THE `S` BLOCK ELEMENTS INTRODUCTION The s block elements of the periodic table are those in which
the last electron enters the outermost s-orbital. 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. The elements of group2 include Beryllium, Magnesium, Calcium,
Strontium, Barium and Radium .These elements are collectively known as alkaline earth metals. The general electronic configuration of s-block elements
is(noble gas)
for alkali metals and(noble gas)
for
alkaline earth metals.
Anomalous properties of Lithium and Beryllium Lithium and beryllium resemble the second element of the
following group. lithium shows similarities to magnesium and beryllium to
aluminum in many of their properties. This type of diagonal similarity is commonly referred to as diagonal relationship in the periodic table. Diagonal relationship is due to the similarities in ionic sizes
and /or charge/radius ratio of the elements.
GROUP 1- ELEMENTS: ALKALI METALS The alkali metals show regular trends in their physical and chemical properties with the increasing atomic number.the atomic ato mic physical and chemical properties are listed below. ELECTRONIC CONFIGURATION
I.
All the electrons have one valence electron,
in
the outside of the noble gas core. II.
The loosely held s-electron in the outermost valence shell of these elements makes them most electropositive metals.
III.
They readily lose electron to give monovalent m+ ions.
IV.
Hence they are Free State in nature.
ATOMIC AND IONIC RADII
With the increase in atomic number, they become larger in size.
The monovalent m+ ions are smaller than the parent atom.
The atomic and ionic radii of the alkali metals increase on moving down the group. IONIZATION ENTHALPY .Theyare considerably low and decreases down the group. .this is because of the effect of larger in size outweighs the increase in nuclear charge, and
GROUP 1- ELEMENTS: ALKALI METALS The alkali metals show regular trends in their physical and chemical properties with the increasing atomic number.the atomic ato mic physical and chemical properties are listed below. ELECTRONIC CONFIGURATION
I.
All the electrons have one valence electron,
in
the outside of the noble gas core. II.
The loosely held s-electron in the outermost valence shell of these elements makes them most electropositive metals.
III.
They readily lose electron to give monovalent m+ ions.
IV.
Hence they are Free State in nature.
ATOMIC AND IONIC RADII
With the increase in atomic number, they become larger in size.
The monovalent m+ ions are smaller than the parent atom.
The atomic and ionic radii of the alkali metals increase on moving down the group. IONIZATION ENTHALPY .Theyare considerably low and decreases down the group. .this is because of the effect of larger in size outweighs the increase in nuclear charge, and
the outermost electron is very well screened from the nuclear charge.
HYDRATION ENTHALPY *The hydration enthalpies of alkali metals ions decrease with increase in ionic size. Li+ > na+ >k+ >Rb+ >Cs+ > Li+ has maximum degree of hydration and for this reason lithium salts are mostly hydrated.
PHYSICAL PROPERTIES All the Alkali metals have low density. The melting and boiling point of the alkali metals
are low.
The alkali metals and their salts impart
characteristics colour to an oxidizing flame. This is because of the heat from the
flame excites the outermost orbital electron to a higher energy level.
WHY CAESIUM AND POTASSIUM USEFUL AS ELECTRODES IN PHOTOELECTIRC CELLS? Alkali metals can be detected by the flame tests and can be determined by flame photometric or atomic absorption spectroscopy. These elements
when irradiated with light, the light energy absorbed maybe sufficient to make an atom lose electron. This property makes caesium and potassium useful as
electrodes in photoelectric cells .
CHEMICAL PROPERTIES Reactivity towards oxygen:
Lithium forms monoxide,sodium forms peroxide and other elements form superoxides. Reactivity towards water
Alkali metals react with water forming hydroxides and hydrogen gas is evolved. Solutions in liquid ammonia Alkali metals dissolve in liquid ammonia
forming deep blue coloured solutions. Explanation In solution ,alkali metal looses the valence
electron.
