Case Study Questions Class 11 Chemistry Chapter 10 s-block Elements
CBSE Class 11 Case Study Questions Chemistry s-block elements. Important Case Study Questions for Class 11 Board Exam Students. Here we have arranged some Important Case Base Questions for students who are searching for Paragraph Based Questions s-block elements.
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CBSE Case Study Questions Class 11 Chemistry s-block Elements
Case Study- 1
The s-block elements of the Periodic Table are those in which the last electron enters the outermost s-orbital. as the s-orbital can accommodate only two electrons, two Groups (1 & 2) belong to the s-block of the Periodic Table. Group 1 of the Periodic Table consists of the elements: Lithium, sodium, potassium, rubidium, caesium and Francium. They are collectively known as the alkali metals. These are so called because they form hydroxides on Reaction with water which are strongly alkaline in nature. The elements of Group 2 include beryllium, magnesium, Calcium, strontium, barium and radium. These elements With the exception of beryllium are commonly known as The alkaline earth metals. These are so called because their Oxides and hydroxides are alkaline in nature and these Metal oxides are found in the earth’s crust. The general electronic configuration of s-block elements is [noble gas] ns1 for alkali metals and [noble gas] ns2 for Alkaline earth metals.
All the alkali metals have one valence electron, ns1 outside the noble gas core. The loosely held s-electron in the outermost Valence shell of these elements makes them the Most electropositive metals. They readily lose Electron to give monovalent M+ Ions. The monovalent ions (M+) are smaller than the parent atom. Hence they Are never found in free state in nature.
The alkali metal atoms have the largest sizes In a particular period of the periodic table. With increase in atomic number, the atom becomes Larger. the atomic and ionic Radii of alkali metals increase on moving down the group i.e., they increase in size while going From Li to Cs.
The ionization enthalpies of the alkali metals Are considerably low and decrease down the Group from Li to Cs. this is because the effect of increasing size outweighs the increasing Nuclear charge, and the outermost electron is very well screened from the nuclear charge.
The hydration enthalpies of alkali metal ions Decrease with increase in ionic sizes. Li+ > Na+ > K+ > Rb+ > Cs+ > Li+ Has maximum degree of hydration and For this reason lithium salts are mostly Hydrated, e.g., LiCl·2H2O
All the alkali metals are silvery white, soft and Light metals. Because of the large size, these Elements have low density which increases down The group from Li to Cs. However, potassium is Lighter than sodium. The melting and boiling Points of the alkali metals are low indicating Weak metallic bonding due to the presence of Only a single valence electron in them. The alkali Metals and their salts impart characteristic Colour to an oxidizing flame. This is because the Heat from the flame excites the outermost orbital Electron to a higher energy level. When the excited Electron comes back to the ground state . Alkali metals can therefore, be detected by The respective flame tests and can be Determined by flame photometry or atomic Absorption spectroscopy. These elements when Irradiated with light, the light energy absorbed May be sufficient to make an atom lose electron. This property makes caesium and potassium Useful as electrodes in photoelectric cells.
The alkali metals are highly reactive due to Their large size and low ionization enthalpy. The Reactivity of these metals increases down the Group.
Reactivity towards air: The alkali metals Tarnish in dry air due to the formation of Their oxides which in turn react with Moisture to form hydroxides. They burn Vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms Peroxide, the other metals form Superoxide. The superoxide O2– Ion is Stable only in the presence of large cations Such as K, Rb, Cs.
Reactivity towards water: The alkali Metals react with water to form hydroxide And dihydrogen.
Reactivity towards dihydrogen: The Alkali metals react with dihydrogen at About 673K (lithium at 1073K) to form Hydrides. All the alkali metal hydrides are Ionic solids with high melting points.
Reactivity towards halogens: The alkali Metals readily react vigorously with Halogens to form ionic halides, M+X– .
Reducing nature: The alkali metals are Strong reducing agents, lithium being the Most and sodium the least powerful.
Solutions in liquid ammonia: The alkali Metals dissolve in liquid ammonia giving Deep blue solutions which are conducting in nature.
[A] MCQ
1) The general electronic configuration of s-block elements is … for alkali metals.
a) [noble gas] ns1
b) [noble gas] ns2
c) [noble gas] ns1np1
d) [noble gas] ns1np2
Ans- a) [noble gas] ns1
2) The general electronic configuration of s-block elements is … for alkaline earth metals.
a) noble gas] ns1
b) [noble gas] ns2
c) [noble gas] ns1np1
d) [noble gas] ns1np2
Ans- b) [noble gas] ns2
3) The atomic and ionic Radii of alkali metals … on moving down the group.
a) constant
b) decrease
c) increase
d) All the above
Ans- c) increase
4) The hydration enthalpies of alkali metal ions … with … in ionic sizes.
a) increase , decrease
b) increase , increase
c) decrease , decrease
d) decrease, increase
Ans – d) decrease, increase
5) Which of the following element is strong reducing agent ?
a) Lithium
b) Sodium
c) Fluorine
d) Helium
Ans- a) lithium
[B]Short Answers
1) Give electronic configuration of alkali metals.
Ans- All the alkali metals have one valence electron, ns1 outside the noble gas core. The loosely held s-electron in the outermost Valence shell of these elements makes them the Most electropositive metals. They readily lose Electron to give monovalent M+ ions. The monovalent ions (M+) are smaller than the parent atom.
Hence they are never found in free state in nature.
2) What are s-block elements ?
Ans- The s-block elements of the Periodic Table are those in which the last electron enters the outermost s-orbital. As the s-orbital can accommodate only two electrons, two Groups (1 & 2) belong to the s-block of the Periodic Table. Group 1 of the Periodic Table consists of the elements: Lithium, sodium, potassium, rubidium, caesium and Francium. They are collectively known as the alkali metals. These are so called because they form hydroxides on Reaction with water which are strongly alkaline in nature. The elements of Group 2 include beryllium, magnesium, Calcium, strontium, barium and radium. These elements With the exception of beryllium are commonly known as the alkaline earth metals.
