Sarvan Kumarhttp://sarvankumar.wordpress.comMy name is Sarvan Kumar, I am post graduate in chemistry (MSc). besides teaching chemistry at various schools and coaching centres I have been giving home tuitions to students for 10 years. I have helped students score good marks in chemistry not only in board examination but also JEE, NEET, SAT, IGCSE and IB examinations. Nearly 90% of my students have scored more than 95% in their CBSE board examination. Moreover they have also secured a seat in prestigious engineering and medical college.
Since every student is different hence my unique style of teach
The homogeneous mixture of two components in which one component is called solute and another component is called solvent.
Types of solution on the basis of physical state of solvent and solute.
Nine types of solution.
Solid+Solid.
Solid+Liquid.
Solid+Gas
Liquid+Solid
Liquid+Liquid
Liquid+Gas
Gas+Solid
Gas+Liquid
Gas+Gas
Define the term concentration.
Amount of the solute present in given amount of solvent is called concentration.
Different types of concentration terms.
1.Molarity.
2.Molality
3.Mole fraction.
4.Mass percentage
5.Concentration in ppm.
Molarity-Number of moles of solute present in one litre of solution.
M=n/V
Where n= Number of moles of solute and V=Volume of solution in liter.
Molality= Number of moles of solute present in one Kg of solvent.
m=n/w
where n= Number of moles of solute and w=mass of solvent in Kg.
Define the term solubility.
The maximum amount of solute that can be dissolved into given amount of solvent at given temperature is called solubility of that solute.
Factors which effect the solubility.
1.Nature of the solvent and solute
2.Temperature.
3.Pressure.
Define the term dissolution.
The process of dissolving solid in liquid due to which concentration of solid increases in solution is called dissolution.
Define the term crystallization.
The process in which solid solute particles collide with each other and get separated out of solution is called crystallization.
Define the term saturated solution.
When rate of dissolution is equal to the rate of crystallization the solution is called saturated solution. In this solution no more solute can be dissolved at given temperature and pressure.
Define the term unsaturated solution.
In this solution more solute can be dissolved at same temperature and pressure.
Conditions of solubility of particular substance in particular solvent.
Like dissolves in like that is a polar solute can dissolved in polar solvent and non polar substance can be dissolved in nonpolar solvent.
Explain the effect of pressure on solubility of solid in a liquid.
Effect of pressure on solubility of solid in liquid is not significant because compressibility of solid and liquid is very low.
Explain the effect of temperature on solubility of solid in a liquid.
Solubility of solid in liquid is reversible process. If dissolution process is endothermic (Δsol H > 0), the solubility should increase with rise in temperature and if it is exothermic (Δsol H < 0) the solubility should decrease according to Le Chateliers Principle.
Explain the effect of pressure on solubility of gas in a liquid.
Solubility of gas in a liquid increases on increasing pressure. Since dissolution of gas in liquid is reversible process so increasing pressure more of gas must be dissolved in liquid to decrease the process so that equilibrium should maintained.
Explain the effect of temperature on solubility of gas in a liquid.
Dissolution process involves dynamic equilibrium and thus must follow Le Chatelier’s principle. As dissolution is an exothermic process, the solubility should decrease with increase of temperature.
Explain Henry’s law and its applications.
The amount of solute that can be dissolved in given amount of solvent at given temperature is directly proportional to the partial pressure of the gas above the liquid
p = KHx
KH is the Henry’s law constant.
X=mole fraction
Applications.
1.When CO2 is packed in soda water or In soft drinks it is packed under high pressure.
2.People suffer from a medical condition called anoxia at high altitude.
It is due to pressure of air at high altitude is low so solubility of oxygen in blood becomes low.
3. Scuba divers suffer from a medical condition called bends when they come out from deep sea water to the surface. It is due to when they come out towards surface pressure gradually decreases and bubble of nitrogen is formed in blood.
Why aquatic species are more comfortable in cold waters rather than in warm waters.
Solubillty of gas increases with decrease in temperature so that aquatic species are more comfortable in cold waters rather than in warm waters.
Define vapour pressure.
In a closed container when rate of evaporation is equal to the rate of condensation the pressure of a vapor above its liquid is called vapor pressure.
Which has higher vapour pressure in solvent and solution if a non volatile solute is added to the solvent.
Vapour pressure of solvent is higher than the vapor pressure of solution because no. of particles of volatile solvent decreases at the surface of liquid in case of solution.
Expain Raoult’s law.
For any solution the partial vapour pressure of each volatile component in the solution is directly proportional to its mole fraction.
Or
When non volatile solute (like solid solute) is added to the solvent vapour pressure of solution is directly proportional to the mole fraction of solvent.
What is colligative properties.
The properties of solution which don’t depend on the nature of the solute but depends on the amount of solute is called colligative properties.
Types of colligative properties.
1.Relative lowering of vapour pressure .
2.Elevation of boiling point.
3.Depression of freezing point.
4.Osmotic pressure.
Derive the equation for relative lowering of vapor pressurewhen non volatile solute is added to the volatile solvent.
According to Raoult’ law
Derive an equation for the calculation of total vapour pressure when two components in solution are volatile.
Why boiling point of solution is higher than that of solventwhen non volatile solute is added to the volatile solvent.
Since vapour pressure of solution is lower than that of solvent hence vapour pressure of solution will be equal to1atm at higher temperature.
Derive an equation for determination of elevation of boiling pointwhen non volatile solute is added to the volatile solvent.
Why freezing point of solution is lower than that of solventwhen non volatile solute is added to the volatile solvent.
Since vapour pressure of solution is lower than that of solvent hence vapour pressure of solution will be equal to vapour pressure of its solid state at lower temperature.
Derive an equation for determination of depression in freezing pointwhen non volatile solute is added to the volatile solvent.
Define osmosis.
The spontaneous phenomena of flowing solvent molecules from lower concentration solution to higher concentration solution is called osmosis.
What is the reason of osmosis.
Since vapor pressure of lower concentration is higher than higher concentration side.
Define osmotic pressure.
The extra pressure which is applied on the higher concentration solution side to prevent osmosis is called osmotic pressure.
How can we calculate osmotic pressure.
A raw mango is placed in concentrated salt solution, it shrivels into pickle why.
Raw mango loses water via osmosis.
Why wilted flowers revive when placed in fresh water?
Osmosis occurs due to vapor pressure difference and water molecules flow from fresh water to flower.
Carrot that has become limp because of water loss into the atmosphere can be placed into the water making it firm once again.
Osmosis occurs due to vapor pressure difference and water molecules flow from fresh water to carrot.
Define isotonic solution.
Two solution having same osmotic pressure at constant temperature is called isotonic solution.
Explain Hypertonic and Hypotonic solution.
Two solutions which are separated by SPM in which one solution has osmotic pressure lower than the other, the lower osmotic pressure solution is called hypotonic solution and higher osmotic pressure solution is called hypertonic solution,
Why sodium chloride solution, called normal saline solution having concentration of 0.9% (mass/mass) is safe to inject intravenously.
The osmotic pressure of the fluid inside the blood cell is equal to 0.9% (mass/mass)
When blood cells placed in a solution containing more than 0.9% (mass/volume) sodium chloride, it would shrink why?
Because water molecules flow out from blood cells to sodium chloride solution via osmosis .
When blood cells placed in a solution containing lower than 0.9% (mass/volume) sodium chloride, it would swell why?
Because water molecules get inside to blood cells from sodium chloride solution via osmosis.
Explain reverse osmosis and application of it.
When extra pressure which is applied to the higher concentration solution is higher than the osmotic pressur the process is reversed. Now water molecules passes from higher concentration solution to lower concentration this process is called reverse osmosis.
We can apply reverse osmosis to reduce the salt level from salty water.
Define ideal and non ideal solution.
Ideal solution–Solutions which obey Raoult’s law over the entire range of concentration are known as ideal solutions.
Solutions which don’t obey Raoult’s law over the entire range of concentration are known non ideal solutions.
What are the features of ideal solution?
3.If the two components of solution are A and B than intermolecular force of attraction between A-B is nearly equal to the A-A and B-B.
4.Examples.n-hexane and n-heptane, bromoethane and chloroethne benzene and toluene .
Types of non ideal solution.
Non Ideal solution have two types.
1.Non Ideal solution having +ve deviation.
3 .If the two components of solution is A and B than intermolecular force of attraction between A-B is lesser than the A-A and B-B.
4. Example.Mixture of alcohol and acetone
1.Non Ideal solution having -ve deviation
3. If the two components of solution is A and B than intermolecular force of attraction between A-B is higher than the A-A and B-B.
4. Example.Mixture of chloroform and acetone.
Define Azeotropic mixture.
The mixture of the two components which boil at same temperature. These type of mixture can’t be separated by fractional distillation.
Types of azeotropic mixture.
a.Maximum boiling azeotropic mixture.
Mixture of the two components which boilng point is higher than the either of the two components. This type of mixture shows large –ve deviations from Raoult’s law. Example. 68% nitric acid and 32% water by mass.
b.Minimum boiling azeotropic mixture.
Mixture of the two components which boilng point is lower than the either of the two components.This type of mixture shows large +ve deviations from Raoult’s law.
Example 95% by volume of ethanol in water.
What do you mean by Van’t Hoff factor.
Factor which makes experimental and theoretical value of colligative property equal.
Q.1.Group number 13-18 elements are known as p-block elements why?
Ans.The last electron of these elements goes into p-orbitals.
Q.2.What is the valance shell electronic configuration of p-block elements.
Ans.ns2np1-6
Q.3.1st member of group 13-17 elements show different behavior from rest of the elements why?
Ans.It is due to small size and high electronegativity of 1st member of group 13-17 elements.
Q.4.1st member of group 13-17 elements can show maximum covalency four but other elements can expand their covalency greater than four why?
Or
While boron forms only [BF4]–, aluminium gives [AlF6]3– ion why?
Ans.1st member of group 13-17 elements has only four outermost orbitals one ns and three np orbitals hence they can show maximum covalency four but other elements have vacant d orbitals hence hence they can expand their covalency greater than four.
Q.5.The first member of group 13-17 elements can form multiple bonds to itself ( e.g., C=C, C≡C, N≡N) and to other elements like C=O, C=N, C ≡N, N=O). But heavier member don’t why?
Ans.It is due to small size and high electronegativity of 1st member of group 13-17 elements but p-orbital of heavier member is so large and diffuse to effective overlapping.
Q.6.What is inert pair effect?
Ans.When 2 electrons of outermost ns orbital are reluctant towards bonding and elements show oxidation state two unit less than the group oxidation state ,this effect is called inert pair effect.
Group 13 elements :boron family
B,Al,Ga,In, Tl
Q.1.Boron has unusually high melting point why?
Ans.Due to very strong crystalline lattice arises from its small size, boron has unusually high melting point.
Q.2. Gallium can exist in liquid state during summer why?
Ans. It is due to very low melting point of Gallium.
Q.3.Boron forms only covalent compound why?
Ans.It is due to very high first three ionization enthalpies of boron it is unable to form +3 ions hence it forms mainly covalent compound.
Q.4.Aluminium is electropositive element and can form ionic compound why?
Ans.It is due to larger size and lower ionization enthalpies of Al can form Al3+ion and can form ionic compound.
Q.5.What are the oxidation states of group 13 elements with or without inert pair effect.
Ans.Without inert pair effect +3
With inert pair effect +1
Q.6.Stability of +3 oxidation state decreases top to bottom and +1 increases top to bottom.
Ans .As we go down the group electronegativity decreases hence energy required to unpair two ns electron is not compensated by energy release in formation of two extra bond.
Q.7.Which is more stable in BCl3 and TlCl3.
Ans.Since stability of +3 oxidation stability decreases top to bottom as energy required to unpair two ns electron is not compensated by energy release in formation of two extra bonds hence the BCl3 is more stable than TlCl3.
Q.8.Atomic radius of Ga is less than that of Al why?.
Ans. Presence of additional 10 d-electrons in Ga which produce poor screening effect and increased effective nuclear charge decreases the atomic size.
Q.9.Ionisation enthalpy of Ga is higher than Al why?
Ans. Atomic size of Ga is lower than Al hence Ionisation enthalpy of Ga is higher than Al
Q.10.Ionisation enthalpy of Tl is higher than In why.
Ans Higher ionization energy compensate the high effective nuclear charge arises due to poor shielding effect of d and f electrons in Tl.
Q.11.The order of ionisation enthalpies of group 13 elements increases in following order why ?
Ans. As the elements losing its electron +ve charge increases thus size decreases hence ionization enthalpy increases.
Q.12.Electronegativity of alkali metals first decreases from B and Al and then increases slightly why?
Ans.It is due to irregularities in atomic size and presence of d and f electrons in heavier member.
Q.13.In trivalent state compounds of 13 elements are Lewis acid why?
Ans. In trivalent state compounds of 13 elements are electron deficient hence they are Lewis acid example BF3, BCl3 etc.
Q.14. Arrange the Lewis basic strength of group 13 trivalent compound.
Ans. The tendency to behave as Lewis acid decreases with the increase in the size down the group. BI3< BBr3< BCl3< BF3
Q.15. Aluminum is very less reactive towards oxygen why?
Ans. Aluminum forms a thin oxide layer on the surface which protects the metal from further attack.
Q.16.AlCl3 exists in dimer why?
Ans. Al atom in molecule is electron deficient due to presence of only six electrons on its valence shell.
Q.17.Arrange acidic property of group 13 oxides.
Ans.Acidic property of group 13 oxide decreases down the group.
Boron trioxide – acidic
Aluminum and gallium oxides – amphoteric
Indium and thallium – basic
Q.18.White fumes appear around the bottle of anhydrous aluminum chloride give reason.
Ans.Anhydrous aluminum chloride liberates HCl gas on hydrolysis with atmospheric moisture moist HCl appears white in colour.
Q.19.Boron is unable to form BF63-why?
Ans. Boron has only four outermost orbital one ns and three np orbitals hence it cannot expand covalency greater than 4.
Q.20.What is the heating effect of borax?
Q.21.What is borax bead test?
Ans.This test is used for identify transition elements. When borax is converted into metaborate of a transition elements it gives characteristic color.
Q.22.The structure of boric acid H3BO3.
Ans .It has a layer structure in which planar BO3 units are joined by hydrogen bonds.