Both the cation and the electron combine
with ammonia to form ammoniated cation and ammoniated electron. This ammoniated electron is responsible for
the blue colour. REDUCING NATURE Alkali metals are strong reducing agents. This is due to their greater ease to loose
electrons They have large value of negative reduction
potentials Lithium is the strongest reducing agent due
to its high hydration energy. General Characteristics Of the Compounds of the Alkali Metals Oxides and Hydroxides
Lithium forms mainly its oxide, Li2O (also some peroxide Li2O2)
Sodium forms peroxide, NaO2
Pottasium, Rubidium and Caesium form superoxides, MO2
As the size of the metal ion increases, the increasing stability of the peroxide or superoxideis due to stabilization of large anions by larger cations through lattice energy effects.
The superoxides are also paramagnetic.
Sodium peroxide is widely used as an oxidizing agent in inorganic chemistry.
Halides
M.P and B.P always follow the trend: fluoride>chloride>bromide>iodide.
All these halides are soluble in water.
The low solubility of LiF in water is due to its high lattice enthalpy whereas the low solubility of CsI is die to smaller hydration enthalpy if its two ions
Other halides of lithium are soluble in ethanol, acetone and ethylacetate.
Salts and Oxo-Acids
Lithium carbonate is not so stable to heat
This is because lithium being very small in size polarizes a large CO
2+ 3
ion leading to the formation of more
stable Li2O and CO.
Its Hydrogencarbonate does not exist as a solid.
Anomalous Properties of Lithium.
Anomalous behavior due to: 1.Very small size of Li atom and ion. 2.High Polarizing effect (i.e. charge/radius ratio)
This results in increased covalent characters of lithium compounds, responsible for their solubility in organic solvents.
Lithium also shows diagonal relationship to magnesium.
Differences between Lithium and other Alkali metals:
Li is harder and its M.P and B.P are higher than the other alkali metals.
It is least reactive, stronger reducing agent than other alkali metals. Its combustion in air leads to formation of monoxide, Li2O and nitride, Li3N unlike other alkali metals
LiCl is deliquescent and crystallizes as a hydrate, LiCl.2H 2O whereas other alkali metal chlorides don’t form hydrates.
Li Hydrogencarbonate is not obtained in solid form while all other elements form solid hudrogencarbonates.
On reacting with ethyne, it DOES NOT form ethynide unlike other alkali metals.
LiNO3 when heated gives Lithium oxide, Li2O whereas other alkali metal nitrates decompose to give their corresponding nitrites. 4LiNO3 → 2Li2O + 4NO2 + O2
2NaNO3 → 2NaNO2 + O2
LiF and Li2O are less soluble in water than fluorides and oxides of other alkali metals (comparatively)
Points of similarities between Lithium and Magnesium
Similarity arises because of their similar sizes between them. Main points of similarity:
Both are harder and lighter than other elements in their respective groups.
Both react slowly with water. Their oxides and hydroxides are much less soluble, their hydroxides decompose easily on heating and both form nitride, Li 3N and Mg3N2 by direct combination with Nitrogen
Their oxides don’t combine with excess oxygen to form
Superoxide
Their carbonates decompose easily on heating to form oxides and CO2. Solid Hydrogencarbonates are not formed by lithium and magnesium.
Both their chlorides dissolve in ethanol
Their chlorides are deliquescent and crystallize from aqueous solution as hydrates, LiCl.2H2O amd MgCl2.8H2O.
SODIUM CARBONATE :(i) Sodium carbonate is generally prepared by solvay process. (ii) In this process advantage is taken of the low solubility of sodium hydrogencarbonate whereby it gets precipitated in the reaction sodium chloride with ammonium hydrogen carbonate. (iii) By passing carbon dioxide to conc.solution of NaCl saturated with ammonia where ammonium carbonate followed by ammonium carbonate are formed.
