1st PUC Chemistry Question Bank Chapter 10 s-Block Elements

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Karnataka 1st PUC Chemistry Question Bank Chapter 10 s-Block Elements

Question 1.
What are s-block elements?
Answer:
The s-block elements are those in which the last electron enter into the outermost s-orbital.

Question 2.
Write the general electronic configuration s-block elements.
Answer:
[Noble gas]ns^1-2

Group 1 Elements: Alkalimetals

Question 1.
Write the general electronic configuration of alkali metals.
Answer:
[Noble gas] ns¹

Question 2.
What is alkali?
Answer:
A water soluble base is called alkali.

Question 3.
What are alkalimetals? Why they are called so?
Answer:
The elements of group 1 except hydrogen are called alkali metals. They are called so because these metals form hydroxides when they react with water, which are strongly alkaline in nature.

Question 4.
Name the radio active alkali metal.
Answer:
Francium (t_1/2 = 21 minutes)

Question 5.
Name the element which is not alkali metal present in group 1.
Answer:
Hydrogen

Question 6.
Name the first alkali metal of group 1 elements.
Answer:
Lithium

Question 7.
What is diagonal relationship between the elements in the periodic table?
Answer:
A relationship within the periodic table by which certain elements in the second period have a close chemical similarity with diagonal element in the next group of the third period is called diagonal relationship.
Example:

Question 8.
What is the cause for the diagonal relationship between the elements in the periodic table?
Answer:
Due to similarity is sizes and charge / radius ratio of the elements.

Question 9.
Give examples which shows diagonal relationship in the periodic table.
Answer:
Li and Mg, Be and Al, B and Si.

Question 10.
Write the electronic configuration of alkali metals.
Answer:

Question 11.
How does ionisation enthalpy of alkali metal changes down the group? Give reason.
Answer:
Ionisation enthalpy of alkali metals decreases down the group because the size of metal atom increases due to the addition of new shells.

Question 12.
How does hydration enthalpy of alkali metals changes down the group? Give reason.
Answer:
Hydration enthalpy of alkali metal ions decrease down the group due to increase in their ionic sizes.
Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺

Question 13.
Li⁺ ion has high degree of hydration. Why?
Answer:
Due to its small ionic size.

Question 14.
Lithium salts are mostly hydrated. Why?
Answer:
Because Li⁺ has maximum degree of hydration.

Question 15.
Alkali metals have low density and decreases down the group. Give reason.
Answer:
Low density of alkali metals is due to large size. Density of alkali metal decreases down the group due to increase in their size.

Question 16.
Alkali metals are soft metals. Why?
Answer:
Due to weak metallic bonding.

Question 17.
Melting and boiling points of alkali metals are low. Why?
Answer:
Due to weak metallic bone ing.

Question 18.
The alkali metals and their salts impart characteristic colour to an non-luminous flame. Explain.
Answer:
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, there is an emission of radiation in the visible region.

Question 19.
How do you detect alkali metals in the laboratory?
Answer:
By flame tests, flame photometry or atomic absorption spectroscopy.

Question 20.
Caesium and potassium metals are used as electrodes in photoelectric cells. Explain.
Answer:
When caesium and potassium metals are irradiated with light, the light energy absorbed is sufficient to make an atom o lose electron. This property makes caesium and potassium useful as electrodes in photoelectric cells.

Question 21.
Alkali metals are highly reactive. Why?
Answer:
Dae to their large size and low ionisation enthalpy.

Question 22.
Explain the reactivity of alkali metals towards air.
Answer:
(i) Alkali- metals tarnish in dry air due to the formation of oxides, which in turn react with moisture to form hydroxides.
(ii) They bum vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide and other alkali metals forms super oxides.

Note: Super oxide (O₂^–) ion is stable only in the presence of large cations such as K⁺, Rb⁺ and Cs⁺

Question 23.
How does lithium reacts with nitrogen of air?
Answer:
Lithium reacts with nitrogen of air to form lithium nitride (Li₃N)
6Li + N₂ → 2Li₃N

Question 24.
Alkali metals are normally kept in kerosene oil. Why?
Answer:
Due to high reactivity of alkali metals with air and water.

