Class 10 Science · Chapter 3

Metals and Non-metals

344 Class 10 Science questions on Metals and Non-metals from the NCERT textbook, NCERT Exemplar and CBSE board papers 2017–2026, each with its answer and a step-by-step solution.

CBSE Class 10 Science previous year papers, with answers →

Questions · page 17 of 18

1. Q321 CBSE 2019
   How ​is the method of ​extraction of metals ​high up ​in the ​reactivity ​series different from that for metals in the middle ? Why can the same process not be applied for them ? Name the process used for the extraction of ​these metals.
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   Middle metals are reduced from oxides by carbon; high metals are extracted by electrolysis of molten compounds because carbon cannot reduce their oxides.

   Metals in the middle of the reactivity series (zinc, iron, lead) are obtained by first converting their sulphide or carbonate ores into oxides, by roasting or calcination, and then reducing the oxides with a reducing agent such as carbon, for example $\text{ZnO}+\text{C}⟶\text{Zn}+\text{CO}$. Metals high up in the series (potassium, sodium, calcium, magnesium, aluminium) are obtained by the electrolysis of their molten chlorides or oxides, the metal being deposited at the cathode. The same process cannot be applied because these metals are very reactive: they have more affinity for oxygen than carbon has, so carbon cannot reduce their oxides. The method of extraction depends on the reactivity of the metal. The process used is electrolytic reduction; the metal obtained can then be purified by electrolytic refining.

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2. Q322 CBSE 2019
   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.
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   The ore is a carbonate such as zinc carbonate; it is calcined to the oxide, which carbon reduces to the metal.

   The ore is a carbonate ore: a carbonate gives off carbon dioxide with dilute hydrochloric acid, seen as brisk effervescence (${\text{ZnCO}}_3+2\text{HCl}\to {\text{ZnCl}}_2+{\text{H}}_2\text{O}+{\text{CO}}_2$). Example: zinc carbonate, ${\text{ZnCO}}_3$ (calamine). Steps to obtain the metal from the enriched ore: (i) Calcination: the ore is heated strongly in a limited supply of air, which converts it into the metal oxide: ${\text{ZnCO}}_3\stackrel{\text{heat}}{\to }\text{ZnO}+{\text{CO}}_2$. (ii) Reduction: the metal oxide is heated with carbon (coke), which reduces it to the metal: $\text{ZnO}+\text{C}\to \text{Zn}+\text{CO}$.

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3. Q323 CBSE 2019
   1. (a)​List ​in tabular form three chemical properties on the ​basis of which we can differentiate between ​a metal and a non-metal.
   2. (b)Give reasons for the following: (i) Most metals conduct electricity well. (ii) The ​reaction of iron (III) oxide [${\text{Fe}}_2{\text{O}}_3$] ​with heated aluminium is used to ​join cracked machine parts.
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   Metals and non-metals differ in oxides, acid reactions and ions; metals conduct through free electrons; the thermit reaction gives molten iron that joins parts.

   (a) Three chemical properties that distinguish a metal from a non-metal (the table is set out in step 5): Reaction with oxygen: a metal forms a basic oxide, for example $\text{MgO}$; a non-metal forms an acidic or a neutral oxide, for example ${\text{SO}}_2$ or $\text{CO}$. Reaction with dilute acids: a metal such as zinc displaces hydrogen from dilute acids; a non-metal does not displace hydrogen from dilute acids. Ions formed: a metal forms positive ions by losing electrons, for example ${\text{Na}}^{+}$; a non-metal forms negative ions by gaining electrons, for example ${\text{Cl}}^{-}$. (b) (i) Metals have free electrons: the outer electrons of metal atoms are loosely held and easily lost, so they move freely through the metal and carry the electric current. (ii) Aluminium is more reactive than iron, so it displaces iron from iron(III) oxide: ${\text{Fe}}_2{\text{O}}_3+2\text{Al}\to 2\text{Fe}+{\text{Al}}_2{\text{O}}_3+\text{heat}$. The reaction is so highly exothermic that the iron formed is molten; the molten iron fills the crack between the machine parts and joins them when it solidifies (thermit welding).

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4. Q324 CBSE 2019
   1. (a)Name two metals which are obtained ​from their ores by ​simple heating.
   2. (b)Differentiate between calcination ​and ​roasting, taking ​examples of zinc ores.
   3. (c)​What is thermit reaction ​? ​State its significance.
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   Mercury and copper come from simple heating; calcination heats carbonate ores, roasting heats sulphide ores; the thermit reaction yields molten iron to join rails.

