Metals and Non-Metals PDF

0  Metals and Non-Metals Introduction All the naturally occuring elements are broadly grouped into two categories, metals and non-metals Some common metals are gold (Au), silver (Ag), lead (Pb), Zinc (Zn), aluminium (Al), magnesium (Mg), nickel (Ni), iron (Fe), etc. S...

0  Metals and Non-Metals Introduction All the naturally occuring elements are broadly grouped into two categories, metals and non-metals Some common metals are gold (Au), silver (Ag), lead (Pb), Zinc (Zn), aluminium (Al), magnesium (Mg), nickel (Ni), iron (Fe), etc. Some common non-metals are nitrogen (N), oxygen (O), hydrogen (H), carbon (C), sulphur (S), phosphorus (P), Chlorine (Cl) etc. Elements An element is a substance that is made entirely from one type of Atom Example: Hydrogen, Helium, Oxygen Metals A metal is an element, whose atoms readily lose electrons to form positive ions (cations), (except hydrogen) and form metallic bonds with other metal atoms and ionic bonds with non-metallic atoms. Metals are electropositive elements as they form positive ions by losing (donating) electrons. Among all the know 118 elements, most of them are metals. Major metals in the earth’s crust are : Aluminium (Al), Iron (Fe), Calcium (Ca), Sodium (Na), Potassium (K) and Magnesium (Mg). Metals Occurrence of Metals Physical Properties of Metals Activity Series of Metals Chemical Properties of Metals Uses of Metals Non-metals Occurrence of Non-metals Physical Properties of Non-metals Chemical Properties of Non-metals Uses of Non-metals Comparative study of Metals and Non-metals Interactions in Metals and Non-metals Chemical Bond Causes of Chemical Bond Formation Types of Chemical Bonds Ionic Bond Examples of Ionic Bond Nature and Structure of Ionic Compounds Characteristics of Ionic Compounds Metals in Nature Ores 1  Metallurgy : Extraction of Metals Extraction of Various Metals According to their reactivity Alloys Preparation of Alloys Properties of Alloys Some Important Alloys Corrosion Essentials of Corrosion Corrosion of iron (Rusting of Iron) Quick Recap Assignment Occurrence of Metals Metals are found mainly in the earth’s crust, both in the free state as well as in the combined state. A) Combined State: Highly reactive metals are present in the form of their compounds as carbonates, sulphates, oxides etc in the reactivity series, all the metals above copper are found in the combined form as sodium is present in the form of sodium chloride, sodium nitrate, sodium carbonate etc. B) Native/Free State: Less reactive metals are generally found in their elemental form in the free state or in native state as Gold (Au); Silver (Ag), Plantinum (Pt) etc. are found as such in nature. Physical Properties of Metals There are some properties which are helpful in the recognition of metals are: A) Physical State : Most of the metals are hard solids at room temperature. Example: Iron (Fe), Silver (Ag) etc. Reason: The atoms of metals are rigidly or tightly packed in the crystal. Exception : (a) Sodium (Na) and Potassium (K) metals are soft. (b) Hg (Mercury) is a liquid metal. B) Melting and Boiling Point: Metals generally have high melting points and high boiling points. Example: Beryllium (Be) ha s1287oC [m.p.] and 2472oC [b.p.]. Reason: Due to the rigid packing of atoms in the crystal, atoms are held together by strong forces of attraction. Exception: Na and K have low melting and low boiling points. Caesium and gallium have such low melting points that they even melt on keeping on one’s palm. C) Lustre: The substances with shiny and reflecting surfaces are generally metals. Metals in their form have lustre. Example: Gold (Au) has a shiny surface while wood does not have lustre. Reason: Due to electronic vibrations of free electrons. 2 D) Malleability: Metals can be beaten into thin sheets. This property is called malleability. Example: Aluminium, Gold, Silver can be beaten into this sheets i.e., aluminium foils, gold foils and silver foils. (Gold and silver are most malleable) Reason: Due to the elastic nature of metals i.e., they can bear the stress. Exception: Zinc (Zn) is non-malleable at most of temperatures however, it is interesting to know that it becomes malleable between 100-150oC and again becomes malleable above 210oC. E) Ductility: Metals can be drawn into chains or wires. This property of metals is called ductility. Example: Gold (Au) and Silver (Ag) chain, Cable wires (copper wires). Gold is the most ductile metal. Reason: Due to the elastic nature of metals. Exception: Zinc. F) Thermal Conductivity: Transfer of heat from one end of a substance to the another end in known as thermal conductivity. Metals can transfer heat from one end to other end so they are called thermal conductors. Example: When utensils are heated from bottom the whole pot turns hot. It means that the heat is transferred from bottom to all other parts of the utensil. Silcer and copper are the best conductors of heat. Exception: Mercury (Hg) and Lead (Pb) are poor conductors of heat. G) Electrical Conductivity: Have you ever noticed that to switch on a TV, we insert the power cable of television into an electric socket, and TV starts working. It is just because the thin wires of metal inside the cables transfer electricity from electric socket to TV. This property of metals is called electrical conductivity. On removing metal wire from cable, we cannot