3) Explain the trend of ionisation enthalpy of alkali metals.
Ans- The ionization enthalpies of the alkali metals Are considerably low and decrease down the Group from Li to Cs. This is because the effect of increasing size outweighs the increasing Nuclear charge, and the outermost electron is very well screened from the nuclear charge.
[C]Long Answers
1) Write physical properties of alkali metals .
Ans- All the alkali metals are silvery white, soft and Light metals. Because of the large size, these Elements have low density which increases down The group from Li to Cs. However, potassium is Lighter than sodium. The melting and boiling Points of the alkali metals are low indicating Weak metallic bonding due to the presence of Only a single valence electron in them. The alkali Metals and their salts impart characteristic colour to an oxidizing flame. This is because the Heat from the flame excites the outermost orbital Electron to a higher energy level. When the excited Electron comes back to the ground state . Alkali metals can therefore, be detected by The respective flame tests and can be Determined by flame photometry or atomic Absorption spectroscopy. These elements when Irradiated with light, the light energy absorbed May be sufficient to make an atom lose electron. This property makes caesium and potassium Useful as electrodes in photoelectric cells.
2) Explain chemical properties of alkali metals.
Ans- The alkali metals are highly reactive due to Their large size and low ionization enthalpy. The Reactivity of these metals increases down the Group.
Reactivity towards air: The alkali metals Tarnish in dry air due to the formation of Their oxides which in turn react with Moisture to form hydroxides. They burn Vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms Peroxide, the other metals form Superoxide. The superoxide O2– Ion is Stable only in the presence of large cations Such as K, Rb, Cs.
Reactivity towards water: The alkali Metals react with water to form hydroxide And dihydrogen.
Reactivity towards dihydrogen: The Alkali metals react with dihydrogen at About 673K (lithium at 1073K) to form Hydrides. All the alkali metal hydrides are Ionic solids with high melting points.
Reactivity towards halogens: The alkali Metals readily react vigorously with Halogens to form ionic halides, M+X– .
Reducing nature: The alkali metals are Strong reducing agents, lithium being the Most and sodium the least powerful.
Solutions in liquid ammonia: The alkali Metals dissolve in liquid ammonia giving Deep blue solutions which are conducting in nature.
Case Study- 2
Lithium metal is used to make useful alloys, for example with lead to make ‘white metal’ Bearings for motor engines, with aluminium to make aircraft parts, and with magnesium to make armour plates. It is used in Thermonuclear reactions. Lithium is also used to make electrochemical cells. Sodium is used To make a Na/Pb alloy needed to make PbEt4 and PbMe4. These organolead compounds were Earlier used as anti-knock additives to petrol, But nowadays vehicles use lead-free petrol. Liquid sodium metal is used as a coolant in Fast breeder nuclear reactors. Potassium has a vital role in biological systems. Potassium Chloride is used as a fertilizer. Potassium Hydroxide is used in the manufacture of soft Soap. It is also used as an excellent absorbent of carbon dioxide. Caesium is used in devising Photoelectric cells.
Points of Difference between Lithium and other Alkali Metals –
i) Lithium is much harder. Its m.p. and b.p. are higher than the other alkali metals.
ii) Lithium is least reactive but the strongest Reducing agent among all the alkali metals. On combustion in air it forms mainly Monoxide, Li2O and the nitride, Li3N unlike Other alkali metals.
iii) LiCl is deliquescent and crystallises as a Hydrate, LiCl.2H2O whereas other alkali Metal chlorides do not form hydrates.
iv) Lithium hydrogencarbonate is not Obtained in the solid form while all other Elements form solid hydrogencarbonate.
v) Lithium unlike other alkali metals forms No ethynide on reaction with ethyne.
vi) Lithium nitrate when heated gives lithium Oxide, Li2O, whereas other alkali metal Nitrates decompose to give the Corresponding nitrite.
vii) LiF and Li2O are comparatively much less Soluble in water than the corresponding Compounds of other alkali metals.
Sodium carbonate is generally prepared by Solvay Process. In this process, advantage is Taken of the low solubility of sodium Hydrogencarbonate whereby it gets Precipitated in the reaction of sodium chloride with ammonium hydrogencarbonate. The Latter is prepared by passing CO2 to a Concentrated solution of sodium chloride Saturated with ammonia, where ammonium Carbonate followed by ammonium Hydrogencarbonate are formed. The equations For the complete process may be written as
Sodium hydrogencarbonate crystal Separates. these are heated to give sodium Carbonate.
The most abundant source of sodium chloride is sea water which contains 2.7 to 2.9% by Mass of the salt. In tropical countries like India, Common salt is generally obtained by Evaporation of sea water. Approximately 50 Lakh tons of salt are produced annually in India by solar evaporation. Crude sodium Chloride, generally obtained by crystallisation Of brine solution, contains sodium sulphate, Calcium sulphate, calcium chloride and Magnesium chloride as impurities. Calcium Chloride, CaCl2, and magnesium chloride, MgCl2 are impurities because they are Deliquescent (absorb moisture easily from the Atmosphere). To obtain pure sodium chloride, The crude salt is dissolved in minimum amount Of water and filtered to remove insoluble Impurities. The solution is then saturated with Hydrogen chloride gas. Crystals of pure Sodium chloride separate out. Calcium and Magnesium chloride, being more soluble than Sodium chloride, remain in solution.
Sodium Hydroxide (Caustic Soda), NaOH is generally prepared Commercially by the electrolysis of sodium Chloride in Castner-Kellner cell. A brine Solution is electrolysed using a mercury Cathode and a carbon anode. Sodium metal Discharged at the cathode combines with Mercury to form sodium amalgam. Chlorine Gas is evolved at the anode.