   (a) Mercury and copper. Cinnabar, $\text{HgS}$, heated in air gives mercuric oxide, which on further heating gives mercury; copper glance, ${\text{Cu}}_2\text{S}$, heated in air gives copper. (b) Calcination is heating a carbonate ore strongly in limited air (or in the absence of air), which drives out carbon dioxide: ${\text{ZnCO}}_3(s)\stackrel{\text{heat}}{\to }\text{ZnO}(s)+{\text{CO}}_2(g)$. Roasting is heating a sulphide ore strongly in excess air, which gives off sulphur dioxide: $2\text{ZnS}(s)+3{\text{O}}_2(g)\stackrel{\text{heat}}{\to }2\text{ZnO}(s)+2{\text{SO}}_2(g)$. Both change the ore into zinc oxide. (c) The thermit reaction is the reaction of iron(III) oxide with aluminium, in which the more reactive aluminium displaces iron: ${\text{Fe}}_2{\text{O}}_3(s)+2\text{Al}(s)\to 2\text{Fe}(l)+{\text{Al}}_2{\text{O}}_3(s)+\text{heat}$. It is so highly exothermic that the iron is produced in the molten state. Significance: it is used to join railway tracks and cracked machine parts.

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5. Q325 CBSE 2019
   (a) Write chemical ​equations ​for the following reactions:

   1. (i)​Calcium ​metal reacts with ​water.
   2. (ii)Cinnabar is heated in the presence of air.
   3. (iii)Manganese ​dioxide is heated with aluminium powder. (b) What are alloys ? List two properties of alloys.
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   Ca gives calcium hydroxide and hydrogen; cinnabar is roasted to HgO, then Hg; Al displaces Mn; an alloy is a harder, corrosion-resistant mixture.

   (a) (i) $\text{Ca}+2{\text{H}}_2\text{O}⟶{\text{Ca(OH)}}_2+{\text{H}}_2$ (ii) $2\text{HgS}+3{\text{O}}_2\stackrel{\text{heat}}{\to }2\text{HgO}+2{\text{SO}}_2$ (on further heating, $2\text{HgO}\stackrel{\text{heat}}{\to }2\text{Hg}+{\text{O}}_2$) (iii) $3{\text{MnO}}_2+4\text{Al}⟶3\text{Mn}+2{\text{Al}}_2{\text{O}}_3+\text{heat}$ (b) An alloy is a homogeneous mixture of two or more metals, or of a metal and a non-metal. Two properties: alloys are harder and stronger than the pure metals, and they resist corrosion better. (They also have lower melting points and lower electrical conductivity; any two properties are accepted.)

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6. Q326 CBSE 2019
   What would ​you observe on adding zinc ​granules ​to ​freshly prepared ferrous sulphate solution ? ​Give reason for ​your answer.
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   The pale green solution turns colourless and iron deposits on the zinc, because zinc is more reactive than iron and displaces it.

   Observation: the pale green colour of the ferrous sulphate solution fades and the solution becomes colourless, and a black (dark grey) coating is deposited on the zinc granules. Reason: zinc is more reactive than iron (it lies above iron in the reactivity series), so it displaces iron from ferrous sulphate. Colourless zinc sulphate forms and the displaced iron deposits on the zinc: $\text{Zn(s)}+{\text{FeSO}}_4\text{(aq)}\to {\text{ZnSO}}_4\text{(aq)}+\text{Fe(s)}$

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7. Q327 CBSE 2019
   What ​is observed ​after about 1 hour of adding the ​strips ​of copper and aluminium separately ​to ferrous sulphate solution ​filled in ​two beakers ? Name the reaction if any change in colour is noticed. Also, write chemical equation for ​the reaction.
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   Copper shows no change; aluminium turns the solution colourless with an iron deposit, a displacement reaction.

   Copper strip: no change is observed - the solution stays pale green and nothing is deposited on the copper, because copper is less reactive than iron and cannot displace it. Aluminium strip: the pale green colour of the solution fades and the solution becomes colourless, and a greyish-black deposit of iron forms on the aluminium strip, because aluminium is more reactive than iron. The reaction in the aluminium beaker is a displacement reaction: $2\text{Al}(s)+3{\text{FeSO}}_4(aq)⟶{\text{Al}}_2({\text{SO}}_4{)}_3(aq)+3\text{Fe}(s)$

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8. Q328 CBSE 2019
   What would a student report nearly ​after 30 ​minutes of placing ​duly ​cleaned strips of aluminium, copper, iron and zinc in freshly ​prepared iron ​sulphate solution taken in four beakers ?
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   Aluminium and zinc turn the solution colourless with an iron deposit; copper and iron show no change.