Q.23.What is the heating effect of boric acid?
Ans.
Q.24.Boric acid is weak monobasic acid why?
Ans.Boric acid accept OH– from water in turn release H+ ion, thus orthoboric acid behaves like monobasic acid.
Q.25.What is inorganic benzene
Ans .B3N3H6 is known as inorganic benzene.
Q.26.Describe the structure of diborane.
Ans. Diborane contains four terminal B-H bonds and two bridge (B-H-B) bonds .The four terminal B-H bonds lie in one plane and two centre-two electron bonds bonds while the two bridge (B-H-B) bonds present above and below this plane and three centre–two electron bonds.
Q.27.What do you mean by banana bond in diborane.
Ans. Two bridge (B-H-B) bonds present above and below this plane and three centre–two electron bonds are known as banana bond.
Q.28.What are the different uses of B and Al.
Ans.1.Boron fibers in making bullet proof vest.
2 Metal borides are in nuclear industry.
3.Borax and boric acid are used in making heat resistant glasses.
4. Aq. orthoboric acid as mild antiseptic.
5. In packing
6. Utensil making,
7.Construction, aeroplane and transportation industry.
Group 14 elements: Carbon family
C,Si,Ge,Sn,Pb
Q.30.What is the common oxidation state of group 14 elements?
Ans. +4 and +2 are the common oxidation states of these elements. Carbon also show negative oxidation state.
Q.31.Compounds in +4 oxidation state are generally covalent in nature why?
Ans. Since the sum of the first four ionization enthalpies of group 14 elements is very high hence these compounds are mainly covalent.
Q.32.What are the oxidation state of group 14 elements without inert pair effect.
Ans +2=With inert pair effect
+4=without inert pair effect
Q.33.Stability of +2 oxidation state increases top to bottom and +4 decreases top to bottom why?
Or
PbX2 is more stable than PbX4 why?
Or
GeX4 is more stable than GeX2
Or
In heavier members of group 14 elements the tendency to show +2 oxidation state increases in the sequence Ge<Sn<Pb.why?
Or
Carbon and silicon mostly show +4 oxidation state why?
Ans. As we go down the group electronegativity decreases hence energy required to unpair two ns electrons is not compensated by energy release in formation of two extra bond.Thus stability of+2 oxidation stae increases and that of +4 decreases.
Q.34.Lead compound in +4 state are strong oxidising agents why?
Ans. +4 oxidation state of Pb is unstable due to inert pair effect hence tend to decrease this oxidation state into stable +2 oxidation state hence works as oxidizing agent.
Q.35. Sn in +2 state is a reducing agent why?
Ans.Sn can show both +2 and +4 oxidation state hence in +2 oxidation state it works as reducing agent.
Q.36.Maximum covalency of carbon is 4 but other elements can extend their covalency greater than 4 why?
Or
The species like, SiF62–, [GeCl6]2–,[Sn(OH)6]2- of heavier group 14 elements known but carbon is unable to form these types of compound why?
Or
Halides of heaviour elements(SiCl4) undergo hydrolysis but CCl4 not why?
Ans. Carbon has only four outermost orbitals one 2s and three 2p orbitals hence they can show maximum covalency four but other elements have vacant d orbitals hence hence they can expand their covalency greater than four.
Q.37.Give the example of acidic,amphoteric and neutral oxide of group14 elements
Ans. CO2, SiO2 and GeO2 –acidic
SnO2 and PbO2 – amphoteric SnO PbO
CO –neutral
Q.38.PbI4 does not exist why?
Ans. Energy released during formation of Pb—I bond cannot unpair 6s2 electrons and excite of them to higher orbital to have four unpaired electrons around lead atom.
Q.39.[SiF6]2– is known whereas [SiCl6]2– not why?
Ans.(i) Due to larger size of Cl– ions they cannot be accommodated around Si4+
(ii)Due to lower electronegativity of Cl– Interaction between lone pair of chloride ion and Si4+ is not very strong.
Q.40.Carbon show different behaviour than rest of the members of its Group why?
Ans. It is due to following reasons of carbon atoms.
Smaller size
High electronegativity
Higher ionisation enthalpy
Unavailability of d orbitals
Q.41.What is catenation property.
Ans.Tendency of atoms to link with one another through covalent bonds to form chains and rings called catenation.
Q.42.Catenation property of carbon is very high why?
Ans.This is due to strong C—C bond catenation property of carbon is very high.
Q.43.Tendency to show catenation property by group 14 elements decreases down the group why?
Ans.The size and electronegativity decreases down the group hence tendency to show catenation property by group 14 elements decreases down the group. Lead does not show catenation. C > > Si > Ge = Sn.
Q.44.Carbon is able to show allotropic forms why?
Ans.Due to high catenation property and ability to form bond by carbon atoms it show allotrope.
Q.45.Name the all crystalline allotropes of carbon.
Ans.(i)Diamond
(ii)Graphite
(iii) Fullerenes
Q.46.Describe the structure of Diamond.
Ans. Diamond has crystalline structure in which each carbon is sp3 hybridised and linked to four other carbon in tetrahedral manner.
Q.47.Diamond is used as an abrasive for sharpening hard tools why?
Or
Diamond is covalent, yet it has high melting point why?
Ans.Diamond is a hardest substance on the earth because bond dissociation enthalpy of extended covalent bonding between carbon atoms is very high hence use as abrasive.
Q.48.What are the other uses of diamond.
Ans.(i) In jewelry
(ii) In making dyes
(ii) In the manufacture of tungsten filaments for electric light bulbs.
Q.49.Describe the structure of Graphite.
Ans (i) Crystalline layered structure.
(ii) Each layer is joined with each other by van der Waals forces.
(iii) Each layer is composed of planar hexagonal rings of carbon atoms.
(iv)Each carbon is sp2 hybridised in hexagonal ring.
(v) Fourth electrons of carbon are delocalized over the whole sheet.
Q.50.Graphite conducts electricity why?
Ans. Since each carbon is sp2 hybridised in hexagonal ring hence fourth electrons of carbon are delocalized over the whole sheet thus graphite is good conductor.
Q.51.Graphite is very soft and slippery why?
Or
Graphite is used as a dry lubricant in machines running at high temperature, where oil cannot be used as a lubricant.
Ans.Due to layered structure graphite cleaves easily between the layers and, therefore it is very soft and slippery.
Q.52.How is fullerene made?
Ans.By heating graphite in an electric arc in the presence of inert gases.
Q.53.Describe the structure of Fullerenes.
Ans. Fullerene contains mainly C60 molecule with traces of C70.
Q.54.Describe the structure of C60 molecule(Buckminsterfullerene)
Ans.(i) It contains 20 six- membered rings and 12 five membered rings.
(ii) A six membered ring is fused with six or five membered rings but a five membered ring can only fuse with six membered rings.
(iii)All the carbon atoms are sp2 hybridised
(iv) The fourth electron of each carbon is delocalised in molecular orbitals.
Q.55.C60 molecule is also called Buckminsterfullerene why?
Ans. C60 molecule has a shape like soccer ball and called Buckminsterfullerene.
Q.56.Which Property of Buckminsterfullerene give its aromatic character?
Ans. The fourth electron of each carbon is delocalised in molecular orbitals.
Q.57.Fullerenes are the only pure form of carbon why?
Ans Fullerenes have smooth structure without having ‘dangling’ bonds.
Q.58. of graphite is taken as zero why?
Ans. Graphite is thermodynamically most stable allotrope of carbon and, therefore of graphite is taken as zero.
Q.59.Give some examples of impure form of carbon.
Ans. Carbon black, coke, and charcoal.
Q.60.How is Carbon black obtained?
Ans. Carbon black is obtained by burning hydrocarbons in a limited supply of air.
Q.61.How is charcoal obtained?
Ans.Charcoal is obtained by heating wood in the absence of air.
Q.62 How is coke obtained?
Ans. Coke is obtained by heating coal in the absence of air.
Q.63.What are the uses of carbon?
Ans.(i) Composites which contain graphite are used in products such as tennis rackets, fishing rods, aircrafts and canoes.
(ii) Graphite is used as electrode in batteries and industrial electrolysis due to its good conducting nature.
(iii) Charcoal is used in gas masks to adsorb poisonous gases due to its porous nature.
(iv) Carbon black is used as black pigment in black ink and as filler in automobile tyres.
(v) Coke is used as a fuel and largely as a reducing agent in metallurgy.
(vi) Diamond in jewellery.
Q.64.What is producer gas?
Ans.Mixture of CO and N2 known as producer gas.
Q.65.What is water gas/synthesis gas/syn gas?
Ans.Mixture of CO and H2 known as producer gas.
Q.66.What are metal carbonyls?
Ans.CO contains a lone pair at carbon and acts as donor atom to the metal and forms metal carbonyl.
Q.67.CO is highly poisonous why?
Ans.CO forms complex with hemoglobin which is more stable than the oxygen-hemoglobin complex prevents hemoglobin in the red blood corpuscles from carrying oxygen round the body and ultimately death.
Q.68.Name the buffer system which helps blood to maintain pH.
Ans.H2CO3/HCO3–
Q.69.What is photosynthesis?
Ans.The process by which green plants convert atmospheric CO2 into carbohydrates called photosynthesis.
Q.70.What is green house effect?
Ans. Increasing amount of CO2 in atmosphere produce during combustion of fossil fuels and decomposition of limestone for cement raise the temperature of the atmosphere called green house effect.
Q.71.What is dry ice?
Ans.Solid CO2 is known as dry ice. Dry ice is used as a refrigerant for ice-cream and frozen food.
Q.72. What are the uses of gaseous CO2?
Ans (i)As and non-supporter of combustion it is used as fire extinguisher.
(ii) In soft drinks.
Q.73.Draw the resonating structure of CO2
Q.74 What is silica?
Ans.Silicon dioxide(SiO2) is known as silica. Example Quartz, cristobalite and tridymite are crystalline forms of silica.
Q.75.Describe the structure of silica.
Ans.Silicon dioxide is a covalent, three-dimensional network solid in which each silicon atom is covalently bonded in a tetrahedral manner to four oxygen atoms.
Q.76.Silica in its normal form is almost nonreactive why?
Ans. It is due to very high Si—O bond enthalpy.
Q.77.What are the different uses of silica?
Ans (i) Quartz is used as a piezoelectric material
(ii) Silica gel is used as a drying agent.
(iii) Kieselghur, an amorphous form of silica is used in filtration plants.
Q.78.What are silicones ?
Ans.silicones are organosilicon polymers formed by –(R2SiO)-as repeated unit.
Q.79.What are the uses of silicones?
Ans. Silicones are used as sealant, greases, electrical insulators and for water proofing of
fabrics.
(ii) In surgical and cosmetic plants.
Q80 What are silicates?
Ans.The basic structural unit of silicates is SiO4 in which silicon atom is bonded to four oxygen atoms in tetrahedron manner.Example feldspar, zeolites, mica and asbestos.
Q81.Give the two examples of man made silicate.
Ans.Glass and cement.
Q82 What are Zeolites?
Ans. If some Aluminum atoms replace few silicon atoms in three-dimensional network of silicon dioxide the structure formed is aluminosilicate a types of zeolite.
Q 83.What are different uses of zeolites?
Ans.(i)ZSM-5 a zeolite used in used to convert alcohols directly into gasoline.
(ii)Hydrated zeolites are used for softening of “hard” water
Why are group 1 and group 2 elements called s-block elements?
The last electron of these elements enters the outermost s-orbital that is why they are called s-block elements.
Group1 elements are collectively known as alkali metals why?
They form hydroxide in water,which are alkaline in nature ,that is why they are called alkali metals.
Group 2 elements are collectively known as alkaline earth metals why?
Because their oxides and hydroxides are alkaline in nature and these metal oxides are found in the earth’s crust.
What is general electronic configuration of s-block elements?
ns1-2
GROUP 1 ELEMENTS ( ALKALI METALS)
Li, Na, K,Rb,Cs,Fr
Q.1.Alkali metals are normally kept in kerosene oil why?
Ans.Alkali metals are highly reactive towards air and water hence they are normally kept in kerosene oil.
Q.2.Alkali metals are never found free in nature why?
Ans. Due to very large size and presence of only one electron in outermost shell they easily lose their valence electron and from M+ ion.
Q.3.The atomic and ionic radii of alkali metals increase on moving down why?
Ans. As moving down the group number of shell increases, that is why atomic and ionic radii of alkali metals increase top to bottom.
Li<Na<K<Rb<Cs
Q.4.The ionization enthalpies of the alkali metals are considerably low and decrease down the group from Li to Cs.
Ans.Alkali metals are very large hence ionization enthalpies of these metals are very low. Since atomic size increases down the group thus ionization enthalpy decreases down the group.
Q.5.The hydration enthalpies of alkali metal ions decrease with increase in ionic sizes why?
Ans.Atomic size increases down the group hence hydration enthalpy decreases down the group.
Li+> Na+ > K+ > Rb+ > Cs+
Q.5. Li+ has maximum degree of hydration why?
Or
Why are lithium salts commonly hydrated and those of the other alkali ions usually anhydrous?
Ans. Li+ is smallest in size thus lithium salts are mostly hydrated example LiCl.2H2O
Q.6.Alkali metals have very low density why?
Ans. Alkali metals are very large hence they have very low density.
Q.7.Why does density of alkali metals increases down the group Why ?
Ans. Because mass increases down the group.
Q.8.Potassium is lighter than sodium why ?
Ans. In case of K electron filling takes place in 4s without 3d thus volume increases so much and density decrease.
Q.9.The melting and boiling points of the alkali metals are low why?
Ans.Weak metallic bonding in alkali metal crystal due to the presence of only a single valence electron in them.
Q.10.Alkali metals give flame test why?
Ans.Outermost orbital electron is excited to higher energy level .When the excited electron comes back to the ground state, there is emission of radiation in the visible region. Therefore they can be detected by the respective flame tests.
Q.11. Cs and K are used in photoelectric cell why?
Ans.The light energy absorbed by these metals and lose electron because outermost electron is loosely bonded with nucleus.
Q.12.Alkali metals are very reactive why?
Ans. Due to very large size and low ionization enthalpy alkali metals are highly reactive.
Q.13.Alkali metals tarnish in air why?
Ans. Alkali metals form oxide with air which further reacts with water and form hydroxide.