A] MCQ
1) NaOH Sodium hydroxide is generally prepared Commercially by the electrolysis of … in Castner-Kellner cell.
a) NaCl
b) Na2CO3
c) NaHCO3
d) NaNH2
Ans- a) NaCl
2) … is used in the manufacture of soft Soap.
a) Sodium Hydroxide
b) Potassium Hydroxide
c) Aluminium hydroxide
d) Beryllium hydroxide
Ans- b) Potassium Hydroxide
3) … is used in devising Photoelectric cells.
a) Hydrogen
b) Lithium
c) Caesium
d) Helium
Ans- c) Caesium
4) … compounds were Earlier used as anti-knock additives to petrol.
a) Organomagnesium
b) Organosilicon
c) Organochloride
d) Organolead
Ans- d) Organolead
5) The sodium amalgam is treated with water to gives ….
a) NaOH
b) Na2CO3
c) NaHCO3
d) NaNH2
Ans- a) NaOH
[B]Short Answers
1) How sodium hydroxide is prepared ?
Ans- Sodium Hydroxide (Caustic Soda), NaOH is generally prepared Commercially by the electrolysis of sodium Chloride in Castner-Kellner cell. A brine Solution is electrolysed using a mercury Cathode and a carbon anode. Sodium metal Discharged at the cathode combines with Mercury to form sodium amalgam. Chlorine Gas is evolved at the anode.
2) Write short note on sodium chloride ?
Ans- The most abundant source of sodium chloride is sea water which contains 2.7 to 2.9% by Mass of the salt. In tropical countries like India, Common salt is generally obtained by Evaporation of sea water. Approximately 50 Lakh tons of salt are produced annually in India by solar evaporation. Crude sodium Chloride, generally obtained by crystallisation Of brine solution, contains sodium sulphate, Calcium sulphate, calcium chloride and Magnesium chloride as impurities. Calcium Chloride, CaCl2, and magnesium chloride, MgCl2 are impurities because they are Deliquescent (absorb moisture easily from the Atmosphere). To obtain pure sodium chloride, The crude salt is dissolved in minimum amount Of water and filtered to remove insoluble Impurities. The solution is then saturated with Hydrogen chloride gas. Crystals of pure Sodium chloride separate out. Calcium and Magnesium chloride, being more soluble than Sodium chloride, remain in solution.
3) Give preparation of sodium carbonate ?
Ans- Sodium carbonate is generally prepared by Solvay Process. In this process, advantage is Taken of the low solubility of sodium Hydrogencarbonate whereby it gets Precipitated in the reaction of sodium chloride with ammonium hydrogencarbonate. The Latter is prepared by passing CO2 to a Concentrated solution of sodium chloride Saturated with ammonia, where ammonium Carbonate followed by ammonium Hydrogencarbonate are formed. The equations For the complete process may be written as
[C]Long Answers
1) Write the uses of alkali metals .
Ans- Lithium metal is used to make useful alloys, for example with lead to make ‘white metal’ Bearings for motor engines, with aluminium to make aircraft parts, and with magnesium to make armour plates. It is used in Thermonuclear reactions. Lithium is also used to make electrochemical cells. Sodium is used To make a Na/Pb alloy needed to make PbEt4 and PbMe4. These organolead compounds were Earlier used as anti-knock additives to petrol, But nowadays vehicles use lead-free petrol. Liquid sodium metal is used as a coolant in Fast breeder nuclear reactors. Potassium has a vital role in biological systems. Potassium Chloride is used as a fertilizer. Potassium Hydroxide is used in the manufacture of soft Soap. It is also used as an excellent absorbent of carbon dioxide. Caesium is used in devising Photoelectric cells.
2) What are the differences between Lithium and other Alkali Metals.
Ans- Points of Difference between Lithium and other Alkali Metals –
i) Lithium is much harder. Its m.p. and b.p. are higher than the other alkali metals.
ii) Lithium is least reactive but the strongest Reducing agent among all the alkali metals. On combustion in air it forms mainly Monoxide, Li2O and the nitride, Li3N unlike Other alkali metals.
iii) LiCl is deliquescent and crystallises as a Hydrate, LiCl.2H2O whereas other alkali Metal chlorides do not form hydrates.
iv) Lithium hydrogencarbonate is not Obtained in the solid form while all other Elements form solid hydrogencarbonate.
v) Lithium unlike other alkali metals forms No ethynide on reaction with ethyne.
vi) Lithium nitrate when heated gives lithium Oxide, Li2O, whereas other alkali metal Nitrates decompose to give the Corresponding nitrite.
vii) LiF and Li2O are comparatively much less Soluble in water than the corresponding Compounds of other alkali metals.
Case Study- 3
Alkaline earth elements have two electrons in the s -orbital of the valence shell . Their general electronic configuration may be represented as [noble gas] ns 2 . Like alkali metals, the compounds of these elements are also predominantly ionic.
The atomic and ionic radii of the alkaline earth metals are smaller than those of the corresponding alkali metals in the same periods. This is due to the increased nuclear charge in these elements. Within the group, the atomic and ionic radii increase with increase in atomic number.
The alkaline earth metals have low ionization enthalpies due to fairly large size of the atoms. Since the atomic size increases down the group, their ionization enthalpy decreases. The first ionisation enthalpies of the alkaline earth metals are higher than those of the corresponding Group 1 metals. This is due to their small size as compared to the corresponding alkali metals. It is interesting to note that the second ionisation enthalpies of the alkaline earth metals are smaller than those of the corresponding alkali metals.
Like alkali metal ions, the hydration enthalpies of alkaline earth metal ions decrease with increase in ionic size down the group. Be2+> Mg2+ > Ca2+ > Sr2+ > Ba2+ The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions. Thus, compounds of alkaline earth metals are more extensively hydrated than those of alkali metals, e.g., MgCl2 and CaCl2 exist as MgCl2.6H2O and CaCl2· 6H2O while NaCl and KCl do not form such hydrates.