   Beakers with aluminium and zinc: the pale green colour of the iron sulphate solution fades and the solution becomes colourless, and a greyish-black deposit of iron forms on the strips. Beakers with copper and iron: no change - the solution stays pale green and the strips are unaffected. Reason: aluminium and zinc are more reactive than iron and displace it: $2\text{Al}+3{\text{FeSO}}_4\to {\text{Al}}_2({\text{SO}}_4{)}_3+3\text{Fe}$ and $\text{Zn}+{\text{FeSO}}_4\to {\text{ZnSO}}_4+\text{Fe}$. Copper is less reactive than iron, and iron cannot displace itself, so no reaction occurs in those beakers.

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9. Q329 CBSE 2019
   1. (a)Arrange the ​following metals ​in the increasing order of their reactivities : Copper, Zinc, ​Aluminium and Iron
   2. (b)List two observations ​you would record in your notebook 30 minutes after adding iron filings to copper ​sulphate ​solution.
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   Reactivity increases as $\text{Cu}<\text{Fe}<\text{Zn}<\text{Al}$; iron filings get a reddish-brown copper coat and the blue solution turns light green.

   (a) Copper, iron, zinc, aluminium (from the least to the most reactive), that is, $\text{Cu}<\text{Fe}<\text{Zn}<\text{Al}$. (b) (i) A reddish-brown coating of copper is deposited on the iron filings. (ii) The blue colour of the copper sulphate solution fades and it turns light green. (Iron is more reactive than copper and displaces it: $\text{Fe}(\text{s})+{\text{CuSO}}_4(\text{aq})\to {\text{FeSO}}_4(\text{aq})+\text{Cu}(\text{s})$.)

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10. Q330 CBSE 2019
    A pale ​green solution of ferrous sulphate was taken in four separate test tubes marked I, II, III and IV. ​Pieces of Cu, Zn and Al were dropped ​in test tubes II, III and IV respectively. In which case(s)

    1. (a)the colour of ferrous ​sulphate solution will ​match with the colour in test tube (I) ? Give reason.
    2. (b)the colour of ferrous ​sulphate solution will fade and black mass will be deposited on the surface of the ​metal ​?
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    Tube II (copper) stays pale green like tube I; in tubes III (zinc) and IV (aluminium) the colour fades and black iron deposits.

    (a) Test tube II (copper). Copper is less reactive than iron, so it cannot displace iron from ferrous sulphate; no reaction takes place and the solution stays pale green, like the solution in test tube I. (b) Test tubes III (zinc) and IV (aluminium). Zinc and aluminium are more reactive than iron and displace it: $\text{Zn}+{\text{FeSO}}_4\to {\text{ZnSO}}_4+\text{Fe}$ and $2\text{Al}+3{\text{FeSO}}_4\to {\text{Al}}_2({\text{SO}}_4{)}_3+3\text{Fe}$. The colourless zinc and aluminium sulphates replace the pale green ferrous sulphate, so the colour fades, and the iron set free is deposited as a black mass on the metal.

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11. Q331 CBSE 2018
    1. (a)Write ​the steps involved in the extraction of pure metals in the middle of the activity series from their carbonate ores.
    2. (b)How is copper extracted from its sulphide ore ? ​Explain the various steps ​supported by chemical equations. Draw labelled diagram ​for ​the electrolytic ​refining of copper.
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    Carbonate ores are calcined, reduced and refined; copper glance is roasted, self-reduced to copper and refined electrolytically.

