Case Study Questions Class 11 Chemistry Chapter 3 Classification of Elements and Periodicity in Properties
CBSE Class 11 Case Study Questions Chemistry Classification of Elements and Periodicity in Properties. Important Case Study Questions for Class 11 Board Exam Students. Here we have arranged some Important Case Base Questions for students who are searching for Paragraph Based Questions Classification of Elements and Periodicity in Properties.
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CBSE Case Study Questions Class 11 Chemistry Classification of Elements and Periodicity in Properties
Case –I
[A] MCQ
1)… was the first to consider the idea of trends among properties of elements.
a) Johann Dobereiner
b) John Alexander Newlands
c) Demitri Mendeleev
d) Henry Moseley
Ans- a) Johann Dobereiner
2) Which of the following system of classifying elements given by Newlands ?
a) Octaves
b) Triads
c) Periods
d) Group
Ans- b) Triads
3) Which of the following system given by Dobereiner ?
a) Triads
b) Periods
c) Octaves
d) Group
Ans- c) Octaves
4) …credited with the development of the Periodic Table.
a) Johann Dobereiner
b) John Alexander Newlands
c) Henry Moseley
d) Demitri Mendeleev.
Ans- d) Demitri Mendeleev
5) Which of the following element Predicted by Mendeleev ?
a) Aluminium
b) Silicon
c) Germanium
d) Oxygen
Ans – c) germanium
[B] Short Answers
1) What is periodic Law of Elements ?
Ans- It states as The properties of the elements are aperiodic function of their atomic weights.
2) Explain – Dobereiners Triads
Ans- Johann Dobereiner in early 1800’s was the first to consider the idea of trends among properties of elements. By 1829 he noted a similarity among the physical and chemical properties of several groups of three elements (Triads). In each case, he noticed that the middle element of each of the Triads had an atomic weight about half way between the atomic weights of the other two.
3) What is Newlands Octaves ?
Ans – John Alexander Newlands arranged the elements in increasing order of their atomic weights and noted that every eighth element had properties similar to the first element. The relationship was just like every eighth note that resembles the first in octaves of music.
[C] Long Answers
1) How were the elements arranged in Dmitri Mendeleev’s periodic table?
Ans- Mendeleev arranged elements in horizontal rows and vertical columns of a table in order of their increasing atomic weights in such away that the elements with similar properties occupied the same vertical column or group. In particular, Mendeleev relied on the similarities in the empirical formulas and properties of the compounds formed by the elements. He realized that some of the elements did not fit in with his scheme of classification if the order of atomic weight was strictly followed. He ignored the order of atomic weights, thinking that the atomic measurements might be incorrect, and placed the elements with similar properties together. For example, iodine with lower atomic weight Than that of tellurium (Group VI) was placed In Group VII along with fluorine, chlorine, Bromine because of similarities in properties.
2) Why did Mendeleev leave gaps in his periodic table ?
Ans – To keep the Primary aim of arranging the elements of Similar properties in the same group, he Proposed that some of the elements were still Undiscovered and, therefore, left several gaps in the table. For example, both gallium and Germanium were unknown at the time Mendeleev published his Periodic Table. He left The gap under Aluminium and a gap under Silicon, and called these elements Eka-Aluminium and Eka-Silicon. Mendeleev Predicted not only the existence of gallium and Germanium, but also described some of their General physical properties. These elements Were discovered later.