Q.14.What is the nature of oxide of alkali metals.
Ans. Lithium forms monoxide, sodium forms peroxide and K, Rb and Cs forms superoxide
4 Li +O →2Li2O (oxide)
2Na +O2 →Na2O2 (peroxide)
M+O2 →MO2 (superoxide)
(M = K, Rb, Cs)
Q.15.Lithium reacts directly with nitrogen of air to form the nitride, Li3N why?
Ans.It is due to very small size of Li it shows anomalous behavior.
Q.16.What is the oxidation state of K in KO2?
Ans +1
Q.17.Li is least reactive with water but other elements react explosively with water why?
Ans. It is due to small size and very hydration energy of lithium it is least reactive with water.
Q.18.All the alkali metal hydrides have high melting points why?
Ans. Alkali metals are ionic solid that is why they have high melting point.
Q.19.Lithium salts are mainly covalent why?
Ans. Li+ is small in size hence distortion of electron cloud(polarization power) is maximum hence its salt is covalent.
Q.20.In all lithium salts LiI is most covalent and why?
Ans. I– having large size thus polarization power is maximum.
Q.21.Lithium is strongest reducing agent why?
Ans. EO value depends on the following factors.
M (s) → M(g) (Sublimation enthalpy)
M(g) → M(g)+ (Ionization enthalpy)
M+(g) + H2O → M+(aq) Hydration enthalpy
Due to small size of lithium ,it has the highest hydration enthalpy resulting low negative Eo value and its high reducing power.
Q.22.Alkali metals give blue color in liquid ammonia.
or
Alkali metals are paramagnetic in liquid ammonia why?
Ans. M +(x + y) NH3 →[M(NH3 )x ]+ + [e(NH3 )y ]−
Due to presence of ammoniated electron alkali metals give blue color in liquid ammonia solution.
Q.23 Alkali metals in liquid ammonia liberate hydrogen on standing give reaction.
Ans. M+ (am) + e− + NH3 (1)→MNH2(am) +½H2(g)
Q.24.What are the different uses of lithium alloys?
Ans. Li+Pb= ‘White metal’ bearings for motor engines,
Li+Al = Aircraft parts
Li+Mg= Armour plates which is used in thermonuclear reactions.
Q.25.What are the uses of other alkali metal.
Ans 1.Liquid sodium metal is used as a coolant in fast breeder nuclear reactors.
2. Na/Pb alloy needed to make PbEt4 and PbMe4 which is used as anti knock additives to petrol earlier.
3. Potassium hydroxide is used in manufacturing of soft soap.
4. Potassium chloride is used as a fertilizer.
5. Cesium is used in photoelectric cells.
Q.26.What is the reason behind the increasing stability of the peroxide or superoxide, as the size of the metal ion increases ?
Ans.It is due to the stabilisation of large anions by larger cations through lattice energy effects.
Q.27. Why is KO2 paramagnetic?
Ans. An unpaired electron is present antibonding π*2p orbital of O2-
Q.28.The alkali metal halides, MX, (X=F,Cl,Br,I) are all high melting crystalline solids Why?
Ans. Alkali metals halides are ionic compound hence have very high melting point.
Q.29.The melting and boiling points of alkali metals decreases in following order fluoride > chloride > bromide> iodide why?
Ans.Metallic bond strength decreases as the size of halide ion increases.
Q.30. All alkali metal halides are soluble in water why.
Ans. Due to weak metallic bonding in alkali metal lattice enthalpy is lower than hydration enthalpy.
Q.31.Solubility of LiF in water is low why?
Ans.High lattice enthalpy of LiF due to small size of F.
Q.32.Solubility of CsI in water is low why?
Ans.Low hydration enthalpy of CsI due to large size of Cs.
Q.33.LiCl, LiBr,and LiI are not only soluble in water but also soluble in organic solvent like acetone,ethanol etc,but LiF is almost insoluble in water.
Ans Difference in lattice enthalpy and hydration enthalpy of other litium halides is higher than LiF, hence lithium halides are soluble in water. Due to larger size of Cl,Br and I their lithium halides have higher covalent character ,hence they are soluble in organic solvent also.
Q.34.The mobilities of the alkali metal ions in aqueous solution are Li+ < Na+ < K+ < Rb+ < Cs+
Ans.As the size of alkali metal ions increases hydration extent of these ions decreases thus mass of hydration ion decreases thus mobility increases.
Q.35.LiI is more soluble in KI in ethanol why?
Ans.LiI is more covalent than than KI due to high polarization power of Li in compare with K.Thus LiI is more soluble in KI in ethanol.
Q.36.Sodium is found to be more useful than potassium why?
Or
What are the biological function of Na and K.
Ans.Na+ ions are found in the blood plasma and also in interstitial fluid which surrounds the cell thus help in variety of biological function.
1.Transmission of nerve signal.
2.Regulate the flow of water across cell membrane.
3.Transport sugar and amino acids into cell.
K+ ions are found inside cell and help in following biological functions.
1.They activate many enzymes
2.Participate in the oxidation of glucose to produce ATP.
From above observation we can say that Sodium is found to be more useful than potassium.
Q.37.Alkali metals are prepared by electrolysis of their fused chlorides why?
Ans.In aq alkali metals chlorides H+ ions are present along with alkali metals ions,and Eo value of H+ is higher than alkali metal ions hence H2 instead of the alkali metals produced at cathode.
Q.38.A solution of Na2CO3 is alkaline why ?
Ans. In aq solution forms NaOH which is strong base.
CO32- + H2O → HCO3– +OH–
Q.39.Define oxoacids.
Ans.Oxo-acids are those in which the acidic proton is on a hydroxyl group with an oxo group attached to the same atom e.g., carbonic acid.
Q.40.The stability of the carbonates and hydorgencarbonates of alkali metal increases from top to bottom in group why?
Ans.As the size of alkali metal ions increases the polarization power of these ions to CO32- and HCO3– decreases hence thermal stability increases.
Q.41.Lithium carbonate is not so stable to heat why?
Or
Why is Li2CO3 decomposed at a lower temperature whereas Na2CO3 at higher
temperature?
Ans. smaller Li+ ions polarize CO32 most hence it decomposes into Li2O and CO2.
Q.42. Why does Li show anomalous behavior?
Ans It is due to
(i) Exceptionally small size of its atom
(ii) High polarising power
(iii)High charge/radius ratio.
Q.43.List some anomalous behavior of Li.
Ans. (i) m.p. and b.p.of Li are higher than the other alkali metals.
(ii) Lithium is least reactive but the strongest reducing agent among all the alkali metals.
(iii) Li forms monoxide, Li2O
(iv)Li forms nitride Li3N unlike other alkali metals.
(v) LiCl is more hydrated in compare with other alkali metals.
(vi) Solid Lithium hydrogencarbonate doesn’t exist while all other elements form solid hydrogencarbonates.
(vii) Lithium unlike other alkali metals forms covalent compound.
Q.44.What is the reason of diagonal relationship between Li and Mg.
Ans Due to similar charge/radius ratio they show similar behavior.
Q.45.What are the points of similarities between Li and Mg.
Ans (i)Both lithium and magnesium are harder than other elements in the respective groups.
(ii) Both Lithium and magnesium react slowly with water.
(iii) Both form nitride Li3N and Mg3N2, by direct combination with nitrogen.
(iv) Both form monoxide with oxygen Li2O and MgO.
(v) The carbonates of lithium and magnesium decompose easily on heating to form the oxides and CO2.
(vi) Solid hydrogencarbonates are not formed by lithium and magnesium.
(vii) Both LiCl and MgCl2 are soluble in ethanol.
(viii) Both LiCl and MgCl2 are deliquescent and crystallise from aqueous solution as hydrates, LiCl·2H2O and MgCl2·8H2O.
Q.46.Potassium carbonate cannot be manufactured by Solvay process why?
Ans.Potassium hydrogencarbonate is too soluble to be precipitated by the addition of ammonium hydrogencarbonate to a saturated solution of potassium chloride.
Q.47.What is the heating effect of washing soda (Na2CO3.10H2O)?
Q.48.What are the different uses of washing soda .
Ans.1 Water Softening,Laundring process.
2 In paper ,Textiles and Paint industries.
3 To manufacture soap,glass and borax.
Q.49. What are the uses of NaCl?
Ans.1Common salt for domestic purpose.
2 To prepare Na2O2.MaOH etc.
Q.50.What are the uses of NaOH (Caustic Soda)?
Ans. (i) To manufacture soap paper, artificial silk etc
(ii) In petroleum refining
(iii) In the purification of bauxite,
(iv) As a laboratory reagent.
Q.50.What are the uses of Sodium hydrogencarbonate (baking soda)
Ans (i) As antiseptic.
(ii) It is used in fire extinguishers.
Q.51.What are the common physical and chemical properties of alkali metals?
Ans.Physical properties.
(i) Crystalline solid
(ii) Silvery white, soft and light metals.
(iii) Low density
(iv) Low melting and boiling point
(v) Give flame test
Chemical properties
(i) Very highly reactive
(ii)+1 is common oxidation state
(iii) Form monoxide ,peroxide ,super oxide with oxygen
(iv) Form strong base on water
(v) Form halides with halogen
(vi) All are very strong reducing agent.
(vii) Their ammonical solution is deep blue and paramagnetic
Q.52.What is the oxidation state of Na in Na2O2.
Ans . (+1) In peroxide oxidation state of oxygen is -1 hence
2(x) -2=0 thus x=1
Q.53.Sodium is less reactive than potassium why?
Ans.Lower ionization enthalpy of K makes it more reactive than sodium.
Q.54.Alkali and alkaline earth metals can not be obtained by chemical
reduction methods why?
Ans. Alkali and alkaline earth metals are themselves strong reducing agent due to low ionization enthalpies thus reducing agents better than alkali metals are not available.
Q.55.Why are K and Cs rather than Li used in photoelectric cell?
Ans.Atomic size of K and Cs is much larger than Li thus light can easily emits electron from K and Cs.
Q.56.Arrange thermal stability of carbonate and hydrogencarbonate of alkali metals in increasing order.
Ans.Li2CO3 <Na2CO3< K2CO3< Rb2CO3< Cs2CO3
LiHCO3 <NaHCO3< KHCO3< RbHCO3< CsHCO3
GROUP 2 ELEMENTS ( ALKALINE EARTH METALS)
Be,Mg,Ca,Sr,Ba,Ra
Q.1.The atomic and ionic radii of alkaline earth metals increase on moving down why?
Ans As moving down the group number of shell increases, that is why atomic and ionic radii of alkaline earth metals increases top to bottom.
Be<Mg<Ca<Sr<Ba
Q.2.The first ionization enthalpy of the alkali metals is lower than corresponding alkaline earth metal why?
Ans.Alkali metals are larger than alkaline earth metals hence ionization enthalpies of these metals are lower than corresponding alkaline earth metals.
Q.3.The 2nd ionization enthalpy of the alkaline earth metals are lower than corresponding alkali metal why?
Ans.After removing of one electron from alkali metals they are converted into stable inert gas configuration hence removing of second electron becomes difficult.
Q.4.The hydration enthalpies of alkaline metal ions decrease with increase in ionic sizes why?
Ans.Atomic size increases down the group hence hydration enthalpy decreases down the group.
Be2+> Mg2+ > Ca2+ > Sr2+ > Ba2+
Q.5.Hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions.?
Ans. Ionic size of alkaline earth metals are lower than alkali metal ions hence the hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions.
Q.6.The alkaline earth metals are harder than the alkali metals why?
or
The melting and boiling points of these metals are higher than the corresponding alkali metals
Ans. Inter atom metallic bonding is stronger in case of alkaline earth metals due to their lower size.
Q.7.Strontium and barium impart characteristic colours to the flame why?
Ans. Flame excites electrons to higher energy levels and when they come back to the ground state, energy is emitted in the form of visible light.
Q.8.Beryllium and magnesium do not give flame test why?
Ans 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.
Q.9.The alkaline earth metals are less reactive than the alkali metals?
Ans.Ionisation enthalpy of alkaline earth metals is higher than alkali metal hence alkaline earth metals are less reactive than the alkali metals.
Q.10.Beryllium and magnesium are kinetically inert to oxygen and water why?
Ans.Be and Mg form of an oxide film on their surface hence further reaction is not possible.
Q.11.Alkaline earth metals are strong reducing agents why?
Ans.It is due to large negative values of their reduction potentials.
Q.12.Reducing power of alkaline earth metal is less than those of their corresponding alkali metals why?
Ans.Alkali metals have larger –ve reduction potential due to their lower ionization enthalpies.
Q.13.The alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions why?
Ans. Alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions due to formation of ammoniated ions.
M +(x + y)NH3 → [M(NH3 )x ]2+ + 2[e(NH3 )y ]−
Q.14.What are the uses of different alkaline earth metals?
Ans.1.Cu +Be alloy is used in prepare high strength springs
2. Metallic beryllium is used for making windows of X-ray tubes.
3. Mg-Al alloys are used in air-craft construction
4. Magnesium (powder and ribbon) is used in flash powders and bulbs, incendiary bombs and signals
5. A suspension of magnesium hydroxide in water (called milk of magnesia) is used as antacid in medicine.
6. Magnesium carbonate is an ingredient of toothpaste.
7. Calcium is used in the extraction of metals from oxides which are difficult to reduce with carbon.
8. Radium salts are used in radiotherapy, for example, in the treatment of cancer.
Q.15.Alkaline earth metals compounds are less ionic than the corresponding compounds of alkali metals.
Ans .It is due to increased nuclear charge and smaller size of alkaline earth metals they distort electron cloud higher than the alkali metals.
Q.16.The fluorides of alkaline earth metals are relatively less soluble than the chlorides why?
Ans.The fluorides of alkaline earth metals have higher lattice energy than the chlorides.
Q.17.The solubility of alkaline earth metal hydroxides in water increase down the group why?
Ans.Both lattice enthalpy and hydration enthalpy decreases down the group. OH– is common hence size of the cation decides the lattice enthalpy. Since lattice enthalpy decreases much more than the hydration enthalpy with increasing ionic size, the solubility increases as we go down the group.
Q.18.Solubility of alkaline earth metal carbonates and sulphates in water decreases down the group Why?
Ans.The size of anions(CO32-and SO42-) are 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 will decrease as found for alkaline earth metal carbonates and sulphates.
