction, more reactive metal displaces the less reactive metal form its salt.
Metal A + Salt of metal B → Salt of metal A + metal B
Example:-
1. Iron displaces copper from copper sulphate solution.
Fe + CuSO4 → FeSO4 + Cu
2. Aluminium displaces copper form copper sulphate solution.
2Al + 3CuSO4 → Al2(SO4)3 + 3Cu
3. Zinc displaces copper form copper sulphate solution.
Zn + CuSO4 → ZnSO4 + Cu
4. When copper is dipped in the solution of silver nitrate, it displaces silver and forms copper nitrate.
Cu + 2AgNO3 → Cu(NO3)2 + 2Ag
In the reaction, copper is more reactive than silver and hence, displaces silver form silver nitrate solution.
5. Silver metal does not react with copper sulphate solution because silver is less reactive than copper and not able to displace copper form its salt solution.
Ag + CuSO4 → No reaction
Similarly, when gold is dipped in the solution of copper nitrate, no reaction takes place because copper is more reactive than gold.
Au + CuSO4 → No reaction
and, no reaction takes place when copper is dipped in the solution of aluminium nitrate because copper is less reactive than aluminium.
Cu + Al(NO3)3 → No reaction
Metal Oxides
Chemical properties
Metal oxides are basic in nature. The aqueous solution of metal oxides turns red litmus blue.
1. Reaction of metal oxides with water
Most of the meta oxides are insoluble in water. Alkali metal oxides are soluble in water. Alkali metal oxides give strong base when dissolved in water.
Example:-
1. Sodium oxide gives sodium hydroxide when react with water.
Na2O + H2O → 2NaOH
2. Potassium oxide gives potassium hydroxide when reacts with water.
K2O + H2O → 2KOH
2. Amphoteric Oxides
Some metal oxides such as aluminium oxide, zinc oxide, etc. show both acidic as well as basic behaviour. Such metal oxides which react with both acids as well as bases to produce salts and water are known as amphoteric oxides. It reacts with base like acid and reacts with an acid like a base.
Example:- Reaction of Metal oxides with acids
1. Aluminium oxide reacts with hydrochloric acid and produces a salt aluminium chloride and water.
Al2O3 + 6HCl → 2AlCl3 + 3H2O
2. Zinc oxide gives zinc chloride and water on reaction with hydrochloric acid.
ZnO + 2HCl → ZnCl2 + H2O
Reaction of Metal oxides with bases
3. Aluminium oxide reacts with sodium hydroxide produces sodium aluminate and water.
Al2O3 + 2NaOH → 2NaAlO2 + H2O
4. Zinc oxide reacts with sodium hydroxide produces sodium zincate and water.
ZnO + 2NaOH → Na2ZnO2 + H2O
3. Solubility of metal oxide in water
Most metal oxides are insoluble in water but some of these dissolve in water to form alkalis. The dissolving of sodium oxide and potassium oxide in water gives sodium hydroxide alkalis and potassium hydroxide alkalis respectively.
Na2O (s) + H2O (l) → 2NaOH (aq)
K2O (s) + H2O (l) → 2KOH (aq)
Non - Metals
Non-metals are the elements that do not conduct electricity and are neither malleable nor ductile.
Example:- 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 the elements which form negative ions by gaining an electron. Thus, non - metals are also known as Electromagnetic elements.
Physical Properties of non-metals
1. Non-Metals are not hard and shiny
Non-metals are not hard rather they are generally soft. But the diamond is an exception. It is hardest and shiny naturally occurring substance. Iodine is lustrous having a shiny surface.
2. Non-metals are lustrous
Non-metals have a dull appearance. Diamond and iodine are exceptions.
3. Non-metals may be solid, liquid or gas.
4. Non-metals are not sonorous.
Non-metals are not sonorous, i.e., they do not produce a typical sound on being hit.
5.Non-metals are bad conductors of heat and electricity
Non-metals are bad conductors of heat and electricity. Graphite, which is allotrope of carbon is a good conductor of electricity and is an exception.
6. Non-metals are not malleable.
7. Non-metals are not ductile.
8. Non-metals have low melting and boiling point
Non-metals generally have low melting and boiling point, except diamond
9. Most of the non-metals have low density.
10. Non-metals are in many colors.
Chemical properties of Non-metals
1. Reaction of Non-metals with oxygen
Non-metals form corresponding oxide when reacting with oxygen.