Case –II
We must bear in mind that when Mendeleev developed his Periodic Table, chemists knew nothing about the internal structure of atom. However, the beginning of the 20th century witnessed profound developments in theories about sub-atomic particles. In 1913, the English physicist, Henry Moseley observed regularities in the characteristic X-ray spectra of the elements. A plot of ν (whereν is frequency of X-rays emitted) against atomic number (Z ) gave a straight line and not the plot of ν vs atomic mass. He thereby showed that the atomic number is a more fundamental property of an element than its atomic mass. Mendeleev’s Periodic Law was, therefore, accordingly modified. This is known as the Modern Periodic Law and can be stated as : The physical and chemical properties of the elements are periodic functions of their atomic numbers.Numerous forms of Periodic Table have been devised from time to time. Some forms emphasise chemical reactions and valence, whereas others stress the electronic configuration of elements. A modern version, the so-called “long form” of the Periodic Table of the elements , is the most convenient and widely used. The horizontal rows (which Mendeleev called series) are called periods and the vertical columns, groups. Elements having similar outer electronic configurations in their atoms are arranged in vertical columns, referred to as groups or families. According to the recommendation of International Union of Pure and Applied Chemistry (IUPAC), the groups are numbered from 1 to 18 replacing the older notation of groups IA … VIIA, VIII, IB … VIIB and 0. There are altogether seven periods. The period number corresponds to the highest principal quantum number (n) of the elements in the period. The first period contains 2 elements. The subsequent periods consists of 8, 8, 18, 18 and 32 elements, respectively. The seventh period is incomplete and like the sixth period would have a theoretical maximum (on the basis of quantum numbers) of 32 elements. In this form of the Periodic Table, 14 elements of both sixth and seventh periods (lanthanoids and actinoids, respectively) are placed in separate panels at the bottom. the IUPAC has made recommendation that until a new element’s discovery is proved, and its name is officially recognised, a systematic nomenclature be derived directly from the atomic number of the element using the numerical roots for 0 and numbers 1-9. The roots are put together in order of digits which make up the atomic number and “ium” is added at the end.Groupwise Electronic Configurations Elements in the same vertical column or group have similar valence shell electronic configurations, the same number of electrons in the outer orbitals, and similar properties. theoretical foundation for the periodic classification. The elements in a vertical column of the Periodic Table constitute a group or family and exhibit similar chemical behaviour. This similarity arises because these elements have the same number and same distribution of electrons in their outermost orbitals. We can classify the elements into four blocks viz., s-block, p-block, d-block and f-block depending on the type of atomic orbitals that are being filled with electrons. Two exceptions to this categorisation. Strictly, helium belongs to the s-block but its positioning in the p-block along with other group 18 elements is justified because it has a completely filled valence shell (1s) and as a result, exhibits properties characteristic of other noble gases. The other exception is hydrogen. It has only one s-electron and hence can be placed in group 1 (alkali metals). It can also gain an electron to achieve a noble gas arrangement and hence it can behave similar to a group 17 (halogen family) elements. Because it is a special case, we shall place hydrogen separately at the top of the Periodic Table.
[A] MCQ
1) In 1913, the English physicist, ….observed regularities in the characteristic X-ray spectra of the elements.
a) Johann Dobereiner
b) John Alexander Newlands
c) Demitri Mendeleev
d) Henry Moseley
Ans- d) Henry Moseley
2) Horizontal row in periodic table called ..
a) Group
b) Period
c) Triad
d) Octave
Ans- b) Period
3) Vertical Column in periodic table called ..
a) Group
b) Period
c) Triad
d) Octave
Ans- a) Group
4) According to Modern Periodic Law the physical and chemical properties of the elements are periodic functions of their ….
a) Atomic mass
b) Atomic numbers
c) Atomic structure
d) Atomic size
Ans- b) Atomic numbers
5) What is IUPAC name of element having atomic number 107.
a) Unnilpentium
b) Unnilhexium
c) Unnilseptium
d) Unniloctium
Ans- c) Unnilseptium
[B]Short Answers
1) How did Moseley interpret the result of X-ray spectra experiment?
Ans- Henry Moseley observed regularities in the characteristic X-ray spectra of the elements. A plot of ν (whereν is frequency of X-rays emitted) against atomic number (Z ) gave a straight line and not the plot of ν vs atomic mass. He thereby showed that the atomic number is a more fundamental property of an element than its atomic mass.
2) How to name elements according to IUPAC ?
Ans- The IUPAC has made recommendation that until a new element’s discovery is proved, and its name is officially recognised, a systematic nomenclature be derived directly from the atomic number of the element using the numerical roots for 0 and numbers 1-9. The roots are put together in order of digits which make up the atomic number and “ium” is added at the end.