2NH3 + H2O+ CO2 ---> (NH4)2CO3 (NH4)2CO3 + H2O + CO2 ---> 2NH4HCO3 NH4HCO3 + NaCl ---> NH4Cl + NaHCO3
PROPERTIES :(i) it is also called washing soda. (ii) sodium carbonate is a white crystalline solid which exists as a decahydrate. (iii) on heating, loses its water of crystallisation to form monohydrate. (iv) above 373 K, the monohydrate becomes completely anhydrous & changes to a white powder called soda ash. 373k Na2CO3.10H2O ----> Na2CO3.H2O + 9H2O <373K Na2CO3.10H2O -----> Na2CO3 + H2O
USES :(i) used in water softening, laundering and cleaning. (ii) used in manufacture of glass, soap and caustic soda. (iii) in paper, paints and textile industries.
SODIUM CHLORIDE :-
process: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 seperate out.Calcium and magnesium chloride, being more soluble than sodium chloride, remains in solution. USES:(i) used as a common salt or table salt for domestic pupose. (ii) used for the preparation of sodium peroxide, sodium hydroxide and sodium carbonate.
SODIUM HYDROXIDE-
*IT IS PREPARED BY THE ELECTROLYSIS OF AQUEOUS SOLUTION OF BRINE USING MERCURY CATHODE AND GRAPHITE ANODE
*IT IS PREPARED BY THE ELECTROLYSIS OF AQUEOUS SOLUTION OF BRINE USING MERCURY CATHODE AND GRAPHITE ANODE
*THE ELCECTROLYSIS IS CARRIED OUT IN A SPECIAL DESIGNED CELL CALLED CASTNER-KELLNER CELL.
*Na METAL DISCHARGED AT THE CATHODE FORMS Na AMALGAM WITH MERCURY CHLORINE GAS AT ANNODE THE FOLLOWING REACTION TAKES PLACE
-at the cathode Na+ + e-
Na-amalgam
Na + amalgam → Na - amalgam
ANODE-
Cl - e- → cl ,
2Na-amalgam +2H2O
PROPERTIES-
-
cl +cl → cl2 2NaOH +2 Hg+H2
(i)NaOH IS A WHITE, TRANSLUCENT SOLID.
(ii)MELTING POINT IS 591 K.
(iii)IT DISSOLVES IN WATER TO FORM STRONGLY ALKALINE SOLUTION.
(iv)IT REACTS WITH CO2 TO FORM NA2CO3.
(v)ITS AQUEOUS SOLUTION IS CORROSIVE AND SOAPY TO TOUCH.
USES
*USED IN THE MANUFACTURE OF PAPER ,ARTIFICIAL SILK ETC.
*IN PETROLEUM REFINING.
*IN THE TEXTILE INDUSTRY.
*FOR THE PREPARATION OF PURE FATS AND OILS.
*AS A LABORATORY REAGENT.
SODIUM HYDROGENCARBONATE :(i) Is commonly known as baking soda because it decomposes on heating to generate bubbles of carbon dioxide. (ii) It is made by saturating a solution of sodium carbonate with carbon dioxide. Na2CO3 + H2O + CO2 -----> 2 NaHCO3 (iii) it is mild antiseptic for skin infections. USES:(i) used in fire extinguishers.
GROUP 2 ELEMENTS: ALKALINE EARTH METALS :-
(i) The group 2 elements comprise Beryllium, Magnesium, Calcium, Strontium, Barium and Radium. (ii) The first element Beryllium differs from trhe rest of the members and shows diagonal relationshipt to Aluminium.
ELECTRONIC CONFIGURATION :(i) These electrons have two electrons in s- orbital of the valence shell. (ii) Their general electronic configuration is ns2 . (iii) The compounds of these elements are also predominantly ionic.
IONIZATION ENTHALPIES :(i) The alkaline earth metals have low ionization enthalpies due to fairly large size of the atom. (ii) Since the atomic size increases down the group, their ionization enthaslpy decreases.
HYDRATION ENTHALPIES :(i) The hydration enthalpies of alkaline earth metal ions decrease with increase in ionic size down the group. (ii) So the compounds of Alkaline Earth Metals are more extensively hydrated than those of alkali metals.