Non-metal + Oxygen → Non-metallic oxide
Example:-
1. When carbon reacts with oxygen, carbon dioxide is formed along with the production of heat.
C + O2 → CO2 + Heat
2. 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.
2C + O2 → 2CO + Heat
3. Sulphur gives sulphur dioxide when reacting with oxygen. Sulphur catches fire when exposed to air.
S + O2 → SO2
4. When hydrogen reacts with oxygen it gives water.
2H2 + O2 → 2H2O
Non-metallic Oxide
Non-metallic oxides are acidic in nature. The solution of non-metal oxides turns blue litmus red.
Example:-
1. Carbon monoxide gives carbonic acid when dissolved in water.
CO2 + H2O → H2CO3
2. Sulphur dioxide gives sulphurous acid when dissolved in water.
SO2 + H2O → H2SO3
3. Sulphur dioxide gives sulphuric acid when reacts with oxygen.
2SO2 + O2 → 2SO3
SO3 + H2O → H2SO4
2. Reaction of Non-metal with chlorine
Non-metal gives respective chloride when they react with chlorine gas.
Non-metal + Chlorine → Non-metal chloride
Example:- Hydrogen gives hydrogen chloride and phosphorous gives phosphorous trichloride when reacting with chlorine.
H2 + Cl2 → 2HCl
P4 + 2Cl2 → 4PCl3
3. Reaction of Non-metals with hydrogen
Non-metals reacts with hydrogen to form covalent hydrides.
Non-metal + Hydrogen → Covalent Hydride
Example:-
1. Sulphur combines with hydrogen to form a covalent hydride is called Hydrogen sulphide.
H2 + S → H2S
2. Nitrogen combines with hydrogen in presence of an iron catalyst to form covalent hydride ammonia.
N2 + 3H2 → 2NH3
**Note:- 1. Non-metals do not react with water or steam to evolve hydrogen gas.
2. 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 and Negative charged ions 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 (Cation):- A positive ion is formed because of the loss of electrons by an atom.
Example:- 1. Sodium forms sodium ion because of the loss of one electron.
2. Magnesium forms positive ion because of the loss of two electrons.
Negative ion (Anion):- A negative ion is formed because of the gain of an electron.
Example:- 1. Chlorine gains one electron in order to achieve a stable configuration. After the gain of one electron, chlorine gets one negative charge over it forming chlorine ion.
Why cations or anions form:-
Noble gasses are stable elements. Every elements other than noble gasses loss, gain or share electrons to make themselves stable. Because of loss or gain of electrons elements become charged (cation or anion).
To understand cation and anion, we have to understand electronic configuration of elements and their valencies.
Valency:-
The number of valance electrons present in the outer most shells of an atom is known as valency.
Example:- Electronic configuration of Sodium (Na) = 2 8 1
Their are three shells in sodium atom and the outer most shell has 1 electron can be shared, so valence electron of sodium is 1.
- If outer most shell has 1, 2, 3, or 4 electrons, these can be given in sharing of electrons. So 1, 2, 3, or 4 will be valance electrons.
- If outermost shell has 5, 6 or 7 electrons can not be given in sharing of electrons as these need electrons to complete their octet.
Required valance electrons for outer most shell having 6 electrons = 8 - 6 = 2
Required valance electrons for outer most shell having 7 electrons = 8 - 7 = 1
How cations and anions form
Difference between number of protons in nucleus of an atom and number of electrons create positive or negative charge on the atom and make ion.
Example:- 1. A sodium atom has one electron in its outermost shell. If it losses the electron from its M-shell then its L-shell now becomes the outermost shell and that has a stable octet. The nucleus of this atom still has 11 protons but the number of electrons has become 10. So there is a net positive charge giving us a sodium cation Na+.
2. A chlorine atom has 7 electrons in its outermost shell. It requires 1 electron to complete its octet. If it gains 1 electron then its M-shell become the outermost shell that has a stable octet.
The nucleus of this atom still has 17 protons but the number of electron has become 18. So there is a net negative charge giving us a chlorine anion Cl-.
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 or Electrovalent Bond.