3) Explain Groupwise electronic configuration in periodic table .
Ans- Groupwise Electronic Configurations Elements in the same vertical column or group have similar valence shell electronic configurations, the same number of electrons in the outer orbitals, and similar properties. theoretical foundation for the periodic classification. The elements in a vertical column of the Periodic Table constitute a group or family and exhibit similar chemical behaviour. This similarity arises because these elements have the same number and same distribution of electrons in their outermost orbitals.
[C] Long Answers
1) Modern periodic table is classified in how many blocks? What are the exception ?
Ans- The elements classify into four blocks viz., s-block, p-block, d-block and f-block depending on the type of atomic orbitals that are being filled with electrons. Two exceptions to this categorisation. Strictly, helium belongs to the s-block but its positioning in the p-block along with other group 18 elements is justified because it has a completely filled valence shell (1s) and as a result, exhibits properties characteristic of other noble gases. The other exception is hydrogen. It has only one s-electron and hence can be placed in group 1 (alkali metals). It can also gain an electron to achieve a noble gas arrangement and hence it can behave similar to a group 17 (halogen family) elements. Because it is a special case, we shall place hydrogen separately at the top of the Periodic Table.
2) How modern period table is arranged ?
Ans-A modern version, the so-called “long form” of the Periodic Table of the elements , is the most convenient and widely used. The horizontal rows (which Mendeleev called series) are called periods and the vertical columns, groups. Elements having similar outer electronic configurations in their atoms are arranged in vertical columns, referred to as groups or families. According to the recommendation of International Union of Pure and Applied Chemistry (IUPAC), the groups are numbered from 1 to 18 replacing the older notation of groups IA … VIIA, VIII, IB … VIIB and 0. There are altogether seven periods. The period number corresponds to the highest principal quantum number (n) of the elements in the period. The first period contains 2 elements. The subsequent periods consists of 8, 8, 18, 18 and 32 elements, respectively. The seventh period is incomplete and like the sixth period would have a theoretical maximum (on the basis of quantum numbers) of 32 elements. In this form of the Periodic Table, 14 elements of both sixth and seventh periods (lanthanoids and actinoids, respectively) are placed in separate panels at the bottom.
Case –III
The s-Block Elements The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) which have ns1and ns2 outermost electronic configuration belong to the s-Block Elements. They are all reactive metals with low ionization enthalpies. They lose the outermost electron(s) readily to form 1+ ion (in the case of alkali metals) or 2+ ion (in the case of alkaline earth metals). The metallic character and the reactivity increase as we go down the group. Because of high reactivity they are never found pure in nature. The compounds of the s-block elements, with the exception of those of lithium and beryllium are predominantly ionic. The p-Block Elements comprise those belonging to Group 13 to 18 and these together with the s-Block Elements are called the Representative Elements or Main Group Elements. The outermost electronic configuration varies from ns2np1 to ns2np6 in each period. At the end of each period is a noble gas element with a closed valence shell ns2np6 configuration. All the orbitals in the valence shell of the noble gases are completely filled by electrons and it is very difficult to alter this stable arrangement by the addition or removal of electrons. The noble gases thus exhibit very low chemical reactivity. Preceding the noble gas family are two chemically important groups of non-metals. They are the halogens (Group 17) and the chalcogens (Group 16).The non-metallic character increases as we move from left to right across a period and metallic character increases as we go down the group. These are the elements of Group 3 to 12 in the centre of the Periodic Table. These are characterised by the filling of inner d orbitals by electrons and are therefore referred to as d-Block Elements. These elements have the general outer electronic configuration (n-1)d1-10ns0-2 . They are all metals. They mostly form coloured