CHEMICAL PROPERTIES :(a) Reactivity towards air and water :(b) (i) Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface. (ii) Calcium, Strontium and Barium are readily attacked by air to form the oxide and nitrite. (iii) They also react with water with increasing vigour even in cold to form hydroxides. (b) Reactivity towards the halogens:(i) All the alkaline earth metals combine with halogens at elavated temperatures forming their halides. M + X2 ----> MX2 (X= F,Cl, Br,I) (b) Reactivity towards the halogens:(i) All the alkaline earth metals combine with halogens at elavated temperatures forming their halides. M + X2 ----> MX2 (X= F,Cl, Br,I) (c) Reactivity towards hydrogen :(i) all the elements except Beryllium combine with hydrogen upon heating to form their hydrides. (d) Reducing Nature :(i) Alkaline earth metals are strong are reducing age(b) Reactivity towards the halogens:(i) All the alkaline earth metals combine with halogens at elavated temperatures forming their halides. M + X2
----> MX2 (X= F,Cl, Br,I)
(c) Reactivity towards hydrogen :(i) all the elements except Beryllium combine with hydrogen upon heating to form their hydrides. (d) Reducing Nature :(i) Alkaline earth metals are strong are reducing age
corresponding alkali metals. (iii) Beryllium has less negative value compared to other alkaline earth metals. (iv) Its reducing nature is due to large hydration energy. (e) Solution in liquid ammonia :(i) The alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions forming ammoniated ions. USES :(i) Beryllium is used in the manufacture of alloys. (ii) Copper- Beryllium alloys are used in the preparation of high strength springs. (iii) Magnesium forms alloys with aluminium, zinc, manganese and tin. (iv) Magnesium-aluminium alooys being light in mass are used in aircraft construction. (v) Calcium is used in the extraction of metals from oxides which are difficult to reduce with carbon. (vi) Radium salts are used in radiotherapy.
GENERAL CHARACTERISTICS (Compounds of Alkaline Earth Metals)
2+
1) Valence- M (Group II elements) 2) Compounds are predominantly ionic (less ionic than alkali metal compounds) - because of increased nuclear charge and smaller size. 2+ 2+ 3) Be and Mg oxides- more covalent (compared to Ca, Sr and Ba oxides) Oxides and Hydroxides 1) alkaline earth metals- burn in oxygen to form monoxide , having rock-salt structure [Exception- BeO; covalent in nature] 2) High enthalpies of formation 3) Very stable to heat 4) Oxides- basic in nature. React with water forming sparingly soluble hydroxides [Exception- BeO; amphoteric] 5) Solubility, thermal stability and basic character- increase with increasing atomic number.The basic character is of hydroxide is due to the low ionisation enthalpies . 6) Hydroxides- less basic and less stable (compared to alkali metal hydroxides) [Exception- Be(OH)2; Amphoteric]
Halides
1) ionic in nature [Exception- Beryllium halides; covalent and soluble in organic solvents] 2) Beryllium chloride- chain structure in solid state.
Vapour phase- chloro-bridged dimer ; dissociates into linear monomer at high temp. 3) Tendency to form halide hydrates decreases down the
group. Salts of Oxoacids a) Carbonates 1) Insoluble in water 2) Precipitated by Sodium or (NH4)2CO3 3) Solubility decreases as atomic number of the metal ion increases. 4) Decompose on heating to give CO2 and oxide. 5) The solubility of carbonates in water decreases down the family 6) Thermal stability increases with increasing cationic size. 7) Exception- Beryllium carbonate ; unstable, kept in atmosphere of CO2 b) Sulphates 1) White solids and stable to heat 2) BeSO4 and MgSO4 – soluble in water, as greater 2+ 2+ hydration enthalpies of Be and Mg overcome the lattice enthalpy.