Q.19.Arrange the thermal stability of carbonates and sulphates in decreasing order.
Ans.As the size of ion increase, ability to polarize the negative part decreases hence thermal stability increases.
Thus BeCO3 < MgCO3 < CaCO3 < SrCO3< BaCO3
BeSO4 < MgSO4 < CaSO4 < SrSO4< BaSO4
Q.20.Beryllium carbonate is unstable and can be kept only in the atmosphere of CO2 why?
Ans. Being very small size it polarize CO32- most hence it is unstable and can be kept only in the atmosphere of CO2
Q.21.Arrange the thermal stability of alkaline earth metal hydroxide in decreasing order.
Ans.Be(OH)2 < Mg(OH)2 <Ca(OH)2< Sr(OH)2 <Ba(OH)2
Q.22.Arrange the basic strength of alkaline earth metal hydroxide in decreasing order.
Ans. Be(OH)2 < Mg(OH)2 <Ca(OH)2< Sr(OH)2 <Ba(OH)2
Q.23.Why are alkaline earth metal hydroxide are less basic, less stable and less soluble than alkali metal hydroxides?
Ans. Lower size, higher ionization enthalpy and higher lattice enthalpy of alkaline earth metals hydroxide are responsible for their less basic and less stable than alkali metal hydroxides?
Q.24.One of the alkaline earth metal hydroxide are amphoteric in nature name that hydroxide.
Ans.Be(OH)2 is amphoteric in nature because it reacts both with alkali and acid.
Q.25.Except Beryllium halides all other halides of alkaline earth metals are ionic in nature why?
Ans.Be is small atom hence it distorts electron cloud towards it most hence it forms mainly covalent halides.
Q.26.Draw the structure of beryllium chlorides in solid state and in vapour state.
Ans. In solid state it exists in polymeric form with monomer BeCl2,and in vapour state it forms chloro-bridged dimer.
Solid State
Vapour state
Q.27.The tendency to form halide hydrates gradually decreases (for example, MgCl2·8H2O, CaCl2·6H2O, SrCl2·6H2O and BaCl2·2H2O) down the group.
Ans.Since size of the alkaline earth metal decreases top to bottom hence hydration strength decreases down the group.
Q.28. List some anomalous behavior of Be.
Ans 1.Compunds of Be are mainly covalent.
2.Be can show maximum coordination number 4 but other elements can show coordination number more than 4.
3. The oxide and hydroxide of beryllium, the group, are amphoteric in nature.
Q.29.List some similar behavior of Be and Al.
Ans. Due to similar charge/radius ratio beryllium resembles aluminum in some ways.
(i) Both Be and Al are not readily attacked by acids.
(ii)Both Beryllium and aluminium hydroxide dissolve in excess of alkali to give a beryllate ion, [Be(OH)4]2– just as aluminum hydroxide gives aluminates ion, [Al(OH)4]–.
(iii) in vapour state the chlorides of both beryllium and aluminum exist in bridged chloride structure
(iv)Both the chlorides are soluble in organic solvents and are strong Lewis acids.
Q.30.Define the process slaking of lime.
Ans.The addition of limited amount of water breaks the lump of lime. This process is known as slaking of lime.
Q.31.What are the uses of quick lime(CaO).
Ans.(i)Used in manufacturing cement.
(ii) Used in manufacturing sodium carbonate from caustic soda.
(iii) In the purification of sugar
(iv)In the manufacture of dye stuffs.
Q.32.What do you mean by lime water and milk of lime?
Ans.The aqueous solution of Ca(OH)2 is called lime water and a suspension of slaked lime in water is called milk of lime.
Q.33.When CO2 is passed through lime water turns milky why?
Ans Due to formation of calcium carbonate.
Q.34.What are the uses of Calcium hydroxide(slaked lime)?
Ans. (i) Used in the preparation of mortar.
(ii)Used in white wash due to its disinfectant nature.
(iii)Used in making glass and preparation of bleaching.
Q.35.What are the uses of CaCO3 ?
Ans.(i) As an antacid
(ii) Mi ld abrasive in tooth paste.
(iii) A constituent of chewing gum.
(iv) Filler in cosmetics.
(V) Used in the manufacture of high quality paper.
Q.36.What do you mean by ‘dead burnt plaster’.
Ans. It is anhydrous calcium sulphate having no water of crystallisation formed after heating plaster (CaSO4 .1/2H2O ) of Paris above 393K.
Q.37.What are the uses of plaster of Paris(CaSO4 .1/2H2O)?
Ans. (i)In the building industry.
(ii) As plasters.
(iii) In ornamental work.
(iv)For making casts of statues and busts.
Q.38.Cement is is also called Portland why?
Ans. It resembles with the natural limestone in the Isle of Portland, England.
Q.39.What are the composition of cement.
Ans. The averagecomposition of Portland cement is : CaO, 50-60%; SiO2, 20-25%; Al2O3, 5-10%; MgO, 2-3%; Fe2O3, 1-2% and SO3, 1-2%.
Q.40.What do you mean by setting of cement?
Ans.When water is mixed with cement it gives give a hard mass, due to the hydration of the molecules of the constituents and their rearrangement.
Q.41.Why gypsum is added to cement?
Ans.Gypsum slow down the process of setting of the cement so that it gets sufficiently hardened.
Q.42.What are the biological use s of alkaline earth metals.
Ans (i)Chlorophyll contains magnesium.
(ii)Ca is present in bones and teeth.
(iii) Ca plays important roles in neuromuscular function,interneuronal transmission, cell membrane integrity and blood coagulation.
Hydrogen is placed separately in the periodic table why?
Hydrogen resembles both with alkali metals and halogens .At same time it also differs in many ways with those elements. Hydrogen is extremely small and H+doest not exist freely hence hydrogen is placed separately in the periodic table.
Properties to which hydrogen resembles with alkali metals.
1.Like alkali metals hydrogen can lose one electron to form unipositive ions.
2. Like alkali metals, hydrogen forms sulphide, oxides and halides.
3.Outermost electronic configuration of alkali metals(ns1) is similar to electronic configuration of hydrogen(1s1)
Properties to which hydrogen differs from alkali metals.
Ionisation enthalpy of hydrogen is very high unlike alkali metals which have low ionization enthalpies.
2. Hydrogen has not metallic properties.
Properties to which hydrogen resembles with halogens.
1.Like halogens hydrogen is also short by one electron to the corresponding noble gas configuration.
2. In term of ionization enthalpy, hydrogen is similar with halogens.
3. Like halogens, hydrogen forms a diatomic molecule.
How many isotopes of hydrogen are known.
Which hydrogen is also known as heavy hydrogen.
Deuterium is also known as heavy hydrogen.
Are chemical properties of different isotopes are same? Explain in which property they differ from each other.
All isotopes of hydrogen have same electronic configuration hence they have almost the same chemical properties but their rates of reactions are different due to their different enthalpy of bond dissociation. Physical properties of these isotopes are different due their large mass differences.
What is water gas ?
The mixture of CO and H2 is called water gas.
What is syn gas?
Water gas (CO+ H2 ) is also known as synthesis gas (syn gas) because it is used for synthesis methanol and hydrocarbons.
What is ‘coal gasification’ ?
The process of producing ‘syn gas’ from coal is called ‘coal gasification’.
Hydrogen is inert at room temperature why or Atomic hydrogen is produced at a high temperature in an electric arc or under ultraviolet radiations.
The bond dissociation enthalpy of H-H is very high thus hydrogen gas is inert at room temperature.
Describe the different ways by which hydrogen combines with other element.
(i) Loss of an electron to give H +
(ii) Gain of an electron to form H–
(iii) Forms single covalent bond by sharing electrons.
Uses of Hydrogen?
(i) It is used as a rocket fuel.
(ii) In preparation of organic compound like methanol etc.
(iii) In preparation of inorganic compound like NH3 ,metal hydrides etc.
(iv) In hydrogenation of vegetable oils into vanaspati.
(v) In metallurgical processes, to reduce heavy metal oxides to metals.
(vi) In fuel cells to generate electrical energy.
Three types of hydrides.
(i) Ionic or saline or salt like hydrides.
(ii) Covalent or molecular hydrides.
(iii) Metallic or non-stoichiometric hydrides.
(i) Ionic or saline or saltlike hydrides.
These hydrides are formed with s block elements like Na ,K and Ca.
Exampes :NaH,KH,CaH2 etc.
Properties.
(i) Mainly ionic but some covalents hydrides are also available.
(ii) Crystalline, non-volatile and non conducting in solid state.
(ii) Covalent or molecular hydrides.
These hydrides are formed with P block elements like C,N.F etc
CH4, NH3, H2O and HF
Properties
(i) They are covalent hydrides.
(ii) Volatile compounds
(iii) Metallic or non-stoichiometric hydrides.
These are formed by many d-block and f-block elements. LaH2.87, YbH2.55, TiH1.5–1.
Properties.
(i) Nonstoichiometric, being deficient in hydrogen.
(ii) Crystalline, non-volatile and nonconducting in solid state.
Pd,Pt can use source of energy why?
These elements can accommodate large number of hydrogen thus they used as its
hydrogen storage and as a source of energy.
Different types of molecular hydrides
(i)Electron-deficient
(ii) Electron-precise
(iii) electron-rich
(i) Electron-deficient
Group 13 hydrides are electron-deficient compounds.like BH3.AlH3 etc.They all are electron acceptors that are Lewis acids.
(ii) Electron-precise
All elements of group 14 form such compounds example CH4 etc.
(iii) Electron-rich hydrides
Group 15-17 hydrides are electron-rich hydrides.Example: NH3 ,H2O, HF.They all are electron donors that are Lewis bases.
Water has high freezing point, high boiling point, high heat of vaporisation and high heat of fusion in comparison to H2S and H2Se why?
It is due to the presence of extensive hydrogen bonding between water molecules.
Which property of water is responsible for moderation of the climate and body temperature of living beings.
The high heat of vaporisation and heat capacity are responsible for moderation of the climate and body temperature of living beings.
What is the structure of ice?
Ice has crystalline structure in which each oxygen atom is surrounded tetrahedrally by four other oxygen.
Water acts as Bronsted acid and base both give reactions to prove it.
What is hard water and soft water?
Water which contains hydrogencarbonate, chloride and sulphate of calcium and magnesium called hard water. Water free from soluble salts of calcium and magnesium is called soft water.Hard water does not give lather with soap. Soft water gives lather with soap easily.
Hard water is unsuitable for laundry also why?
Hard water forms scum/precipitate with soap hence unsuitable for laundry.
Hard water is harmful for boilers why?
On boiling there is deposition of salts in the form of scale. This reduces the efficiency of the boiler.
Types of hardness of water?
Hardness of water is of two types:
(i) Temporary hardness
Temporary hardness is due to the presence of magnesium and calcium hydrogen carbonates
(ii) Permanent hardness
Permanent hardness is due to the presence of magnesium and calcium chlorides and sulphates.
30% solution of H2O2 is known as 100 volume hydrogen peroxide. It means 1 ml of 30 % H2O2 gives 100ml of oxygen at STP.
What are the uses of H2O2?
(i) Used as a hair bleach and as a mild disinfectant.
(ii) Perhydrol is H2O2 used as antiseptic
(ii) In manufacturing sodium perborate and per-carbonate, which are used in high quality detergents.
(iii) In the synthesis of hydroquinone, tartaric acid, pharmaceuticals (cephalosporin) etc.
(iv)Used as a bleaching agent for textiles, paper pulp, leather, oils etc.
What is heavy water?
D2O, used as a moderator in nuclear reactors.
What are the advantage of using hydrogen as fuel?
(i)On a mass for mass basis dihydrogen can release more energy than other sources of fuel
(ii) Pollutants in combustion of dihydrogen are less than other sources of fuel.
What are the disadvantages of using hydrogen as fuel cell?
1. Mass of a cylinder containing compressed dihydrogen of will be very high.
2.Expensive insulated tanks will be used to convert hydrogen gas into liquid.
What will be only possible pollutants of fuel cell?
The only pollutants will be the oxides of dinitrogen.
What is the basic principle of hydrogen economy?
To transport and store of energy in the form of liquid or gaseous dihydrogen. Advantage of hydrogen economy is that energy is transmitted in the form of dihydrogen and not as electric power.
Definition no 1: Addition of oxygen to an element or a compound called oxidation.
Definition:2 The removal of hydrogen from a substance called oxidation.
Definition 3 : Removal of electropositive element from substance is called oxidation.
Definition 4: Addition of electronegative element to a substance is called oxidation.
Definitions of Reduction.
Definition no1: Addition of hydrogen to an element or a compound is called reduction
Definition:2 The removal of oxygen from a substance is called reduction
Definition 3: Addition of electropositive element from substance is called reduction.
Definition 4: Removal of electronegative element from substance is called reduction.
Define redox reactions.
When oxidation and reduction occur simultaneously that reaction is called redox reaction.
Define oxidation and reduction reactions in terms of electron transfer.
loss of electrons called oxidation and the gain of electrons is called reduction.
Define oxidizing agent and reducing agent.
Compound which oxidises other and reduces itself is called oxidizing agent.
Compound which reduces other and oxidises itself is called reducing agent.
Define Oxidation number.
Oxidation state of an element in a compound assuming there is a complete transfer of electron from a less electronegative atom to a more electonegative atom and determined by a set of rules.
Define oxidation,reduction,oxidizing agent ,reducing agent and redox reactions in terms of oxidation number.
Oxidation: Increase in oxidation number of the element in the given substance.
Reduction: Decrease in oxidation number of the element in the given substance.
Oxidising agent: A reagent which increases the oxidation number of an element in a given substance and decreases itself.
Reducing agent: A reagent which lowers the oxidation number of an element in a given substance and increases itself.
Redox reactions: Reactions in which there is increase in oxidation number of an element and also decrease in oxidation number takes place is called redox reduction.
What are the rules for calculation of oxidation number?
1. The oxidation number of an atom in free elements is zero. Hydrogen in H2, sulphur in S8, phosphorus in P4 oxygen in O3 is zero.
2.Fluorine has oxidation number -1
3. Oxidation number of oxygen is -2 in all compounds except in peroxides, superoxides and oxygen fluorides. Oxidation number in peroxide is -1, in superoxides -1/2, in OF2 +2 and in O2F2 is +1.
