Mix Examples - Metals and Non-metals Questions in English

(N/A) Sodium $(Na)$ is a highly reactive metal that reacts vigorously with oxygen and moisture in the air,so it is stored in kerosene to prevent reaction.
$(b)$ Iodine $(I_2)$ is a non-metal that possesses a characteristic lustrous appearance and is purple-black in colour.
$(c)$ Gallium $(Ga)$ and Caesium $(Cs)$ have very low melting points and melt when held in the palm of a human hand.
$(d)$ Lead $(Pb)$ and Mercury $(Hg)$ are metals that are poor conductors of heat compared to other metals like copper or silver.

(N/A) Ionic compounds are formed by the transfer of electrons between atoms,resulting in the formation of positively charged cations and negatively charged anions.
These oppositely charged ions are held together by very strong electrostatic forces of attraction.
Because of these strong inter-ionic forces,a significant amount of energy is required to break the crystal lattice and overcome these forces,which results in high melting and boiling points.

(A) The metal $M$ is Aluminium $(Al)$.
The oxide $Al_{2}O_{3}$ is amphoteric in nature,meaning it reacts with both acids and bases.
Reaction with dilute sulphuric acid:
$Al_{2}O_{3}(s) + 3H_{2}SO_{4}(aq) \rightarrow Al_{2}(SO_{4})_{3}(aq) + 3H_{2}O(l)$
Reaction with dilute sodium hydroxide solution:
$Al_{2}O_{3}(s) + 2NaOH(aq) + 3H_{2}O(l) \rightarrow 2Na[Al(OH)_{4}](aq)$ (or Sodium aluminate $2NaAlO_{2} + H_{2}O$)

(N/A) $Na$ reacts vigorously with cold water.
$2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)$
$(b)$ $Mg$ does not react with cold water but reacts with hot water.
$Mg(s) + 2H_2O(l) \rightarrow Mg(OH)_2(aq) + H_2(g)$
$(c)$ $Fe$ does not react with cold or hot water but reacts with steam.
$3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$

(N/A) The element found in bones is Calcium $(Ca)$. The oxide of this metal used in the cement industry is Calcium oxide $(CaO)$,also known as quicklime.
When $CaO$ reacts with water,it forms Calcium hydroxide $(Ca(OH)_2)$,which is basic in nature and turns red litmus blue.
The chemical reaction is: $CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq)$.
$(b)$ Among the given metals,Magnesium $(Mg)$ reacts only with hot water.
The products formed are Magnesium hydroxide $(Mg(OH)_2)$ and Hydrogen gas $(H_2)$.
The chemical reaction is: $Mg(s) + 2H_2O(l) \rightarrow Mg(OH)_2(aq) + H_2(g)$.

(N/A) $(i)$ $3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$
$(ii)$ $Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g)$
$(iii)$ $2Cu(s) + O_2(g) \xrightarrow{\Delta} 2CuO(s)$

(N/A)

Metals	Non-Metals
$1$. Burn in air to form metal oxides which are basic in nature.	$1$. Burn in air to form non-metal oxides which are acidic in nature.
$2$. React with acids to form salt and hydrogen gas.	$2$. They are electron acceptors,so they cannot produce hydrogen gas from acids.
$3$. Metals form ionic compounds with non-metals. $e.g., NaCl$.	$3$. Non-metals form covalent compounds with other non-metals. $e.g., CCl_4$.

(N/A) $(i)$ Sodium $(Na)$ and potassium $(K)$ are highly reactive metals that react vigorously with moisture and oxygen in the air,so they are stored in kerosene to prevent reaction.
$(ii)$ Nickel $(Ni)$ and chromium $(Cr)$ are added to iron $(Fe)$ to produce stainless steel,which is resistant to corrosion.
$(iii)$ Gold $(Au)$ and silver $(Ag)$ are the most malleable and ductile metals,allowing them to be beaten into thin sheets and drawn into thin wires.

