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Some Basic Concepts of Chemistry — Chemistry STD 11 Science — Question

Rajasthan BoardEnglish MediumSTD 11 ScienceChemistrySome Basic Concepts of Chemistry4 Marks
Question
The uncertainty in the experimental or the calculated values is indicated by mentioning the number of significant figures. Significant figures are meaningful digits which are known with certainty plus one which is estimated or uncertain. The uncertainty is indicated by writing the certain digits and the last uncertain digit. there are certain rules for determining the Number of significant figures. These are Stated below:

  • All non-zero digits are significant. For Example in 285cm, there are three Significant figures and in 0.25 mL, there are two significant figures.
  • Zeros preceding to first non-zero digit are not significant. such zero indicates the position of decimal point. thus, 0.03 has one significant figure and 0.0052 has two significant figures.
  • Zeros between two non-zero digits are significant. thus, 2.005 has four Significant figures.
  • Zeros at the end or right of a number are significant, provided they are on the right side of the decimal point. For example, 0.200 g has three significant figures. But, if otherwise, the terminal zeros are not significant if there is no decimal point.
Precision refers to the closeness of various measurements for the same quantity. However, accuracy is the agreement of a particular valueto the true value of the result.

LAWS OF CHEMICALCOMBINATIONS- The combination of elements to form compounds is governed by the following five basic laws.

  1. Law of Conservation of Mass-This law was put forth by Antoine Lavoisierin 1789. He performed careful experimental studies for combustion reactions and reached to the conclusion that in all physical andchemical changes, there is no net change inmassduring the process. Hence, he reachedto the conclusion that matter can neither becreated nor destroyed. This is called ‘Law ofConservation of Mass’.
  2. Law of Definite Proportions-This law was given by, a French chemist, Joseph Proust. He stated that a given compound always contains exactly the same proportion of elements by weight.
  3. Law of Multiple Proportions-This law was proposed by John Dalton. According to this law, if two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element, are in the ratio of small whole numbers. For example, hydrogen combines with oxygen to form two compounds, namely, water and hydrogen peroxide.

Hydrogen + Oxygen→ Water

2g                 16g       18g

Hydrogen + Oxygen → Hydrogen Peroxide

    2g                  32g        34g

Here, the masses of oxygen (i.e., 16 g and 32 g), which combine with a fixed mass of hydrogen (2g) bear a simple ratio, i.e., 16:32 or 1:2.

  1. Gay Lussac’s Law of Gaseous Volumes-This law was given by Gay Lussac in 1808. Heobserved that when gases combine or are produced in a chemicalreaction they do so in asimple ratio by volume,provided all gases are at the same temperature and pressure.
  2. Avogadro’s Law – In 1811, Avogadro proposed that equal volumes of all gases at the same temperature and pressure should contain equal number of molecules.
In 1808, Dalton published ‘A New System of Chemical Philosophy’, in which he proposed the following :

  1. Matter consists of indivisible atoms.
  2. All atoms of a given element have identical properties, including identical mass. Atoms of different elements differ in mass.
  3. Compounds are formed when atoms of different elements combine in a fixed ratio.
  4. Chemical reactions involve reorganisati on of atoms. These are neither created nor destroyed in a chemical reaction.
  1. … refers to the closeness of variousmeasurements for the same quantity.
  1. Accuracy
  2. Reliability
  3. Precision
  4. Uncertainty
  1. Law of Conservation of mass was put forth by ….in 1789.
  1. Joseph Proust
  2. Antoine Lavoisier
  3. Joseph Louis
  4. Gay Lussac
  1. Which of the following number has twosignificant figures.
  1. 0.0052
  2. 052
  3. 52
  4. 0052
  1. … is the agreement of a particular valueto the true value of the result.
  1. Accuracy
  2. Reliability
  3. Precision
  4. Uncertainty
  1. Law of Multiple Proportions proposed by....
  1. Joseph Proust
  2. Antoine Lavoisier
  3. Joseph Louis
  4. John Dalton
✓

Answer

  1. (c) Precision
  1. (b) Antoine Lavoisier
  1. (a) 0.0052
  1. (a) Accuracy
  1. (d) John Dalton

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The existing large number of organic compounds and their ever-increasing numbers has made it necessary to classify them on the basis of their structures. Organic compounds are broadly classified as open-chain compounds which are also called aliphatic compounds. Aliphatic compounds further classified as homocyclic and heterocyclic compounds. Aromatic compounds are special types of compounds. Alicyclic compounds, aromatic compounds may also have heteroatom in the ring. Such compounds are called heterocyclic aromatic compounds. Organic compounds can also be classified on the basis of functional groups, into families or homologous series. The members of a homologous series can be represented by general molecular formula and the successive members differ from each other in a molecular formula by a $- CH _2$ unit.