4.The oxidation number of hydrogen is +1 in all of its compounds except metallic hydrides. In metallic hydrides it is +1.
5.The oxidation number of an ion is equal to the charge on it.
6.The oxidation number of group 1 elements(Li,Na,K,Rb,Cs) is +1 and the oxidation number of group 2 elements(Be,Mg,Ca,Sr,Ba) is +2.
7.For complex ions, the algebraic sum of oxidation numbers of all the atoms is equal to the net charge on the ion.
8.For neutral molecules the algebraic sum of oxidation numbers of all the atoms is equal o zero.
Balance the following equation by oxidation number method in acidic medium.
Step1.Calculate the oxidation number of each elements taking part in reaction.
Step 2. Balance the atoms which oxidation number changed during the reaction.
Step.3 Calculate the increase or decrease in the oxidation number per atom which oxidation number changed during the reaction.
Step.4 If these changes are not equal then multiply by suitable number so that these become equal.
Step.5 Balance oxygen if unbalance.
Balance the following equation by ion electron method method/half reaction method in acidic medium.
Step1.Calculate the oxidation number of each elements taking part in reaction.
Step 2. Write individual oxidation and reduction reaction.
Step 3:Balance the atoms which oxidation number changed during the reaction.
Step 4: Calculate the increase or decrease in the oxidation number per atom which oxidation number changed during the reaction and add it in higher oxidation state side.
Step5. Add two reactions so that added electron must be cancelled.
\
Step6: Balance Oxygen atoms
Rule for balancing oxygen in acidic medium
Add H+ twice of the unbalance oxygen in that side in which higher number of oxygen is present and H2O equal number of unbalance oxygen in another side.
Rule for balancing oxygen in basic medium
Add H+ equal number of the unbalance oxygen in that side in which lower number of hydrogen is present.
Rule for balancing hydrogen in basic medium
Add OH– equal number of the unbalance oxygen in that side in which lower number of oxygen is present and H2O equal number of unbalance hydrogen in another side.
Rule for balancing hydrogen in basic medium
Add OH– equal number of the unbalance oxygen in that side in which higher number of oxygen is present and H2O equal number of unbalance hydrogen in another side.
Define electrolytic cell.
The cell which convert electrical energy into chemical energy is called electrolytic cell.
Define Galvanic cell.
The cell which convert chemical energy into electrical energy is called electrolytic cell.
Define electrode potential.
Potential of individual electrode is called electrode potential.
Dehine Standard Electrode Potential.
At 298K when the concentration of each species taking part in the electrode reaction is unity reaction is carried out at 298K, then the potential of each electrode is said to be the Standard Electrode Potential .
Rate of transfer of molecules from solid into liquid and of reverse transfer from liquid into solid is equal at normal melting point or normal freezing point of the substance.
Liquid -vapour equilibrium.
When rate of evaporation is equal to rate of condensation at normal boiling point of
the liquid.
Solid-vapour equilibrium.
Rate of transfer of molecules from solid into vapour and of reverse transfer from vapour into solid is equal.
Equilibrium involving dissolution of solid or gases in liquids.
For dissolution of solid in liquids.
The rate of dissolution of solid substance = rate of crystallization of solid substance.
For dissolution of solid in liquids.
Define Henry’s Law.
The amount of solute that can be dissolved in given amount of solvent at given temperature in mole fraction is directly proportional to the partial pressure of the gas.
Define equilibrium/Dyanamic equilibrium.
Actual reaction doesn’t stop but it proceeds from both sides with same speed that is rate of forward reaction is equal to rate of backward reaction.
What is Law of mass action?
According to law of mass action rate of reaction is directly proportional to the active massif reactants raised to the power equal to stochiometric coefficient.
In equilibrium expressions, we use activity (a) instead of concentration. For most practical cases:
For gases → activity ≈ partial pressure
For solutions → activity ≈ concentration
For pure solids and pure liquids → activity = 1 (unity)
🔹 Why is activity of a solid = 1?
Constant density
In a solid, particles are tightly packed.
The “effective concentration” (number of particles per unit volume) does not change.
No effect on equilibrium
Even if you take more or less solid, its concentration remains constant.
So it does not affect the equilibrium position.
Defined reference state
Activity is defined relative to a standard state.
For a pure solid, its standard state = itself → so activity = 1.
Equilibrium constant (Kc)
What is KP.
KP is equilibrium constant for reactions involving gaseous reactants and products. In place of concentration partial pressures are used.
What is the relationship between and KP and KC.
Define heterogeneous and homogenous equilibrium reaction.
Equilibrium is a state of chemical reaction in which the rates of backward and forward reactions are equal.
Reaction in which equilibrium exists between reactants and products having more than one phase is called heterogeneous equilibrium.
Factors effecting equilibrium constant.
(i) Equilibrium constant is independent of initial concentrations of the reactants and products.
(ii) Equilibrium constant depends on temperature only.
How can we predict the extent of reaction by using equilibrium constant.
(i) If value of Kc is very large that is greater than 103 than, the reaction is nearly to completion. The concentration of products is much larger than that of the reactants at equilibrium.
(ii) If value of Kc is very small that is lesser than 10-3 than, the backward reaction is favored. Concentration of reactants is much larger than that of products that is the reaction proceeds to a very small extent in forward reaction
(iii) If value of Kc is in between 103 and 10-3 than reaction proceeds toward equilibrium that is the reactants and products are comparable.
Define reaction quotient, Qc or Qp.
Reaction quotient is calculated similar as equilibrium constant but concentration of reactants and products are not necessarily equilibrium values.
How can we predict the direction of reaction by using reaction quotient.
(i) If Qc > Kc, the reaction is in reverse direction.
(ii) If Qc < Kc, the reaction is in forward direction
(iii) If Qc = Kc, the reaction mixture is at equilibrium.
When the particular reaction is non spontaneous, spontaneous and equilibrium
Define Le Chatelier’s principle.
According to this principle when a chemical equilibrium is disturbed by a change of temperature, pressure, concentration or presence of catalyst equilibrium shifts in that direction in which these changes can be nullified.
(i)Effect of concentration.
(a)Addition of reactant shifts the equilibrium in forward direction
(b)Addition of products shifts the equilibrium in backward direction.
(c)Removal of reactant shifts the equilibrium in backward direction.
(d)Removal of products shifts the equilibrium in forward direction.
(ii)Effect of temperature:
(a)Increase the temperature to an exothermic reaction shifts the equilibrium on backward direction.
(b)Increase the temperature to an endothermic reaction shifts the equilibrium on forward direction.
(iii)Effect of pressure:
Increase the pressure shifts the equilibrium in that direction in which lower number of moles of gaseous reactants or products are present.
(iv) Effect of catalyst:
Catalyst does not affect the equilibrium but it helps the reaction to achieve equilibrium faster.
What is electrolyte?
Substances which dissociates into constitutes cation and anion in aq. solution is called electrolyte.
Define the types of electrolyte.
There are two types of electrolytes strong electrolyte and weak electrolyte.
Strong electrolyte:Types of electrolyte which dissociates almost completely.
Examples: All salts(sodium chloride), Strong acids HCl,HNO3, H2SO4,H3PO4 ,HClO4,HBr,HI. Strong bases-Group 1 and group 2 elements hydroxides. Like NaOH ,LiOH ,KOH etc.
Weak electrolytes : Types of electrolyte which does not dissociates completely.
Examples: Weak acids like acetic acid,HNO2,HF etc. Weak bases like NH4OH.
Condition in which ionic equilibrium condition is established.
Ionic equilibrium condition is established in case of weak electrolyte, between dissociated ions and the unionized ions.
Define acids and bases according to Arrhenius theory.
Acids are substances which give H+( hydrogen ions) in water. and bases are substances that give OH– ( hydroxide ions) in water
What is the limitation of Arrhenius concept of acids and bases.
(i) Arrhenius concept is limited to only aq. Solution.
(ii) It does not explain the acidity and basicity of those compounds which do not contain hydrogen H+ and OH–.
What are the Bronsted and Lowry concept of acids and bases.
Acids are the substance which donate proton(H+) and bases the substance which accept proton.
Define conjugate acid and base pairs.
The acid-base pair differs only by a proton(H+) is called a conjugate acid-base pair.
HCl and Cl– are conjugate acid base pairs. Similarly H2O and H3O+ are another conjugate acid base pairs.
Species which accepts electron pair are called Lewis acids and which donates an electron pair are called Lewis bases. Strong acids have weak conjugate bases and vice versa.
Examples of Lewis acids.
(i) Some +ve charge species are Lewis acids.
Like H+, Fe3+ etc.
(ii) Electron deficient molecules are Lewis acids.
Like BF3, B2H6 ( BH3 ) , AlCl3 etc.
Examples of Lewis bases.
(i)-ve charge species are Lewis bases
Cl–,OH– etc.
(ii) Electron rich molecules are Lewis bases.
Like H2O, NH3 etc.
Give Two reactions to prove amphoteric nature of water.
Pure water also show amphoteric behavior how?
Ionic product of water.
Because concentration of water (55.55 M) is very high and doesn’t change during reaction
How can we determine the acidic, basic and neutral aq solution
[H3O+] > [OH– ]= Acidic
[H3O+] = [OH– ]= Neutral
[H3O+] < [OH–]= Basic
Define pH scale.
Prove pH+ pOH= 14
Ionization constant (Ka) for weak acid.
Ionization constant (Kb) for weak base
What is pKa and pKb.
pKa=-log[Ka] and pKb=-[log Kb]
Relation between Ka and Kb for conjugate acid base pair.
Determine the factors effecting acidic strength.
(i)Larger the value of ionization constant Ka higher the acidic strength.
(ii) In a given group acidic strength increases as the bond length increases.
Example: HI>HBr>HCl>HF
(iii) If the difference between electronegativity increases acidic strength increases.
Example: CH4 < NH3 < H2O < HF
What is common ion effect.
Common Ion effect is based on the Le Chatelier’s principle. If any common ion is added equilibrium shifts in that direction in which this increased concentration is consumed.
What is difference between hydration and hydrolysis of salt?
Salt is dissociated into constituent ions in aq.solution and remains as hydrated ions called hydration.
When ions interact with water to form corresponding acids or bases called hydrolysis.
Examples:
(i) Salts of weak acid and strong base e.g.CH3COONa.
(ii) Salts of strong acid and weak base e.g. NH4Cl
(iii) salts of weak acid and weak base, e.g.CH3COONH4
Salt of strong acid + strong base
Example: NaCl No hydrolysis
pH = 7 (neutral)
2. Salt of weak acid + strong base
Example: CH₃COONa
✔ Solution is basic (pH > 7)
3. Salt of strong acid + weak base
Example: NH₄Cl
✔ Solution is acidic (pH < 7)
4. Salt of weak acid + weak base
Example: NH₄CH₃COO
✔ pH depends on Ka vs Kb
If Ka > Kb → acidic
If Kb > Ka → basic
If Ka = Kb → neutral
What is the formula for determining pH of aq.solution containing salt of weak acid and weak base, e.g.CH3COONH4
pH = 7 + ½ (pKa – pKb)
Define Buffer Solutions.
The solutions which resist change in pH value on the addition of small amounts of acid or alkali are called Buffer Solutions.
uffer solutions are solutions that resist change in pH when small amounts of acid or base are added.
There are 2 main types of buffer solutions
1. Acidic Buffer
Made from: weak acid + its salt (with strong base)
Examples:
CH₃COOH + CH₃COONa
H₂CO₃ + NaHCO₃
✔ pH < 7 (acidic range) ✔ Common range: pH 3–6
Formula (Henderson–Hasselbalch equation):
2. Basic Buffer (Alkaline Buffer)
Made from: weak base + its salt (with strong acid)
Examples:
NH₄OH + NH₄Cl
✔ pH > 7 (basic range) ✔ Common range: pH 8–11
Formula:
Then convert:
Quick Revision Table
Type
Components
pH
Acidic Buffer
Weak acid + its salt
< 7
Basic Buffer
Weak base + its salt
> 7
What is solubility product Ksp.
Equilibrium is established between the undissolved sparingly soluble salt in water
Substance which is taken under observation is called system , rest of the universe is called surrounding.
Explain the different types of system with example.
There are three types of the system.
(i) Open system.
The type of system in which both matter and energy is exchanged between system and surroundings.
Example: Substance in open container.
(ii) Closed system.
Example: Substance in close container having thermoconducting wall.
The type of system in which only energy is exchanged between system and surroundings.
(iii) Isolated system.
The type of system in which neither matter nor energy is exchanged between system and surroundings.
Example: Substance in insulated vessel (Thermos flask)
The state of the system
In thermodynamics, the state of a system means the condition of a system at a particular moment, described by certain measurable properties.
Definition
The state of a system is the condition of the system specified by state variables such as pressure, temperature, volume, and composition.
State Variables (State Functions)
These properties determine the state of a system:
Temperature (T)
Pressure (P)
Volume (V)
Number of moles (n)
Density
Internal energy (U)
f these variables are known, the state of the system is completely defined.
Example
Consider a gas in a container:
Temperature = 300 K
Pressure = 1 atm
Volume = 2 L
These values describe the state of the system.
Important Points
When any state variable changes, the system goes to a new state.
The change from one state to another is called a thermodynamic process.
Example: Heating a gas changes temperature, pressure, or volume, so the state changes.
Short Definition for Students
State of a system: “The set of values of thermodynamic properties such as pressure, temperature, and volume that describe the condition of a system at a given time.”
Define state function or state variables and path function.
Those properties of the system which depend only on the state of the system and not on how it is reached to that state are called state function.
Examples:
State Functions (Thermodynamic Properties)
1. Basic Thermodynamic State Variables
Temperature (T)
Pressure (P)
Volume (V)
Density (ρ)
Number of moles (n)
Concentration
2. Energy Functions
Internal Energy (U)
Enthalpy (H)
Helmholtz Free Energy (A or F)
Gibbs Free Energy (G)
3. Other Thermodynamic State Functions
Entropy (S)
Heat Capacity (C)
Molar heat capacity
Specific heat
Those properties of the system which depend on how the change is carried out are called path function.
Examples: :Heat energy(q),Work done(w)
Define internal energy (U).
The total energy contained by the system is called internal energy.