(N/A) $(i)$ Iron does not burn in air even on strong heating. At high temperature,it combines with oxygen of the air to form $Fe_{3}O_{4}$.
$3Fe(s) + 2O_{2}(g) \xrightarrow{\text{Heat}} Fe_{3}O_{4}(s)$
$(ii)$ Calcination of lead carbonate involves heating the ore in the absence of air.
$PbCO_{3}(s) \xrightarrow{\text{Heat}} PbO(s) + CO_{2}(g)$
$(iii)$ This is a thermite reaction where aluminium acts as a reducing agent.
$Cr_{2}O_{3}(s) + 2Al(s) \rightarrow Al_{2}O_{3}(s) + 2Cr(l) + \text{Heat}$

(N/A) The most reactive element is $B$ because it displaces both $A$ and $C$ from their respective compounds.
$(b)$ The least reactive element is $C$ because it is displaced by both $A$ and $B$ in the given reactions.
$(c)$ These are displacement reactions,specifically single displacement reactions.

(N/A) $(i)$ The order of decreasing reactivity is: Sodium > Aluminium > Copper > Gold.
$(ii)$ When aluminium powder is heated with manganese dioxide, a thermite reaction occurs, producing manganese metal and aluminium oxide: $4Al(s) + 3MnO_2(s) \xrightarrow{\Delta} 3Mn(l) + 2Al_2O_3(s) + \text{Heat}$.

(N/A) $(i)$ Sodium is so soft that it can be cut with a knife.
$(ii)$ It has a low density compared to most metals.
$(iii)$ It has a low melting point.

(N/A) $(i)$ Due to the strong electrostatic force of attraction between the oppositely charged ions,ionic compounds exist as solids. They are generally brittle and break into pieces when pressure is applied.
$(ii)$ Ionic compounds are generally soluble in water because the polar water molecules break the ionic bonds.
$(iii)$ Ionic compounds do not conduct electricity in the solid state because the ions are held in a fixed lattice. However,they conduct electricity in the aqueous or molten state because the ions become free to move.

(N/A) $(i)$ Mercury $(Hg)$ is a metal that does not stick to glass due to its high surface tension and weak adhesive forces with glass.
$(ii)$ Aluminium $(Al)$ is the metal commonly used in thermite welding because it acts as a strong reducing agent when reacting with metal oxides.
$(iii)$ Zinc oxide $(ZnO)$ is amphoteric in nature,meaning it reacts with both acids and bases to form salt and water.

(N/A) Sodium is a highly reactive metal with a very high affinity for oxygen. Carbon is a weaker reducing agent than sodium,so it cannot reduce sodium oxide to sodium metal.
$(b)$ Sodium is obtained from molten sodium chloride $(NaCl)$ by the process of electrolytic reduction (electrolysis). During this process,sodium ions $(Na^+)$ are reduced at the cathode to form metallic sodium $(Na)$.

(N/A) $(i)$ The insoluble impurities that settle down at the bottom of the anode during electrolytic refining are known as 'anode mud'. The cathode is made of a thin strip of pure metal.
$(ii)$ During the electrolytic refining of copper:
At anode: $Cu(s) \rightarrow Cu^{2+}(aq) + 2e^-$
At cathode: $Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)$

(N/A) Two metals purified by electrolytic refining are $Copper$ $(Cu)$ and $Zinc$ $(Zn)$.
In the electrolytic refining process:
$1$. The $Anode$ is made of an impure block of the metal.
$2$. The $Cathode$ is made of a thin strip of pure metal.
$3$. The $Electrolyte$ is an aqueous solution of a salt of the metal being refined.
The pure metal is deposited on the $Cathode$.

(N/A) Magnesium reacts with very dilute nitric acid to produce magnesium nitrate and hydrogen gas:
$Mg(s) + 2HNO_3(aq) \longrightarrow Mg(NO_3)_2(aq) + H_2(g)$
$(b)$ Aluminium powder reacts with iron$(III)$ oxide in a thermite reaction to produce iron and aluminium oxide:
$Fe_2O_3(s) + 2Al(s) \longrightarrow 2Fe(l) + Al_2O_3(s) + \text{Heat}$
$(c)$ Zinc sulphide is roasted in the presence of oxygen to produce zinc oxide and sulphur dioxide gas:
$2ZnS(s) + 3O_2(g) \xrightarrow{\Delta} 2ZnO(s) + 2SO_2(g)$

(N/A) The two categories for this reaction are:
$(i)$ Displacement reaction (specifically,a thermite reaction).
$(ii)$ Redox reaction (reduction-oxidation reaction).
$(b)$ Oxidation is defined as the addition of oxygen to a substance or the removal of hydrogen from a substance.
In the given reaction:
- Aluminium $(Al)$ is oxidised because it gains oxygen to form $Al_{2}O_{3}$.
- Iron$(III)$ oxide $(Fe_{2}O_{3})$ is the oxidising agent because it provides oxygen to aluminium.