1. The successive members of a homologous series differ by which mass of amu? (1)
2. Does Pyridine, pyrrole, thiophene are all heteroaromatic compounds (1)
3. Difference between heterocyclic and homocyclic compound. (2)
OR
Is tetrahydrofuran is aromatic compounds? (2)
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The existing large number of organic compounds and their ever-increasing numbers has made it necessary to classify them on the basis of their structures. Organic compounds are broadly classified as open-chain compounds which are also called aliphatic compounds. Aliphatic compounds further classified as homocyclic and heterocyclic compounds. Aromatic compounds are special types of compounds. Alicyclic compounds, aromatic compounds may also have heteroatom in the ring. Such compounds are called heterocyclic aromatic compounds. Organic compounds can also be classified on the basis of functional groups, into families or homologous series. The members of a homologous series can be represented by general molecular formula and the successive members differ from each other in a molecular formula by a $- CH _2$ unit.

1. The successive members of a homologous series differ by which mass of amu?
OR
Is tetrahydrofuran is aromatic compounds?
2. Does Pyridine, pyrrole, thiophene are all heteroaromatic compounds
3. Difference between heterocyclic and homocyclic compound.
→
The molecular orbital theory is based on the principle of a linear combination of atomic orbitals. According to this approach when atomic orbitals of the atoms come closer, they undergo constructive interference as well as destructive interference giving molecular orbitals, i.e., two atomic orbitals overlap to form two molecular orbitals, one of which lies at a lower energy level (bonding molecular orbital). Each molecular orbital can hold one or two electrons in accordance with Pauli's exclusion principle and Hund's rule of maximum multiplicity. For molecules up to $N _2$, the order of filling of orbitals is:
Image
Bond order $=\frac{1}{2}$ [bonding electrons - antibonding electrons]
Bond order gives the following information:
i. If bond order is greater than zero, the molecule/ion exists otherwise not.
ii. Higher the bond order, higher is the bond dissociation energy.
iii. Higher the bond order, greater is the bond stability.
iv. Higher the bond order, shorter is the bond length.

1. Arrange the following negative stabilities of $CN , CN ^{+}$and $CN ^{-}$in increasing order of bond.
2. The molecular orbital theory is preferred over valence bond theory. Why?
3. Ethyne is acidic in nature in comparison to ethene and ethane. Why is it so?
OR
Bonding molecular orbital is lowered by a greater amount of energy than the amount by which antibonding molecular orbital is raised. Is this statement correct?
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Lewis dot structures, in general, do notrepresent the actual shapes of the molecules.In case of polyatomic ions, the net charge ispossessed by the ion as a whole and not by aparticular atom. It is, however, feasible toassign a formal charge on each atom. Theformal charge of an atom in a polyatomicmolecule or ion may be defined as thedifference between the number of valenceelectrons of that atom in an isolated or freestate and the number of electrons assignedto that atom in the Lewis structure. It isexpressed as :Generally the lowest energystructure is the one with the smallestformal charges on the atoms. The formalcharge is a factor based on a pure covalentview of bonding in which electron pairsare shared equally by neighbouring atoms. The octet rule, though useful, is not universal.It is quite useful for understanding thestructures of most of the organic compoundsand it applies mainly to the second periodelements of the periodic table. There are threetypes of exceptions to the octet rule.

  • The incomplete octet of the central atom
  • Odd-electron molecules
  • The expanded octetFrom the Kössel and Lewis treatment of theformation of an ionic bond, it follows that theformation of ionic compounds wouldprimarily depend upon:
  • The ease of formation of the positive andnegative ions from the respective neutralatoms;
  • The arrangement of the positive andnegative ions in the solid, that is, thelattice of the crystalline compound.