Ionic Compounds
The compounds formed in this manner by transfer of electrons from a metal to a non-metal are known as ionic compounds or electrovalent compounds.
Example:- In sodium chloride, sodium is a metal and chlorine is a non-metal. 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.
Properties of Ionic compounds
- 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 compounds facilitate the passage of electricity by moving towards opposite electrodes.
- Ionic compounds conduct electricity in the molten state.
Occurrence and Extraction of Metals
Sources of metals:- 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 mineral from which a metal can be profitably extracted are called Ores.
Gangue:- Ores mined from the earth are usually contaminated with large amounts of impurities such as soil, sand, etc. are called gangue.
- 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, 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 categorized 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.
- 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. First of all, concentration is done to remove impurities from ores. The concentration of ores is also known as enrichment of ores. Gravity separation, Electromagnetic separation, Froth floating 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.
Calcenation:- Heating of carbonate ores in the limited supply of air to convert them into oxides is known as Calcination.
Difference between Calcination and Roasting
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.
Extraction of Mercury:
Cinnabar (HgS) is first heated in air. This turns HgS (Mercury Sulphide) 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:
Copper glance (Cu2S) is roasted in the presence of air. Roasting turns copper glance into Copper (I) Oxide. Copper oxide is then heated in the absence of air. This reduces copper (I) 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 blend (ZnS) and Calamine (ZnCO3) are ores of zinc. Zinc blends is roasted to be converted into zinc oxide.
Calamine 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:
Hematite (Fe2O3) ore is heated with carbon to be reduced to iron metal.
Extraction of Lead:
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 aluminum ( a reducing agent) to be reduced to metal.
Example:- Manganese dioxide and copper oxide are reduced to respective metals when heated with aluminum.
Thermite Reaction
Some displacement reactions are highly exothermic, the amount of heat evolved is so large that the metals are produced in the molten state. Such a reaction is known as thermite reaction.
Example:- Ferric oxide, when heated with aluminum, is reduced to iron metal. In this reaction, a lot of heat is produced.
Use of Thermite Reactions:- The thermite reaction is used in the welding of electric conductors, iron joints, cracked machine parts, 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 form 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 form the process of electrolytic reduction are pure.
Refining or Purification of metals
Metals extracted form various methods contains some impurities, thus, they are required to be obtained. 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 form a lump of impure metal. In this, impure metal is used as anode and pure metal 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 form 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.
Corrosion
Most of the metals keeps 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 carbonates, etc. As a result, the metal gets eaten up. This process is called Corrosion.
Rusting of Iron
Rusting of iron is the most common form of corrosion. When iron comes in contact with moisture present in the air, the upper layer of iron turns into iron oxide (Fe2O3.32H2O) is reddish-brown in colour and is known as Rust. The phenomenon is called Rusting of Iron.
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. This can be done by-
1. painting
2. Greasing or Oiling
3. Galvanization
4. Electroplating
5. Alloying
Galvanisation:- A method to protecting steel and iron from rusting by coating them with a thin layer of zinc. This process is known as Galvanisation.
Electroplating:- Electroplating is basically the process of plating a metal onto the other by hydrolysis mostly to prevent corrosion of metal or for decorative purposes.
Alloying:- Pure metals are not used to make articles. So there are mixed some other substances to make them hard and strong and causing changing in metal's properties. This process is called alloying.
Alloys
The homogeneous mixture of two or more metals, or a metal and a non-metal is called alloy.
Some examples of alloys are-
1. Brass - 80% Cu + 20% Zn
2. Bronze - 90 % Cu + 10 % Sn
3. Solder - 50% Pb + 50% Sn
4. Duralumin - 95% Al + 4% Cu + 0.5% Mg + 0.5% Mn
5. Steel - 99.95% Fe + 0.05% C
6. Stainless steel -74% Fe + 18%Cr + 8%Ni
7. Magnesium alloy - 95% Al + 5%Mg
8.German Silver - 60% Cu + 20% Zn + 20% Ni
9. 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.
10. Amalgams - An alloy in which mercury (Hg) is present. For example Sodium amalgams [Na(Hg)] and Zinc amalgams [Zn(Hg)].
Properties of 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. For example Solder [50% Pb +50% Sn] has lower melting point than Pb and Sn.
- The electrical conductivity of alloys is lower than the constituent metals.
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