(N/A) The reaction which is highly exothermic and in which metal is produced in the molten state is known as the thermit reaction.
Consider the following displacement reaction:
$Fe_{2}O_{3}(s) + 2Al(s) \xrightarrow{\text{Heat}} Al_{2}O_{3}(s) + 2Fe(l) + \text{Heat}$
In this reaction,when iron $(III)$ oxide $(Fe_{2}O_{3})$ is heated with aluminium powder $(Al)$,the amount of heat evolved is so large that the iron $(Fe)$ metal is produced in a molten state.
This molten iron flows into the gaps of the railway tracks or cracked machine parts and,upon cooling,solidifies to join them together.

(N/A) Formation of $MgO$:
$Mg (2, 8, 2) \to Mg^{2+} + 2e^-$
$O (2, 6) + 2e^- \to O^{2-}$
$Mg^{2+} + O^{2-} \to MgO$
$(b)$ $(i)$ Cation: Magnesium ion $(Mg^{2+})$.
$(ii)$ Anion: Oxide ion $(O^{2-})$.
$(c)$ Properties of ionic compounds:
$(i)$ They are generally crystalline solids and are soluble in water.
$(ii)$ They have high melting and boiling points due to strong electrostatic forces of attraction.
$(iii)$ They conduct electricity in the molten state or in aqueous solution because the ions are free to move.

(N/A) Aluminium forms a strong,non-porous,and protective layer of aluminium oxide $(Al_2O_3)$ on its surface upon exposure to air. This layer does not peel off easily and prevents further oxidation or corrosion of the underlying metal.
$(b)$ Dilute $HNO_3$ is a strong oxidising agent. When it reacts with metals like zinc,it oxidises the hydrogen gas produced during the reaction into water $(H_2O)$ and itself gets reduced to nitrogen oxides (such as $N_2O$,$NO$,or $NO_2$).
$(c)$ Aluminium has a much higher affinity for oxygen than carbon does. Therefore,carbon cannot effectively remove oxygen from aluminium oxide $(Al_2O_3)$ to produce aluminium metal. Consequently,electrolytic reduction is used instead of carbon reduction.

(A) Calcium starts floating when added to water because the bubbles of hydrogen gas formed during the reaction stick to the surface of the metal,making it buoyant.
$(b)$ Nitric acid $(HNO_{3})$ is a strong oxidizing agent. When it reacts with metals,it oxidizes the hydrogen gas produced to water $(H_{2}O)$ and itself gets reduced to nitrogen oxides $(N_{2}O, NO, NO_{2})$.
$(c)$ The chemical equation is: $3Fe(s) + 4H_{2}O(g) \rightarrow Fe_{3}O_{4}(s) + 4H_{2}(g)$.
The compound obtained is iron($II$,$III$) oxide,also known as ferroferric oxide $(Fe_{3}O_{4})$.

(N/A) An alloy is a homogeneous mixture of two or more metals,or a metal and a non-metal,which cannot be separated into its components by physical methods.
Properties that make alloys more useful than pure metals:
$(i)$ Alloys are often harder and stronger than their constituent pure metals.
$(ii)$ They are more resistant to corrosion.
$(iii)$ Their electrical conductivity is generally lower than that of pure metals.
$(iv)$ Their melting point is lower than that of pure metals.
Examples:
$(a)$ Brass (an alloy of $Cu$ and $Zn$) and Bronze (an alloy of $Cu$ and $Sn$) are not good conductors of electricity,whereas pure copper is an excellent conductor used in electrical circuits.
$(b)$ Solder (an alloy of $Pb$ and $Sn$) has a very low melting point,making it ideal for welding electrical wires together.

(N/A) Carbonate ores are subjected to calcination.
Example: Zinc carbonate $(ZnCO_{3})$ is calcined to form zinc oxide $(ZnO)$.
$ZnCO_{3}(s) \xrightarrow{\text{Heat}} ZnO(s) + CO_{2}(g)$
$(b)$ The calcined ore is obtained in the form of a metal oxide. It can be reduced to the metal using a reducing agent like carbon (coke) or by electrolytic reduction.
For the example of zinc oxide $(ZnO)$:
$ZnO(s) + C(s) \rightarrow Zn(s) + CO(g)$
$(c)$ Two metals used as reducing agents are Aluminum $(Al)$ and Sodium $(Na)$.