The alkaline earth metals, in general, are silvery white, lustrous and relatively soft but harder than the alkali metals. Beryllium and magnesium appear to be somewhat greyish. The melting and boiling points of these metals are higher than the corresponding alkali metals due to smaller sizes. The trend is, however, not systematic. Because of the low ionisation enthalpies, they are strongly electropositive in nature. The electropositive character increases down the group from Be to Ba. Calcium, strontium and barium impart characteristic brick red, crimson and apple green colours respectively to the flame. In flame the electrons are excited to higher energy levels and when they drop back to the ground state, energy is emitted in the form of visible light. The electrons in beryllium and magnesium are too strongly bound to get excited by flame. Hence, these elements do not impart any colour to the flame. The flame test for Ca, Sr and Ba is helpful in their detection in qualitative analysis and estimation by flame photometry. The alkaline earth metals like those of alkali metals have high electrical and thermal conductivities which are typical characteristics of metals.
Chemical Properties- The alkaline earth metals are less reactive than the alkali metals. The reactivity of these elements increases on going down the group.
i) Reactivity towards air and water: Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface. Magnesium is more electropositive and burns with dazzling brilliance in air to give MgO and Mg3N2. Calcium, strontium and barium are readily attacked by air to form the oxide and nitride.
ii) Reactivity towards the halogens: All the alkaline earth metals combine with halogen at elevated temperatures forming their halides.
iii) Reactivity towards hydrogen: All the elements except beryllium combine with hydrogen upon heating to form their hydrides, MH2. BeH2, however, can be prepared by the reaction of BeCl2 with LiAlH4
iv) Reactivity towards acids: The alkaline earth metals readily react with acids liberating dihydrogen. M + 2HCl → MCl2 + H2
v) Reducing nature: Like alkali metals, the alkaline earth metals are strong reducing agents. This is indicated by large negative values of their reduction potentials. However their reducing power is less than those of their corresponding alkali metals. Beryllium has less negative value compared to other alkaline earth metals
vi) Solutions in liquid ammonia: Like alkali metals, the alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions forming ammoniated ions.
From these solutions, the ammoniates, [M(NH3)6]2+can be recovered.
Beryllium is used in the manufacture of alloys. Copper -beryllium alloys are used in the preparation of high strength springs. Metallic beryllium is used for making windows of X-ray tubes. Magnesium forms alloys with aluminium, zinc, manganese and tin. Magnesium-aluminium alloys being light in mass are used in air-craft construction. Magnesium (powder and ribbon) is used in flash powders and bulbs, incendiary bombs and signals. A suspension of magnesium hydroxide in water (called milk of magnesia) is used as antacid in medicine. Magnesium carbonate is an ingredient of toothpaste. Calcium is used in the extraction of metals from oxides which are difficult to reduce with carbon. Calcium and barium metals, owing to their reactivity with oxygen and nitrogen at elevated temperatures, have often been used to remove air from vacuum tubes. Radium salts are used in radiotherapy, for example, in the treatment of cancer.
[A] MCQ
1) The atomic and ionic radii of the alkaline earth metals are … than those of the corresponding alkali metals in the same periods.
a) smaller
b) bigger
c) different
d) None of above
Ans- a) smaller
2) Within the group, the atomic and ionic radii of alkaline earth metals … with … in atomic number.
a) increase , decrease
b) increase , increase
c) decrease , increase
d) decrease , decrease
Ans- b) increase , increase
3) Alkaline earth elements have … electrons in the s -orbital of the valence shell.
a) Zero
b) One
c) Two
d) Three
Ans- c) Two
4) Ionization enthalpy …. down the group of alkaline earth metals.
a) first increases then decreases
b) first decreases then increases
c) increase
d) decreases
Ans- d) decreases
5) The hydration enthalpies of alkaline earth metal ions … with … in ionic size down the group.
a) increase , decrease
b) increase , increase
c) decrease , increase
d) decrease , decrease
Ans- c) decrease , increase
[B]Short Answers
1) Give the electronic configuration of alkaline earth metals .
Ans- Alkaline earth elements have two electrons in the s -orbital of the valence shell . Their general electronic configuration may be represented as [noble gas] ns 2 . Like alkali metals, the compounds of these elements are also predominantly ionic.
2) Write short note on Ionization Enthalpies of alkaline earth metals.
Ans- The alkaline earth metals have low ionization enthalpies due to fairly large size of the atoms. Since the atomic size increases down the group, their ionization enthalpy decreases. The first ionisation enthalpies of the alkaline earth metals are higher than those of the corresponding Group 1 metals. This is due to their small size as compared to the corresponding alkali metals. It is interesting to note that the second ionisation enthalpies of the alkaline earth metals are smaller than those of the corresponding alkali metals.
3) Explain the trends in hydration enthalpy of alkaline earth metals.
Ans- Like alkali metal ions, the hydration enthalpies of alkaline earth metal ions decrease with increase in ionic size down the group. Be2+> Mg2+ > Ca2+ > Sr2+ > Ba2+ The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions. Thus, compounds of alkaline earth metals are more extensively hydrated than those of alkali metals, e.g., MgCl2 and CaCl2 exist as MgCl2.6H2O and CaCl2· 6H2O while NaCl and KCl do not form such hydrates.
[C]Long Answers
1) Write physical properties of alkaline earth metals.
Ans- The alkaline earth metals, in general, are silvery white, lustrous and relatively soft but harder than the alkali metals. Beryllium and magnesium appear to be somewhat greyish. The melting and boiling points of these metals are higher than the corresponding alkali metals due to smaller sizes. The trend is, however, not systematic. Because of the low ionisation enthalpies, they are strongly electropositive in nature. The electropositive character increases down the group from Be to Ba. Calcium, strontium and barium impart characteristic brick red, crimson and apple green colours respectively to the flame. In flame the electrons are excited to higher energy levels and when they drop back to the ground state, energy is emitted in the form of visible light. The electrons in beryllium and magnesium are too strongly bound to get excited by flame. Hence, these elements do not impart any colour to the flame. The flame test for Ca, Sr and Ba is helpful in their detection in qualitative analysis and estimation by flame photometry. The alkaline earth metals like those of alkali metals have high electrical and thermal conductivities which are typical characteristics of metals.
2) Explain the chemical properties of alkaline earth metals.