U=Kinetic energy + Potential energy
Internal Energy Includes
Translational kinetic energy of molecules
Rotational energy
Vibrational energy
Electronic energy
Nuclear energy
Intermolecular potential energy
Kinetic energy=translational energy + rotational +vibrational.
Internal energy does not include:
Macroscopic kinetic energy
Macroscopic potential energy
Thermodynamic Processes (JEE / NEET)
1. Isothermal Process
Temperature remains constant.
Condition:
For an ideal gas:
PV = constant
Key Points
Heat exchange occurs.
Internal energy change = 0 (for ideal gas).
Example Slow expansion of gas in contact with a heat reservoir.
2. Adiabatic Process
No heat exchange between system and surroundings.
Condition:
Equation:
Key Points
Temperature changes.
Work done comes from internal energy.
Example Rapid compression of gas.
3. Isobaric Process
Pressure remains constant.
Condition:P=constant
Work done:
Key Points
Volume and temperature change.
Example Heating gas in a cylinder with movable piston.
4. Isochoric Process (Isovolumetric)
Volume remains constant.
Condition:
Key Points
No work done
Heat supplied increases internal energy.
Example Heating gas in a rigid container.
Other Processes (Sometimes Asked)
Cyclic Process
System returns to initial state.
Example Heat engine cycle.
Reversible Process
Infinitely slow process.
System always in equilibrium.
Irreversible Process
Fast process.
System not in equilibrium.
Example Free expansion of gas.
What is adiabatic system?
The types of the system in which there is no heat exchange between the system and surroundings through its boundary is called adiabatic system.
What is the difference between adiabatic system and isolated system.
Adiabatic system is a specified condition
Mention the sign convention for different thermodynamics function.
Energy absorbed by the system ‘q’=+ve
Energy released by the system ‘q’ =-ve
Work done by the system ‘w’ = -ve
Work done on the system ‘w’ = +Ve
Define first Law of thermodynamics.
The First Law of Thermodynamics is based on the principle of conservation of energy. It states that energy can neither be created nor destroyed; it can only be converted from one form to another. Internal energy change of the system is equal to the energy provide to the system and work done by the system.
Obtain a formula for mechanical work (pressure- volume work) against constant external pressure .
What do you mean by free expansion?
Expansion of a gas in vacuum when is called free expansion.
No work during free expansion of an ideal gas whether the process is reversible or irreversible.
Define reversible and irreversible process.
The process which occur infinitely slowly such that system and surrounding are always in equilibrium with each other. The process can be reversed at any moment by an infinitesimal change called reversible process. An irreversible process can be defined as a process in which the system and the surroundings do not return to their original condition once the process is initiated
Change in internal energy of a system when
(i)Work done by the system or on the system and heat is absorbed by the system.
(ii)If a process is carried out at constant volume and heat is absorbed by the system.
(iii) For isothermal process.
(iv) For adiabatic system.
What is the equation for calculating work done for isothermal irreversible process.
Obtain equation for the work done for the isothermal reversible process.
Define enthalpy.
A thermodynamic function, the enthalpy H is equal to H = U + pV .
U=Internal energy of the system
P=Pressure of the system.
V=Volume of the system.
What will be the change in enthalpy during the reaction at constant pressure.
What is the equation for calculating change in enthalpy at constant pressure for a reaction involving gases.
What is extensive property and intensive properties of a system?
Those properties which depends on the amount of matter present in the system are called extensive property.Example: Mass, volume, internal energy, enthalpy.
Complete List of Common Extensive Properties
Mass (m)
Volume (V)
Number of moles (n)
Internal Energy (U)
Enthalpy (H)
Entropy (S)
Gibbs Free Energy (G)
Helmholtz Free Energy (A or F)
Total Charge
Heat Capacity (C)
Total Energy
Total Momentum
Total Electric Charge
Surface Area
Length
Weight
Total Magnetic Moment
Total Heat Content
Those properties which do not depend on the amount of matter present in the system are known as intensive properties.Example:Temperature, density, pressure etc.
Complete List of Common Intensive Properties
Temperature (T)
Pressure (P)
Density (ρ)
Concentration
Refractive Index
Boiling Point
Melting Point
Specific Heat Capacity
Molar Heat Capacity
Molar Volume
Viscosity
Surface Tension
Thermal Conductivity
Electrical Conductivity
Specific Gravity
Color
Odor
Hardness
Molar Entropy
Molar Enthalpy
Molar Gibbs Energy
Define heat capacity.
Heat requires by the substance to increase the temperature by one 0C is called heat capacity.
Define specific heat capacity.
Heat requires by the unit mass of the substance to increase the temperature by one 0C is called specific heat capacity.
Molar heat capacity is the amount of heat required to raise the temperature of 1 mole of a substance by 1 Kelvin (or 1°C).
What is the relation between heat capacity at constant volume (Cv) and heat capacity at constant pressure Cp)?
What is bomb calorimeter?
A calorimeter in which the chemical reaction is carried out at constant volume. Hence there is no work done.
How can we calculate the enthalpy change for reaction?
Define standard enthalpy of reaction.
The standard enthalpy of reaction is the enthalpy change that occurs on a system when one mole of matter is transformed by a chemical reaction under standard condition.
Define standard enthalpy of fusion
Enthalpy of fusion is the change in enthalpy when one mole of solid is melted in liquid.
Define standard enthalpy vaporization
Amount of heat required to vaporize one mole of liquid at constant temperature.
Define standard enthalpy of sublimation
The change in enthalpy when one mole of a solid substance is converted into vapour state at constant temperature.
Define standard enthalpy of formation
Enthalpy change during the formation of one mole of a compound from its constituent element in their standard state.
If an element is already in its most stable form at 1 bar pressure, its enthalpy of formation is defined as zero.
This is just a reference point in thermodynamics.
Examples
because these are elements in their standard states.
Define enthalpy of combustion
Enthalpy change during combustion of one mole of a substance when all the reactants and products are in standard state.
Define enthalpy of atomisation.
(i )Enthalpy change during breaking of one mole of bond of a compound completely into gaseous state.
(ii)Enthalpy of atomization is also called enthalpy of sublimation in case of solid element.
Define bond dissociation enthalpy.
The bond dissociation enthalpy is the change in enthalpy when one mole of covalent bond is broken to form product in gaseous state.
Define mean bond enthalpy.
For the polyatomic molecules mean bond enthalpy is defined as total bond dissociation enthalpies divided by number of bonds.
Define Hess’s Law of constant heat summation.
When a reaction takes place in several steps then its standard enthalpy is the sum of the standard enthalpies of the intermediate reactions.
Define enthalpy of solution
Enthalpy change when one mole of substance is dissolved in given amount of solvent at give temperature.
Define lattice enthalpy.
The enthalpy change when one mole of an ionic compound dissociates into it constituent ions in gaseous state.
Define enthalpy of hydration.
The energy released when new interacting are made between the gaseous ions and water molecules is called hydration enthalpy.
What is the relationship between enthalpy of solution,lattice enthalpy,hydration enthalpy.
For ionic compound
Criteria for solubility of ionic compounds in water
Main Solubility Criterion
An ionic compound dissolves in water when:
Meaning:
Hydration energy ≥ lattice energy
Then dissolution becomes energetically favorable.
What is Born -Haber cycle.
With the help of Born-Haber Cycle we determine the lattice enthalpy of the substance.
Heat of Neutralisation
1. Definition
Heat of neutralisation is the enthalpy change when 1 mole of water is formed from the reaction of an acid and a base.
2. Standard Value (Most Important Point)
3. Strong vs Weak (Very Important)
Case 1: Strong acid + Strong base
Example: HCl + NaOHΔH=−57.1kJ/mol
Case 2: Weak acid + Strong base
Example: CH₃COOH + NaOH
Case 3: Strong acid + Weak base
Example: HCl + NH₄OH
Same logic → less heat released.
Case 4: Weak + Weak
Least heat evolved.
4. Key Formula (Numerical)
5. Limiting Reagent Concept (Very Important)
Heat depends on moles of water formed, not volume.
Example:
If H+ = 0.1 mol
OH− = 0.05 mol
→ Water formed = 0.05 mol
6. Temperature Rise Formula
Trap 1
Students think more volume = more heat ❌ ✔ Depends on moles reacting
Trap 2
Using -57.1 for weak acids ❌ ✔ Only for strong + strong
Trap 3
Forgetting diprotic acids
Example:
Define spontaneous reaction.
The reaction which occurs itself without any external force is called spontaneous reaction.
Examples:Burning of carbon in oxygen.Cooling of hot water.
What do you mean by spontaneity of chemical reaction?
Spontaneity means feasibility of particular chemical reaction that is that reaction is possible or not.
What is the criteria for spontaneous process?
Define entropy(s)
The degree of randomness or disorderness is called entropy.
How can we calculate the change in entropy of a reversible reaction?
Where q=enthalpy change during chemical reaction or heat released or absorbed by the system.
What is the relation between change in entropy of the system ,change in entropy of surrounding )?
What is the relation between change in entropy of the system. change in enthalpy of system ) and Gibbs energy change ?
In that case in which randomness increases the change in entropy is taken as +ve and when randomness decreases the change in entropy is taken as –ve.
What is the formula for calculating equilibrium constant Kc?
When the particular reaction is non spontaneous, spontaneous and equilibrium according to .
Predict whether the reaction may occur spontaneously or not in following cases.
ALL FORMULAS OF ENTROPY CHANGE (ΔS)
All important formulas for an adiabatic process (very useful for JEE/NEET)
The forces of attraction and repulsion between constituent particles of matter a molecule or atom is called intermolecular force of attraction.
Different types of attractive intermolecular forces of attraction (van der Waals forces).
van der Waals forces are of three types.
(i)London forces/dispersion force-Between non polar molecules.
(ii)Dipole-dipole forces- Between polar molecules.
(iii)Dipole-induced dipole forces- Between non polar and polar molecules.
H-bonding-A particular type of dipole-dipole interaction acts between hydrogen of one molecule and F or O or N of other molecule.
What is thermal energy.
Energy of the substance arises due to motion of it constituent atoms or molecules.
Thermal energy of substance in their different physical states decreases in following order.
Gas>Liquid>Solid
Define Boyle’s Law.
At constant temperature, the pressure of a given mass of gas is inversely proportional to its volume.
Define isotherm.
Each curve corresponds to a different constant temperature in pressure, p vs. Volume, v is called isotherm.
Draw Graph of pressure, p vs.1/V of a gas at different constant temperatures.
Define Charle’s law.
At constant pressure the volume of a given mass of a gas is directly proportional to its absolute temperature.
Draw Volume vs. Temperature graph at constant pressure.
Define isobar.
Each line in Volume vs Temperature graph at constant pressure is called isobar.
Define absolute zero.
Hypothetical temperature at which volume of gases are considered zero is called absolute zero.This temperature is -273.15oC.
Explain Gay Lussac- Law.
At constant volume, pressure of a given mass of a gas is directly proportional to the temperature.
Define isochore.
Each line in Pressure vs. Temperature graph at constant volume is called isochore.
Define Avogadro Law.
Equal volumes of all gases under the same temperature and pressure contain equal number of molecules.
What do you mean by STP.
STP means standard temperature and pressure. Standard temperature is 273.15 K (0oC) and standard pressure is 1 bar .At STP molar volume of an ideal gas is 22.7 L mol–1.
Define ideal gas.
Gas which obeys all types of gas laws Boyle’s law, Charles’ law and Avogadro law strictly is called an ideal gas.
What is ideal gas equation.
Combination of four gas laws gives ideal gas equation.
Define partial pressure.
The pressure exerted by the individual gas in mixture of gases is called partial pressure.
Define Dalton’s Law of partial pressure.
The total pressure exerted by the mixture of gases is equal to the sum of the partial
pressures of individual gases.
PTotal = PA+PB+PC……………
Let at the temperature T, three gases A,B and C enclosed in the volume V, exert partial pressure pA, pB and pC respectively. then,
Waht is the relation between partial pressure mole fraction, total pressure and partial pressure.
Define aqueous tension.
Pressure exerted by saturated water vapour is called aqueous tension.
What is the different assumptions or postulates of the kinetic molecular theory of gases.
(i)The volume of the molecules is negligible in comparison to the empty space between them.
(ii) There is no force of attraction between the particles of a gas at ordinary temperature and pressure.
(iii) Particles of a gas are always in constant and random motion in straight lines
(iv)Pressure exerted is due to collision of the particles with the walls of the container
(v) Collisions of gas molecules are perfectly elastic thus the total energy of molecules before and after the collision remains constant.
(vi)The average kinetic energy of the gas molecules is directly proportional to the absolute temperature.
Which assumption of kinetic theory explains the great compressibility of gases.
The volume of the molecules is negligible in comparison to the empty space between them.
Which assumption of kinetic theory explains that gases expand and occupy all the space available to them.
There is no force of attraction between the particles of a gas at ordinary temperature and pressure.
Why do gases deviate from the ideal behaviour?
It is due to following reasons.
(i) The volume of the molecules is negligible in comparison to the empty space between them.
(ii) There is no force of attraction between the particles of a gas at ordinary temperature and pressure.
What are the conditions under which gases deviate from ideality?
Gases deviate from ideal behaviour when compressibility factor Z>1or Z<1.This is possible at high pressure and low temperature when molecules of gases are very close to each other and intermolecular forces become significant.
What are the conditions under which gases follow ideal gas equation and very close to ideality?
Gases behave as ideal gas when compressibility factor Z=1.This is possible at very low pressure and high temperature when molecules of gases are very far to each other and intermolecular forces are negligible.
What are the real gas/van der Waals equation.
What is the significance of ‘a’ and ‘ b’ in real gas equation.?
’a’ is measure of magnitude of attractive forces between the molecules .Higher the value of ‘a ‘easier’ the liquefaction. ‘b’ is effective size of molecules, higher the value of ‘b’ smaller is the compressible volume.
What is the unit of van der Waals ‘a’ and ‘b’.
Unit of a=L2 mol-2atm.
Unit of b=Lmol-1
What is compressibility factor Z.
Z determine the deviation of gas from ideal behaviour.
Define Boyle point/ Boyle temperature.
The temperature at which a real gas obeys ideal gas law over an appreciable range of pressure is called Boyle temperature or Boyle point.
How can compressibility factor is determined id VReal and VIdeal are given.