(A) Since the ore releases $SO_2$ gas upon heating in air,it is a sulphide ore.
The steps of metallurgy are as follows:
$(i)$ Concentration of ore: Usually done by the froth floatation process for sulphide ores.
$(ii)$ Roasting: The concentrated sulphide ore is heated strongly in the presence of excess air to convert it into metal oxide.
$(iii)$ Reduction: The metal oxide is reduced to the metal by heating with a reducing agent like carbon or by self-reduction.
$(b)$ $(i)$ $Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)$ : This reaction will take place because $Zn$ is more reactive than $Cu$ and can displace it from its salt solution.
$(ii)$ $Fe(s) + ZnSO_4(aq) \rightarrow FeSO_4(aq) + Zn(s)$ : This reaction will not take place because $Fe$ is less reactive than $Zn$ and cannot displace $Zn$ from its salt solution.

(N/A) $(i)$ Lead is a metal that is not lustrous.
$(ii)$ Mercury is a metal that exists as a liquid at room temperature.
$(iii)$ Gallium and Caesium are metals that have very low melting points.
$(iv)$ Iodine is a non-metal that is lustrous.
$(v)$ Graphite (an allotrope of carbon) is a non-metal that is a good conductor of electricity.

(N/A) $(i)$ Calcium starts floating because the bubbles of hydrogen gas formed stick to its surface.
$Ca(s) + 2H_2O(l) \longrightarrow Ca(OH)_2(aq) + H_2(g)$
$(ii)$ In dry air,sodium reacts with oxygen to form sodium oxide. In moist air,it reacts with water vapor to form sodium hydroxide.
$4Na(s) + O_2(g) \longrightarrow 2Na_2O(s)$ (In dry air)
$2Na(s) + 2H_2O(l) \longrightarrow 2NaOH(aq) + H_2(g) + \text{Heat}$ (In moist air)
$(iii)$ Iron displaces copper from copper sulphate solution due to higher reactivity.
$Fe(s) + CuSO_4(aq) \longrightarrow FeSO_4(aq) + Cu(s)$
$(iv)$ Potassium reacts violently with cold water,releasing hydrogen gas and heat.
$2K(s) + 2H_2O(l) \longrightarrow 2KOH(aq) + H_2(g) + \text{Heat}$
$(v)$ Carbon dioxide dissolves in water under high pressure to form carbonic acid.
$CO_2(g) + H_2O(l) \longrightarrow H_2CO_3(aq)$

(N/A) $(i)$ Zinc reacts with sulphuric acid to form zinc sulphate and hydrogen gas: $Zn(s) + H_{2}SO_{4}(aq) \rightarrow ZnSO_{4}(aq) + H_{2}(g)$
$(ii)$ Zinc reacts with hydrochloric acid to form zinc chloride and hydrogen gas: $Zn(s) + 2HCl(aq) \rightarrow ZnCl_{2}(aq) + H_{2}(g)$
$(iii)$ Zinc is less reactive than aluminium,so it cannot displace aluminium from aluminium chloride: $Zn(s) + AlCl_{3}(aq) \rightarrow \text{No reaction}$
$(iv)$ Zinc reacts with sodium hydroxide to form sodium zincate and hydrogen gas: $Zn(s) + 2NaOH(aq) \rightarrow Na_{2}ZnO_{2}(aq) + H_{2}(g)$
$(v)$ Zinc reacts with concentrated nitric acid to form zinc nitrate,nitrogen dioxide,and water: $Zn(s) + 4HNO_{3}(conc.) \rightarrow Zn(NO_{3})_{2}(aq) + 2NO_{2}(g) + 2H_{2}O(l)$

(N/A) $(i)$ Ionic compounds have high melting and boiling points because a considerable amount of energy is required to break the strong inter-ionic electrostatic forces of attraction.
$(ii)$ Ionic compounds are soluble in water because the polar water molecules interact with the ions,overcoming the lattice energy and causing them to dissociate.
$(iii)$ They are solids and somewhat hard because of the strong electrostatic force of attraction between the oppositely charged ions,which holds them in a rigid crystal lattice structure.
$(iv)$ Ionic compounds conduct electricity in the molten state because the heat overcomes the electrostatic forces,allowing the ions to become free to move and carry charge.
$(v)$ Metals have a tendency to lose electrons to attain a stable,completely filled valence shell configuration (noble gas configuration).