turn on our electrical appliances. Example: Silver, Copper, Iron etc. Reason: Due to the presence of free electrons. H) Sonority: Metals produce a ringing sound when struck with hard material, this property of metals is called sonority and therefore metals are called ‘sonorous’. Example: Wind chime of metal sticks or hand bell. I) Strength: Metals are strong i.e., they can hold heavy weights without snapping (breaking). Example: Iron metal in different forms, used for different purposes as in making bridges, vehicles, railway lines, buildings etc. Exception: Sodium and potassium are soft metals. J) Density: Metals usually have high densities i.e., they are heavy substances. Example: Iron and manganese have high densities i.e., 7.8 g/cm3 and 7.4 g/cm3 respectively. Reason: Due to rigid structures, number of atoms and their mass per unit volume is more. Exception: Sodium and potassium have low densities i.e., 0.97 g/cm3 and 0.86 g/cm3 respectively. 3  K) Colour: Usually, metals are of silvery/grey or golden yellow in colur. Example: Silver (Ag), Aluminium (Al). Exception: Copper (Cu) (Reddish brown) L) Alloy Formation: Alloys are homogeneous mixtures of two or more metals (sometimes small traces of non-metals are also present). Example: Brass is a mixture of Copper (Cu) and Zince (Zn). Activity Series of Metals Metals, according to their reactivity, are arranged in a specific arrangement in order of their decreasing reactivity which is termed as activity series of metals. In the activity series, the most reactive metal is potassium (K) which is placed at the top of series while Platinum (Pt) is the least reactive metal which is placed at the bottom of the series. Activity series is also called the reactivity series. Hydrogen is although a non-metal, but it is included in the series because it can form a positive ion. Characteristics Regarding Activity Series of Metals Electropositive character decreases down the series. Metals that are placed above hydrogen can displace hydrogen from water and dilute acids. Metals that are placed below hydrogen cannot displace hydrogen from water and dilute acids. A high placed metal can displace a low placed metal from its salt solution. A low placed metal cannot displace a high placed metal from its salt solution. Oxides of metals K, Na, Ca, Mg and Al cannot be reduced by reducing agents like H2, C or CO. Oxides of metals (below aluminium) can be reduced by reducing agents like H2, C or CO. Oxides (and nitrates) of very less reactive metals like Hg, Au can be reduced by heating or by strong heating. Chemical Properties of Metals When metals go for chemical reactions, they behave uniquiely. Chemical properties of metals are as follows: Metals are electropositive elements Metals are very reactive. Metals are electropositive elements i.e., metals tend to lose electrons easily and form positively charged ions called cations. Example: (a) Na → Na+ + e− (loss of one electron) (b) Mg → Mg 2+ + 2e− (loss of two electrons) The electropositive nature allows metals to form compounds with other elements easily. A) Reaction of Metals with Oxygen Most of the metals combine with oxygen to form metal oxides. Sodium react with oxygen at room temperature to form sodium oxide which on reaction with water gives alkali. Metal + oxygen → Metal oxide Metal oxides are basic oxides since they form bases (or alkalis i.e., water soluble bases) when dissolved in water and turn red litmus blue. Metal oxide + Water → Base 4  Some metals also form amphoteric oxides that can react with acids as well as bases. Examples: i) Reaction of Sodium with Oxygen Sodium combines with oxygen at room temperature to form sodium oxide which on reaction with water gives sodium hydroxide (alkali). 4Na(s) + O2 (g) → 2Na2 (s) Sodium Oxygen Sodium oxide (basic) Na2 O(s) + H2 O(l) → 2NaOH(aq) Sodium oxide Water Soluble base (alkali) Sodium hydroxide ii) Reaction of Magnesium with Oxygen Mg does not react with oxygen at room temperature. On heating, Mg burns in air with intese white light and produces large amount of heat to form MgO, which on reaction with water gives magnesium hydroxide base. ∆ 2Mg(s) + O2 (g) → 2MgO(s) Magnesium Oxygen Magnesium oxide MgO + H2 O → Mg(OH)2 (s) Magnesium Water Magnesium hydroxide oxide (base) iii) Reaction of Copper with Oxygen Copper is less reactive metal and does not burn in air on heating. However, on prolonged and strong heating, copper reacts with oxygen and forms copper (II) oxide (also called cupric oxide) which on reaction with water forms copper (II) hydroxide. Strong heating 2Cu + O2 → 2CuO Copper Oxygen Copper (II) oxide CuO + H2 O → Cu(OH)2 Copper (II) Water Copper (II) hydroxide oxide (base) Order of Reactivity is : K > Na > Mg > Cu B) Reaction of Metals with Water Hydrogen gas is produced when metals react with water. Second product will be the metal oxide or the hydoxide. Highly reactive metals form their respective hydroxides while less reactive metals form their respective oxides. Metal + Water → Metal hydroxide + Hydrogen gas Metal + Water → Metal oxide + Hydrogen gas Examples: i)Reaction of Potassium with Water Potassium reacts vigorously with cold water at room temperature and gives aqueous potassium hydroxide and hydrogen gas. room 2K + 2H2 O → 2KOH + H2 ↑ (Potassium) (Cold water) temperature (Potassium (Hydrogen) hydroxide) 5  ii) Reaction of Sodium with Water Sodium