Question 25.
What is the oxidation state of K in KO₂?
Answer:
+ 1

Question 26.
How does alkali metals react with water? Give general reaction.
Answer:
Alkali metals react with water to form alkali metal hydroxide with the liberation of dihydrogen gas.
2M + 2H₂O → 2MOH + H₂
where M = alkali metal

Question 27.
Although E° of Li is most negative than sodium, lithium reacts with water less vigorously than sodium. Why?
Answer:
Because Li has small size and high hydration energy.

Question 28.
Explain the reactivity of alkali metals with dihydrogen.
Answer:
Alkali metals react with dihydrogen at about 673 K (lithium 1073 K) to form ionic solids called alkali metal hydrides. .
2M + H₂ → 2M⁺H^–

Question 29.
How does alkali metals react with halogens?
Answer:
Alkali metals react with halogens vigorously to form ionic halides M⁺X^–.

Question 30.
Lithium halide is covalent. Why?
Answer:
Due to high polarisaiton capability of lithium ion.

Question 31.
What is polarisation?
Answer:
The distortion of electron cloud of the anion by the cation is called polarisation.

Question 32.
Lithium iodide is most covalent nature than other lithium halides. Why?
Answer:
Since anion with large size (I^–) can be easily distorted, among halides. So lithium iodide is the most covalent in nature.

Question 33.
Lithium is powerful reducing agent than other alkali metals. Why?
Answer:
Due to small size of lithium ion, its hydration enthalpy is very high. Therefore Li has more negative E° (SRP) value and hence it has high reducing power.
Note: Lesser E° (SRP), stronger reducing agent.

Question 34.
Alkali metal solutions in ammonia are blue and conducting in nature. Explain.
Answer:
The blue colour of alkalil metal solutions in ammonia is due to ammoniated electrons which absorb energy corresponding to red region of the visible light for their excitation to higher energy levels. The transmitted light is blue which imparts blue to the solution.
The electrical conductivity of alkali metal solutions in ammonia is due to both ammoniated cations and ammoniated electrons.

Question 35.
Write the chemical reaction which occurs when ammoniated alkali metal solutions on standing.
Answer:
Ammoniated alkali metal solutions on standing liberates hydrogen gas resulting in the formation of amide.
M⁺_am + e^– + NH₃(l) → MNH_2(am) + \(\frac { 1 }{ 2 }\)H₂(g)
where am = solution in ammonia.

Question 36.
Give any four uses of alkali metals.
Answer:
1. 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
• with magnesium to make armour plates, used in thermonuclear reactions.

2. Lithium is also used to make electrochemical cells.

3. Sodium is used to make a Na/Pb alloy needed to make tetraethyl lead (PbEt₄) and tetramethyl lead (PbMe₄). These organolead compounds were earlier used as anti-knock additives to petrol, but nowadays vehicles use lead-free petrol.

4. Liquid sodium metal is used as a coolant in fast breeder nuclear reactors.

5. 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.

6. Caesium rubidium are used in photoelectric cells.

General Characteristics of the Compounds of the Alkali Metals

Question 1.
Name the type of oxides formed when alkali metals are heated with air.
Answer:
When heated with excess of air, lithium forms monoxide, Li₂O; sodium forms peroxide, Na₂O₂ whereas potassium, rubidium and caesium form super oxides having general formula MO₂.

Question 2.
Stability of peroxide or super oxide increases as sizes of the metal ion increases. Why?
Answer:
Due to stabilisation of large anions by large cations through lattice energy effects.