    (a) After the ore is concentrated, a metal in the middle of the activity series is extracted from its carbonate ore in three steps, in this order: (i) Calcination: the carbonate ore is heated strongly in limited air and changes into the metal oxide, for example ${\text{ZnCO}}_3\stackrel{\text{heat}}{\to }\text{ZnO}+{\text{CO}}_2$. (ii) Reduction: the metal oxide is heated with a reducing agent such as carbon, for example $\text{ZnO}+\text{C}\to \text{Zn}+\text{CO}$. (iii) Purification (refining) of the impure metal, usually by electrolytic refining. (b) Copper glance, ${\text{Cu}}_2\text{S}$, is heated in air (roasting), and part of it is changed into copper(I) oxide: $2{\text{Cu}}_2\text{S}+3{\text{O}}_2\stackrel{\text{heat}}{\to }2{\text{Cu}}_2\text{O}+2{\text{SO}}_2$. On further heating, the oxide reacts with the remaining sulphide and copper is formed (self-reduction): $2{\text{Cu}}_2\text{O}+{\text{Cu}}_2\text{S}\stackrel{\text{heat}}{\to }6\text{Cu}+{\text{SO}}_2$. The impure copper is refined electrolytically (labelled diagram in the figure): the electrolyte is acidified copper sulphate solution, the anode is a block of impure copper connected to the positive terminal, and the cathode is a thin strip of pure copper connected to the negative terminal. Copper dissolves from the anode, $\text{Cu}\to {\text{Cu}}^{2+}+2e^{-}$, and pure copper is deposited on the cathode, ${\text{Cu}}^{2+}+2e^{-}\to \text{Cu}$; the insoluble impurities settle below the anode as anode mud.

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12. Q332 CBSE 2018
    1. (a)List ​in tabular form any three chemical properties on ​the basis of which ​metals and non-metals are differentiated.
    2. (b)​State ​two ways ​to prevent the ​rusting of iron.
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    Metals and non-metals differ in their oxides, reaction with dilute acids and ions formed; rusting is prevented by painting, oiling or galvanising.

    (a) Metals versus non-metals (chemical properties): 1. Nature of oxides: Metals form basic oxides, e.g. $\text{MgO}$ (a few, such as ${\text{Al}}_2{\text{O}}_3$ and $\text{ZnO}$, are amphoteric); non-metals form acidic oxides, e.g. ${\text{SO}}_2$ and ${\text{CO}}_2$, or neutral oxides, e.g. $\text{CO}$. 2. Reaction with dilute acids: Metals above hydrogen in the activity series displace hydrogen from dilute acids, e.g. $\text{Zn}+2\text{HCl}\to {\text{ZnCl}}_2+{\text{H}}_2$; non-metals do not displace hydrogen from dilute acids. 3. Ions formed: Metals lose electrons to form positive ions (they are electropositive), e.g. $\text{Na}\to {\text{Na}}^{+}+e^{-}$; non-metals gain electrons to form negative ions (they are electronegative), e.g. $\text{Cl}+e^{-}\to {\text{Cl}}^{-}$. Other correct chemical properties, such as ionic chlorides of metals against covalent chlorides of non-metals, also earn credit. (b) Rusting of iron can be prevented by: 1. painting, oiling or greasing the iron, which keeps air and moisture away from its surface; 2. galvanising, i.e. coating the iron with a thin layer of zinc. Chrome plating, tin plating or making alloys such as stainless steel also prevent rusting.

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13. Q333 CBSE 2018
    Distinguish ​between metals and non-metals ​on ​the basis ​of ​

    1. (i)two physical ​properties and ​
    2. (ii)​three chemical properties.
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    Metals are malleable, ductile conductors forming basic oxides and positive ions; non-metals are poor conductors forming acidic or neutral oxides and negative ions.

    (i) Physical properties: 1. Malleability and ductility: metals are malleable and ductile; non-metals are neither malleable nor ductile (solid non-metals are brittle). 2. Conduction: metals are good conductors of heat and electricity; non-metals are poor conductors (except graphite, which conducts electricity). (ii) Chemical properties: 1. Oxides: metals form basic oxides (some, such as aluminium oxide and zinc oxide, are amphoteric); non-metals form acidic or neutral oxides. 2. Nature: metals are electropositive, losing electrons to form positive ions; non-metals are electronegative, gaining electrons to form negative ions. 3. Dilute acids: metals above hydrogen in the reactivity series displace hydrogen from dilute acids; non-metals do not displace hydrogen from dilute acids. Other correct properties also earn credit.

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14. Q334 CBSE 2018
    1. (a)Explain why the method of extraction ​of metals high up in the reactivity ​series is different from the method of extraction ​of metals in ​the middle of the series. Why cannot the same process be applied for them ?
    2. (b)With the help of equation explain the process ​of extraction ​of sodium.
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    Carbon cannot reduce the oxides of highly reactive metals, so they are electrolysed in molten form; sodium forms at the cathode from molten NaCl.