Ans- The alkaline earth metals are less reactive than the alkali metals. The reactivity of these elements increases on going down the group.
i) Reactivity towards air and water: Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface. Magnesium is more electropositive and burns with dazzling brilliance in air to give MgO and Mg3N2. Calcium, strontium and barium are readily attacked by air to form the oxide and nitride.
ii) Reactivity towards the halogens: All the alkaline earth metals combine with halogen at elevated temperatures forming their halides.
iii) Reactivity towards hydrogen: All the elements except beryllium combine with hydrogen upon heating to form their hydrides, MH2. BeH2, however, can be prepared by the reaction of BeCl2 with LiAlH4.
iv) Reactivity towards acids: The alkaline earth metals readily react with acids liberating dihydrogen. M + 2HCl → MCl2 + H2
v) Reducing nature: Like alkali metals, the alkaline earth metals are strong reducing agents. This is indicated by large negative values of their reduction potentials. However their reducing power is less than those of their corresponding alkali metals. Beryllium has less negative value compared to other alkaline earth metals
vi) Solutions in liquid ammonia: Like alkali metals, the alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions forming ammoniated ions.
From these solutions, the ammoniates, [M(NH3)6]2+ can be recovered.
3) Give the uses of alkaline earth metals.
Ans- Beryllium is used in the manufacture of alloys. Copper -beryllium alloys are used in the preparation of high strength springs. Metallic beryllium is used for making windows of X-ray tubes. Magnesium forms alloys with aluminium, zinc, manganese and tin. Magnesium-aluminium alloys being light in mass are used in air-craft construction. Magnesium (powder and ribbon) is used in flash powders and bulbs, incendiary bombs and signals. A suspension of magnesium hydroxide in water (called milk of magnesia) is used as antacid in medicine. Magnesium carbonate is an ingredient of toothpaste. Calcium is used in the extraction of metals from oxides which are difficult to reduce with carbon. Calcium and barium metals, owing to their reactivity with oxygen and nitrogen at elevated temperatures, have often been used to remove air from vacuum tubes. Radium salts are used in radiotherapy, for example, in the treatment of cancer.
Case Study – 4
The dipositive oxidation state (M2+) is the Predominant valence of Group 2 elements. The Alkaline earth metals form compounds which Are predominantly ionic but less ionic than the Corresponding compounds of alkali metals. This is due to increased nuclear charge and Smaller size. The oxides and other compounds Of beryllium and magnesium are more covalent Than those formed by the heavier and large Sized members (Ca, Sr, Ba). The general Characteristics of some of the compounds of Alkali earth metals are described below.
Oxides and Hydroxides: The alkaline Earth metals burn in oxygen to form the Monoxide, MO which, except for BeO, have Rock-salt structure. The BeO is essentially Covalent in nature. The enthalpies of formation Of these oxides are quite high and consequently They are very stable to heat. BeO is amphoteric While oxides of other elements are ionic in Nature. All these oxides except BeO are basic In nature and react with water to form sparingly Soluble hydroxides.
MO + H2O → M(OH) 2
The solubility, thermal stability and the Basic character of these hydroxides increase With increasing atomic number from Mg(OH) 2 To Ba(OH) 2. The alkaline earth metal Hydroxides are, however, less basic and less Stable than alkali metal hydroxides. Beryllium Hydroxide is amphoteric in nature as it reacts With acid and alkali both.
Be(OH) 2 + 2OH– → [Be(OH)4]2–
Beryllate ion
Be(OH) 2 + 2HCl + 2H2O → [Be(OH)4]Cl2
Halides: Except for beryllium halides, all Other halides of alkaline earth metals are ionic In nature. Beryllium halides are essentially Covalent and soluble in organic solvents. Beryllium chloride has a chain structure in the Solid state as shown below: In the vapour phase BeCl2 tends to form a Chloro-bridged dimer which dissociates into the Linear monomer at high temperatures of the Order of 1200 K. The tendency to form halide Hydrates gradually decreases down the group. The dehydration Of hydrated chlorides, bromides and iodides Of Ca, Sr and Ba can be achieved on heating; However, the corresponding hydrated halides Of Be and Mg on heating suffer hydrolysis. The Fluorides are relatively less soluble than the Chlorides owing to their high lattice energies.
Salts of Oxoacids: The alkaline earth Metals also form salts of oxoacids. Some of These are :
Carbonates: Carbonates of alkaline earth Metals are insoluble in water and can be Precipitated by addition of a sodium or Ammonium carbonate solution to a solution Of a soluble salt of these metals. The solubility Of carbonates in water decreases as the atomic Number of the metal ion increases. All the Carbonates decompose on heating to give Carbon dioxide and the oxide. Beryllium Carbonate is unstable and can be kept only in The atmosphere of CO2. The thermal stability Increases with increasing cationic size.
Sulphates: The sulphates of the alkaline earth Metals are all white solids and stable to heat. BeSO4, and MgSO4 are readily soluble in water; The solubility decreases from CaSO4 to BaSO4. The greater hydration enthalpies of Be2+ and Mg2+ ions overcome the lattice enthalpy factor And therefore their sulphates are soluble in Water.
Nitrates: The nitrates are made by dissolution Of the carbonates in dilute nitric acid. Magnesium nitrate crystallises with six Molecules of water, whereas barium nitrate Crystallises as the anhydrous salt. This again Shows a decreasing tendency to form hydrates With increasing size and decreasing hydration Enthalpy. All of them decompose on heating to Give the oxide like lithium nitrate.
Beryllium, the first member of the Group 2 Metals, shows anomalous behaviour as Compared to magnesium and rest of the Members. Further, it shows diagonal Relationship to aluminium which is discussed Subsequently.
i) Beryllium has exceptionally small atomic And ionic sizes and thus does not compare Well with other members of the group. Because of high ionisation enthalpy and Small size it forms compounds which are Largely covalent and get easily hydrolysed.
ii) Beryllium does not exhibit coordination Number more than four as in its valence Shell there are only four orbitals. The Remaining members of the group can have A coordination number of six by making Use of d-orbitals.
iii) The oxide and hydroxide of beryllium, Unlike the hydroxides of other elements in The group, are amphoteric in nature.