ions, exhibit variable valence (oxidation states), paramagnetism and oftenly used as catalysts. However, Zn, Cd and Hg which have the electronic configuration, (n-1) d10ns2 do not show most of the properties of transition elements. In a way, transition metals form a bridge between the chemically active metals of s-block elements and the less active elements of Groups 13 and 14 and thus take their familiar name “Transition Elements”.The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids, Ce(Z = 58) – Lu(Z = 71) and Actinoids, Th(Z = 90) – Lr (Z = 103) are characterised by the outer electronic configuration (n-2)f 1-14 (n-1)d 0–1ns2 . The last electron added to each element is filled in f- orbital. These two series of elements are hence called the Inner- Transition Elements (f-Block Elements). They are all metals. Within each series, the properties of the elements are quite similar. The chemistry of the early actinoids is more complicated than the corresponding lanthanoids, due to the large number of oxidation states possible for these actinoid elements. Actinoid elements are radioactive. Many of the actinoid elements have been made only in nanogram quantities or even less by nuclear reactions and their chemistry is not fully studied. The elements after uranium are called Transuranium Elements. The elements can be divided into Metals and Non-Metals. In contrast, non-metals are located at the top right hand side of the Periodic Table. The elements become more metallic as we go down a group; the non- metallic character increases as one goes from left to right across the Periodic Table. Periodic Table show properties that are characteristic of both metals and non- metals. These elements are called Semi-metals or Metalloids.
[A] MCQ
1) Alkali metal and alkaline earth metal belongs to ..
a) S – block
b) P – block
c) D – block
d) F – block
Ans -S – block
2) The metallic character and the reactivity … as we go down the group.
a) Decreases
b) Increases
c) Remains Constant
d) None of Above
Ans – b) Increases
3) Group … Elements known as chalcogens.
a) 12
b) 14
c) 16
d) 18
Ans – c) 16
4) Elements Ce(Z = 58) to Lu(Z = 71) are known as ..
a) Halogens
b) Chalcogens
c) Actinoids
d) Lanthenoids
Ans- d) Lanthenoids
5) The elements after uranium are called … Elements.
a) Halogens
b) Chalcogens
c) Actinoids
d) Transuranium
Ans- d) Transuranium
[B] Short Answers
1) What is metalloid ?
Ans- The elements become more metallic as we go down a group; the non- metallic character increases as one goes from left to right across the Periodic Table. Periodic Table show properties that are characteristic of both metals and non- metals. These elements are called Semi-metals or Metalloids.
2) What are S-block elements ?
Ans- The s-Block Elements The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) which have ns1and ns2 outermost electronic configuration belong to the s-Block Elements. They are all reactive metals with low ionization enthalpies. They lose the outermost electron(s) readily to form 1+ ion (in the case of alkali metals) or 2+ ion (in the case of alkaline earth metals).
3) Give characteristics of nobel gas elements .
Ans- The end of each period is a noble gas element with a closed valence shell ns2np6 configuration. All the orbitals in the valence shell of the noble gases are completely filled by electrons and it is very difficult to alter this stable arrangement by the addition or removal of electrons. The noble gases thus exhibit very low chemical reactivity.
[C] Long Answers
1) Describe in brief – f block elements .
Ans- The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids, Ce(Z = 58) – Lu(Z = 71) and Actinoids, Th(Z = 90) – Lr (Z = 103) are characterised by the outer electronic configuration (n-2)f 1-14 (n-1)d 0–1ns2 . The last electron added to each element is filled in f- orbital. These two series of elements are hence called the Inner- Transition Elements (f-Block Elements). They are all metals. Within each series, the properties of the elements are quite similar. The chemistry of the early actinoids is more complicated than the corresponding lanthanoids, due to the large number of oxidation states possible for these actinoid elements. Actinoid elements are radioactive. Many of the actinoid elements have been made only in nanogram quantities or even less by nuclear reactions and their chemistry is not fully studied. The elements after uranium are called Transuranium Elements.
2) What are d-block elements ? Give Characteristics of d-block elements .