(N/A) Silver: It reacts with sulfur compounds in the air to form a black layer of silver sulfide.
$(b)$ Copper: It reacts with moist carbon dioxide in the air to form a green layer of basic copper carbonate.
$(c)$ Platinum: It is a noble metal and is highly resistant to corrosion.
$(d)$ Nickel or Chromium: These are added to iron to make stainless steel.
$(e)$ Sodium: It is a soft,reactive metal that does not exhibit the typical metallic luster because it reacts rapidly with air and moisture.

(A) $(i)$ The metal $E$ is Sodium $(Na)$.
$(ii)$ Reaction with air: $4Na + O_{2} \rightarrow 2Na_{2}O$.
Reaction with water: $Na_{2}O + H_{2}O \rightarrow 2NaOH$.
$(iii)$ Sodium is obtained from molten sodium chloride $(NaCl)$ by the process of electrolytic reduction (electrolysis). In this process,an electric current is passed through the molten salt.
At cathode: $Na^{+} + e^{-} \rightarrow Na$ (Reduction).
At anode: $2Cl^{-} \rightarrow Cl_{2} + 2e^{-}$ (Oxidation).

(N/A) $(i)$ Gold and platinum are noble metals that are highly lustrous and resistant to corrosion,making them ideal for jewellery.
$(ii)$ Copper is less reactive than hydrogen in the reactivity series,so it cannot displace hydrogen from dilute acids.
$(iii)$ Stainless steel is an alloy of iron with chromium and nickel,which forms a protective oxide layer that prevents rusting.
$(iv)$ Metals are malleable (can be beaten into sheets) and ductile (can be drawn into wires),allowing them to be shaped as required.
$(v)$ $HNO_3$ is a strong oxidizing agent. When it reacts with metals,it oxidizes the produced hydrogen gas $(H_2)$ into water $(H_2O)$ instead of releasing it.

(N/A) $Mg$ has a higher affinity for oxygen than carbon does,meaning $MgO$ is more stable than $CO_2$ at the temperatures required for reduction.
$(b)$ Sodium is obtained by the electrolytic reduction of molten sodium chloride $(NaCl)$.
At cathode: $2Na^+ + 2e^- \rightarrow 2Na$
At anode: $2Cl^- \rightarrow Cl_2 + 2e^-$
$(c)$ Copper is obtained from its sulphide ore $(Cu_2S)$ by roasting followed by self-reduction:
$2Cu_2S + 3O_2 \rightarrow 2Cu_2O + 2SO_2$
$2Cu_2O + Cu_2S \rightarrow 6Cu + SO_2$

(A-D) The coating on silver is silver sulphide $(Ag_2S)$ and on copper is basic copper carbonate $(CuCO_3 \cdot Cu(OH)_2)$.
$(b)$ Galvanisation is a method of protecting steel and iron from rusting by coating them with a thin layer of zinc. The purpose is to prevent corrosion,as the zinc layer acts as a sacrificial anode and protects the underlying metal even if the coating is scratched.
$(c)$ An alloy is a homogeneous mixture of two or more metals or a metal and a non-metal. It is prepared by melting the primary metal and then dissolving the other elements in it in a definite proportion,followed by cooling to room temperature.
$(i)$ When a small quantity of carbon is mixed with iron,it becomes hard and strong.
$(ii)$ When nickel and chromium are mixed with iron,it becomes stainless steel,which is hard and does not rust.

(N/A) $(i)$ Extraction of copper from its sulphide ore $(Cu_2S)$:
$(a)$ Roasting: The sulphide ore is heated in the presence of excess air to convert it into copper$(I)$ oxide.
$2Cu_2S(s) + 3O_2(g) \xrightarrow{\text{heat}} 2Cu_2O(s) + 2SO_2(g)$
$(b)$ Reduction: The copper$(I)$ oxide is then heated with more copper$(I)$ sulphide to obtain copper metal.
$2Cu_2O(s) + Cu_2S(s) \xrightarrow{\text{heat}} 6Cu(s) + SO_2(g)$
$(ii)$ Refining of metals: It is the process of removing impurities from the crude metal obtained after reduction to get pure metal.
Electrolytic refining of copper:
Anode: Impure copper block.
Cathode: Pure copper strip.
Electrolyte: Acidified copper sulphate solution $(CuSO_4 + H_2SO_4)$.