Critical temperature(Tc)- Highest temperature at which gases can be liquefied by increasing pressure above this temperature gases can’t be liquefied.
Critical pressure(Pc)– Pressure of one mole of gas at critical temperature is called critical pressure (pC).
Critical volume (VC) -Volume of one mole of the gas at critical temperature is called critical volume (VC).
What do you mean by permanent gases.?
Gases which has Z>1 that is show continuous positive deviation in Z(compressibility factor) value are permanent gases. These gases have lower value of critical temperature and can’t be liquefied easily.Examples:H2.He.N2 etc.
Give the examples of easily liquefiable gases.
CO2,NH3,H2O etc. are easily liquefiable gases which has having high critical temperature. These gases first show decrease in Z value with increasing pressure, which reaches a minimum value. On further increase in pressure, the value of Z increases continuously.
Which of these gases will liquefy first when we start cooling from 700 K to their critical temperature ?CO2,NH3,H2O, H2.He,N2,O2
H2O> NH3 >CO2> O2 >N2 >H2 >He. Higher the value of critical temperature easier is liquefaction.
Define vapour pressure.
In closed container when rate of evaporation is equal to the rate of condensation (equilibrium condition) the pressure exerted by the vapour of the liquid to the walls of the container is called vapour pressure.
What the boiling terms represent for liquid?
The condition of free vaporization throughout the liquid is called boiling.
Define boiling point?
The temperature at which vapour pressure of liquid is equal to the external pressure is called boiling point at that pressure.
What is normal boiling point.
The temperature at which vapour pressure of liquid is equal to the 1atm is called normal boiling point.
What is standard boiling point.
The temperature at which vapour pressure of liquid is equal to the 1 bar is called standard boiling point
Standard boiling point(99oC) of the liquid is slightly lower than the normal boiling point(100oC) why?
Since 1 bar pressure is slightly less than 1 atm pressure hence standard boiling point of the liquid is slightly lower than the normal boiling point.
The pressure cooker is used for cooking food on hills.
Since atmospheric pressure at high altitude is lower than sea level hence liquid boil at lower temperature on high altitude.Thus pressure cooker is used for cooking food on hills
Define Surface tension.
Surface tension is the force acting per unit length perpendicular to the line drawn on the surface of liquid.
What is the S.I unit of surface tension.
Nm-1
What is the cause of surface tension?
In the bulk of liquid, each molecule is pulled equally in every direction due to presence of similar molecules in all directions, resulting in a net force of zero.The molecules at the surface do not have other molecules on all sides, thus net attractive force is towards the interior of the liquid.
Liquid droplets are perfectly spherical in the gravity free environments why?
Surface tension is responsible for the shape of liquid droplets. It is the surface tension due to which liquids tend to have minimum surface area.
Why do liquid tends to rise (or fall) in the capillary.?
Surface tension gives stretching property to the surface of liquid hence liquid tends to rise (or fall) in the capillary.
Why do particles of soil at the bottom of river remain separated but they stick together when taken out?
It is due to surface tension which reduces surface area of thin film of water on particles of soil.
Heating makes sharp glass edges smooth why?
On heating, the glass melts and the surface of the liquid tends to take the rounded shape at the edges due to surface tension which makes the edges smooth.
Define viscosity.
Internal friction between layers of fluid which resist the flow if liquid is called viscosity.
Define laminar flow.
The flow of a viscous liquid on which particles of the fluid move in parallel layers and there is regular gradation of velocity in passing from one layer to the next layer is called laminar flow.
Obtain a formula for determining viscous force between two layers
Define coefficient of viscosity.
Viscosity coefficient is the viscous force when velocity gradient is unity and the area of contact is also unity.
Describe factors effecting viscosity.
(i)Temperature-higher the temperature lesser the viscosity ,this slower the motion.
(ii) Intermolecular force of attraction.- Hydrogen bonding and van der Waals forces are responsible for viscosity.
Windowpanes of old buildings are thicker at the bottom than at the top why?
The attractive force of attraction which holds the atoms together in a molecule called chemical bond.
What is octet rule?
Atoms can combine either by transfer of valance electrons from one atom to another or by sharing of valance electrons in order to have an octet in their valance shell.
How does ionic bond form explain with example?
Ionic bond forms by transfer of valance electron from one atom to another.
Example:Formation of NaCl
(i)Loss of an electron by Na metal and formation of cations.
(ii) Gain of an electron by Cl nonmetal and formation of anions
(iii)The negative anion and positive cation forms NaCl by electrostatic force of attraction.
Define electrovalent bond.
The electrostatic force of attraction between cation and anion which holds the ion together in a molecule is called electrovalent bond.
Define electrovalency.
The number of electrons loss or gain by elements in particular ionic compound is called electrovalency.
Example:NaCl
Electrovalencey of Na =1
Electrovalencey of Cl =1
Explain formation of covalent bond with example.
Covalent bond forms by sharing of electrons between two atoms.
Example formation of H2O.
formation of covalent bonds
Define covalency.
Number of electrons an atom can share with other atoms is called covalency.
Example: In H2O Covalency of H=1,and covalency of oxygen is 2
The drawback of octet theory.
(i) Octet rule is based upon the chemical interest of noble gases .However some noble gases also form some compound.
(ii)Octet rule does not explain the shape of the molecules.
(iii) It does not explain the relative stability of the molecules.
Describe the factors for formation of ionic compound.
(i) Metals which form cation must have low ionization enthalpy.
(ii)Non metals which form anion must have high electron gain enthalpy.
(iii)Lattice enthalpy of ionic compound should be high.
Define lattice enthalpy.
Energy require to separate one mole of a solid ionic compound into gaseous constituent ions is called lattice enthalpy.
Define bond order.
Number of bonds between two atoms in a molecule is called bond order.
Define bond length.
Bond length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule.
Define covalent radius.
The covalent radius is half of the distance between two similar atoms joined by a covalent bond in a same molecule.
Define van der Waals radius.
Half of the distance between two similar atoms in a separate molecule.
Define bond angle.
Angle between the orbitals containing bonding electron pairs around the central atom in molecule.
Define Bond enthalpy.
Energy required to break one mole of bonds of a particular type between two atoms in a gaseous state.
Define average bond enthalpy.
Average bond enthalpy in a polyatomic molecule can be obtained by dividing total bond dissociation enthalpy by number of bonds broken.
Explain formal charge.
Formal charge =[Total number of valence electrons in the free atom]-[Total number of nonbonding electrons( lone pair electrons)]-1/2[Total number of bonding (shared electrons) electrons]
What are the limitations of octet rule?
(i)The incomplete octet of the central atom.
Incomplete octet
(ii )Odd –electron molecules.
Odd –electron molecules.
(iii)The expanded octet
The expanded octet
Draw Lewis dot structure of following compound.
C2H4
Lewis dot structure of ethene
N2
Lewis dot structure of Nitroogen
H2
Lewis dot structure of Hydrogen
O2
Lewis dot structure of oxygen
O3
Lewis dot structure of ozone
NF3
Lewis dot structure of NF3
CO32-
Lewis dot structure of carbonate ion
H2O
Lewis dot structure of water
CCl4
Lewis dot structure of carbon tetrachloride
PCl5
Lewis dot structure of phosphorus pentachloride
SF6
Lewis dot structure of sulphur hexachloride
H2SO4
Lewis dot structure of sulphuric acid
CH4
Lewis dot structure of methane
NH4+
Lewis dot structure of ammonium ion
SO2
Lewis dot structure of sulphur dioxide
NH3
Lewis dot structure of Ammonia
SF4
Lewis dot structure of sulphur tetrafluooride
ClF3
Lewis dot structure of ClF3
BrF5
Lewis dot structure of BrF5
XeF4
Lewis dot structure of XeF4
NO3–
Lewis dot structure of Nitrate ion
What is Fajan’s Rule?
The partial covalent character of ionic bond was discussed by Fajan’s,by following rules.
(i)The smaller the size of cation and the larger the size of anion the greater the covalent character of particular ionic bond.
(ii)The greater the charge on the cation the greater the covalent character if the ionic bond.
Explain VBT with example.
(i)Covalent bond between two atom forms by overlapping of atomic orbitals.
(ii)Only valence shell orbitals can take part in bonding and must have opposite spin.
(iii)Greater the overlap the stronger is the bond formed between two atoms.
Example : Formation of H2 molecule.
H2
Electronic configuration of one H atom 1s1
Shape of s orbital
Overlapping
Thus H-H forms, having one bond
Explain the formation of following molecules on the basis of VBT.
N2
Electronic configuration of one N atom 1s2,2s2,2p3
Shape of three p orbitals
Overlapping
O2
Electronic configuration of one O atom 1s2,2s2,2p4
Shape of two p orbitals
Overlapping
Thus O=O forms having one sigma and one pi bond.
Mention two types of overlapping and nature of covalent bond.
There are two types of bond formed during overlapping of atomic orbitals.
(i)Sigma bond( Covalent bond which forms by head on overlapping along the intermolecular axis.
Following types of overlapping can form sigma bond
(i)s-s overlapping.
(ii)s-p overlapping.
(iii)p-p overlapping.
(iii )Pi Bond( Covalent bond forms by side on overlapping .
Following types of overlapping can form pi bond.
(iii)p-p overlapping.
Which bond is stronger in sigma and Pi bond?
Sigma bond is stronger bond than Pi bond because in case of sigma bond the overlapping of orbitals takes place to a larger extent.
What are the main postulates of VSEPR theory.
(i) Electron pair Geometry of the molecule is decided by both bond pair and lone pair orbitals ,but Molecular geometry of the molecules is decided by only bond pair orbitals.
(iii)The repulsive interaction of electrons pair decreases in the following order.
What are the steps to find shape of the molecules according to VSEPR theory.
Example: NH3
Step (i)Write Lewis dot structure of NH3.
Step(ii)Chose central atom here it is N
Step(iii)Find bond pair orbitals and lone pair orbitals
bp=3,lp=1
Step(iv)Find the sum of lp and bp
Here lp+bp=4
Now structure is determined by following table.
lp+bp
Structure (Electron pair Geometry)
2
Linear
3
Trigonal planer
4
Tetrahedral
5
Triangular bipyrmidal
6
Square bipyramidal
7
Pentagonal bipyarnidal
VSEPR Theory
Determine the shape of following molecules.
NH3
Central atom=N
l.p=1
b.p=3
l.p+b.p =4
Electron pair Geometry=Tetrahedral
Shape of NH3
Molecular Geometry= Pyramidal
H2O
Central atom=O
l.p=2
b.p=2
l.p+b.p =4
Electron pair Geometry=Tetrahedral
shape of water
Molecular Geometry=Bent/Angular
CCl4
Central atom=c
l.p=0
b.p=4
l.p+b.p =4
Electron pair Geometry/Molecular Geometry =Tetrahedral
BeH2
Central atom=Be
l.p=0
b.p=2
l.p+b.p =2
Electron pair Geometry /Molecular Geometry=Linear
BF3
Central atom=B
l.p=0
b.p=3
l.p+b.p =3
Electron pair Geometry/Molecular geometry=Trigonal planer
PF5
Central atom=P
l.p=O
b.p=5
l.p+b.p =5
Electron pair Geometry/Molecular geometry=Trigonal bipyramidal
SF6
Central atom=S
l.p=0
b.p=6
l.p+b.p =6
Electron pair Geometry/Molecular geometry=Square bipyramidal
CH4
Central atom=C
l.p=O
b.p=4
l.p+b.p =4
Electron pair Geometry/Molecular geometry=Tetrahedral
SF4
Central atom=S
l.p=1
b.p=4
l.p+b.p =5
Electron pair Geometry=Trigonal bipyramidal
Molecular geometry=See-Saw
ClF3
Central atom=Cl
l.p=2
b.p=3
l.p+b.p =5
Electron pair geometry=Trigonal bipyramidal
Molecular Geometry=T shape
XeF4
Central atom=Xe
l.p=2
b.p=4
l.p+b.p =6
Structure=Square bipyramidal
Shape-Square Planar
Define Dipole moment.
The Product of the magnitude of charge and the distance between the centers of the two positive and negative and nonpolar covalent co charge is called dipole moment.
What is Debye(D)
Debye is unit of dipole moment.
What is polar covalent compound and nonpolar covalent compound?
Polar covalent Compound-Compounds having net dipole moment not equal to zero are polar covalent compound.
Non polar covalent Compound-Compounds having net dipole moment equal to zero are nonpolar covalent compound
Chose the polar and nonpolar compound in following compounds.
NF3
Dipole moment of NF3
H2O
Dipole moment of H2O
BeF2
Dipole moment of BeF2
BF3
CH4
Dipole moment of CH4
NH3
Dipole moment of NH3
In NH3 and NF3 which has higher dipole moment?
Dipole moment of NH3 is higher than that of NF3. In NH3 the orbital dipole due to lone pair is in same direction as the resultant dipole moment of N-H bonds ,whereas in NF3 the orbital dipole is in the direction opposite to the resultant dipole moment of three N-F bonds.
Explain the term resonance in a covalent compound.
When a particular compound can be draw in more than two Lewis dot structure in which atomic positions are remain constant and differ on only arrangement of electrons.
Draw the resonating structure of following compounds.
CO2
Resonance Of CO2
O3
Resonance of Ozone
SO3
Resonance of carbonate ion SO3
NO3–
Resonance of nitrate ion
CO32-
Resonance of carbonate ion
Define hybridisation.
The process of intermixing of the orbitals having slightly different energies and formation of new type of orbitals called hybrid orbitals of equivalent energies and shape.
What are the different types of hybridization and their structures.
Hybridisation
Structure
sp
Linear
Sp2
Triagonal planer
Sp3
Tetrahedreal
Sp3d
Triangular bipyramidal
Sp3d2
Square bipyramidal
Using hybridization theory show the formation of following molecules.
H2O
Central atom is Oxygen
Electronic configuration in ground state
Hybridisation among 2s and 2p orbitals
Hybridisation is sp3
Orientation of four sp3 hybrid orbitals in space
Overlapping between hybrid orbital and s orbital of hydrogen and formation of two bonds.
Hybridisation of H2O
BCl3
Central atom is B
Electronic configuration in ground state
Electronic configuration in excited state
Hybridisation among 2s and 2p orbitals
Hybridisation is sp2
Orientation of three sp2 hybrid orbitals in space
Overlapping between hybrid orbital and p orbital of Chlorine and formation of three bonds.