Ans-The elements of Group 3 to 12 in the centre of the Periodic Table. These are characterised by the filling of inner d orbitals by electrons and are therefore referred to as d-Block Elements. These elements have the general outer electronic configuration (n-1)d1-10ns0-2 . They are all metals. They mostly form coloured ions, exhibit variable valence (oxidation states), paramagnetism and oftenly used as catalysts. However, Zn, Cd and Hg which have the electronic configuration, (n-1) d10ns2 do not show most of the properties of transition elements. In a way, transition metals form a bridge between the chemically active metals of s-block elements and the less active elements of Groups 13 and 14 and thus take their familiar name “Transition Elements”.
Case –IV
There are many observable patterns in thephysical and chemical properties of elementsas we descend in a group or move across aperiod in the Periodic Table.Atomic Radius the determination of the atomic sizecannot be precise. In other words, there is no practical way by which the size of an individualatom can be measured. However, an estimateof the atomic size can be made by knowing thedistance between the atoms in the combinedstate. One practical approach to estimate thesize of an atom of a non-metallic element is tomeasure the distance between two atoms whenthey are bound together by a single bond in acovalent molecule and from this value, the“Covalent Radius” For metals, we define theterm “Metallic Radius” which is taken as halfthe internuclear distance separating the metalcores in the metallic crystal. Atomic Radius to refer to both covalent ormetallic radius depending on whether theelement is a non-metal or a metal. Atomic radiican be measured by X-ray or otherspectroscopic methods. The atomic size generallydecreases across a period. It is because within the period the outerelectrons are in the same valence shell and theeffective nuclear charge increases as the atomicnumber increases resulting in the increasedattraction of electrons to the nucleus.Note that the atomic radii of noble gasesAre not considered here. Being monoatomic,Their (non-bonded radii) values are very large.In fact radii of noble gases should be comparednot with the covalent radii but with the van derWaals radii of other elements. The removal of an electron from an atom resultsin the formation of a cation, whereas gain ofan electron leads to an anion. The ionic radiican be estimated by measuring the distancesbetween cations and anions in ionic crystals.In general, the ionic radii of elements exhibitthe same trend as the atomic radii. A cation issmaller than its parent atom because it hasfewer electrons while its nuclear charge remainsthe same. The size of an anion will be largerthan that of the parent atom because theaddition of one or more electrons would resultin increased repulsion among the electronsand a decrease in effective nuclear charge. When we find some atoms and ions whichcontain the same number of electrons, we callthem isoelectronic species. For example,O2–, F–, Na+ and Mg2+ have the same number ofelectrons (10). Their radii would be differentbecause of their different nuclear charges.A quantitative measure of the tendency of anelement to lose electron is given by itsIonization Enthalpy. It represents the energyrequired to remove an electron from an isolatedgaseous atom (X) in its ground state. The ionization enthalpy is expressed inunits of kJ mol–1. We can define the secondionization enthalpy as the energy required toremove the second most loosely boundelectron The first ionization enthalpies of elementshaving atomic numbers up to 60 are plotted then The periodicity of the graph is quitestriking. You will find maxima at the noble gaseswhich have closed electron shells and verystable electron configurations. On the otherhand, minima occur at the alkali metals andtheir low ionization enthalpies can be correlated with their high reactivity. In addition, you willnotice two trends the first ionization enthalpygenerally increases as we go across a periodand decreases as we descend in a group. Electron Gain Enthalpy. when an electron is added to a neutral gaseousatom (x) to convert it into a negative ion, theenthalpy change accompanying the process isdefined as the electron gain enthalpy (∆egh).Electron gain enthalpy provides a measure ofthe ease with which an atom adds an electronto form anion. electron gain enthalpies have largenegative values toward the upper right of theperiodic table preceding the noble gases.The variation in electron gain enthalpies ofelements is less systematic than for ionizationenthalpies. As a general rule, electron gainenthalpy becomes more negative with increasein the atomic number across a period. Theeffective nuclear charge increases from left toright across a period and consequently it willbe easier to add an electron to a smaller atomsince the added electron on an average wouldbe closer to the positively charged nucleus. ElectronegativityA qualitative measure of the ability of an atomin a chemical compound to attract sharedelectrons to itself is called electronegativity.Unlike ionization enthalpy and electron gainenthalpy, it is not a measureable quantity.However, a number of numerical scales ofelectronegativity of elements viz., Pauling scale,Mulliken-Jaffe scale, Allred-Rochow scale havebeen developed. The one which is the most widely used is the Pauling scale. Electronegativity generallyincreases across a period from leftto right (say from lithium tofluorine) and decrease down a group(say from fluorine to astatine) inthe periodic table. Non-metallic elements have strong tendencyto gain electrons. Therefore, electronegativityis directly related to that non-metallicproperties of elements. It can be furtherextended to say that the electronegativity isinversely related to the metallic properties of elements. Thus, the increase inelectronegativities across a period isaccompanied by an increase in non-metallicproperties (or decrease in metallic properties)of elements. Similarly, the decrease inelectronegativity down a group is accompanied by a decrease in non-metallic properties (orincrease in metallic properties) of elements.