The Lattice Enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions. For example, the lattice enthalpy of NaCl is 788 kJ mol–1. This means that 788 kJ of energy is required to separate one mole of solid NaCl into one mole of Na+  (g) and one mole of Cl– (g) to an infinite distance.Bond length is defined as the equilibriumdistance between the nuclei of two bondedatoms in a molecule. Bond lengths aremeasured by spectroscopic, X-ray diffractionand electron-diffraction techniques . The covalent radius is measuredapproximately as the radius of an atom’score which is in contact with the core ofan adjacent atom in a bonded situation.  The vander Waals radius represents the overall sizeof the atom which includes its valence shellin a nonbonded situation.  Bond Angle is defined as the angle between the orbitalscontaining bonding electron pairs around thecentral atom in a molecule/complex ion. Bondangle is expressed in degree which can beexperimentally determined by spectroscopicmethods. It gives some idea regarding thedistribution of orbitals around the centralatom in a molecule/complex ion and hence ithelps us in determining its shape. Forexample H–O–H bond angle in water can berepresented as under:

Bond Enthalpy It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state. The unit of bond enthalpy is kJ mol–1. For example, the H–H bond enthalpy in hydrogen molecule is 435.8 kJ mol–1 . H2 (g) → H(g) + H(g); ∆aH  = 435.8 kJ mol–1. Bond OrderIn the Lewis description of covalent bond,the Bond Order is given by the number ofbonds between the two atoms in amolecule. The bond order, for example in H2(with a single shared electron pair), in O2(with two shared electron pairs) and in N2(with three shared electron pairs) is 1,2,3respectively. A general correlation useful forunderstanding the stablities of moleculesis that: with increase in bond order, bondenthalpy increases and bond lengthdecreases. The concept of resonance was introducedto deal with the type of difficulty experiencedin the depiction of accurate structures ofmolecules like O3. According to the conceptof resonance, whenever a single Lewis structure cannot describe a moleculeaccurately, a number of structures with similar energy, positions of nuclei, bonding and non-bonding pairs of electrons are taken as the canonical structures of the hybrid which describes the molecule accurately.

Thus for O3, the two structures shown above constitute the canonical structures or resonance structures and their hybrid i.e., theIII structure represents the structure of O3more accurately. This is also called resonance hybrid. Resonance is represented by a double headed arrow. In general, it may be stated that

  • Resonance stabilizes the molecule as the energy of the resonance hybrid is lessthan the energy of any single cannonical structure; and,
  • Resonance averages the bond characteristics as a whole. Thus the energy of theO3resonancehybrid is lower than either of the two cannonical froms I and II.
  1. Which of the following techniques used to measure bond length?
  1. Spectroscopic techniques
  2. X-ray diffraction
  3. Electron-diffraction techniques
  4. All the above
  1. The unit of bond enthalpy is …
  1. kJ mol–1
  2. Cal mol-1
  3. Cal mol
  4.  kJ mol
  1. With increase in bond order, bond enthalpy… and bond length ….
  1. Decreases, decreases
  2. Increases, decreases
  3. Increases, increases
  4. Decreases, increases
  1. The …. is measured approximately as the radius of an atom’s core which is in contact with the core of an adjacent atom in a bonded situation.
  1. Ionic radius
  2. Metallic radius
  3. Covalent radius
  4. None of above
  1. … is given by the number of bonds between the two atoms in a molecule.
  1. Bond Order
  2. Bond size
  3. Bond enthalpy
  4. Bond angle
→
A less obvious example of electron transfer is realised when hydrogen combines with oxygen to form water by the reaction:

$2\text{H}_2(\text{g}) + \text{O}_2 (\text{g}) → 2\text{H}_2\text{O} (\text{l})$

Though not simple in its approach, yet we can visualise the H atom as going from a Neutral (zero) state in H2 to a positive state in H2O, the O atom goes from a zero state in O2 To a dinegative state in H2O. It is assumed that There is an electron transfer from H to O and Consequently H2 is oxidised and O2 is reduced. However, as we shall see later, the charge Transfer is only partial and is perhaps better Described as an electron shift rather than a Complete loss of electron by H and gain by O. Two examples of this class Of the reactions are:

$\text{H}_2 (\text{s}) + \text{Cl}_2(\text{g}) → 2\text{HCl} (\text{g})$ And,

$\text{CH}_4 (\text{g}) + 4\text{Cl}_2 (\text{g}) → \text{CCl}_4(\text{l}) + 4\text{HCl (g)}$

In order to keep track of electron shifts in Chemical reactions involving formation of Covalent compounds, a more practical method Of using oxidation number has been Developed. In this method, it is always Assumed that there is a complete transfer of Electron from a less electronegative atom to a More electonegative atom. For example, we Rewrite equations to show Charge on each of the atoms forming part of The reaction:

It may be emphasised that the assumption Of electron transfer is made for book-keeping Purpose only and it will become obvious at a Later stage in this unit that it leads to the simple Description of redox reactions. Oxidation number denotes the Oxidation state of an element in a Compound ascertained according to a set Of rules formulated on the basis that electron pair in a covalent bond belongs Entirely to more electronegative element.  It is not always possible to remember or Make out easily in a compound/ ion, which Element is more electronegative than the other. Therefore, a set of rules has been formulated To determine the oxidation number of an Element in a compound/ ion. We may at this stage, state the rules for the Calculation of oxidation number. These rules are:

1.) In elements, in the free or the uncombined State, each atom bears an oxidation Number of zero. Evidently each atom in H2, O2, Cl2, O3, P4, S8, Na, Mg, Al has the Oxidation number zero.