reacts vigorously with cold water at room temperature and gives aqueous sodium hydroxide and hydrogen gas. room 2Na + 2H2 O → 2NaOH + H2 ↑ Sodium Cold water temperature Sodium Hydrogen hydroxide Potassium, sodium and calcium all react vigorously with water but their reacivity order is K > Na > Ca. iii) Reaction of Magnesium with Water Mg + H2 O → Mg(OH)2 + H2 ↑ Magnesium Hot water Magnesium Hydrogen hydroxide Mg + H2 O → MgO + H2 ↑ Magnesium Steam Magnesium Hydrogen oxide iv) Reaction of Zinc with Water (steam) Zn + H2 O → ZnO + H2 ↑ Zinc Steam Zinc oxide Hydrogen v) Reaction of Iron with steam 3Fe + 4H2 O → Fe3 O4 + 4H2 ↑ Iron Steam Iron oxide Hydrogen Magnesiumm reacts with hot water whereas aluminium, iron or zince reacts with steam to form their respective metallic oxides along with hydrogen gas. Reactivity order is : Mg > Al > Zn > Fe Metals like Cu, Hg, Au and Sn do not react with water (even at high temperatures.) Reactivity order : K > Na > Ca > Mg > Al > Zn > Fe C) Reaction of Metals with Dilute Acids Metal salt and hydrogen gas are produced when a metal reacts with dilute acid. Metal + Dilute mineral acid → Metal salt + hydrogen gas More reactive metals readily evolve hydrogen from acids as compared to the less reactive ones. The rate of displacement of hydrogen from acid depends upon electropositive nature of metals as well as the nature of acid. Highly electropositive metals are good reducing agents and hence can lose their electron rapidly and reduce acids. K > Na > Ca > Mg > Al > Zn Metals, placed above hydrogen in the activity series can displace hydrogen from acids but the metals that are placed below hydrogen cannot displace it from acids. Strong acids like HCl, H2SO4 readily lose their hydrogen (ionisable hydrogen) as compared to that of weak acids like phosphoric acids (H3PO4) etc. as: 2Al + 6HCl → 2AlCl3 + 3H2 ↑ (rapid) Aluminium Strong acid Aluminium Hydrogen chloride 2Al + 2H3 PO4 → 2AlPO4 + 3H2 ↑ (slow) Aluminium Weak acid Aluminium Hydrogen phosphate Examples: 6  i) Reaction of Na with dilute Acid: 2Na + 2HCl(aq) → 2NaCl + H2 ↑ Sodium dil. hydrochloric Sodium Hydrogen acid chloride ‘H2’ is evolved readily and vigorously. ii) Reaction of Magnesium with Dilute Acid: Mg + 2HCl → MgCl2 + H2 ↑ dil. hydrochloric Magnesium Magnesium Hydrogen acid chloride ‘H2’ is evolved quite less readily. iii) Reaction of ‘Al’ with Dilute Acid: 2Al(s) + 6HCl(aq) → 2AlCl3 (aq) + 3H2 (g) ↑ Aluminium dil. hydrochloric Aluminium chloride Hydrogen acid ‘H2’ is evolved comparatively less readily. iv) Reaction of Zinc with Dilute Acid: Zn(s) + 2HCl(aq) → ZnCl2 + H2 ↑ Zinc dil. hydrochloric Zinc chloride Hydrogen acid ‘H2’ is evolved comparatively less readily. v) Reaction of Iron with Dilute Acid: Fe + 2HCl → FeCl2 + H2 ↑ Iron dil. hydrochloric Iron III) chloride Hydrogen acid vi) Reaction of Copper with Dilute Acid: Cu(s) + HCl(aq) → No reaction Copper dil. hydrochloric acid ‘H2’ is not evolved because Cu is less reactive than hydrogen and cannot displace it from acid. Q. Which gas is produced when dilute hydrochloric acid is added to a reactive metal ? Write the chemical reaction when iron reacts with dilute 𝐇𝟐 𝐒𝐎𝟒. Sol. Hydrogen gas is produced when dilute hydrochloric acid is added to a reactive metal. Chemical reaction when iron reacts with dilute H2 SO4 : Fe(s) + H2 SO4 (aq) → FeSO4 (aq) + H2 (g) D) Reaction of Metals with Salt Solutions of less reactive metals: More reactive metal displaces less reactive metal from its salt solution and this type of reaction is called displacement reaction. Actually, when a more reactive metal comes in contact with a solution that contains ions of less reactive metal, it displaces those ions and deposits the less reactive metal in pure metallic form. M1 + Salt solution of M2 → M2 + Salt solution of M1 More Salt solution of reactive less reactive metal Here M1 and M2 represents two different metals. Example: 7  i) Reaction of iron with Copper Sulphate Solution: Iron + Copper (II) sulphate → Iron (II) sulphate + Copper metal Fe(s) + CuSO4 (aq) → FeSO4 (aq) + Cu(s) Iron Blue Green Reddish−brown Copper (II) sulphate solution is blue; iron sulphate solution is light green when dilute. During reaction, the blue colour of solution fades and the iron metal is seen to turn red- brown as the displaced copper gets deposited on it. ii) Reaction of Magnesium with Copper Sulphate Solution: Magnesium being more reactive displaces ‘Cu’ from its salt which gets deposited as pure metal. Mg(s) + CuSO4 (aq) → MgSO4 (aq) + Cu(s) Magnesium Blue Colourless Copper On adding magnesium to blue copper (II) sulphate solution, the blue colour fades as colourless magnesium sulphate is formed and brown bits of copper metal form a precipitate. iii) Reaction of Silver with Copper Sulphate Solution: Silver, being less reactive than copper, cannot displace Cu from its solution. Hence, no reaction takes place. Ag(s) + CuSO4 (aq) → No reaction Silver Blue On adding magnesium to blue copper (II) sulphate solution, the blue colour fades as colourless Uses of Metals Metals are important in every aspect of life. Some uses of metals are mentioned below: A) Metals in Bio-processes Iron (Fe) is a constituent of haemoglobin. Na, Mg, K, Ca etc., are essential metals required as minerals in the living body. B) Metals in daily goods and machines Copper (Cu) and Aluminium (Al) are used for making electrical wires and cables. Iron is used for making stove burners, gutter pipe, railway tracks etc. Steel (Iron and Carbon) is used for making engine parts, utensils, equipments etc. C) Metals in Accessories Gold (Au), Silver (Ag) and Platinum (Pt) are used to make jewellery. D) Metals in Health Equipments and Instruments Mercury (Hg) is used in thermometers, sphygmomanometers and barometers etc. E) Metals in Space Projects Sodium (Na), Titanium (Ti) and Zirconium (Zr) are used in atomic space projects. F) Metals as Nuclear Fuel Uranium (U), Netpunium (Np) etc. are used as nuclear fuel in reactors or weapons. 