    (a) Metals in the middle of the reactivity series (zinc, iron, lead, copper) are moderately reactive and occur mainly as sulphides or carbonates. Their ores are first converted into oxides, by roasting (sulphide ore heated strongly in excess air) or calcination (carbonate ore heated strongly in limited air), and the oxides are then reduced to the metals by heating with carbon, e.g. $\text{ZnO}+\text{C}\to \text{Zn}+\text{CO}$. Metals high up in the series (potassium, sodium, calcium, magnesium, aluminium) are very reactive and have a greater affinity for oxygen than carbon has. So carbon cannot reduce their oxides, and these metals are obtained by electrolytic reduction of their molten compounds (sodium, calcium and magnesium from their molten chlorides, aluminium from its molten oxide). The same process cannot be applied because carbon holds oxygen less strongly than these metals do: heating their oxides with carbon does not set the metal free, so they are obtained by electrolytic reduction. (b) Sodium is extracted by the electrolysis of molten (fused) sodium chloride. When an electric current is passed, ${\text{Na}}^{+}$ ions move to the cathode (negative electrode) and gain electrons to form sodium metal, which collects there; ${\text{Cl}}^{-}$ ions move to the anode (positive electrode) and lose electrons, and chlorine gas is given off. At the cathode: ${\text{Na}}^{+}+e^{-}\to \text{Na}$ At the anode: $2{\text{Cl}}^{-}\to {\text{Cl}}_2+2e^{-}$ Overall: $2\text{NaCl(l)}\to 2\text{Na(l)}+{\text{Cl}}_2\text{(g)}$

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15. Q335 CBSE 2018
    1. (1)A non-metal ‘P’ which is the largest constituent of air, when ​heated with hydrogen in $1:3$ ​ratio in the presence of catalyst ​(Fe) gives a gas Q.
    2. (2)On heating P with ​oxygen an oxide R is ​formed.
    3. (3)If ​this oxide (R) is passed ​into water in the presence of air, an acid S is formed which acts as a ​strong oxidizing agent. (a) Identify P, Q, R and S, and give chemical equations of the reactions involved to justify your answer. (b) To which group and period of modern periodic table does this non-metal P belong ?
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    P is nitrogen, Q ammonia, R nitric oxide and S nitric acid; nitrogen is in group 15, period 2.

    (a) P is nitrogen, ${\text{N}}_2$, the largest constituent of air. Q is ammonia, ${\text{NH}}_3$: ${\text{N}}_2+3{\text{H}}_2\stackrel{\text{Fe, heat}}{\to }2{\text{NH}}_3$ R is nitric oxide, NO: ${\text{N}}_2+{\text{O}}_2\stackrel{\text{heat}}{\to }2\text{NO}$ S is nitric acid, ${\text{HNO}}_3$, a strong oxidising agent: in air, $2\text{NO}+{\text{O}}_2\to 2{\text{NO}}_2$, and then $4{\text{NO}}_2+{\text{O}}_2+2{\text{H}}_2\text{O}\to 4{\text{HNO}}_3$ (Nitrogen dioxide, ${\text{NO}}_2$, is also accepted for R.) (b) Nitrogen (atomic number $7$, electronic configuration $2,5$) belongs to group $15$ and period $2$ of the modern periodic table.

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16. Q336 CBSE 2018
    1. (a)What are amphoteric ​oxides ? Give examples ​of two amphoteric oxides.
    2. (b)What ​is ​Cinnabar ? ​How is metal extracted from Cinnabar ​? Explain ​in brief.
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    Amphoteric oxides such as ${\text{Al}}_2{\text{O}}_3$ and ZnO react with acids and bases; cinnabar, HgS, is roasted to HgO and heated to give mercury.

    (a) Amphoteric oxides are metal oxides that react with both acids and bases to form a salt and water; they show both basic and acidic behaviour. Two examples: aluminium oxide (${\text{Al}}_2{\text{O}}_3$) and zinc oxide ($\text{ZnO}$). For example, ${\text{Al}}_2{\text{O}}_3+6\text{HCl}\to 2{\text{AlCl}}_3+3{\text{H}}_2\text{O}$ and ${\text{Al}}_2{\text{O}}_3+2\text{NaOH}\to 2{\text{NaAlO}}_2+{\text{H}}_2\text{O}$ (sodium aluminate). (b) Cinnabar is the ore of mercury: mercury(II) sulphide, $\text{HgS}$. Extraction: mercury is low in the reactivity series, so it is obtained by heating alone. Cinnabar is first heated strongly in air (roasting), which changes it into mercuric oxide: $2\text{HgS(s)}+3{\text{O}}_2\text{(g)}\stackrel{\text{heat}}{\to }2\text{HgO(s)}+2{\text{SO}}_2\text{(g)}$ On further heating, the mercuric oxide is reduced to mercury: $2\text{HgO(s)}\stackrel{\text{heat}}{\to }2\text{Hg(l)}+{\text{O}}_2\text{(g)}$