Question 3.
Explain the hydrolysis of alkali metal oxides, peroxides and super oxides.
Answer:
The alkali metal oxides, peroxides and super oxides readily undergoes hydrolysis by water to form the corresponding hydroxides.
M₂O + H₂O → 2M⁺ + 20H^–
M₂O₂ + 2H₂O → 2M⁺ + 2OH^– + H₂O₂
2MO₂ + 2H₂O → 2M⁺ + 2OH^– + H₂O₂ + O₂

Question 4.
Why is KO₂ is paramagnetic?
Answer:
Due to the presence of unpaired electron in π*_2p molecular orbital of O^–₂ ion.

Question 5.
The solubility of LiF in water is less. Why?
Answer:
Due to its high hydration enthalpy.

Question 6.
Low solubility of Csl in water is due to what?
Answer:
Due to smaller hydration enthalpy of its two ions.

Question 7.
What is an oxoacid?
Answer:
Oxoacids are those in which the acidic proton is on hydroxyl group with an OXO group attached to the same atom.
Examples: H₂CO₃, H₂SO₄

Question 8.
How does alkali metals reacts with oxoacids and explain the stability of their products.
Answer:
Alkali metals form salts with all oxoacids.
The carbonates (M₂CO₃) of alkali metals are stable upto 1273 K. Li₂CO₃ is less stable and decomposes readily.

All the bicarbonates on gentle heating undergo decomposition to form carbonates with the evolution of CO₂.

Anomalous Properties of Lithium

Question 1.
Give reasons: Li shows amomalous properties from rest of the group from many respects.
Answer:

• Small size of lithium atom and its ion.
• Higher polarization power of Li⁺. (i.e., charge to size ratio)
• High ionisation enthalpy.
• Non-availability of rf-orbitals in its valence shell.
• Strong intermetallic bonding.

Question 2.
Point out the difference between Lithium and other alkali metals.
Answer:

• Lithium is harder than sodium and potassium which are so soft that they can be cut by a knife.
• The melting and boiling points of lithium are comparatively high.
• Lithium forms only monoxide with oxygen, other alkali metals also form peroxides and superoxides.
• Lithium combines with nitrogen to form nitrides, while other alkali metals do not.
• Lithium chloride is deliquescent and crystallizes as a hydrate LiCl.2H₂O, whereas other alkali metal chlorides do not form hydrates.

Question 3.
Give the similarities between lithium and magnesium due to diagonal relationship.
Answer:

• Both form chlorides and bromides that hydrolyse slowly and are soluble in ethanol.
• Both form colourless or slightly coloured crystalline nitrides.
• Both burn in air to give the monoxide only.
• Both form carbonates that decompose on heating.
• Oxides o.f lithium and magnesium do not combine with excess of oxygen to form any superoxide.
• Both lithium and magnesium are harder and lighter than other elements in the respective groups.

Some Important Compounds of Sodium

Question 1.
Describe the manufacture of sodium carbonate by Solvay’s process or Ammonia – soda process.
Answer:
Solvay process consists of the following steps:
Step 1: The sodium hydrogen carbonate is prepared by passing CO₂ to a concentrated solution of sodium chloride (brine), saturated with ammonia.
2NH₃ + H₂O + CO₂ → (NH₄)₂ CO₃
(NH₄)₂ CO₃ + H₂O + CO₂ → 2NH₄HCO₃
2NH₄HCO₃ + NaCl → NH₄Cl + NaHCO₃

Step 2: Sodium hydrogen carbonate is heated to about 300°C to decompose into sodium carbonate and carbon dioxide.

Step 3: Ammonia can be recovered by treating ammonium chloride with calcium hydroxide.
2NH₄Cl (aq) + Ca(OH)₂(s) → 2NH₃(g) + CaCl₂ (aq) + 2H₂O(1)
Note: Solvay process cannot be extended to the manufacture of potassium carbonate because potassium hydrogen carbonate is too soluble to be precipitated by the addition of ammonium hydrogen carbonate to a saturated solution of potassium chloride.