Hybridisationn of BCl3
PCl5
Central atom is P
Electronic configuration in ground state
Electronic configuration in excited state
Hybridisation among 3s ,3p and 3d orbitals
Orientation of five sp3d hybrid orbitals in space
Overlapping between hybrid orbital and p orbitals of chlorine and formation of five bonds.
Hybridisationn of PCl5
SF6
Central atom is S
Electronic configuration in ground state
Electronic configuration in excited state
Hybridisation among 3s ,3p and 3d orbitals
Orientation of five sp3d2 hybrid orbitals in space
Overlapping between hybrid orbital and p orbitals of fluorine and formation of six bonds.
Hybridisation of SF6
CH4
Central atom is Carbon
Electronic configuration in ground state
Electronic configuration in excited state
Hybridisation among 2s and 2p orbitals
Hybridisation is sp3
Orientation of four sp3 hybrid orbitals in space
Overlapping between hybrid orbital and s orbitals of hydrogen and formation of four bonds.
Hybridisation of CH4
NH3
Central atom is N
Electronic configuration in ground state
Hybridisation among 2s and 2p orbitals
Hybridisation is sp3
Orientation of four sp3 hybrid orbitals in space
Overlapping between hybrid orbital and s orbitals of hydrogen and formation of three bonds.
Hybridisation Of NH3
C2H4
Central atoms are two carbon atoms.
Electronic configuration in ground state
Electronic configuration in excited state
Hybridisation among 2s and 2p orbitals
Hybridisation is sp2
Orientation of three sp2 hybrid orbitals in space
Similarly of other carbon atoms
Overlapping
Hybridisation of ethene
C2H2
Central atoms are two carbon atoms.
Electronic configuration in ground state
Electronic configuration in excited state
Hybridisation among 2s and 2p orbitals
Hybridisation is sp
Orientation of two sp hybrid orbitals in space
Similarly of other carbon atoms
Overlapping
What are the main features of molecular orbital theory?
(i)Like atoms are present in the atomic orbitals ,electrons in a molecule are present in a molecular orbitals.
(ii)The atomic orbitals of proper orientation combine to form molecular orbitals.
(iii)Like electron in a atomic orbitals are influenced by one nucleus ,in molecular orbital it is influenced by two or more nuclei.
(iv)The number of molecular orbitals formed are equal to the number of combining orbitals half are called bonding orbitals and another half are called antibonding orbitals.
(v)Bonding molecular orbitals have greater stability because they have lower energy than that of antibonding orbitals.
(vi)Molecualr orbitals are filled according to the aufbau principle,Pauli’s ,Hund’s rule.
What is LACO theory.?
LCAO-linear combination of atomic orbitals.
Let a molecule contain two atoms A and B .Formation of molecular orbitals are due to linear combination of atomic orbitals that can take place by addition and by subtraction of wave functions.
LCAO
What is the sequence of energy levels of molecular orbitals.
Define bond order.
Bond order is defined as half of the difference between the number of electrons present in bonding and antibonding orbitals.
Explain stability of compound using bond order.
Stable molecule-Nb>Na
Unstable molecule- Nb<Na, Nb=Na
When is compound diamagnetic and paramagnetic?
(i)If all the molecular orbitals in a molecule are doubly occupied the substance is diamagnetic and repelled by magnetic field.
(ii) If one or more molecular orbitals are singly occupied the substance is paramagnetic and attracted by magnetic field.
Discuss the bonding of H2, He2, Li2, C2, O2 O2+,O2– on MOT
Draw molecular orbital diagram of ,H2,He2,O2.
H2
Molecular orbital diagram of H2,
He2
Molecular orbital diagram of He2,
O2
Molecular orbital diagram of O2
Define H-bonding.
The attractive force which binds hydrogen atom of one molecule with the F,O,N of another molecule.
Define the two types of H-bonding.
(a)Intermolecular H-bonding.
(b)Intramolecular H-bonding.
(a)Intermolecular H-bonding.
It forms between two different molecules of the same or different compounds.
Example:H-F.H2O etc
(b)Intramolecular H-bonding.
It forms when H-atom and F,O,N is present on the same molecule.
Classification of elements and periodicity in properties
Why do you need to classify elements?
It is very difficult to study chemistry of all elements individually ,because number of elements are very large.
Explain Dobereiner ‘s triads rule.
Dobereiner arranged three elements in increasing order of their atomic weight called triads ,He observed that the middle elements of each triads had an atomic weight equal to the arithmetic mean of the other two.
Example.Li(7) ,Na(23),K(39)
What was the major drawback of Dobereiner ‘s triads rule?
The triads rule valid for limited number of elements.
Explain Newlands Law of octaves.
He arranged the elements in increasing order of their atomic weights and noticed that every eight element have properties similar to the first element.
What was the major drawback of Newland’s octave rule?
The Law was true only for elements up to Ca.
Explain Mendeleev periodic law.
The properties of the elements are periodic function of their atomic weight.
What was the major drawback of Mendeleev periodic table?
Iodine with lower atomic weight than that of tellurium was placed in group VII along with F,Cl.Br.I.
Give special property of Mendeleev periodic table which makes it very special.
He proposed that some of the elements were still discovered and left the several gaps in the table. For example both Galium and Germaimum were unknown at the time Mendeleev published his periodic table. He left the gap under Al and Si and called these elements Eka Aluminium and Eka silicon.
What do you mean by Eka Aluminium and Eka silicon?
He left the gap under Al and Si and called these elements Eka Aluminium(Galium) and Eka silicon(Germanium).
What is modern periodic law?
The properties of the elements are periodic function of their atomic number.
What do you mean by periods and groups in long form of periodic table.
The Horizontal rows are called periods and vertical columns are called groups.
How many groups and periods are in long form of periodic table?
There are eighteen groups (1,2,3,4,5,6,7,8,9,10,11,12,13,14,15,16,17,18) and seven periods (1,2,3,4,5,6,7)
What is the IUPAC name of elements having atomic number 101 and 102?
101-Unnilunbium.
102-unnilnium.
4f and 5f inner transition series of elements are placed in the separately in the periodic table why?
To maintain the structure and preserve the principle of classification by keeping elements with similar properties in a single column.
Why the elements of same group exhibit similar chemical behavior?
Because these elements have the same distribution of electrons in their outermost shell.
Long form of periodic table is classified into four blocks name that groups.
1.S-block elements.
2.p-block elements.
3.d-block elements.
4.f-block elements.
What is the general electronic configuration of the last shell of
1.S-block elements 2.p-block elements 3.d-block elements 4.f-block elements.
1.S-block elements-[ns1-2]
2.p-block elements-[ns1-2np1-6]
3.d-block elements[(n-1)d1-10ns0-2]
4.f-block elements[(n-2)f1-14(n-1)d0-1ns2]
Define following terms.
1.Alkali metals.2.Alakaline earth metals.3.Representative elements or Main group elements.4.Noble gases.5.Halogens.6.Chalogens.7.Transition elements.8.inner transition elements.
1.Alkali metals.-Group 1 elements.
2.Alakaline earth metals. Group 2 metals.
3.Representative elements or Main group elements.-S-block elements and p-block elements.
4.Noble gases.Group 18 elements.
5.Halogens-Group 17 elements.
6.Chalogens-Group 16 elements.
7.Transition elements-d-block elements.
8.inner transition elements-f-block elements.
Why are group 1 elements called alkali metals?
When alkali metals are dissolved in water they form hydroxide which are basic in nature.
Why are group 2 elements called alkaline earth metals.
Hydroxide and oxide of group 2 elements are basic in nature and they are found in earth crust.
Group 18 elements are called Noble gases why?
All the orbitals in the valence shell of the Noble gases are completely filled by electrons and it is very difficult to remove electrons.
d-block elements are called Transition metals why?
d-block elements forma bridge between the chemically active metals of s-block elements and less reactive elements of group 13 and group 14.
What are Lanthanoids and Actinoids.
Lanthanoids-14 elements after La are called Lanthanoids.Last electron goes into 4f orbital.
Electronic configuration is 4f1-145d0-16s2.
Actinoids. 14 elements after Ac are called Actinoids.Last electron goes into 5f orbital.
Electronic configuration is 5f1-146d0-17s2.
What are transuranium elements?
The elements after Uranium are called transuranium elements.
Properties of s-block elements.
1.s-block elements are metals.
2.They are high reactive metals.
3.They have low ionization energy.
4.They lose the outermost electrons readily to form +1 ions(Group 1 elements)+and +2 ion (group2 elements.)
5.The compounds of the s-block elements are ionic except Li and Be.
Properties of the p-block elements.
1.p-block elements contain both metals and non metals.
2.Group 17 and group 16 elements are mainly nonmetals.
3.Group 17 and and Group 16 elements have very high negative electron gain enthalpy.
4.Group 18 elements are noble gases.
Properties of d-block elements.
1.They all are metals.
2.They form colored ions .
3.They have variable oxidation states.
4.They are paramagnetic.
5.They are good catalyst.
Zn,Hg,Cd do not show most of the properties of transition metals due to completely filled d –orbitals.
Properties of f-block elements.
1.They all are metals.
2.With in each series properties of the elements are quite similar.
3.The chemistry of actinoids are more complicated than the corresponding lanthanoids.
4.Actionoids are radioactive elements.
How many types of elements are in periodic table.
There are two types of elements.
1. Metals-78% of all elements are metals. They are located in left side of periodic table.
2. Non metals. They are located in at the top right side of the periodic table.
What are the properties of metals?
1.Metals are usually solid at room temperature.
2.They are good conductors.
3.They have high melting and boiling point.
4.They are malleable and ductile.
What are the properties of non metals?
1.Usually solids or gas at room temperature.
2.Low melting point and boiling point.
Boron and carbons are exceptions.
3.They are poor conductors of heat and electricity.
4.Brittle and are neither malleable and ductile.
5.Metallic character increases top to bottom and decreases left to right.
6.Non metallic decreases top to bottom and increases left to right.
What are semi metals and metalloids?
Elements which show properties of both metals and non metals.
eg. Silicon ,Germanium,Arsenic ,Antimony and Tellurium.
What do you mean by covalent radius and metallic radius.
Covalent radius-It gives the size of an atom of non metal. It is the half of the distance between two atoms when they are bound together by a single bond.
Metallic radius-Half of the distance between two adjacent kernels present at the lattice site in crystal.
Atomic size of elements increases top to bottom why?
As we go down the group number of shell per period increases thus size increases.
Atomic size of elements decreases from left to right in period why?
As we go from left to right in period shell number remains constant and electrons are added to same shell thus attraction between nucleus and outermost electron increases thus size decreases.
The size of anion is larger than its parent atom why?
It is due to addition of one or more electron would result in increased repulsion among the electrons and decrease in effective nuclear charge.
The size of cation is smaller than its parent atom why?
It is due to removal of one or more electron would result in decreased repulsion among the electrons and increases in effective nuclear charge.
The cation with the greater positive charge will have a smaller radius why?
It is due to greater attraction of the electron to the nucleus.
The anion with the greater –ve charge will have large radius why?
It is due to the net repulsion of the electron will more outweigh the nuclear charge and the will expand in size.
Define ionization enthalpy.
The energy required to the remove an electron from an isolated gaseous atom.
Define first Ionization enthalpy and second ionization enthalpy.
The first ionization enthalpy for an element X is the enthalpy change for the reaction.
The second ionization enthalpy for an elements X is the enthalpy change for the reaction.
Why Ionization enthalpies are always positive?
Because energy is always required to remove electron from an atom.
Why second ionization is always higher than the first ionization enthalpy.
It is more difficult to remove an electron from an positively charged ion than from neutral atom.
What is shielding effect or screening effect?
The effective nuclear charge experienced by an valance electron in an atom will be less than the actual charge on nucleus because of the repulsion between the core electrons and valance electrons.
Why 1st ionization enthalpy of elements decreases top to bottom in group why?
It is due to atomic size increases top to bottom in group.
Ionization enthalpy of elements increases in period from left to right why?
Atomic size of elements decreases from left to right in period.
The first ionization enthalpy of boron(B) is slightly higher than Beryllium(Be) why?
In Be electron is removed from fully filled (stable) 2s orbital and in boron the electron is removed from partially filled(less stable) 2p orbital.
The 1st ionization enthalpy of oxygen (o) is lower than that of Nitrogen(N) why?
It is easier to remove a electron from partially filled (2p4)orbital in case of oxygen than that of half orbital (2p3) orbital in case of nitrogen.
What are isoelectronic species?
Atoms and ions which contain the same number of electron is called isoelectronic dpecies.
e.g.O2- =10 electron.F–=10 electron.
Elctron gain enthalpy increases from left to right up to group 17 why?
It is due to the size of atom decreases from left to right in period.
Why negative electron gain enthalpy decreases from top to bottom in group why?
It is due to the size of atom increases from top to bottom in group.
Which oxidation state is shown by group 1 and group 2 elements?
Group 1 elements have only one valance electron hence they show +1 oxidation state Group 2 elements have two valance electron hence they show +2 oxidation state .
Periods 2 elements (top elements from group 13-17) show anomalous behavior why?
I t is due to small size ,high electronegativity and unavailability of d orbitals.
What do you mean by diagonal relationship?
Properties of 2 elements are similar with properties of period 3 elements diagonally.
The 1st member of group 13-17 show anomalous behavior why?
1st member of group has only four valance orbitals one 2s and three 2p orbitals ,whereas the other members have d orbitals hence they can expand their covalency.
B can only form BF4– whereas Al Can form AlF63- why?
It has due to B has only four valance orbitals but Al has d-orbitals thus it can expand its covalency.
The first member of p-block elements display greater ability to form multiple bonds to itself and to the other elements (like C=C,C=N,N=O) but other elements not why?
It is due to first member of p-block element has following properties.
1.Small size.
2.large charge/radius ratio.
3.High electronegativity.
Group 1 elements and halogens are highly reactive why?
Group1 elements have only one valance electron on their last shell hence they can lose electron easily and thus can form cation easily.
Halogens are seven electron on their last shell hence they need only on electron to complete their octet thus can form anion easily.
bond by carbon atoms it show allotrope.
of graphite is taken as zero why?
Sarvan sir- Chemistry For All 