Diagonal Relationship between Beryllium and Aluminium-The ionic radius of Be2+ is estimated to be 31 pm; the charge/radius ratio is nearly the Same as that of the Al3+ ion. Hence beryllium Resembles aluminium in some ways. Some of The similarities are:
i) Like aluminium, beryllium is not readily Attacked by acids because of the presence Of an oxide film on the surface of the metal.
ii) Beryllium hydroxide dissolves in excess of Alkali to give a beryllate ion, [Be(OH)4]2– just As aluminium hydroxide gives aluminate Ion, [Al(OH)4]– .
iii) The chlorides of both beryllium and Aluminium have Cl– Bridged chloride Structure in vapour phase. Both the Chlorides are soluble in organic solvents And are strong Lewis acids. They are used As Friedel Craft catalysts.
iv) Beryllium and aluminium ions have strong Tendency to form complexes, BeF42–, AlF63–.
[A] MCQ
1) The dipositive oxidation state (M2+) is the Predominant valence of … elements.
a) Group 2
b) Group 1
c) Group 17
d) Group 18
Ans- a) Group 2
2) … the first member of the Group 2
metals.
a) Magnesium
b) Beryllium
c) Barium
d) Radium
Ans- b) Beryllium
3) Except for beryllium halides, all
other halides of alkaline earth metals are ionic
in nature.
a) Magnesium halides
b) beryllium halides
c) Calcium halides
d) Radium halides
Ans- c) beryllium halides
4) The ionic radius of Be2+ is estimated to be
… pm.
a) 310
b) 500
c) 50
d) 31
Ans- d) 31
5) Beryllium
carbonate can be kept only in
the atmosphere of …
a) N2
b) H2
c) CO2
d) O2
Ans- c) CO2
[B]Short Answers
1) Explain the anomalous behaviour of beryllium .
Ans- Beryllium, the first member of the Group 2 Metals, shows anomalous behaviour as Compared to magnesium and rest of the Members. Further, it shows diagonal Relationship to aluminium which is discussed Subsequently.
i) Beryllium has exceptionally small atomic And ionic sizes and thus does not compare Well with other members of the group. Because of high ionisation enthalpy and Small size it forms compounds which are Largely covalent and get easily hydrolysed.
ii) Beryllium does not exhibit coordination Number more than four as in its valence Shell there are only four orbitals. The Remaining members of the group can have A coordination number of six by making Use of d-orbitals.
iii) The oxide and hydroxide of beryllium, Unlike the hydroxides of other elements in The group, are amphoteric in nature.
2) Describe the Diagonal Relationship between Beryllium and Aluminium.
Ans- Diagonal Relationship between Beryllium and Aluminium-The ionic radius of Be2+ is estimated to be 31 pm; the charge/radius ratio is nearly the Same as that of the Al3+ ion. Hence beryllium Resembles aluminium in some ways. Some of The similarities are:
i) Like aluminium, beryllium is not readily Attacked by acids because of the presence Of an oxide film on the surface of the metal.
ii) Beryllium hydroxide dissolves in excess of Alkali to give a beryllate ion, [Be(OH)4]2– just As aluminium hydroxide gives aluminate Ion, [Al(OH)4]– .
iii) The chlorides of both beryllium and Aluminium have Cl– Bridged chloride Structure in vapour phase. Both the Chlorides are soluble in organic solvents And are strong Lewis acids. They are used As Friedel Craft catalysts.
iv) Beryllium and aluminium ions have strong Tendency to form complexes, BeF42–, AlF63–.
C]Long Answers
1) Give the explanation on Oxo acid salts of alkaline earth metal elements.
Ans- The alkaline earth Metals also form salts of oxoacids. Some of These are :
Carbonates: Carbonates of alkaline earth Metals are insoluble in water and can be Precipitated by addition of a sodium or Ammonium carbonate solution to a solution Of a soluble salt of these metals. The solubility Of carbonates in water decreases as the atomic Number of the metal ion increases. All the Carbonates decompose on heating to give Carbon dioxide and the oxide. Beryllium Carbonate is unstable and can be kept only in The atmosphere of CO2. The thermal stability Increases with increasing cationic size.
Sulphates: The sulphates of the alkaline earth Metals are all white solids and stable to heat. BeSO4, and MgSO4 are readily soluble in water; The solubility decreases from CaSO4 to BaSO4. The greater hydration enthalpies of Be2+ and Mg2+ ions overcome the lattice enthalpy factor And therefore their sulphates are soluble in Water.
Nitrates: The nitrates are made by dissolution Of the carbonates in dilute nitric acid. Magnesium nitrate crystallises with six Molecules of water, whereas barium nitrate Crystallises as the anhydrous salt. This again Shows a decreasing tendency to form hydrates With increasing size and decreasing hydration Enthalpy. All of them decompose on heating to Give the oxide like lithium nitrate.
2) Give the general characteristics of the compounds of alkaline earth metals.
Ans- The general Characteristics of some of the compounds of Alkali earth metals are described below.
Oxides and Hydroxides: The alkaline Earth metals burn in oxygen to form the Monoxide, MO which, except for BeO, have Rock-salt structure. The BeO is essentially Covalent in nature. The enthalpies of formation Of these oxides are quite high and consequently They are very stable to heat. BeO is amphoteric While oxides of other elements are ionic in Nature. All these oxides except BeO are basic In nature and react with water to form sparingly Soluble hydroxides.
MO + H2O → M(OH) 2
The solubility, thermal stability and the Basic character of these hydroxides increase With increasing atomic number from Mg(OH) 2 To Ba(OH) 2. The alkaline earth metal Hydroxides are, however, less basic and less Stable than alkali metal hydroxides. Beryllium Hydroxide is amphoteric in nature as it reacts With acid and alkali both.