(N/A)

Roasting	Calcination
Heating in the presence of excess air. Done for sulphide ores.	Heating in a limited supply of air. Done for carbonate ores.

$(b)$ $(i)$ Self-reduction: Less reactive metals like mercury $(Hg)$ or copper $(Cu)$ can be reduced by heating alone.
$2HgO \xrightarrow{\text{heat}} 2Hg + O_2$
$2Cu_2O + Cu_2S \xrightarrow{\text{heat}} 6Cu + SO_2$
$(ii)$ Reduction using carbon or displacement: Moderately reactive metals are reduced by heating with carbon or using highly reactive metals like aluminum $(Al)$ in thermite reactions.
$ZnO + C \rightarrow Zn + CO$
$Fe_2O_3 + 2Al \rightarrow 2Fe + Al_2O_3 + \text{Heat}$
$(iii)$ Electrolytic reduction: Highly reactive metals (e.g.,$Na, Mg, Ca$) are extracted by the electrolysis of their molten chlorides because they have a high affinity for oxygen.

(N/A) Metals are chemically reactive in nature. They readily react with oxygen,sulfur,and other elements to form compounds such as oxides,sulfides,halides,carbonates,and sulfates. Due to this high reactivity,they are rarely found in their free (native) state in nature.
$(b)$ Metals are good conductors of heat because they possess free electrons. These free electrons move throughout the metallic lattice,effectively transferring kinetic energy from hotter regions to cooler regions,which facilitates the rapid conduction of heat.

(B) Element $A$ is a non-metal because non-metal oxides are typically acidic or neutral in nature.
$AO$ (e.g.,$CO$ - Carbon monoxide) is a neutral oxide.
$AO_2$ (e.g.,$CO_2$ - Carbon dioxide) is an acidic oxide.
Since metals generally form basic oxides,the formation of neutral and acidic oxides confirms that $A$ is a non-metal.

(N/A) The constituents of the given alloys are as follows:
$(a)$ Brass: It is an alloy of Copper $(Cu)$ and Zinc $(Zn)$.
$(b)$ Bronze: It is an alloy of Copper $(Cu)$ and Tin $(Sn)$.
$(c)$ Solder: It is an alloy of Lead $(Pb)$ and Tin $(Sn)$.

(N/A) Copper is more reactive than silver. In the reactivity series,it is placed above silver. So,copper displaces silver from its salt.
Observation:
$(i)$ The colour of the solution turns bluish green due to the formation of copper nitrate.
$(ii)$ $A$ greyish-white deposit of silver is observed on the surface of the copper wire.
$(b)$ The balanced chemical equation is:
$Cu(s) + 2AgNO_3(aq) \to Cu(NO_3)_2(aq) + 2Ag(s)$

(N/A) Most metal oxides are basic in nature. Metal oxides that are soluble in water are called basic oxides,and their aqueous solutions are known as alkalis.
$(b)$ Metals like magnesium,aluminium,and zinc react with atmospheric oxygen to form a thin,protective oxide layer on their surface. This layer is composed of the metal oxide. Its importance lies in the fact that it prevents further oxidation and corrosion of the underlying metal.
$(c)$ Alkali metals like $Li$ (Lithium),$Na$ (Sodium),and $K$ (Potassium) have a low density and weak metallic bonding due to their high electropositive character,making them soft enough to be cut with a knife.

(N/A) The rate of formation of hydrogen gas bubbles is slow with calcium (less violent) in comparison to that with potassium (more violent). The rate of hydrogen gas evolution is faster in the case of potassium,which generates enough heat to ignite the hydrogen gas.
$Ca + 2H_2O \rightarrow Ca(OH)_2 + H_2 \uparrow$
(The reaction is less vigorous,so the hydrogen gas does not catch fire.)
$(b)$ $(i)$ $Mg$ (Magnesium) or $Mn$ (Manganese).
$(ii)$ $Cu$ (Copper),$Hg$ (Mercury),$Ag$ (Silver),or $Au$ (Gold).
$(iii)$ $Fe$ (Iron),$Zn$ (Zinc),or $Al$ (Aluminium).

(N/A) $(i)$

Electrolytic reduction	Electrolytic refining
It is a process of obtaining metals from their molten chlorides or oxides by electrolysis. The metals are deposited at the cathode,whereas chlorine or oxygen is liberated at the anode.	It is a process of refining impure metals obtained by any reduction process. The impure metal is taken as the anode,pure metal as the cathode,and a metal salt solution as the electrolyte.