Question 2.
What is the composition of washing soda?
Answer:
Na₂CO₃ . 10H₂O

Question 3.
Explain the properties of sodium carbonate.
Answer:
Sodium carbonate is a white, crystalline compound soluble in water (absorbing moisture from the air) but insoluble in alcohol. It forms a strongly alkaline aqueous solution. Sodium carbonate decahydrate (Na₂CO₃ . 10H₂O), is commonly called “washing soda”. It is readily soluble in water. On heating, the decahydrate loses its water of crystallization to form monohydrate. Above 373K, the monohydrate becomes completely anhydrous and changes to a white powder called soda ash.

Question 4.
How does sodium carbonate undergoes hydrolysis with water? Give equation.
Answer:
Sodium carbonate gets hydrolysed by water to form an alkaline solution.
Na₂CO₃ + H₂O → 2NaOH + H₂CO₃
Sodium carbonate solution is basic due to hydrolysis.

Question 5.
Give any four uses of sodium carbonate.
Answer:
Sodium carbonate is used:

1. in water softening, laundering and cleaning.
2. in the manufacture of glass, soap, borax and caustic soda.
3. in paper, paints and textile industries.
4. It is an important laboratory reagent both in qualitative and quantitative analysis.

Question 6.
Which is the main source of sodium chloride?
Answer:
Sea water

Question 7.
What is the percentage of sodium chloride in sea water?
Answer:
2.7 to 2.9%

Question 8.
Explain the production of sodium chloride from sea water.
Answer:
NaCl is obtained from sea water by evaporation in the presence of sun light. Crude sodium chloride contains sodium sulphate, calcium sulphate, calcium chloride and magnesium chloride as impurities. 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.
Note: 36 g NaCl is soluble in 100 g water at 273 K.

Question 9.
Give any two uses of sodium chloride.
Answer:

1. It is used as a common salt or table salt for domestic purpose.
2. It is used in the preparation of Na₂O₂ NaOH and Na₂CO₃.

Question 10.
Describe the manufacture of sodium hydroxide or caustic soda (NaOH) by Castner – Kellner’s process.
Answer:
Caustic soda is manufactured by Castner-Kellner’s electrolytic process. A Brine solution (saturated solution of sodium chloride in water) is electrolysed using a mercury cathode and a carbon anode. Sodium metal discharged at the cathode combines with mercury to form sodium amalgam.
2Na⁺ + 2e^– → 2Na (reduction)
Na + Hg → Na/Hg (amalgam)

Chlorine gas is evolved at the anode.
2Cl^– → Cl₂ + 2e- (oxidation) e^–

The mercury containing dissolved sodium is sent to another chamber called ‘decomposer’ where sodium reacts with water forming sodium hydroxide and hydrogen. Decomposer is packed with graphite blocks as hydrogen is easily liberated over graphite surface.
2Na/Hg + 2H₂O → 2NaOH + H₂ + 2Hg
The solution, which flows out from decomposer, is NaOH solution, which is evaporated to dryness.

Question 11.
Give two properties of sodium hydroxide.
Answer:
Properties: Sodium hydroxide is a deliquescent solid. So, it absorbs moisture and carbon dioxide to form liquid carbonates and bicarbonates. Therefore, sodium hydroxide is always stored in air tight bottles.

Question 12.
Give any four uses of sodium hydroxide.
Answer:
Sodium hydroxide is mainly used for the following purposes:

1. Refining of petroleum
2. Making paper
3. in the textile industries for mercerising cotton fabrics
4. Printing and Dyeing
5. Manufacture of soap
6. Purification of bauxite
7. Refining of oils and fats

Question 13.
Sodium hydrogen carbonate is called baking soda. Why?
Answer:
Because it decomposes on heating to generate bubbles of CO₂.

Question 14.
How do you prepare sodium hydrogen carbonate?
Answer:
Sodium bicarbonate is prepared from purified aqueous sodium carbonate solution by passing carbon dioxide through it.
Na₂CO₃ + H₂O + CO₂ → 2NaHCO₃
The bicarbonate precipitates out. It is separated and dried.

Question 15.
Give any two uses of sodium bicarbonate (baking soda)
Answer:

1. Used as baking powder.
2. Used in treating wool and silk
3. Used as fire extinguisher
4. Used as mild antiseptic for skin infections, etc.