8 G) Metals in Economy Over the period of time, coin has been used for monetary transaction. Au, Ag, Cu etc., have been used for making coins. Therfore, these are also termed as coinage metals. Non-Metals A non metal is an element whose atoms readily gain electrons to form negative ions (anions) and form covalent bon with other non-metals and ionic bonds with metals. Non-metals are electronegative elements as they form negative ions by gaining electrons. Overall there are 18 known non-metals. C(Carbon), N(Nitrogen), O(Oxygen), P(Phosphorus), S(Sulphur), Se(Selenium) are some of the non-metals. Occurrence of Non-metals Non-metals are prresent in the earth’s crust. Oxygen is the most abundant non-metal in earth’s crust that makes about 46.6% by mass of earth’s crust. Silicon is the second most abundant non-metal that forms about 27.7% by mass of earth’s crust. Phosphorus and sulphur also form a considerable part of earth’s crust. Physical Properties of Non-metals Non-metals are usually opposite to metals in characteristics. The important physical properties of non-metals are as follows: A) Physical State: Non-metals are found in all the three states of matter i.e., solid, liquid and gas as: B) Melting and Boiling Point: Non-metals generally have low boiling and melting points than metals. Reason : Due to less compact packing of atoms in them. C) Hardness: Non-metals are generally soft. Reason: Due to non-rigid packing of their atoms. Exception: Diamond (an allotrope of carbon) is the hardest naturally occuring substance. D) Lustre: Non-metals do not show ductility. E) Ductility: Non-metals do not show ductility. F) Malleability: Non-metals are non-malleable. G) Conductivity: Non-metals are bad conductors. (Both thermally and electrically) Reason: Absence of free electrons. Exception: Graphite (an allotrope of carbon) is a good conductor. H) Sonority: Non-metals are non-sonorous. They do not produce sound when hit with an object. I) Strength: Non-metals are not as strong as metals. J) Density: Non-metals have low densities. Example: Nitrogen (1.25 g/cm3 ), Oxygen (1.32 g/cm3 ) K) Colour: Non-metals show different colours as: Sulphur : Yellow Phosphorus : White or Red or Black 9  Graphite : Black Chlorine : Yellowish green Table: Occurrence of some of the non-metals Non-metals Occurrence Carbon As carbon dioxide in air and as carbonates in rocks. Oxygen As oxides and along with other compounds like carbonates. Hydrogen As water and as hydrides. Nitrogen As nitrites and nitrates in soil. Sulphur As Sulphides and Sulphates. Phosphorus As phosphates in rocks. Silicon As silica or SiO2 in the form of sand, quartz. As silicates in the form of mica, asbestos and clay. Chemical Properties of Non-metals A) Reaction of Non-metal with Oxygen Non-metals form respective non-metallic oxides with oxygen. Non-metals + Oxygen → Non-metallic oxide Non-metallic oxides generally are acidic in nature because they form acids on dissolving in water and turns blue litmus paper red. Non-metallic oxide + Water → Acid Some neutral oxides are also obtained which neither react with acids nor with bases. Examples: i) Reaction of Carbon with Oxygen: When carbon reacts with oxygen it forms acidic oxide i.e., carbon dioxide CO2. This is in water forms carbonic acid. A solution of carbon dioxide gas in water turns blue litmus paper red, showing that it is acidic in nature. C(s) + O2 (g) → CO2 ↑ Carbon Oxygen Carbon dioxide (acidic oxide) CO2 (g) + H2 O(l) → H2 CO3 (aq) Carbon dioxide Water Carbonic acid Carbon with limited amount of oxygen gives carbon monoxide which is neutral in nature and hence does not react with acids and bases. 1 C + O → CO ↑ Carbon 2 2 Carbon monoxide (limited) (neutral) Besides carbon monoxide, water (H2 O), nitrogen monoxide or nitric oxide (NO) and nitrous oxide (𝑁2 𝑂) are also some of the examples of neutral oxides. 10  ii) Reaction of Sulphur with Oxygen: Sulphur forms acidic sulphur dioxide with oxygen which in turn gives sulphurous acid on dissolving in water. S + O2 → SO2 ↑ Sulphur Oxygen Sulphur dioxide SO2 + H2 O → H2 SO3 Sulphur Water Sulphurous dioxide acid iii) Reaction of Phosphorus with Oxygen: Phosphorus forms phosphorus pentaoxide with oxygen which gives phosphoric acid on dissolving in water. P4 + 5O2 → 2P2 O5 Phosphorus Oxygen Phosphorus pentaoxide P2 O5 + 3H2 O → 2H3 PO4 Phosphorus Water Phosphoric pentaoxide acid B) Reaction with Water Generally, non-metals do not react with water or steam. This is because non-metals give electrons to reduce hydrogen ions of water into hydrogen gas. Examples: Reaction of Sulphur and Coke with Water: S + H2 O → No reaction Sulphur Water C + H2 O → No reaction Coke Water Coke is the purest form of carbon. C) Reaction of Non-metals with Dilute Acide Non-metals do not react with dilute acids