Be(OH) 2 + 2OH– → [Be(OH)4]2–
Beryllate ion
Be(OH) 2 + 2HCl + 2H2O → [Be(OH)4]Cl2
Halides: Except for beryllium halides, all Other halides of alkaline earth metals are ionic In nature. Beryllium halides are essentially Covalent and soluble in organic solvents. Beryllium chloride has a chain structure in the Solid state .the vapour phase BeCl2 tends to form a Chloro-bridged dimer which dissociates into the Linear monomer at high temperatures of the Order of 1200 K. The tendency to form halide Hydrates gradually decreases down the group. The dehydration Of hydrated chlorides, bromides and iodides Of Ca, Sr and Ba can be achieved on heating; However, the corresponding hydrated halides Of Be and Mg on heating suffer hydrolysis. The Fluorides are relatively less soluble than the Chlorides owing to their high lattice energies.
Case Study – 5
Important compounds of calcium are calcium Oxide, calcium hydroxide, calcium sulphate, Calcium carbonate and cement. These are Industrially important compounds. The large Scale preparation of these compounds and Their uses are described below. Calcium Oxide or Quick Lime CaO – It is prepared on a commercial scale by Heating limestone (CaCO3) in a rotary kiln at 1070-1270 K.
The carbon dioxide is removed as soon as It is produced to enable the reaction to proceed To completion. Calcium oxide is a white amorphous solid. It has a melting point of 2870 K. On exposure To atmosphere, it absorbs moisture and carbon Dioxide.
The addition of limited amount of water Breaks the lump of lime. This process is called Slaking of lime. Quick lime slaked with soda Gives solid sodalime. Being a basic oxide, it Combines with acidic oxides at high Temperature.
Uses: It is an important primary material for Manufacturing cement and is the cheapest Form of alkali. It is used in the manufacture of sodium Carbonate from caustic soda. It is employed in the purification of sugar And in the manufacture of dye stuffs.
Calcium Carbonate – CaCO3 occurs in nature in several Forms like limestone, chalk, marble etc. It can Be prepared by passing carbon dioxide Through slaked lime or by the addition of Sodium carbonate to calcium chloride.
excess of carbon dioxide should be Avoided since this leads to the formation of Water soluble calcium hydrogencarbonate. Calcium carbonate is a white fluffy powder. It is almost insoluble in water. When heated To 1200 K, it decomposes to evolve carbon Dioxide. It is used as a building material in the form of Marble and in the manufacture of quick lime. Calcium carbonate along with magnesium Carbonate is used as a flux in the extraction of Metals such as iron. Specially precipitated CaCO3 Is extensively used in the manufacture Of high quality paper. It is also used as an Antacid, mild abrasive in tooth paste, a Constituent of chewing gum, and a filler in Cosmetics.
Calcium Sulphate (Plaster of Paris),CaSO4·½ H2O – It is a hemihydrate of calcium sulphate. It is Obtained when gypsum, CaSO4·2H2O, is Heated to 393 K.
above 393 K, no water of crystallisation is left and anhydrous calcium sulphate, CaSO4 is Formed. This is known as ‘dead burnt plaster’. It has a remarkable property of setting with Water. On mixing with an adequate quantity Of water it forms a plastic mass that gets into a Hard solid in 5 to 15 minutes.
Uses: The largest use of Plaster of Paris is in the Building industry as well as plasters. It is used For immoblising the affected part of organ where There is a bone fracture or sprain. It is also Employed in dentistry, in ornamental work and For making casts of statues and busts.
Cement: Cement is an important building Material. it was first introduced in England in 1824 by Joseph Aspdin. It is also called Portland cement because it resembles with the Natural limestone quarried in the Isle of Portland, England. Cement is a product obtained by Combining a material rich in lime, CaO with Other material such as clay which contains Silica, SiO2 along with the oxides of Aluminium, iron and magnesium. The average Composition of Portland cement is : CaO, 50-60%; SiO2, 20-25%; Al2O3, 5-10%; MgO, 2-3%; Fe2O3, 1-2% and SO3, 1-2%. For a good Quality cement, the ratio of silica (SiO2) to Alumina (Al2O3) should be between 2.5 and 4 And the ratio of lime (CaO) to the total of the Oxides of silicon (SiO2) aluminium (Al2O3) And iron (Fe2O3) should be as close as possible To 2. The raw materials for the manufacture of Cement are limestone and clay. When clay and Lime are strongly heated together they fuse and React to form ‘cement clinker’. This clinker is Mixed with 2-3% by weight of gypsum (CaSO4·2H2O) to form cement. Thus important Ingredients present in Portland cement are Dicalcium silicate (Ca2SiO4) 26%, tricalcium silicate (Ca3SiO5) 51% and tricalcium Aluminate (Ca3Al2O6) 11%.
Setting of Cement: When mixed with water, The setting of cement takes place to give a hard Mass. This is due to the hydration of the Molecules of the constituents and their Rearrangement. The purpose of adding Gypsum is only to slow down the process of Setting of the cement so that it gets sufficiently Hardened.
Uses: Cement has become a commodity of National necessity for any country next to iron And steel. It is used in concrete and reinforced Concrete, in plastering and in the construction Of bridges, dams and buildings.
Biological importance of magnesium and calcium- An adult body contains about 25 g of Mg and 1200 g of Ca compared with only 5 g of iron And 0.06 g of copper. The daily requirement In the human body has been estimated to be 200 – 300 mg. All enzymes that utilise ATP in phosphate Transfer require magnesium as the cofactor. The main pigment for the absorption of light In plants is chlorophyll which contains Magnesium. About 99 % of body calcium is Present in bones and teeth. It also plays Important roles in neuromuscular function, Interneuronal transmission, cell membrane Integrity and blood coagulation. The calcium Concentration in plasma is regulated at about 100 mgL–1. It is maintained by two hormones: Calcitonin and parathyroid hormone. Do you Know that bone is not an inert and unchanging Substance but is continuously being Solubilised and redeposited to the extent of 400 mg per day in man? All this calcium Passes through the plasma.