Biological Importance of Sodium and Potassium

Question 1.
Give the biological importance of sodium and potassium.
Answer:
Sodium ions participate:

• in the transmission of nerve signals.
• in regulating the flow of water across cell membranes.
• in the transparent of sugars and amino acids into cells.

Potassium ions

• activate many enzymes.
• participate in the oxidation of glucose to produce ATP with sodium.
• Potassium ions are responsible for the transmission of nerve signals.

Group 2 Elements: Alkaline Earth Metals

Question 1.
What are alkaline earth metals? Why are they called so?
Answer:
The elements of group 2 (except beryllium) are called alkaline earth metals. They are called so because their oxides and hydroxides are alkaline and these metal oxides are found in earth’s crust.

Question 2.
Name the radioactive alkaline earth metal.
Answer:
Radium

Question 3.
Write the general electronic configuration of alkaline earth metals.
Answer:
[Noble gas] ns²

Question 4.
Write the electronic configurations of alkaline earth metals.
Answer:

Question 5.
Atomic and ionic radii of the alkaline earth metals are smaller than corresponding alkali metals. Give reason.
Answer:
Due to increased nuclear charge in alkaline earth metals.

Question 6.
How does atomic and ionic radii of alkaline earth metals vary down the group?
Answer:
Atomic and ionic radii increase with increase in atomic number.

Question 7.
Alkaline earth metals have higher first ionisation enthalpies than alkali metals. Why?
Answer:
Due to small size and completely filled s-orbitals.
Note: Second ionisation enthalpies of alkaline earth metals are less than corresponding alkali metals.

Question 8.
How does ionisation enthalpy vary down the group of alkaline earth metals?
Answer:
Ionisation enthalpy of alkaline earth metals decreases down the group because atomic size increases.

Question 9.
Explain the variation of hydration enthalpies of alkaline earth metals.
Answer:
Hydration enthalpies of alkaline earth metals decreases down the group due to increase in size. So hydration enthalpies decreases in the order.
Be²⁺ >Mg²⁺ >Ca²⁺ > Sr²⁺ > Ba²⁺
Hydration enthalpies of alkaline earth metals are greater than alkali metals. Therefore alkaline earth metals are extensively hydrated than those of alkali metals. Thus, MgCl₂ and CaCl₂ exist as MgCl₂.6H₂O and CaCl₂.6H₂O.

Question 10.
Alkaline earth metals impart characteristic colour when heated on a Bunsen flame. Explain.
Answer:
This property is due to the ease of excitation of their valence electrons. In flame the electrons return to other ground state, they emit absorbed energy in the visible light having characteristic wavelengths. Depending upon the wavelength of the light emitted, different colours are imparted to the flame.

Question 11.
Give the colours imparted by Ca²⁺, Sr²⁺, Ba²⁺ and Ra²⁺.
Answer:

Ion	Colour
Ca2+	Brick-red
Sr2+	Crimson red
Ba2+	Apple Green
Ra2+	Carmine – red

Question 12.
Beryllium and magnesium do not impart any colour to the flame. Why?
Answer:
This is because the electrons in beryllium and magnesium are strongly bound (very high ionisation energy) to get excited by flame.

Question 13.
Be and Mg are almost inert to air and water. Why?
Answer:
Because of the formation of oxide layer on their surface.

Question 14.
Name the products formed when powdered Be burns in air.
Answer:
Beryllium oxide (BeO) and Beryllium nitride (Be₃N₂)

Question 15.
Name the products formed when Mg burns in air.
Answer:
Magnesium oxide (MgO) and Magnesium nitride (Mg₃N₂)

Question 16.
How does alkaline earth metals react with halogens?
Answer:
The alkaline earth metals (M = Be, Mg, Ca, Sr, Ba) react with halogens (X = Cl, Br, F, I) at elevated temperature to form halides of the types MX₂.
M + X₂ → MX₂

Question 17.
How do you prepare BeF₂?
Answer:
By thermal decomposition of (NH₄)₂BeF₄.
(NH₄)₂BeF₄ → BeF₂ + 2NH₄F

Question 18.
Give the chemical reaction for the preparation of BeCl₂ from its oxide.
Answer:

Question 19.
Explain the reactivity of alkaline earth metals towards hydrogen.
Answer:
All the alkaline earth metals except ‘Be’ combine with hydrogen directly on heating to form metal hydrides of formula MH₂.