3) Solubility decreases from CaSO4 to BaSO4
c) Nitrates 1) Formed by dissolution of carbonates in dil. HNO3 2) Magnesium nitrate- crystallizes with 6 H2O Barium nitrate- anhydrous 3) As size increases and hydration enthalpy
decreases, tendency to form hydrates decreases. 4) Decompose on heating to give the oxide.
EXERCISES
Question 10.6: Compare the alkali metals and alkaline earth metals with respect to (i) ionization enthalpy (ii) basicity of oxides and (iii) solubility of hydroxides.
Alkali metals (i)
Alkaline earth metals
Ionization enthalpy: (i) Ionization enthalpy: Smaller Lowest ionization atomic size and higher effective enthalpies because of nuclear charge. So, their first their large atomic sizes. ionization enthalpies are higher than that of alkali metals.
However, their second ionization enthalpy is less than the corresponding alkali metals.
(ii) Basicity of oxides: Very basic in nature, due to the highly electropositive nature of alkali metals.
(ii Basicity of oxides: Not as basic ) as those of alkali metals because alkaline earth metals are less electropositive than alkali metals.
(iii )
(ii Solubility of hydroxides: Less i) soluble than those of alkali metals, due to the high lattice energies of alkaline earth metals.
Solubility of hydroxides: More soluble than those of alkaline earth metals.
Question 10.20: The hydroxides and carbonates of sodium and potassium are easily soluble in water while the corresponding salts of magnesium and calcium are sparingly soluble in water. Explain. The atomic size of sodium and potassium is larger than that of magnesium and calcium. Thus, the lattice energies of carbonates and hydroxides formed by calcium and magnesium are much more than those of sodium and potassium. Hence, carbonates and hydroxides of sodium and potassium dissolve readily in water whereas those of calcium and magnesium are only sparingly soluble.
Question 10.32:
Which one of the alkaline earth metal carbonates is thermally the most stable? (a) MgCO3 (b) CaCO3 (c) SrCO3 (d) BaCO3cation present in the carbonate. The increasing order of the cationic size of the given alkaline earth metals is Mg < Ca < Sr < Ba Hence, the increasing order of the thermal stability of the given alkaline earth metal carbonates is MgCO 3 < CaCO3 < SrCO3 < BaCO3
Anomalous behaviour of beryllium Beryllium shows anomalous behavior compared to magnesium
and rest of the second group members. It shows diagonal relationship to aluminum
I.
Beryllium has exceptionally small atomic and ionic sizes
compared with other members of the group. Because of high ionization enthalpy and small size it forms compounds which are largely covalent and get easily hydrolyzed. 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 groups 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 aluminum The ionic radius of
is estimated to be 31 pm; the
charge/radius ratio is nearly the same s that of the
ion.
Hence beryllium resembles aluminum is some ways. some f the similarities are:
I.
Like the aluminum, beryllium is not readily attacked by acids because of the presence of an oxide film on the surface of the metal. Beryllium hydroxide dissolves in excess of alkali to give a
II.
beryllate ion,
just as aluminum hydroxide gives acuminate ion,
III.
.
The chlorides of both beryllium and aluminum have
bridged chloride structure in vapor phase. Both the chlorides are soluble in organic solvents and are strong Lewis acids. They are used as friedel craft catalysts. IV. Beryllium and aluminum ions have strong tendency to form complexes,
,
.
SOME IMPORTANT COMPOUNDS OF CALCIUM Important compounds: Calcium oxide or (quick lime), CaO Calcium hydroxide (slaked lime), Ca(OH)2 Calcium carbonate (limestone), CaCO3 Calcium sulphate (Plaster of Paris), CaSO 4 .1/2H2O Cement
CALCIUM OXIDE, CaO Prepared by heating limestone (CaCO 3) in 1070-1270 K.
CaCO3 CaO +CO2 Since the reaction is reversible, CO2 is removed as soon in order to shift the equilibrium in favour of products. CaO is white amorphous solid. Melting point is 2870 K.
