All questions
15
Q1
Distinguish between metals and non-metals on the basis of their physical properties. Mention at least five properties with one example or exception for each.
Metals are generally lustrous (shining), malleable (can be beaten into thin sheets), ductile (can be drawn into wires), good conductors of heat and electricity, and sonorous (produce ringing sound when struck). Non-metals lack these properties: they are non-lustrous (except iodine), non-malleable, non-ductile, poor conductors of heat and electricity (except graphite), and non-sonorous. Notable exceptions: mercury is the only liquid metal at room temperature; sodium and potassium are soft metals that can be cut with a knife; gallium and caesium have very low melting points — they melt when kept on the palm.
Q2
Physical properties alone are not sufficient to classify elements as metals or non-metals. Justify this statement with at least four specific examples where metals or non-metals show unexpected physical behaviour.
(i) Mercury (a metal) is liquid at room temperature, whereas all other metals are solids. (ii) Sodium and potassium (metals) are so soft they can be cut with a knife, unlike most hard metals. (iii) Gallium and caesium (metals) have such low melting points that they melt when held on the palm. (iv) Iodine (a non-metal) has a shining, lustrous surface — a property typically associated with metals. (v) Graphite (a non-metal allotropic form of carbon) is a good conductor of electricity, unlike other non-metals. (vi) Diamond (a non-metal allotropic form of carbon) is the hardest known natural substance. These exceptions show that properties overlap between the categories, making classification by physical properties alone unreliable.
Q3
How do metals react with oxygen? Taking aluminium as an example, explain the concept of amphoteric oxides. Write balanced chemical equations for the reaction of aluminium oxide with hydrochloric acid and with sodium hydroxide.
Almost all metals combine with oxygen on heating to form metal oxides: Metal + Oxygen → Metal oxide. Most metal oxides are basic in nature, but some metal oxides such as aluminium oxide and zinc oxide react with both acids and bases to produce salt and water — these are called amphoteric oxides. Aluminium burns in air to form aluminium oxide: . Its amphoteric nature is shown by: (i) With HCl: . (ii) With NaOH: (sodium aluminate).
Q4
Describe how different metals react with water. Classify the metals into groups based on their reactivity with cold water, hot water, and steam. Write chemical equations for sodium, calcium, and iron reacting with water/steam.
Metals react with water to produce metal oxide/metal hydroxide and hydrogen gas. Classification: (a) React vigorously with cold water: Potassium and sodium — reaction is violent and exothermic; the evolved hydrogen catches fire. . (b) React with cold water (less vigorous): Calcium — the heat is insufficient to ignite hydrogen; Ca floats due to H₂ bubbles. . (c) React with hot water only: Magnesium — (or Mg(OH)₂). (d) React with steam only: Aluminium and iron. . (e) No reaction with water: Lead, copper, silver, and gold.
Q5
With the help of Activity 3.11, describe how metals react with dilute acids. Explain why hydrogen gas is not evolved when metals (except Mg and Mn) react with dilute nitric acid. What is aqua regia and what makes it special?
Metals react with dilute acids to produce a salt and hydrogen gas: Metal + Dilute acid → Salt + H₂↑. Activity 3.11 shows that Mg, Al, Zn, and Fe react with dilute HCl with decreasing vigour (Mg > Al > Zn > Fe), with Mg giving the highest temperature rise. Copper does not react at all. However, nitric acid (HNO₃) is a strong oxidising agent — it oxidises the H₂ produced to H₂O and is itself reduced to nitrogen oxides (NO, NO₂, N₂O), so H₂ is not evolved. Only Mg and Mn produce H₂ with very dilute HNO₃. Aqua regia is a freshly prepared 3:1 mixture of concentrated HCl and concentrated HNO₃. It is one of the few reagents capable of dissolving gold and platinum.
Q6
With reference to Activity 3.12, explain displacement reactions using the reaction of an iron nail with copper sulphate solution. Write the balanced equation and explain how such reactions help establish the reactivity series.