The hydride of beryllium can also be obtained by the reduction of BeCl₂ with LiAlH₄.

Question 20.
Name the gas liberated when alkaline earth metals react with acids.
Answer:
Dihydrogen gas (H₂)

Question 21.
Alkaline earth metals are weaker reducing agents than alkali metals. Why?
Answer:
Because ionisation energies of alkaline earth metals are higher and their electrode potentials are less negative than corresponding alkali metals.

Question 22.
Although Be has less reduction potential than other alkaline earth metals, it is good reducing agent. Why?
Answer:
Because Be has high hydration enthalpy (due to its small size) and large value of atomisation enthalpy.

Question 23.
What is the colour of the solution when alkaline earth metals are dissolved in ammonia?
Answer:
Blue

Question 24.
Give important uses of alkaline earth metals.
Answer:
1. 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.

2. 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.

3. Calcium is used in the extraction of metals from oxides which are difficult to reduce with carbon.
4. 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.
5. Radium salts are used in radiotherapy, for example, in the treatment of cancer.

General Characteristics of Alkaline Earth Metals

Question 1.
Alkaline earth metal compounds are less ionic than alkali metal compounds. Why?
Answer:
Due to increased nuclear charge and smaller size.

Question 2.
Which oxide of alkaline earth metal is amphoteric and covalent in nature?
Answer:
BeO

Question 3.
Oxides of alkaline earth metals are stable to heat. Why?
Answer:
Because enthalpies of formation of these oxides are quite high.

Question 4.
Explain the amphoteric nature of Beryllium hydroxide.
Answer:
Beryllium hydroxide, Be(OH)2 reacts with acids and alkalies forming amphoteric products.
Be(OH)₂ + 2OH^– → [Be(OH)₄]^2- (beryllate ion)
Be(OH)₂ + 2HCl + H₂O → [Be(OH)₄]Cl₂

Question 5.
Write the polymeric structure of Beryllium chloride.
Answer:

Question 6.
Which alkaline earth metal halide is covalent in nature?
Answer:
BeCl₂

Question 7.
Fluorides of alkaline earth metals are less soluble in water than chlorides. Why?
Answer:
Due to their high lattice energies.

Question 8.
Which carbonate salt of alkaline earth metal is unstable?
Answer:
Beryllium carbonate.

Question 9.
Beryllium and magnesium sulphates are soluble in water. Why?
Answer:
Greater hydration enthalpies of Be²⁺ and Mg²⁺ ions overcome the lattice enthalpy factor. Therefore, beryllium and magnesium sulphates are soluble in water.

Question 10.
Why does the solubility of alkaline earth metal hydroxides in water Increases down the group?
Answer:
Among alkaline earth metal hydroxides, the anion being common, the cationic radius will influence the lattice enthalpy. Since lattice enthalpy decreases much more than the hydration enthalpy with Increasing ionic size, the solubility increases down the group.

Question 11.
Why does the solubility of alkaline earth metal carbonates and sulphates in water decreases down the group?
Answer:
The size of anions being much larger compared to cations, the lattice enthalpy will remain almost constant within a particular group. Since the hydration enthalpies decrease down the group, solubility also decreases for alkaline earth metal carbonates and sulphates.

Anomalous Behavior of Beryllium

Question 1.
Beryllium shows anomalous behaviour as compared other elements of the some group. Why?
Answer:
It is mainly because of,

• Its small size and high polarizing power.
• Relatively high electronegativity and ionization energy as compared to other members.
• Absence of vacant ‘d’-orbitals in its valence shell.