Extremely stable and does not decompose. Absorbs moisture and CO2 on exposure to air.
CaO + H2O Ca(OH)2 CaO +CO2 CaCO3 Slaking of lime: Addition of limited amount of H 2O to break the lump of lime. Quicklime slaked with soda gives solid soda lime. CaO when heated in oxyhydrogen flame emits a brilliant light (lime light). CaO (basic oxide) combines with acidic oxide on heating at high temperature. CaO + SiO2 CaSiO3 (Calcium silicate) 6 CaO + P4O10 2 Ca(PO4)2 (Calcium phosphate)
USES OF CaO: Used in the manufacture of Calcium Carbide. Primary material for manufacturing cement. Used in the purification of sugar and in softening of hard water. Used in the manufacture of dye stuffs. Used in tanning industry and in drying of glasses and alcohol. CaO is used in the preparation of Ca(OH) 2
CALCIUM HYDROXIDE (slaked lime), Ca(OH) 2 Slaking of lime: Prepared by adding limited amount of
water to CaO. CaO + H2O Ca(OH)2
Ca(OH)2 is a white amorphous powder, sparingly
soluble in water. Suspension of slaked lime in water is called milk of lime. Clear aqueous solution of Ca(OH) 2 is lime water. When CO2 is passed through lime water it turns milky due to the formation of calcium carbonate. Ca(OH)2 + CO2 CaCO3 + H2O On passing excess of CO 2, the precipitate dissolves to form calcium hydrogen carbonate. CaCO3 + CO2+ H2O Ca(HCO3)2 The clear solution on heating again gives milkiness due to the decomposition of Ca(HCO3)2 to CaCO3 Milk of lime reacts with chlorine to form hypochlorite,
a constituent of bleaching powder. 2Ca(OH)2 + 2Cl2 CaCl2 + Ca(OCl)2 +2 H2O
USES of Ca(OH)2: Used in white wash due to its disinfectant. Manufacture of bleaching powder. Well known laboratory reagent for the detection of
CO2 Ca(OH)2 + CO2 CaCO3 + H2O Used in the manufacture of calcium hydrogen sulphate, Ca(HSO4)2,which is used in paper industry. Used in the preparation of mortar, a building material.
[Mortar is formed by adding H 2O to a mixture of sand and Ca(OH)2 to form a paste. As it dries, it hardens to form CaCO3] Ca(OH)2 + CO2 CaCO3 + H2O Used in glass making, in tanning industry, and for the purification of sugar.
Calcium carbonate (limestone), CaCO3 Occurrence: limestone, chalk, marble, dolomite. Preparation: by passing CO 2 through Ca(OH)2 or by the
addition of sodium carbonate to calcium chloride. Ca(OH)2 + CO2 CaCO3 + H2O CaCl2 + Na2CO3 CaCO3 + 2 NaCl The product obtained is precipitated chalk. Excess of CO2 would lead to the formation of water soluble calcium hydrogen carbonate. CaCO3 is white fluffy powder, almost insoluble in water. Limestone on calcinations gives CaO and CO 2. CaCO3 CaO +CO2 It reacts with dilute acid to liberate carbon-dioxide. CaCO3 + 2HCl CaCl2 + H2O+ CO2 CaCO3 + H2SO4 CaSO4 + H2O+ CO2
USES OF CaCO3 used in the preparation of cement and lime stone Calcium carbonate with magnesium carbonate is used as flux during smelting iron ores.
used as building material (marble) Precipated chalk is used in medicines and toothpastes
and in the manufacture of high quality paper. Used as an antacid, a constituent in chewing gum and as a filler in cosmetics.
Calcium sulphate (plaster of Paris), CaSO 4. ½ H2O Hemihydrates of calcium sulphate. Obtained by heating gypsum, CaSO 4.2H2O, to 393 K.