[A] MCQ
1) The atomic size generally … across a period.
a) Increases
b) Decreases
c) Remains Constant
d) None of above
Ans- b) decreases
2) The ionization enthalpy is expressed in units of ….
a) kJ mol–1
b) mole kJ-1
c) mole kJ
d) kJ mol-1
Ans- a)kJ mol–1
3) Which of the following is/are numerical scales of electronegativity of elements .a) Pauling scale
b) Mulliken-Jaffe scale
c) Allred-Rochow scale
d) All the above
Ans- d) All the above
4) The … in electronegativity down a group is accompanied by a … in non-metallic properties.
a) Increase , Decrease
b) Decrease , Increase
c) Decrease , Decrease
d) Increase , Increase
Ans- c) Decrease , Decrease
5) Electronegativity generally … across a period from left to right and … down a group in the periodic table.
a) Increase , Decrease
b) Decrease , Increase
c) Decrease , Decrease
d) Increase , Increase
Ans- a) Increase , Decrease
[B] Short Answers
1) Explain the terms – a) Anion b) Cation
Ans – The removal of an electron from an atom results in the formation of a cation, whereas gain of an electron leads to an anion.
2) What are isoelectronic species ? Give an example .
Ans- Some atoms and ions which contain the same number of electrons, we call them isoelectronic species. For example, O 2–, F– , Na+ and Mg2+ have the same number of electrons (10).
3) Define – Electron Gain Enthalpy.
Ans –When an electron is added to a neutral gaseous atom (x) to convert it into a negative ion, the enthalpy change accompanying the process is defined as the electron gain enthalpy (∆egh).
[C] Long Answers
1) Explain Periodic trends in atomic raddi.
Ans- An estimate of the atomic size can be made by knowing the distance between the atoms in the combined state. One practical approach to estimate the size of an atom of a non-metallic element is to measure the distance between two atoms when they are bound together by a single bond in a covalent molecule and from this value, the “Covalent Radius” For metals, we define the term “Metallic Radius” which is taken as half the internuclear distance separating the metal cores in the metallic crystal. Atomic Radius to refer to both covalent or metallic radius depending on whether the element is a non-metal or a metal. Atomic radii can be measured by X-ray or other spectroscopic methods. The atomic size generally decreases across a period. It is because within the period the outer electrons are in the same valence shell and the effective nuclear charge increases as the atomic number increases resulting in the increased attraction of electrons to the nucleus.
2) Describe periodic trends in electronegativity.
Ans-A qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself is called electronegativity. Unlike ionization enthalpy and electron gain enthalpy, it is not a measureable quantity. However, a number of numerical scales of electronegativity of elements viz., Pauling scale, Mulliken-Jaffe scale, All red-Rochow scale have been developed. The one which is the most widely used is the Pauling scale. Electronegativity generally increases across a period from left to right (say from lithium to fluorine) and decrease down a group (say from fluorine to astatine) in the periodic table. Non-metallic elements have strong tendency to gain electrons. Therefore, electronegativity is directly related to that non-metallic properties of elements.