because they cannot displace hydrogen from acids. This is because to replace hydrogen from acids, electrons should be supplied to acids but non- metals are electron acceptor themselves. Example: Reaction of Carbon and Phosphorous with Dilute Hydrochloric Acid: C + HCl → No reaction Carbon (dil) P4 + HCl → No reaction Phosphorus (dil) D) Reaction of Non-metals with Salt Solutions A more reactive non-metal displaces a less reactive non-metal from its salt solution. Examples: 2NaBr(aq) + Cl2 (g) → 2NaCl(aq) + Br2 (l) Sodium bromide Chlorine Sodium chloride Bromine (less reactive) (more reactive) 11 Uses of Non-metals A) Non-metals in Bio-processes CO2 and O2 are gases of life. Oxygen is taken during respiration by all living beings while CO2 is a raw material for photosynthesis in plants. B) Non-metals in Laboratory For the preparation of various chemicals of commercial use, many compounds of non-metals are used as laboratory reagens. For example, sulphur as sulphuric acid (H2 SO4 ), nitrogen as nitric acid (HNO3 ), hydrogen as reducing agent etc. C) Non-metals in Health and Protection Sulphur is used as fungicide. Petroleum jelly is composed of carbon compouns, chlorine is an important disinfectant and insecticide. D) Non-metals in Various Goods Sulphur is used as fungicide. Petroleum jelly is composed of carbon compounds, chorine is an improtant disinfectant and insecticide. E) Non-metals as Fuel Hydrogen is the most efficient fuel. It is used as rocket fuel. Fossil fuels having carbon compounds, such as coal and petroleum, are used as fuel for almost all energy consuming processes. Comparative Study of Metals and Non-metals A) Differences in Physical Properties Metals Non-metals 1. Malleable, ductile and hard. 1. Non-malleable, non-ductile and generally soft. 2. Lustrous sonorous. 2. Non-lustrous and non-sonorous. 3. Generally solids at room 3. Generally liquids or gases at room temperature. temperature. 4. Have high densities and high 4. Have comparatively low densities and low melting and boiling points. melting and boiling points. 5. Have high tensile strength. 5. Have low tensile strength. 6. Used in alloy formation 6. With metals, occasionally form alloys. B) Differences in Chemical Properties Metals Non-metals 1. Lose electrons easily to form cations. 1. Gain electrons easily to form anions. 2. Form basic oxides. Some metals 2. Form acidic oxides. Some non-metals form amphoteric oxides also. form neutral oxides also. 3. Evolve hydrogen from dilute acids. 3. Do not evolve hydrogen gas from water. 12  4. Evolve hydrogen gas from water. 4. Do not evolve hydrogen gas from water. 5. Form ionic chlorides. 5. Form covalent chlorides. 6. Usually do not combine with hydrogen 6. Non-metals form covalent hydrides but some metals form their respective ionic with hydrogen. hydrides with hydrogen. Interactions in Metals and Non-metals When metals react with non-metals, they form ionic compounds. On the other hand, when non- metals react with other non-metals, they form covalent compounds. When atoms of elements combine to form molecules, a force of attraction is developed between atoms which holds them together. This force is called a ‘Chemical bond’. Chemical Bond The force of attraction that holds two or more atoms, ions etc. together in different chemical species is called a chemical bond. Causes of Chemical Bond Formation A) To attain minimum energy and maximum stability Atoms combine with one another to attain stability by foming molecules. Molecules are stable species and have less kinetic energy. B) To attain nearest inert gas electronic configuration Generally, atoms of inert gases are non reactive or negligibly reactive and hence are stable. Most of the atoms tend to become stable by attaining electronic configuration of inert gases (8 electrons in valence shell). The atoms that have 8 electrons in the valence shell are stable, (except He, Be etc., that have 2 electrons in valence shell). This is in accordance with Lewis Octet Rule. Types of Chemical Bonds Atoms combine to form molecules in different manners. On the basis of electronic configurations of atoms and the process by which the atoms attain stable configurations, chemical bonds may be of various types as: i) Ionic bond or electrovalent bond (formed by the transfer of electrons) (non-directional in nature). ii) Covalent bond or electron pair bond (formed by bidirectional sharing of electrons). iii) Co-ordinate bond (formed by unidirectional sharing of electrons). 13 Ionic Bond The bond that is formed by the complete transfer of elctron from one atom to another atom is known as ionic bond or the force that hlods atoms together through complete transfer of electrons is called ionic bond. An ionic bond is formed when one atom can donate electrons to attain inert gas configuration and the another atom can gain electrons to attain inert gas configuration. Generally, an ionic bond is formed between a metal and a non-metal. The number of electrons, lost or gained by an atom during the ionic bond formation, is known as electrovalency. Explanation: Actually, metals have a tendency to lose electrons while non-metals have a tendency to gain electrons to attain stability after attaining inert gas configuration. Metals form cations while non-metals form anions. The positively charged cation and negatively charged anion attract one another and the strong force of