In Activity 3.12, an iron nail dipped in copper sulphate solution (blue) gets coated with a reddish-brown layer of copper metal, while the blue solution fades to pale green. A copper wire in iron sulphate solution shows no change. The reaction is: . Iron, being more reactive than copper, displaces copper from its salt solution. Copper cannot displace iron because it is less reactive. Such displacement reactions provide direct comparative evidence to arrange metals in the order of decreasing reactivity — a more reactive metal can always displace a less reactive metal from its salt solution.
Q7
What is the reactivity series? List the metals in decreasing order of reactivity from potassium to gold. State the significance of the position of hydrogen in the series and give one practical application of the reactivity series.
The reactivity series is a list of metals arranged in decreasing order of their reactivity, established through displacement experiments and reactions with oxygen, water, and acids. The series (most to least reactive): K > Na > Ca > Mg > Al > Zn > Fe > Pb > [H] > Cu > Hg > Ag > Au. Hydrogen is included in the series as a reference because metals above hydrogen can displace it from dilute acids (producing H₂ gas), while metals below hydrogen cannot. Practical application: the extraction method for a metal is chosen based on its position in the series — highly reactive metals (K, Na, Ca, Mg, Al) need electrolytic reduction; moderately reactive metals (Zn, Fe, Pb, Cu) are reduced using carbon; least reactive metals (Ag, Au) are found in the free state.
Q8
Explain the formation of ionic compounds taking the examples of sodium chloride (NaCl) and magnesium chloride (MgCl₂). Show the electron transfer using electron-dot notation and state what type of ions are formed.
Ionic (electrovalent) compounds form when a metal atom transfers its valence electrons to a non-metal atom, creating oppositely charged ions held together by strong electrostatic forces. Na (electronic configuration 2,8,1) loses its one valence electron to form Na⁺ (2,8), attaining the stable octet of neon. Cl (2,8,7) gains that one electron to form Cl⁻ (2,8,8), attaining the stable octet of argon. Na⁺ and Cl⁻ ions attract to form NaCl — the ionic compound does not exist as discrete molecules but as an aggregate of ions. For MgCl₂: Mg (2,8,2) loses two valence electrons to become Mg²⁺ (2,8). Two chlorine atoms each gain one electron to become two Cl⁻ ions (2,8,8). The Mg²⁺ cation and two Cl⁻ anions form MgCl₂. The cations are Na⁺ and Mg²⁺; anions are Cl⁻.
Q9
List the general properties of ionic compounds. Explain why ionic compounds have high melting and boiling points, why they conduct electricity in the molten state but not in the solid state, and why they are generally soluble in water but not in organic solvents.
Properties of ionic compounds: (i) They are solids at room temperature and are hard but brittle — they break on applying pressure. (ii) They have high melting and boiling points because strong electrostatic forces of attraction between oppositely charged ions require a large amount of energy to overcome. (iii) In the solid state, ions are held in fixed positions in a rigid lattice and cannot move — hence solids do not conduct electricity. In the molten (liquid) state, heat energy overcomes the electrostatic forces, freeing the ions to move towards opposite electrodes — hence they conduct electricity. In aqueous solution too, ions dissociate and conduct electricity. (iv) They are soluble in polar solvents like water (which separates the ions) but insoluble in non-polar organic solvents like petrol and kerosene.
Q10
Define the terms: mineral, ore, and gangue. Explain how the occurrence of metals in nature is related to their position in the reactivity series. Give examples of metals found in the free (native) state and in the combined state.
A mineral is a naturally occurring element or compound in the earth's crust. An ore is a mineral from which a metal can be extracted profitably (i.e., it contains a high percentage of the metal). Gangue refers to the earthy impurities like soil, sand, and silt present in mined ores. Low-reactivity metals (Au, Ag, Pt) are found in the free/native state because they do not react easily. Copper and silver also occur as sulphides or oxides. Medium-reactivity metals (Zn, Fe, Pb) occur as oxides, sulphides, or carbonates. Highly reactive metals (K, Na, Ca, Mg, Al) are never found free — they always exist as compounds because they readily react with other elements.