2.) For ions composed of only one atom, the Oxidation number is equal to the charge On the ion. Thus Na+ Ion has an oxidation Number of +1, Mg2+ ion +2 , Fe3+ ion, +3, Cl –  Ion, –1, O2 – ion, –2; and so on. In their Compounds all alkali metals have Oxidation number of +1, and all alkaline Earth metals have an oxidation number of +2. Aluminium is regarded to have an Oxidation number of +3 in all its Compounds.

3.) The oxidation number of oxygen in most Compounds is –2. However, we come across Two kinds of exceptions here. One arises In the case of peroxides and superoxides, The compounds of oxygen in which oxygen Atoms are directly linked to each other. While in peroxides (e.g., H2 O2, Na2 O2), each Oxygen atom is assigned an oxidation Number of –1, in superoxides (e.g., K O2, Rb O2) each oxygen atom is assigned an Oxidation number of –(½). The second Exception appears rarely, i.e. when oxygen Is bonded to fluorine. In such compounds e.g., oxygen difluoride (OF2) and dioxygen difluoride (O2F2), the oxygen is assigned an oxidation number of +2 and +1, respectively. The number assigned to oxygen will depend upon the bonding state of oxygen but this number would now be a positive figure only.

4.) The oxidation number of hydrogen is +1, Except when it is bonded to metals in binary Compounds (that is compounds containing Two elements). For example, in LiH, NaH, And Ca H2, its oxidation number is –1.

5.) In all its compounds, fluorine has an Oxidation number of –1. Other halogens (Cl, Br, and I) also have an oxidation number Of –1, when they occur as halide ions in Their compounds. Chlorine, bromine and Iodine when combined with oxygen, for Example in oxoacids and oxoanions, have Positive oxidation numbers.

6.) The algebraic sum of the oxidation number Of all the atoms in a compound must be Zero. In polyatomic ion, the algebraic sum Of all the oxidation numbers of atoms of The ion must equal the charge on the ion. Thus, the sum of oxidation number of three Oxygen atoms and one carbon atom in the Carbonate ion, (CO3) 2– must equal –2.

A term that is often used interchangeably With the oxidation number is the oxidation State. Thus in CO2, the oxidation state of Carbon is +4, that is also its oxidation number And similarly the oxidation state as well as Oxidation number of oxygen is – 2. This implies That the oxidation number denotes the Oxidation state of an element in a compound.

The oxidation number/state of a metal in a Compound is sometimes presented according To the notation given by German chemist, Alfred Stock. It is popularly known as Stock Notation. According to this, the oxidation Number is expressed by putting a Roman Numeral representing the oxidation number In parenthesis after the symbol of the metal in The molecular formula. Thus aurous chloride And auric chloride are written as Au(I)Cl and Au(III)Cl3. Similarly, stannous chloride and Stannic chloride are written as Sn(II) Cl2 and Sn(IV)Cl4. This change in oxidation number Implies change in oxidation state, which in Turn helps to identify whether the species is Present in oxidised form or reduced form. Thus, Hg2(I) Cl2 is the reduced form of Hg(II) Cl2.

  1. H atom goes from a … state in H2 to a positive state in H2O in water formation.
  1. Neutral
  2. Positive
  3. Negative
  4. All the above
  1. In oxidation number method, there is a complete transfer of electron from a …. electronegative atom to a … electonegative atom.
  1. more, less
  2. less, more
  3. non, more
  4. non, less
  1. Oxidation number of Mg2 + ion is:
  1. -2
  2. -1
  3. +2
  4. +1
  1. In Na2O2 each oxygen atom is assigned an oxidation number of …
  1. +1
  2. -2
  3. +2
  4. -1
  1. The algebraic sum of the oxidation number of all the atoms in a compound must be…
  1. 0
  2. 1
  3. 2
  4. -2

 