2 (CaSO4.2H2O) 2 (CaSO4). H2O +3 H2O Dead burnt plaster: Above 393 k, no water of crystallization is left and anhydrous sulphate, CaSO4 is formed. It is so called because it loses the property of setting with water.
USES of Plaster of Paris. On mixing with water it changes into plastic mass and solidifies due to rehydration. This is called setting of Plaster of Paris. Used for producing moulds for industries such as pottery, ceramics. For setting broken or fractured bones and sprain in the body. For making statues, models, and other decorative material. Employed in dentistry Question corner 1. What happens when
(i) (ii) (iii) Ans
Quicklime is heated with silica Chlorine reacts with slaked lime Calcium nitrate is heated? (i)
On heating quicklime with silica, calcium silicate is formed. CaO + SiO2 CaSiO3 (ii) The reaction of chlorine with slaked lime forms bleaching powder. 2Ca(OH)2 + 2Cl2 CaCl2 + Ca(OCl)2 +2 H2O (iii) On heating, calcium nitrate decomposes to give NO2 gas along with O2 gas. 2Ca(NO3)2 2CaO + 4NO2 +O2 2. Describe two important uses of quick lime. ( NCERT 10.18) 3. Describe the importance of the following: (NCERT 10.21) Limestone (i) Plaster of Paris (ii) 4. What is plaster of Paris? How is it prepared? 5. Give four industrial uses of lime and limestone. 6. What is the difference between quick lime, slaked lime, and lime water? 7. How is bleaching powder prepared? 8. Give the formula for dolomite. Ans: CaCO3. MgCO3 9. Gypsum on heating to 393 K gives ______. Ans: Plaster of Paris 10. What is the formula of bleaching powder and its chemically known term? Ans: Ca(OCl)Cl or CaCl2 + Ca(OCl)2 and is known as chlorohypochlorite as it is a mixed salt of HCl and HOCl.
CEMENT It is an important building material.It was introduced in 1824 by Joseph Aspdin.It is called as portland cement.
It is a product obtained by combining rich in lime. The average composition of portland cement : CaO= 50-60%;SiO2=20-25%;Al2O3=5-10%;MgO=23%;Fe2O3=1-2%; SO3=1-2% Quality cement the ratio of silica to alumina should be 2.5-4 and lime to the total oxides of silicon,aluminium and iron should be close to 2. When clay and lime are strongly heated together they fuse and react to form cement clinker and is mixed with 2-3%of gypsun (CaSO4.2H2O) to form cement.Thus, inghredients of portland cement are dicalcium silicate 26%,tricalcium silicate 51%,and tricalcium aluminate 11%.
Setting of Cement: 1.When the cement is mixed with the water it give a hard mass.This is due to the hydration of the molecules of the constituent. 2.The purpose of adding gypsum is to slow down the process of setting of cement so that it gets sufficiently hardened.
Uses of Cement: 1.It is next to iron and steel of comodity of national necessity for any country. 2.It is used in concrete and reinforced concrete, in plastering and in the construction of bridges,dams and buildings. REVISION EXERCISES 1.When alkali metal dissolves in liquid ammonia ,it gives different colours. Explain. a)The dilute solutions of alkali metals are blue.This is due to ammoniated electrons. +
-
M+( X+Y) NH3 →[ M(NH3)x] +[ e( NH3 )Y]
When conc. Increases ammoniated electron get bound by free electron and colour changes to bronze. 2.potassium carbonate cannot be prepared by solvay process.why ? a)potassium carbonate being more soluble than sodium carbonate cannot be precipitated whenCO2 passed through a conc. Soln . of KCl. 3.Lithium carbonate is decomposed at a lower temperature whereas sodium carbonate at a higher temperature.why? a) litium is very small in size polarises a large carbonate ion and is not stable to heat.
4.The hydroxides and carbonates of sodium and potassium are easily soluble in water while those of magnesium and calcium are less soluble. Explain. a)The lattice enthalpies of sodium and potassium are are lower due to the large size. Thats why they are easily soluble in water. 5.LiF