Question 2.
Give the anomalous properties of beryllium.
Answer:
1. Beryllium forms covalent compounds whereas other members form ionic compounds. (This is due to large ionisation enthalpy and small ionic size of beryllium).

2. Beryllium does not exhibit coordination more than four (as it has four orbitals in the valence shell).
The remaining members of the group can have coordination number of six (by making use of d-orbitals).

3. The oxide and hydroxide of beryllium are amphoteric in nature where as other alkaline earth metal oxides and hydroxides are basic in nature.

Question 3.
Be shows diagonal relation (similar properties) with Al. Why?
Answer:
Because charge / radius ratio of Be²⁺ and Al³⁺ ion are same.

Question 4.
Give the similarities of Be and AI on the basis of diagonal relationship.
Answer:
1. Like aluminium, Be is not attacked by acids because of the presence of an oxide film on the surface of metal (called passivity).

2. Both beryllium and aluminium hydroxides dissolved in sodium hydroxide (alkali) to form beryliate ion, [Be(OH)₄]^2- and aluminate ion [Al(OH)₄]^– respectively.

3. Both BeCl₂ and AICI₃ have bridged chloride polymeric structure. Both are lewis acids and used as Friedel crafts catalysts.

4. Both beryllium and aluminium ions have strong tendency to form complexes BeF^2-₄ , AlF^3-₆

Some Important Compounds of Calcium

Question 1.
Give the preparation of quicklime (calcium oxide).
Answer:
It is prepared commercially by heating limestone (CaCO₃) in a rotary kiln at 1070-1270 K.
CaCO₃ ⇌ CaO + CO₂ : ΔH = + ve
This is a reversible reaction. It is endothermic in the forward direction. To ensure that the limestone is completely converted to lime, carbon dioxide is allowed to escape from the kiln. This makes carbon dioxide unavailable for the reverse reaction.

Question 2.
What is slaking of lime?
Answer:
When water is added to lime, it becomes hot and cracks to form a white powder. This process is called slaking of lime.

Question 3.
How is soda lime obtained?
Answer:
Quick lime slaked with soda gives soda lime.

Question 4.
How is CaO oxide reacts with SiO₂ and P₄O₁₀?
Answer:
CaO, being a basic oxide reacts with acidic oxides like SiO₂ and P₄O₁₀ at high temperature.

Question 5.
Give two uses of quick lime.
Answer:

• used in the manufacture of cement
• used in the manufacture of sodium carbonate from caustic soda
• used in the purification of sugar.

Question 6.
How is calcium hydroxide prepared?
Answer:
Calcium hydroxide prepared by adding water to quick lime (CaO).
CaO + H₂O → Ca(OH)₂
Note: Calcium hydroxide is also called as slaked lime.

Question 7.
When carbon dioxide is passed through lime water, it turns milky. Why? Give the chemical reaction.
Answer:
Due to the formation of insoluble calcium carbonate.
Ca(OH)₂ + CO₂ → CaCO₃ + H₂O

Question 8.
What is lime water?
Answer:
Aqueous solution of calcium hydroxide is called lime water.

Question 9.
What is milk of lime?
Answer:
The suspension of slaked lime in water is called milk of lime.

Question 10.
The precipitate of calcium carbonate dissolves by passing excess CO₂. Why? Write the chemical reaction.
Answer:
Due to the formation of soluble calcium hydrogen carbonate.
CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂

Question 11.
How does milk of lime reacts with chlorine? Write the chemical reaction.
Answer:
Milk of lime reacts with chlorine to form hypochlorate.

Question 12.
Give uses of calcium hydroxide, (slaked lime)
Answer:
Calcium hydroxide is used
(a) in the preparation of morter, a building material.
(b) in white wash due to its disinfactant nature.
(c) in glass making, tanning industry.
(d) in the preparation of bleaching powder and for purification of sugar.