→

The orbital wave function or $\psi$ for an electronin an atom has no physical meaning. It issimply a mathematical function of thecoordinates of the electron. However, fordifferent orbitals the plots of correspondingwave functions as a function of r (the distancefrom the nucleus) are different. According to the German physicist, MaxBorn, the square of the wave function(i.e.,$\psi^2$) at a point gives the probability densityof the electron at that point. Boundary surface diagrams of constantprobability density for different orbitals give afairly good representation of the shapes of theorbitals. In this representation, a boundarysurface or contour surface is drawn in spacefor an orbital on which the value of probabilitydensity $\mid\psi\mid2$ is constant. In principle manysuch boundary surfaces may be possible.However, for a given orbital, only thatboundary surface diagram of constantprobability density* is taken to be goodrepresentation of the shape of the orbital whichencloses a region or volume in which theprobability of finding the electron is very high,say, 90%.

In hydrogen atom, electron has the same energy when it is in the2s orbital as when it is present in 2p orbital.The orbitals having the same energy are calleddegenerate. The 1s orbital in a hydrogenatom, as said earlier, corresponds to the moststable condition and is called the ground stateand an electron residing in this orbital is moststrongly held by the nucleus.

An electron inthe 2s, 2p or higher orbitals in a hydrogen atomis in excited state.The filling of electrons into the orbitals ofdifferent atoms takes place according to theaufbau principle which is based on the Pauli’sexclusion principle, the Hund’s rule ofmaximum multiplicity and the relativeenergies of the orbitals. Theaufbausprinciple states : In the ground state of theatoms, the orbitals are filled in order oftheir increasing energies. In other words,electrons first occupy the lowest energy orbitalavailable to them and enter into higher energyorbitals only after the lower energy orbitals arefilled.The number of electrons to be filled in variousorbitals is restricted by the exclusion principle,given by the Austrian scientist Wolfgang Pauli(1926). According to this principle : No twoelectrons in an atom can have the sameset of four quantum numbers. Pauliexclusion principle can also be stated as : “Onlytwo electrons may exist in the same orbitaland these electrons must have oppositespin.” This means that the two electrons canhave the same value of three quantum numbersn, l and ml, but must have the opposite spinquantum number.Hund’s Rule of Maximum Multiplicity rule deals with the filling of electrons into the orbitals belonging to the same subshell. It states : pairing ofelectrons in the orbitals belonging to thesame subshell (p, d or f) does not take placeuntil each orbital belonging to thatsubshell has got one electron each i.e., itis singly occupied.

The distribution of electrons into orbitals of anatom is called its electronic configuration.If one keeps in mind the basic rules whichgovern the filling of different atomic orbitals,the electronic configurations of different atomscan be written very easily.The electronic configuration of differentatoms can be represented in two ways. Forexample :

  1. sa pbdc…… notation
  2. Orbital diagram

  1. …at a point gives the probability density of the electron at that point.

  1. $\psi\times2$

  2. $-\psi^2$

  3. $\psi$

  4. $\psi^2$

  1. Only …. electrons may exist in the same orbital and these electrons must have opposite spin.
  1. One
  2. Two
  3. Three
  4. Four
  1. …deals with the filling of electrons into the orbitals belonging to the same subshell.
  1. Hund’s Rule of Maximum Multiplicity rule
  2. Pauli’s exclusion principle
  3. Aufbau principle
  4. Werner Heisenberg
  1. Electrons first occupy the …. energy orbital available to them and enter into … energy orbitals.
  1. Lowest, Higher
  2. Higher, Lowest
  3. Middle, Higher
  4. Higher, Middle
→
The ionic character of metallic halides tends toward covalent nature as per Fajan's rule. Such covalent halides behave as non-metal in their higher oxidation states. The property to hydrolyse to give oxy-acids of the element and corresponding hydro halogen acid for most non-metallic elements proceeds exceptionally in the way, keeping oxidation number of element and halide sam in oxo-acids.
Non-polar halides are immiscible in water, as they do not show hydrolysis, but halides of some elements with empty d-orbital undergo hydrolysis. Stability of halides of the higher state is governed by the inert-pair effect.