attraction developed between the ions with opposite charge in known as ionic bond. The compounds containing ionic bonds are called compounds. Examples of Ionic Bond A) Formation of Sodium Chloride (NaCl) Sodium ( 11Na) ∶ Electronic configuration : 2 , 8, 1 −1e− 2, 8, 1 → 2, 8 [Ne gas configuration] Chlorine ( 17Cl) ∶ Electronic configuration : 2, 8, 7 Chlorine can gain one electron to be stable and form negatively charged ion +1e− 2, 8, 1 → 2, 8 [Ne gas configuration] Positively charged ions are called cations while negatively charged ions are known as anions. Electron-dot structure: B) Formation of Magnesium Chloride (𝐌𝐠𝐂𝐥𝟐 ): Magnesium ( 12Mg) ∶ Electronic configuration : 2, 8, 2 Magnesium can lose two electrons to become stable or to attain inert gas electronic, configuration. −2e− 2, 8, 2 → 2, 8 [Ne gas configuration] Chlorine( 17Cl) ∶ Electronic configuration : 2, 8, 7 Chlorine can gain one electron to become stable or attain inert gas electronic configuration 14  +1e− 2, 8, 7 → 2, 8, 8 [Ar gas configuration] Electron-dot structure: Characteristics of Ionic Compounds A) Physical State: Ionic compounds are solids at room temperature. Explanation : Since ionic compounds exist as clusters of ion pairs, there are strong forces of attraction between them which form strong crystals that cannot be liquefied easily and hence ionic compounds are solids. B) Hardness: Ionic compounds are hard. Explanation: Due to strong electrostatic forces of attraction between the constituent atoms, ionic compounds are hard. C) Conductivity: In the solid state, the ionic compounds are non-conductors but in the molten (fused) or in the aqueous state, they conduct electricity. Explanation: In the solid state, the ions are held so strongly by electrostatic forces of attraction that they remain stationary in ionic crystal and hence do not conduct electricity, but in the fused state or in the aqueous state, the ions are available for conductivity. D) High melting point and boiling point: Ionic compounds have high melting and boiling point. Explanation : Due to strong electrostatic forces of attraction, the constituent ions have rigid packing which results in high melting and boiling point of ionic compounds. Ionic Compound Melting Point (K) Boiling Point (K) NaCl 1074 1686 LiCl 887 1600 CaCl2 1045 1900 MgCl2 981 1685 E) Solubility: Ionic compounds are generally soluble in water. Explanation: Ionic compounds dissolve in water (polar solvent) because of the attraction of positive and negative charges. These water molecules arrange around ions of compound in such a way that each cation is surrounded by the negative pole and each anion is surrounded by positive pole of water molecule. This phenomenon is called hydration and the energy released during this process is known as hydration energy. Solvation energy in further dissociation of ionic compound into ions. 15 Metals in Nature In nature, metals are generally found either in the native state or in the combined state. The native state is also called free state or metallic state. Less reactive metals are generally present in the free elemental state. These are found usually associated with rocks or alluvial materials like clay, sand etc. Sometimes metals are found as lumps which are called ‘nuggets’. Example: Silve, gold, platinum etc. In the combined state, metals form minerals which contain various valuable substances and impurities along with metals. The minerals which are present as mixture of metallic compounds and extractable impurities are known as ores and the compound of a metal found in nature is called mineral. The impurities from the desired metals are removed by various methods called extraction methods. The complete process of extraction of metals is known as metallurgy. Ores The minerals from which the metals can be conveniently and economically extracted are know as ‘Ores’. Type Name of Ore Composition (a) Oxides Haematite Fe2 O3 Magnetite Fe3 O4 Bauxite Al2 O3. 2H2 O (b) Carbonate Limestone CaCO3 Dolomite CaCO3. MgCO3 Calamine ZnCO3 (c) Sulphates Epsom Salt MgSO4. 7H2 O Gypsum CaSO4. 2H2 O (d) Sulphides Galena PbS Zinc blende ZnS Iron pyrite FeS2 Extraction of Metals “Metallurgy is the process of extraction of metals in pure state from their respective ores and refining them for use.” Extraction of Metals according to reactivity The highly reactive metals are generally extracted by ‘electrolysis’ due to strong bonding they have with other components of ore. So, cannot be reduced by heating with carbon. 16  Example: Potassium (K), Sodium (Na), Calcium (Ca), Magnesium (Mg) etc., are extracted by electrolysis. The moderately reactive metals are generally extracted by reduction process. This reduction is done with reducing agents like coke (C) etc. Example: Zinc (Zn), Iron (Fe), Lead (Pb) etc., are extracted by this method. Less reactive metals are extracted from their oxides by heating alone. Example: Copper (Cu), Mercury (Hg). Very less reactive metals are present in free state in nature in the metallic form. Example: Silver (Ag), Gold (Au), Platinum (Pt). Extraction of Various Metals According to their Reactivity A) Mining of Ore: Generally ores are found deep inside the earth but some may be present only a few metres under earth’s surface. Mining is the process to take out the ores from mines. B) Sizing of Ores: Mined ores are found generally in form of big lumps which are