Q11
Distinguish between roasting and calcination. Why are carbonate and sulphide ores converted to their oxides before reduction? Write a balanced chemical equation for the roasting of zinc sulphide and the calcination of zinc carbonate, followed by reduction to zinc metal.
Roasting is the process of heating a sulphide ore strongly in the presence of excess air to convert it into its oxide. Calcination is heating a carbonate ore strongly in limited air to convert it into its oxide. Ores are converted to oxides first because it is easier to obtain a metal from its oxide than directly from its sulphide or carbonate — oxides are more readily reduced by carbon. Roasting of ZnS: . Calcination of ZnCO₃: . Reduction of ZnO: .
Q12
Describe the extraction of metals low in the activity series, taking mercury (from cinnabar) and copper (from copper glance) as examples. Write all the chemical equations involved.
Metals low in the activity series (below Cu) are very unreactive, and their oxides can be reduced to metals simply by heating. Mercury is extracted from its sulphide ore cinnabar (HgS) in two steps: (1) — roasting in air converts cinnabar to mercuric oxide. (2) — mercuric oxide decomposes on further heating to give mercury. Copper from copper glance (Cu₂S): (1) — roasting. (2) — copper(I) oxide reacts with remaining sulphide to produce copper metal.
Q13
Explain the thermit reaction with a balanced chemical equation and state one important application. Why cannot metals at the top of the activity series (like Na, Mg, Al) be reduced from their oxides using carbon?
The thermit reaction is a highly exothermic displacement reaction in which aluminium powder reduces iron(III) oxide to molten iron: . The enormous heat generated melts the iron, which is then used to join railway tracks or cracked machine parts. Metals at the top of the activity series (K, Na, Ca, Mg, Al) have a higher affinity for oxygen than carbon has. Therefore, carbon cannot pull oxygen away from their oxides — the reduction does not occur. Instead, these metals are obtained by electrolytic reduction of their molten compounds (e.g., molten chlorides).
Q14
Describe the process of electrolytic refining of copper with a labelled diagram explanation. What happens to the soluble and insoluble impurities during this process?
Electrolytic refining is used to obtain highly pure metals from impure metal obtained after reduction. For copper refining: the anode (+) is a thick block of impure copper, the cathode (-) is a thin strip of pure copper, and the electrolyte is an acidified copper sulphate (CuSO₄) solution. On passing electric current: at the anode, pure copper oxidises and dissolves as Cu²⁺ ions into the electrolyte: . At the cathode, Cu²⁺ ions from the electrolyte gain electrons and deposit as pure copper: . Soluble impurities dissolve into the solution, while insoluble impurities settle down below the anode as anode mud.
Q15
What is corrosion? Describe the conditions necessary for the rusting of iron using Activity 3.14. List any four methods of preventing corrosion and explain the principle of galvanisation. Also, define alloys and give two examples of alloys with their composition and one specific property each.
Corrosion is the slow degradation of a metal due to reaction with substances in its environment (moisture, acids, etc.). Activity 3.14 shows that iron rusts only when BOTH air (oxygen) and water (moisture) are present — nails in a tube with water and air rust, nails in boiled water covered with oil (no dissolved air) do not rust, and nails with anhydrous CaCl₂ (no moisture) do not rust. Prevention methods: (i) painting, (ii) oiling/greasing, (iii) galvanisation (coating with zinc), (iv) chrome plating, (v) anodising, (vi) alloying. Galvanisation coats iron with a thin layer of zinc — even if the zinc coating is scratched, zinc (being more reactive) corrodes preferentially and still protects the underlying iron. An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal. Example 1: Stainless steel — iron + nickel + chromium; hard and does not rust. Example 2: Solder — lead + tin; has a low melting point, used for welding electrical wires.