1. How does halide undergo hydrolysis to give oxy-acids of underlined element $PCl _3$ ?
2. Out of $NCl _3$ and $BCl _3$ undergoes hydrolysis to form oxy-acids? Write the chemical reaction for the correct answer.
3. Out of $PbCl _4, PbF _4, PbI _4$ and $PbBr _4$ which one doesn't exist?
OR
Non-Polar halides are immiscible in water. Why?
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Read the passage given below and answer the following questions from (i) to (v).
When covalent bond is formed betweentwo similar atoms, for example in H2,O2, Cl2,N2Or F2, the shared pair of electrons is equally Attracted by the two atoms. As a result electronPair is situated exactly between the twoIdentical nuclei. The bond so formed is calledNonpolar covalent bond. As a result of polarisation, the moleculePossesses the dipole moment which can be defined as the productOf the magnitude of the charge and theDistance between the centres of positive andNegative charge. It is usually designated by aGreek letter ‘µ’. Mathematically, it is expressedAs follows :Dipole moment (µ) = charge (Q) × distance ofSeparationDipole moment is usually expressed inDebye units (D). The conversion factor is1 D = 3.33564×10–30 C mWhere C is coulomb and m is meter. Just as all the covalent bonds haveSome partial ionic character, the ionicBonds also have partial covalentCharacter. The partial covalent character of ionic bonds was discussed by Fajans in terms of the following rules:
  • The smaller the size of the cation and theLarger the size of the anion, the greater theCovalent character of an ionic bond.
  • The greater the charge on the cation, theGreater the covalent character of the ionic bond.
  • For cations of the same size and charge,The one, with electronic configuration(n-1)d0ns0, typical of transition metals, isMore polarising than the one with a nobleGas configuration, ns2 np6, typical of alkali and alkaline earth metal cations.
Sidgwick and Powell in 1940, proposed a simple theoryBased on the repulsive interactions of theElectron pairs in the valence shell of the atoms.It was further developed and redefined byNyholm and Gillespie (1957).The main postulates of VSEPR theory areAs follows:
  • The shape of a molecule depends uponThe number of valence shell electron pairs(bonded or nonbonded) around the centralAtom.
  • Pairs of electrons in the valence shell repelone another since their electron clouds arenegatively charged.
  • These pairs of electrons tend to occupySuch positions in space that minimiseRepulsion and thus maximise distanceBetween them.
  • The valence shell is taken as a sphere withThe electron pairs localising on theSpherical surface at maximum distanceFrom one another.
  • A multiple bond is treated as if it is a singleElectron pair and the two or three electronPairs of a multiple bond are treated as aSingle super pair.
  • Where two or more resonance structuresCan represent a molecule, the VSEPRModel is applicable to any such structure.
The arrangement of electron pairs and the atoms around the central atom can be : linear,Trigonal planar, tetrahedral, trigonal-Bipyramidal and octahedral. Valence bond theory was introduced byHeitler and London (1927) and developedFurther by Pauling and others. A discussionOf the valence bond theory is based on the knowledge of atomic orbitals, electronicConfigurations of elements.partialmerging of atomic orbitals is called overlappingof atomic orbitals which results in the pairingof electrons. The extent of overlap decides thestrength of a covalent bond. according toorbital overlap concept, the formation of acovalent bond between two atoms results bypairing of electrons present in the valence shellhaving opposite spins. When orbitals of two atoms come close to formbond, their overlap may be positive, negativeor zero depending upon the sign anddirection of orientation of amplitude of orbitalwave function in space. Positive andnegative sign on boundary surface diagramsin the show the sign (phase) of orbitalwave function and are not related to charge.Orbitals forming bond should have same sign(phase) and orientation in space. This is calledpositive overlap. The criterion of overlap, as the main factorfor the formation of covalent bonds appliesuniformly to the homonuclear/heteronucleardiatomic molecules and polyatomic molecules.
  1. Dipole moment is usually expressed in….
  1. Debye
  2. Centimeter
  3. Columbs
  4. Ergs
  1. 1D = .....
  1. 33564×10–28Cm
  2. 3.3564×10–30Cm
  3. 33564×10–32Cm
  4. 33564×10–34Cm
  1. Valence bond theory was introduced by ….
  1. Pauling and lewis
  2. Nyholm and Gillespie
  3. Heitler and London
  4. Sidgwick and Powell
  1. Pair is situated exactly between the two Identical nuclei the bond so formed is called …. covalent bond.
  1. Unipolar
  2. Bipolar
  3. Polar
  4. Nonpolar
  1. Pairs of electrons in the valence shell … one another since their electron clouds are negatively charged.
  1. Attract
  2. Repel
  3. Both a) & b)
  4. None if above
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Read the passage given below and answer the following questions from (i) to (v).