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17. Q337 CBSE 2018
    A student added few ​pieces of aluminium metal to two test tubes A and B containing aqueous solutions ​of iron ​sulphate and copper sulphate. In the second part of her experiment, she ​added iron metal to another test tubes C and D containing aqueous solutions of aluminium sulphate and copper sulphate. In ​which test ​tube or test tubes will she observe colour change ? On the basis of this experiment, state which one is the most reactive ​metal and ​why.
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    Colour changes in tubes A, B and D, not C; aluminium is the most reactive metal, as it displaces both iron and copper.

    Colour change is observed in test tubes A, B and D (no change in C). A: the pale green iron sulphate solution becomes colourless, as aluminium displaces iron. B: the blue copper sulphate solution becomes colourless, as aluminium displaces copper. D: the blue copper sulphate solution turns pale green, as iron displaces copper. Aluminium is the most reactive metal, because it displaces both iron and copper from their salt solutions, whereas iron cannot displace aluminium from aluminium sulphate (tube C). Order of reactivity: $\text{Al}>\text{Fe}>\text{Cu}$.

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18. Q338 CBSE 2017
    The atomic number ​of ​an element is 19. (a) Write the electronic configuration of this element and determine

    1. (i)the valency of this ​element, and
    2. (ii)whether this element is a ​metal or ​a non-metal. (b) Write the formula ​of the oxide of this element. (c) Is this element more reactive or less reactive than Na ​(atomic number 11) ? Justify your answer giving example.
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    The element is potassium ($2,8,8,1$): valency 1, a metal, oxide ${\text{K}}_2\text{O}$, and more reactive than sodium.

    (a) Electronic configuration: 2, 8, 8, 1 (the element is potassium, K). (i) Its outermost shell has one electron, which it loses, so its valency is 1. (ii) It is a metal, since it forms a positive ion by losing its outer electron. (b) The formula of its oxide is ${\text{K}}_2\text{O}$ (potassium, valency 1, with oxygen, valency 2). (c) It is more reactive than Na (2, 8, 1). Its valence electron is in the fourth shell, farther from the nucleus than sodium's, which is in the third, so it is held less strongly and lost more easily. Example: potassium reacts with cold water even more vigorously than sodium, $2\text{K}+2{\text{H}}_2\text{O}\to 2\text{KOH}+{\text{H}}_2$, and it stands above sodium in the reactivity series.

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19. Q339 CBSE 2017
    The atomic number of ​an element ​M ​is 19.

    1. (a)Write electronic ​configuration and valency of this element.
    2. (b)Is ‘M’ a metal or a non-metal ? ​Give reason in support of your ​answer.
    3. (c)Write the formula and nature (acidic/basic) ​of ​the oxide of M.
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    M has configuration $2,8,8,1$ and valency 1, is a metal, and forms the basic oxide ${\text{M}}_2\text{O}$.

    (a) Electronic configuration of M: $2,8,8,1$ (K, L, M and N shells). Valency: $1$. (b) M is a metal: it has only one electron in its outermost shell, which it loses readily to form the positive ion ${\text{M}}^{+}$, as metals do. (c) Formula of the oxide: ${\text{M}}_2\text{O}$ (M is potassium, so the oxide is ${\text{K}}_2\text{O}$). Nature: basic; it dissolves in water to form the alkali $\text{MOH}$.

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20. Q340 CBSE 2017
    The atomic number of an element is 20.

    1. (a)Write its ​electronic configuration and determine its valency.
    2. (b)Is ​it ​a ​metal or a non-metal ?
    3. (c)Write the formula of its chloride.
    4. (d)Is it ​more reactive or less reactive than Mg (atomic ​number ​12) ? Give reason ​for your answer.
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    The element is calcium ($2,8,8,2$): valency 2, a metal, chloride ${\text{CaCl}}_2$, and more reactive than magnesium.

    (a) Electronic configuration: K, L, M, N $=2,8,8,2$. It has $2$ valence electrons, which it loses, so its valency is $2$. (b) It is a metal (it is calcium, in group 2). (c) Its chloride is ${\text{CaCl}}_2$. (d) It is more reactive than magnesium. Magnesium ($2,8,2$) and this element are both in group 2, but this element has one more shell, so its valence electrons are farther from the nucleus and are lost more easily.

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