[A] MCQ
1) Quick Lime is prepared on a commercial scale by heating … in a rotary kiln at 1070-1270 K.
a) CaCO3
b) Ca3Al2O6
c) CaSO42H2O
d) CaO
Ans- a) CaCO3
2) An adult body contains about … of Ca
a) 600g
b) 1200g
c) 1800g
d) 2400g
Ans- b) 1200g
3) The calcium Concentration in plasma is regulated at about …
a) 10
b) 50
c) 100
d) 500
Ans- c) 100
4) It was first introduced in England in 1824 by ….
a) Edgar Dobbs
b) Egor Cheliev
c) James Parker
d) Joseph Aspdin
Ans- d) Joseph Aspdin
5) Molecular Formula of plaster of paris is …
a) CaSO42H2O
b) CaSO4½ H2O
c) Ca2SiO4
d) Ca3Al2O6
Ans- b) CaSO4·½ H2O
[B]Short Answers
1) State the biological importance of magnesium and calcium.
Ans- Biological importance of magnesium and calcium- An adult body contains about 25 g of Mg and 1200 g of Ca compared with only 5 g of iron And 0.06 g of copper. The daily requirement In the human body has been estimated to be 200 – 300 mg. All enzymes that utilise ATP in phosphate Transfer require magnesium as the cofactor. The main pigment for the absorption of light In plants is chlorophyll which contains Magnesium. About 99 % of body calcium is Present in bones and teeth. It also plays Important roles in neuromuscular function, Interneuronal transmission, cell membrane Integrity and blood coagulation. The calcium Concentration in plasma is regulated at about 100 mgL–1. It is maintained by two hormones: Calcitonin and parathyroid hormone.
2) Give the preparation and uses of calcium Sulphate (Plaster of Paris).
Ans- It is a hemihydrate of calcium sulphate. It is Obtained when gypsum, CaSO4·2H2O, is Heated to 393 K.
above 393 K, no water of crystallisation is left and anhydrous calcium sulphate, CaSO4 is Formed. This is known as ‘dead burnt plaster’. It has a remarkable property of setting with Water. On mixing with an adequate quantity Of water it forms a plastic mass that gets into a Hard solid in 5 to 15 minutes.
Uses: The largest use of Plaster of Paris is in the Building industry as well as plasters. It is used For immoblising the affected part of organ where There is a bone fracture or sprain. It is also Employed in dentistry, in ornamental work and For making casts of statues and busts.
3) Give the preparation and uses of Calcium Oxide or Quick Lime.
Ans- It is prepared on a commercial scale by Heating limestone (CaCO3) in a rotary kiln at 1070-1270 K.
The carbon dioxide is removed as soon as It is produced to enable the reaction to proceed To completion. Calcium oxide is a white amorphous solid. It has a melting point of 2870 K. On exposure To atmosphere, it absorbs moisture and carbon Dioxide.
The addition of limited amount of water Breaks the lump of lime. This process is called Slaking of lime. Quick lime slaked with soda Gives solid sodalime. Being a basic oxide, it Combines with acidic oxides at high Temperature.
Uses: It is an important primary material for Manufacturing cement and is the cheapest Form of alkali. It is used in the manufacture of sodium Carbonate from caustic soda. It is employed in the purification of sugar And in the manufacture of dye stuffs.
[C]Long Answers
1) Give the preparation and uses of Calcium Carbonate, CaCO3.
Ans- CaCO3 occurs in nature in several Forms like limestone, chalk, marble etc. It can Be prepared by passing carbon dioxide Through slaked lime or by the addition of Sodium carbonate to calcium chloride.
excess of carbon dioxide should be Avoided since this leads to the formation of Water soluble calcium hydrogencarbonate. Calcium carbonate is a white fluffy powder. It is almost insoluble in water. When heated To 1200 K, it decomposes to evolve carbon Dioxide. It is used as a building material in the form of Marble and in the manufacture of quick lime. Calcium carbonate along with magnesium Carbonate is used as a flux in the extraction of Metals such as iron. Specially precipitated CaCO3 Is extensively used in the manufacture Of high quality paper. It is also used as an Antacid, mild abrasive in tooth paste, a Constituent of chewing gum, and a filler in Cosmetics.
2) What is cement ?
Ans- Cement is an important building Material. it was first introduced in England in 1824 by Joseph Aspdin. It is also called Portland cement because it resembles with the Natural limestone quarried in the Isle of Portland, England. Cement is a product obtained by Combining a material rich in lime, CaO with Other material such as clay which contains Silica, SiO2 along with the oxides of Aluminium, iron and magnesium. The average Composition of Portland cement is : CaO, 50-60%; SiO2, 20-25%; Al2O3, 5-10%; MgO, 2-3%; Fe2O3, 1-2% and SO3, 1-2%. For a good Quality cement, the ratio of silica (SiO2) to Alumina (Al2O3) should be between 2.5 and 4 And the ratio of lime (CaO) to the total of the Oxides of silicon (SiO2) aluminium (Al2O3) And iron (Fe2O3) should be as close as possible To 2. The raw materials for the manufacture of Cement are limestone and clay. When clay and Lime are strongly heated together they fuse and React to form ‘cement clinker’. This clinker is Mixed with 2-3% by weight of gypsum (CaSO4·2H2O) to form cement. Thus important Ingredients present in Portland cement are Dicalcium silicate (Ca2SiO4) 26%, tricalcium silicate (Ca3SiO5) 51% and tricalcium Aluminate (Ca3Al2O6) 11%.
Setting of Cement: When mixed with water, The setting of cement takes place to give a hard Mass. This is due to the hydration of the Molecules of the constituents and their Rearrangement. The purpose of adding Gypsum is only to slow down the process of Setting of the cement so that it gets sufficiently Hardened.
Uses: Cement has become a commodity of National necessity for any country next to iron And steel. It is used in concrete and reinforced Concrete, in plastering and in the construction Of bridges, dams and buildings.