The presence of positive charge on thenucleus is due to the protons in the nucleus.As established earlier, the charge on the proton is equal but opposite to that of electron.Atomic number (Z) = number of protons inthe nucleus of an atom = number of electrons in a nuetral atom. protons and neutrons present in thenucleus are collectively known as nucleons.The total number of nucleons is termed asmass number (A) of the atom.
mass number (A) = number of protons (Z) + number of neutrons (n).
Isobars are the atoms with same massnumber but different atomic number forexample, 64C and 714N. On the other hand, atomswith identical atomic number but differentatomic mass number are known as Isotopes.For example,considering of hydrogen atom again, 99.985% of hydrogen atoms contain only one proton.This isotope is called protium (11H). Rest of thepercentage of hydrogen atom contains two otherisotopes, the one containing 1 proton and 1neutron is called deuterium (21D, 0.015%)and the other one possessing 1 proton and 2neutrons is called tritium (13T )..the studiesof interactions of radiations with matter haveprovided immense information regarding thestructure of atoms and molecules. Neils Bohrutilised these results to improve upon themodel proposed by Rutherford. Twodevelopments played a major role in theformulation of Bohr’s model of atom. Thesewere:
  1. Dual character of the electromagneticradiation which means that radiations possess both wave like and particle likeproperties, and
  2. Experimental results regarding atomicspectra.
James Maxwell (1870) was the first to givea comprehensive explanation about theinteraction between the charged bodies andthe behaviour of electrical and magnetic fieldson macroscopic level. He suggested that whenelectrically charged particle moves underaccelaration, alternating electrical and magnetic fields are produced and transmitted.These fields are transmitted in the forms ofwaves called electromagnetic waves orelectromagnetic radiation.radiations are characterised by theproperties, namely, frequency (ν) and wavelength $(\lambda).$The SI unit for frequency (ν) is hertz(Hz, s–1), after Heinrich Hertz. It is defined asthe number of waves that pass a given pointin one second.Wavelength should have the units of lengthand as you know that the SI units of length ismeter (m). Since electromagnetic radiationconsists of different kinds of waves of muchsmaller wavelengths, smaller units are used.In vaccum all types of electromagneticradiations, regardless of wavelength, travel atthe same speed, i.e., 3.0 × 108m s–1 (2.997925× 108 ms–1, to be precise). This is called speedof light and is given the symbol ‘c‘. Thefrequency (ν), wavelength $(\lambda)$ and velocity of light(c) are related by the following equation .
$\text{c}=\text{v}\lambda$
The other commonly used quantityspecially in spectroscopy, is the wavenumber.It is defined as the number of wavelengthsper unit length. Its units are reciprocal ofwavelength unit, i.e., m–1. However commonlyused unit is cm–1
  1. The presence of positive charge on the nucleus is due to the …. in the nucleus.
  1. Protons
  2. Neutrons
  3. Electron
  4. Nucleons
  1. Atomic Number is denoted by:
  1. A
  2. Z
  3. N
  4. M
  1. Atomic Mass number is denoted by:
  1. M
  2. Z
  3. N
  4. A
  1. … are the atoms with same mass number but different atomic number.
  1. Isotopes
  2. Allotropes
  3. Isobars
  4. None of above
  1. Atoms with identical atomic number but different atomic mass number are known as ..
  1. Isotopes
  2. Allotropes
  3. Isobars
  4. None of above
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Read the passage given below and answer the following questions from (i) to (v).
When a liquid evaporates in a closed container, molecules with relatively higher kinetic energy escape the liquid surface into the vapour phase and number of liquid molecules from the vapour phase strike the liquid surface and are retained in the liquid phase. It gives rise to a constant vapour pressure because of an equilibrium in which the number of molecules leaving the liquid equals the number returning to liquid from the vapour. We say that the system has reached equilibrium state at this stage. However, this is not static equilibrium and there is a lot of activity at the boundary between the liquid and the vapour. Thus, at equilibrium, the rate of evaporation is equal to the rate of condensation. It may be represented by
$\text{H}_2\text{O}_{(\text{l})}\rightleftharpoons\text{H}_2\text{O}_{(\text{vap})}$
The double half arrows indicate that the processes in both the directions are going on simultaneously. The mixture of reactants and products in the equilibrium state is called an equilibrium mixture.
Equilibrium can be established for both physical processes and chemical reactions. The reaction may be fast or slow depending on the experimental conditions and the nature of the reactants. When the reactants in a closed vessel at a particular temperature react to give products, the concentrations of the reactants keep on decreasing, while those of products keep on increasing for some time after which there is no change in the concentrations of either of the reactants or products. This stage of the system is the dynamic equilibrium