Question 13.
Give the occurrence of calcium carbonate.
Answer:
Calcium carbonate occurs in nature in several forms like lime stone, chalk, marble etc.

Question 14.
Give two methods of preparation of calcium carbonate.
Answer:
1. Calcium carbonate is prepared by passing carbon dioxide through slaked lime.
Ca(OH)₂ + CO₂ → CaCO₃ + H₂O
Note: Excess of CO₂ should be avoided because this leads to the formation of water soluble calcium hydrogen carbonate. Or Calcium bicarbonate.

2. CaCO₃ can also be obtained by the addition of sodium carbonate to calcium chloride.
CaCl₂ + Na₂CO₃ → CaCO₃ + 2NaCl

Question 15.
What is the action of heat on calcium carbonate?
Answer:
When calcium carbonate is heated strongly to 1200 K, it undergoes decomposition to form CaO and CO₂.

Question 16.
How does CaCO₃ reacts with HCI and H₂SO₄? Give examples.
Answer:
CaC03 reacts with dilute acids like HCl, H₂SO₄ etc. to liberate carbon dioxide.
CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂

Question 17.
Give the uses of calcium carbonate.
Answer:

• Calcium carbonate is mostly used in the construction industry as a building material such as limestone, marble or chalk. It can also be used as an ingredient of cement.
• Calcium carbonate in the form of chalk is used as filler in the paper industry. It is used as a cheap way to make bright opaque paper.
• Limestone is used in decoration and construction. The main use being its contribution to make mortar, it is used in bonding bricks, concrete blocks, stones, and tiles.
• Limestone is used in the steel industry, ornamental stone, chemical and optical uses and as mineral specimens.
• The limestone can also be used to make glasses.
• It is also used as an antacid, mild abrasive in tooth paste, a constituent of chewing gum, and a filler in cosmetics.

Question 18.
What is plaster of paris? How it is obtained?
Answer:
Plaster of paris is a hemihydrate of calcium sulphate with composition CaSO₄ . \(\frac { 1 }{ 2 }\) H₂O. Plaster of paris is obtained by heating gypsum (CaSO₄.2H₂O) at 393 K.

Question 19.
What is dead burnt plaster?
Answer:
When gypsum is heated above 393 K, anhydrous calcium sulphate is formed. This is called dead burnt plaster.

Question 20.
Give uses of plaster of paris.
Answer:

• In surgery for setting broken or fractured bones.
• For making casts for statues, in dentistry, for surgical instruments, and toys, etc.
• In making black board chalks.
• In construction industry.

Question 21.
What is cement? Give the average composition of Portland cement.
Answer:
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:

Components of cement	Approximate composition
CaO	50-60%
SiO2	20-25%
Al2O3	5-10%
MgO	3%
Fe2O3	1-2%
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.

Question 22.
How is cement obtained? Give the important ingredients present in Portland cement.
Answer:
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. Important ingredients present in Portland cement are dicalcium silicate (Ca₂SiO₄) 26%, tricalcium silicate (Ca₃SiO₅) 51% and tricalcium aluminate (Ca₃Al₂O₆) 11%.

Question 23.
Explain the process of setting of cement.
Answer:
Cement sets when mixed with water by a complex series of chemical reactions. 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.

Question 24.
Give uses of cement.
Answer:

• Buildings (floors-concrete and reinforced concrete (RCC), beams, columns, roofing, piles, bricks, mortar, panels, plaster)
• Transportation (roads, pathways, crossings, bridges, sleepers, viaducts, tunnels, stabilization, runways, parking)
• Water supply (pipes, culverts, kerbing, drains, canals, weirs, dams, tanks, pools)
• Civil works (piers, docks, retaining walls, silos, warehousing, poles, pylons, fencing)
• Agriculture (buildings, processing, housing, feedlots, irrigation).

Biological Importance of Magnesium and Calcium

Question 1.
Give the biological importance of magnesium and calcium.
Answer:

• 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.
• Calcium plays important role in neuromuscular function, interneuronal transmission, cell membrane integrity and blood coagulation.