$(ii)$ Minerals are naturally occurring elements or compounds in the Earth's crust. Ores are those minerals from which metals can be extracted profitably and conveniently.
$(iii)$ An alloy is a homogeneous mixture of two or more metals or a metal and a non-metal. An amalgam is a specific type of alloy in which one of the components is mercury $(Hg)$.

(N/A) Sulphur $(S)$ is a non-metal because:
$1.$ It is brittle in its solid state and possesses low melting and boiling points.
$2.$ It is an electronegative element that forms anions $(S^{2-})$ and acts as a poor conductor of heat and electricity.
$3.$ It reacts with oxygen to form acidic oxides,which dissolve in water to produce acids.
Chemical equation:
$(a) S(s) + O_2(g) \rightarrow SO_2(g)$
$(b) SO_2(g) + H_2O(l) \rightarrow H_2SO_3(aq)$
These properties are characteristic of non-metals,confirming that sulphur is a non-metal.

(N/A) Magnesium $(Mg)$ is considered a metal due to the following reasons:
$1.$ Magnesium is a hard solid with a high melting and boiling point,which are characteristic physical properties of metals.
$2.$ Magnesium atoms lose two electrons to form electropositive $(Mg^{2+})$ ions and act as good conductors of heat and electricity.
$3.$ Magnesium reacts with oxygen to form a basic oxide,which dissolves in water to form a base. This is a characteristic chemical property of metals.
Chemical equations:
$(a) 2Mg(s) + O_2(g) \rightarrow 2MgO(s)$
$(b) MgO(s) + H_2O(l) \rightarrow Mg(OH)_2(aq)$
Since magnesium exhibits these metallic properties,it is classified as a metal.

(A) We can store copper sulphate in a silver vessel because silver is less reactive than copper,so no displacement reaction occurs. However,we cannot store silver nitrate solution in a copper vessel because copper is more reactive than silver and will displace silver from its salt solution.
$Cu(s) + 2AgNO_{3}(aq) \rightarrow Cu(NO_{3})_{2}(aq) + 2Ag(s)$
$(b)$ Nitric acid $(HNO_{3})$ is a strong oxidizing agent. When zinc reacts with dilute $HNO_{3}$,the hydrogen gas produced is immediately oxidized to water $(H_{2}O)$,while the nitric acid itself gets reduced to nitrogen oxides like $N_{2}O$,$NO$,or $NO_{2}$.
$(c)$ Tin is less reactive than zinc. Zinc is more reactive and can easily react with organic acids present in food,potentially contaminating it. Tin provides a protective,non-reactive layer that preserves the food.

(N/A) $(i)$ Carbon acts as a reducing agent. In the reactivity series,zinc is less reactive than carbon,so carbon can displace oxygen from zinc oxide to form zinc metal:
$ZnO(s) + C(s) \rightarrow Zn(s) + CO(g)$
However,calcium is much more reactive than carbon and has a very high affinity for oxygen. Therefore,carbon cannot reduce calcium oxide to calcium.
$(ii)$ Copper and aluminium are used for electricity transmission because they possess high electrical conductivity and low electrical resistance,which minimizes energy loss during transmission.

(N/A) $(i)$ Alloys generally have higher electrical resistivity than their constituent pure metals. This property allows them to produce more heat when current passes through them. Additionally,they do not oxidize (burn) readily at high temperatures.
$(ii)$ Gold is a noble metal and is highly resistant to corrosion. It does not react with oxygen,moisture,or atmospheric gases,which is why it retains its luster for years.
$(iii)$ Iron reacts with oxygen and moisture present in the air to form hydrated iron$(III)$ oxide,commonly known as rust. Painting the grill creates a protective barrier that prevents the iron from coming into contact with air and moisture,thereby preventing rusting.

(B) Magnesium reacts with hot water to produce magnesium hydroxide and hydrogen gas $(Mg + 2H_2O \rightarrow Mg(OH)_2 + H_2)$.
As the reaction proceeds, small bubbles of hydrogen gas are produced.
These bubbles of hydrogen gas stick to the surface of the magnesium ribbon.
Due to the buoyancy provided by these gas bubbles, the magnesium ribbon starts floating on the surface of the water.