The chemical equilibrium may be classified in three groups.
  1. The reactions that proceed nearly to completion and only negligible concentrations of the reactants are left. In some cases, it may not be even possible to detect these experimentally.
  2. The reactions in which only small amounts of products are formed and most of the reactants remain unchanged at equilibrium stage.
  3. The reactions in which the concentrations of the reactants and products are comparable, when the system is in equilibrium.
The equilibrium involving ions in aqueous solutions which is called as ionic equilibrium.
Solid-Liquid Equilibrium Ice and water kept in a perfectly insulated thermos flask (no exchange of heat between its contents and the surroundings) at 273K and the atmospheric pressure are in equilibrium state and the system shows interesting characteristic features. We observe that the mass of ice and water do not change with time and the temperature remains constant. However, the equilibrium is not static. The intense activity can be noticed at the boundary between ice and water. Molecules from the liquid water collide against ice and adhere to it and some molecules of ice escape into liquid phase. There is no change of mass of ice and water, as the rates of transfer of molecules from ice into water and of reverse transfer from water into ice are equal at atmospheric pressure and 273 K. It is obvious that ice and water are in equilibrium only at particular temperature and pressure. For any pure substance at atmospheric pressure, the temperature at which the solid and liquid phases are at equilibrium is called the normal melting point or normal freezing point of the substance. The system here is in dynamic equilibrium and we can infer the following:
  1. Both the opposing processes occur simultaneously.
  2. Both the processes occur at the same rate so that the amount of ice and water remains constant.
Solid – Vapour Equilibrium Let us now consider the systems where solids sublime to vapour phase. If we place solid iodine in a closed vessel, after sometime the vessel gets filled up with violet vapour and the intensity of colour increases with time. After certain time the intensity of colour becomes constant and at this stage equilibrium is attained. Hence solid iodine sublimes to give iodine vapour and the iodine vapour condenses to give solid iodine. The equilibrium can be represented as,
$\text{l}_2(\text{solid})\rightleftharpoons\text{l}_2(\text{vapour})$
Other examples showing this kind of equilibrium are,
$\text{Camphor}_{(\text{solid})}\rightleftharpoons\text{Camphor}_{(\text{vapour})}$
$\text{NH}_4\text{CI}_{(\text{solid})}\rightleftharpoons\text{NH}_4\text{CI}_{(\text{vapour})}$
The equilibrium Involving Dissolution of Solid in Liquids Only a limited amount of salt or sugar can dissolves in a given amount of water at room temperature. If we make a thick sugar syrup solution by dissolving sugar at a higher temperature, sugar crystals separate out if we cool the syrup to the room temperature. We call it a saturated solution when no more of solute can be dissolved in it at a given temperature. The concentration of the solute in a saturated solution depends upon the temperature. In a saturated solution, a dynamic equilibrium exits between the solute molecules in the solid state and in the solution: Sugar (solution) Sugar (solid), and the rate of dissolution of sugar = rate of crystallisation of sugar. Equality of the two rates and dynamic nature of equilibrium has been confirmed with the help of radioactive sugar. If we drop some radioactive sugar into saturated solution of non-radioactive sugar, then after some time radioactivity is observed both in the solution and in the solid sugar. Initially there were no radioactive sugar molecules in the solution but due to dynamic nature of equilibrium, there is exchange between the radioactive and non-radioactive sugar molecules between the two phases. The ratio of the radioactive to non- radioactive molecules in the solution increases till it attains a constant value.
  1. Which of the following symbol represents equilibrium.
  1. $\rightleftharpoons$
  2. $\leftrightarrows$
  3. $\nLeftrightarrow$
  4. $\uparrow\downarrow$
  1. When there is no change in the concentrations of either of the reactants or products, this stage of the system is the …
  1. Static equilibrium
  2. Dynamic equilibrium
  3. Physical equilibrium
  4. Chemical equilibrium
  1. A … solution means no more of solute can be dissolved in it at a given temperature.
  1. Unsaturated
  2. Supersaturated
  3. Saturated
  4. None of these.
  1. The equilibrium involving ions in aqueous solutions which is called as …
  1. Static equilibrium
  2. Dynamic equilibrium
  3. Physical equilibrium
  4. Ionic equilibrium
  1. The concentration of the solute in a saturated solution depends upon the …
  1. Solvent
  2. Pressure
  3. Temperature
  4. System
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