crushed into small pieces and then powdered. C) Concentration of Ore/enrichment of ore/dressing of Ore: The removal of undesired foreign impurities i.e., gangue from ore is called concentration of ore. Many physical and chemical methods are used for the concentration of ore. D) Oxidation of Ore into Oxide Ore: The oxidation of ore is an essential step during metallurgical process because the reduction of metal oxides is comparatively easier than the reduction of other compounds of metals (as carbonate, nitrate etc.) and hence, metallic compounds are first converted into their respective oxides and then reduced to metals. Conversion of ore into oxide ore can be done by either of the following methods: 17  a) Calcination: In this process, the ore is heated in absence of air or limited supply of air strongly, below its fusion temperature (meltin point). This process removes organic matter and moisture from the ores. It is generally carried out in the case of carbonate ores or hydrated oxide ores. Calcination is done in reverberatory furnace. i) Hydrated oxide ores undergo calcination in order to remove moisture as– ∆ Al2 O3. 2H2 O → Al2 𝑂3 + 2H2 O Bauxite Alumina Water ∆ 2Fe2 O3. 3H2 O → 2Fe2 O3 + 3H2 O Iron (III)oxide Iron(III) Water trihydrate oxide ii) Carbonate ores undergo calcination in order to remove carbon dioxide as – ∆ CaCO3 → CaO + CO2 Calcium carbonate Calcium Carbon Limestone oxide dioxide ∆ CaCO3. MgCO3 → CaO + MgO + 2CO2 ↑ Dolomite Calcium Magnesium Carbon oxide oxide dioxide ∆ ZnCO3 → ZnO + +CO2 ↑ Zinc carbonate Zinc Carbon Calamine oxide dioxide b) Roasting: In this method, ore is heated either alone or with some material but the excess of oxygen (air) is the essential condition. The ore is heated below its melting point (fusion temperature). This process is employed when oxidation of the ore is required. It is generally carried out for sulphide ores. i) 2PbS + 3O2 → 2PbO + 2SO2 Galena Oxygen Lead (II) oxide Sulphur dioxide ii) 2Cu2 𝑆 + 3O2 → 2Cu2 O + 2SO2 Chalcocite Oxygen Copper (I) Sulphur oxide dioxide iii) 2ZnS + 3O2 → 2ZnO + 2SO2 Zinc blende Oxygen Zinc (II) oxide Sulphur dioxide Roasting is carried out in a special type of reverberatory furnace. In roasting, moisture is given out along with the impurities of sulphur and phosphorus. After roasting the ore becomes porous that helps in reduction of ore in obtaining desired metal. E) Reduction of Oxide Ore into metal: The metal oxides are finally converted into metals by reduction process. Depending upon the reactivity of metals, reduction is done in different ways as: a) Smelting (Reduction with Carbon): In this process, the roasted or calcined ore is mixed with suitable quantity of coke or charcoal (which act as reducing agent) and is heated to a high temperature above its melting point. Example: i) PbO + C → Pb + CO ↑ Plumbus Carbon Lead Carbon oxide monoxide 18  PbO + CO → Pb + CO2 ↑ Plumbus Carbon Lead Carbon oxide monoxide dioxide ii) Fe2 O3 + 3C → 2Fe + 3CO ↑ Haematite Carbon Iron Carbon monoxide Fe2 O3 + 3CO → 2Fe + 3CO2 ↑ Haematite Carbon Iron Carbon monoxide dioxide b) Thermite process: It is the technique, to reduce metal oxide using more reactive metal powder as fuel. Aluminium, magnesium, titanium are some metals which are used as fuel in thermite process. In this process, a mixture of concentrated oxide ore and metal powder (i.e., thermite) is taken in a steel crucible and kept on sand. A mixture of magnesium powder and barium peroxide (called ignition mixture) is used to ignite the reaction mixture. A large amount of hat is evolved during the raction which melts the metal. Example: Cr2 O3 (s) + 2Al(s) → 2Cr(l) + Al2 O3 (s) Chromium (III) Aluminium Chromium Alumina oxide 3Mn3 O4 + 8Al → 9Mn(l) + 4Al2 O3 (s) Chromium (II,III) Aluminium Mangnese Alumina oxide Fe2 O3 (s) + 2Al(s) → 2Fe(l) + Al2 O3 (s) + Heat This process is also known as Gold-Schmidt aluminothermic reduction as aluminium is most commonly used as thermite in the process. c) Electrolytic reduction: Highly reactive metals like Na, K, Mg, Ca, Al, etc, are reduced by electrolysis of their respective oxides, hydroxides of chloride in molten state. On passing electric current into the molten solution, metal is liberated at cathode while impurities are settled down as anode mud generally. Example: Reduction of sodium from NaCl (by electrolysis) as– NaCl(molten) → Na+ + Cl− At cathode: Na+ + e− → Na Metal At anode: 2Cl− − e− → Cl2 ↑ d) Refining/Purification of Metal: The reduced metals obtained are generally impure which may be associated with following types of impurities as – i) Uncharged (not reduced) ore. ii) Other metals that are produced by simultaneous reduction of their compounds originally present in the ore. iii) Non-metals like silicon, carbon, phosphorous etc. iv) Slag, flux etc., which is present in residual condition. These impurities can be removed by “refining of metals”. Purification of copper (electrorefining of copper) i) Dissociation of electrolyte (𝐂𝐮𝐒𝐎𝟒 solution) CuSO4 → Cu2+ + SO2− 4 19  ii) Reaction at cathode Positively charged copper ions (Cu2+ ) from CuSO4 solution go to the cathode, get reduced and deposit as pure Cu metal. Cu2+ + 2e− → Cu Copper ions Electrons Copper atom (from electrolyte) (from cathode) (at cathode) iii) Reaction at anode Copper atoms of impure anode lose two electrons each to anode and form copper ions Cu2+ , which go into electrolyte as – Cu → Cu2+ + 2e− Copper atoms Copper ions Electrons (from impure anode) (goes to electrolyte) (to anode) Important NCERT Questions Q1. Reverse of the following chemical reaction is. not possible : Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s) Justify this statement with reason. Sol. If a strip of zinc metal is put in copper sulphate solution, then the blue colour of copper sulphate fades gradually due to the formation of colourless zinc sulphate solution and reddish- brown copper metal is deposited on zinc strip. CuSO4(aq) + Zn(s) → ZnSO4(aq) + Cu(s) Copper sulphate Zinc Zinc sulphate Copper (Blue solution) (Colourless solution) (Reddish brown) In this reaction, zinc metal being more reactive than copper displaces copper from copper sulphate solution. If however, a strip of copper metal is placed in zinc sulphate solution, then no reaction occurs. This is because copper metal is less reactive than zinc metal and hence, cannot displace zinc from its salt solution. Cu(s) + ZnSO4(aq) → No reaction Copper Zinc sulphate Q2. Name a metal which: a) is the best conductor of heat. b) has a very low melting point. c) does not react with oxygen even at high temperature. d) is most ductile. Sol. a) Metal which is the best conductor of heat is silver. b) Gallium has a very low melting point. c) Silver and gold do not react with oxygen even at high temperature. d) Gold is the most ductile metal. Q3. What is meant by amphoteric oxides? Choose the amphoteric oxides from the following: Na2 O, ZnO, CO2 , Al2 O3 , H2 O 20 Ans. Amphoteric oxides are those which show acidic as well as basic character, i.e., they react with bases as well as acids. ZnO and Al2O3 are amphoteric oxides. Q4. a) Why does calcium start floating when it reacts with water? Write the balanced chemical equation of the reaction. b) Name two metals which do not react with water. Ans. a) Calcium reacts with cold water to form calcium hydroxide and hydrogen gas. Room Ca(s) + 2H2 O(l) → Ca(OH)2(aq) + H2(g) temperature The bubbles of hydrogen gas produced stick to the surface of calcium and hence, it starts floating on the surface of water. b) Gold and silver do not react with water. Q5. Give reason: a) Aluminium is a reactive metal but is still used for packing food articles. b) Calcium starts floating when water added to it. Ans. a) Aluminium is a strong and cheap metal. It is also a good conductor of heat. But it is highly reactive. When it is exposed to moist air, its surface is covered with a thin impervious layer of aluminium oxide (Al2 O3 ). This layer does not allow moist air to come in contact with the fresh metal and hence, protects the metal underneath from further damage or corrosion. Thus, after the formation of this protective layer of Al2 O3 , aluminium becomes resistant to corrosion. It is because of this reason that although aluminium is a highly reactive metal, it is still used in food packaging. b) Refer to answer 7(a) Q6. An ore on treatment with dilute hydrochloric acid produces brisk effervescence. Name the type of ore with one example. What steps will be required to obtain metal from the enriched ore? Also write the chemical equations for the reactions involved in the process. Ans. The ore on treatment with dilute hydrochloric acid produces brisk effervescence hence, it must be a carbonate ore. Calamine (ZnCO3 ) is an important carbonate ore of zinc. Steps required to obtain metal from the enriched carbonate ore: a) Comversion of rhe carbonate ore into metal oxide: This done by calcination (for carbonate ores). Calcination is the process of heating the ore strongly in the absence or limited supply of air. The zinc carbonate on heating decomposes to form zinc oxide as shown: Heat ZnCO3(s) → ZnOs + CO2(g) Zinc carbonate (Absence of air) (Calamine−ore of Zn) 21  b) Reduction of the metal oxide to metal: As zinc is moderately reactive, zinc oxide cannot be reduced by heating alone. Hence, it is reduced to zinc by using a reducing agent such as carbon. Heat ZnO(s) + 𝐶(𝑠) → Zn(s) + CO(g) Zinc oxide Zinc Carbon monoxide The reduction of metal oxides by heating with coke is called smelting. Q7. Write balanced chemcial equations to explain what happens, when i) Mercuric oxide is heated. ii) Mixture of cuprous oxide and cuprous sulphide is heated. iii) Aluminium is reacted with manganese dioxide. iv) Ferric oxide is reduced with aluminium. v) Zinc carbonate undergoes calcination. Ans. i) On heating, mercuric oxide decomposes to give mercury and oxygen. Heat 2HgO(s) → 2Hg (l) + O2(g) ii) On heating mixture of cuprous oxide and cuprous sulphide, copper and sulphur dioxide are produced. Heat 2Cu2 O(s) + Cu2 S(s) → 6Cu(s) + SO2(g) iii) When aluminium is heated with manganese dioxide, manganese and aluminium oxide are formed. Heat 3MnO2(s) + 2Al(s) → 2Fe(l) + Al2 O3(s) iv) Ferric oxide reacts with aluminium to produce aluminium oxide and iron. Heat Fe2 O3(s) + 2Al(s) → 2Fe(l) + Al2 O3(s) v) On calcination, zinc carbonate produces zinc oxide and carbon dioxide. Calcination ZnCO3(s) → ZnO(s) + CO2(g) Q8. Draw a schematic diagram of the various steps involved in the extraction of metals from ores for metals of medium reactivity and for metals of low reactivity. Ans. Various steps involved in the extraction of a metal from its ore followed by refining of the metal is called ‘metallurgy’. The steps involved are summarised as follows: 22