Periodic Properties of Elements

Periodic Properties

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Periodic Table

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Section 1 Development of Periodic Table 1.1 Dobereigner's law of triads ................................................................................ 3 1.2 Newlans’s Law of Octaves .................................................................................. 4 1.3 Lothar Meyer's atomic volume curve .............................................................. 6 1.4 Mendeleev’s Periodic Law and Periodic Table .............................................. 7 1.5 Modern Classification ........................................................................................ 10 1.5 Classification of Elements ................................................................................ 13

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1 Development of Periodic Table You must have visited a library. There are thousands of books in a large library. In spite of this if you ask for a particular book, the library staff can locate it easily. How is it possible? In library the books are classified into various categories and sub-categories. They are arranged on shelves accordingly. Therefore location of books becomes easy. Electrons are filled in various shells and subshells in a fairly regular fashion. Therefore, properties of elements are repeated periodically. Such trends in their physical and chemical properties were noticed by chemists in the nineteenth century and attempts were made to classify elements on their basis long before structure of atom was known.

1.1 Dobereigner's law of triads A German Chemist, Dobereigner, in 1829, reported a significant relation be-tween atomic weights and properties of elements. He reported that there were several groups of three closely related elements, called triads, in which the atomic weight of the middle element was almost equal to the arithmetic mean of the other two. This relationship of elements was called Dobereigner's law of triads.

Doberigner's laws of triads states that in a chemically similar group of three elements, the atomic weight of the middle element is almost equal to the arithmetic mean of the other two. For example, in the following table in triad I, the atomic weight of the middle element sodium (Na) and the mean of the atomic weights of lithium (Li) and potassium ( K ) is the same. Similarly, in triad II, the atomic weight of strontium (Sr) and the mean of the atomic weights of calcium (Ca) and barium (Ba) are approximately equal. Similarly, sulphur, selenium, tellurium and chlorine, bromine, iodine form such triads. Further, it was shown that the elements could be classified into groups consisting of more than three elements. For example, Fluorine was added to the triad {chlorine, bromine, iodine and magnesium was added to the triad {calcium, strontium, barium.

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Limitations of Dobereigner's law of triads This law suffered from following drawbacks. 1. The law could not be applied to all the chemically similar elements even when they could form triads. For example, in the triad of copper (63.5), silver(108) and gold(197) and also in triad of zinc(65), cadmium(112.5) and mercury (200) were chemically similar, the law failed i.e. the atomic weight of the middle element was not even close to the arithmetic mean of the other two elements . 2. The atomic weights of some elements were faulty. So the validity of the law was questionable.

1.2 Newlands' law of octaves An Englishman, Newlands, in 1864, used the analogy of musical octaves to classify the elements. He showed that when elements were arranged in the ascending order of their atomic weights, the eighth element, like the eighth note in music, resembled the _rst , the ninth element resembled the second and so on. The main idea in his classi_cation was the repetition of properties of the elements after certain interval when elements were arranged in the increasing order of their atomic weights. This was an example of periodicity in properties. This type of periodicity is also observed in nature. Days and nights repeat, sunrise and sunset repeat, seasons of the year repeat, the heights of the tide repeat, swinging pendulum returns to its original position in each swing, as the years pass the fashions repeat, even the history repeats itself. In short, periodicity is the order of nature. Newlands'law of octaves can be stated as follows. If the elements are arranged in the increasing order of atomic weights, the eighth element starting from any given element is a kind of repetition of the first like the eighth note of octave in music. EDUDIGM 1B Panditya Road, Kolkata 29

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Newlands' arrangement of elements was as follows. (1) H (2) Li (3) Be (4) B (5) C (6) N (7) O (8) F (9) Na (10) Mg (11) Al (12) Si (13) P (14) S (15) Cl (16) K (17) Ca (18) Ti (19) Cr (20) Mn (21) Fe (22) Co, Ni (23) Cu | etc. Newlands noticed that elements belonging to the same class usually appeared in the same column. Following table shows Newlands' column wise arrangement of elements. Lihium Li Sodium Na Potassium K

Beryllium Be Magnesium Mg Calcium Ca

Boron B Aluminium Al Titanium Ti

Carbon C Silicon Si Chromium Cr

Nitrogen N Phosphorus P Manganese Mn

Oxygen O Sulphur S Iron Fe

Fluorine F Chlorine Cl Cobalt,Nickle Co,Ni

Limitations of Newlands' law of octaves This law had following limitations. 1. Some of the elements had faulty atomic weights. So they did not _t in the order. For example, the atomic weight of beryllium was initially reported as 14.5. So beryllium was placed after boron. As a result, the properties of beryllium did not match with elements in the same column. Later the atomic weight of beryllium was corrected as 9.4 and its place in the order was also changed. 2. Newlands included some elements in the column that looked out of place due to their properties. For example, Newlands placed B, Al, Ti in the same column, C, Si, Cr in the same column and so on. Here Ti and Cr looked out of place considering the properties of B, Al and C, Si. 3. The law of octaves was found to be applicable to elements with low atomic weights. For example, the law was applicable to B, Al but not to Ti in the same column, it was applicable to C, Si but not to Cr in the same column. 4. The number of elements in each octave was not the same. For example, cobalt and nickel were placed together. This increased the number of elements in that octave. 5. Newlands did no leave gaps for undiscovered elements. For example, ele-ments with atomic weights 44, 68 and 72 were discovered later and were given their proper places in the respective columns after their discovery.

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1.3 Lothar Meyer's atomic volume curve Amidst 1969-70, a German scientist Lothar Meyer made an important contribution to the work of classification of elements. He used a physical property viz. atomic volume for the classification of elements. Lothar Meyer plotted a graph of atomic volume (ordinate) against atomic weight (abscissa). (Today such a graph is plotted for atomic volume against atomic number.) Lothar Meyer's atomic volume curve showed that the elements which occupy similar positions on the curve show similar chemical properties. Lothar Meyer made following observations in his graph. 1. The light alkali metals occupy the crests (except lithium) while the heavy transition elements are found in the troughs of the curves. 2. The halogen elements are on the ascending part and the alkaline earth metals were on the descending part of the curve. 3. The metallic elements are generally placed on the descending part while nonmetallic elements were placed on the ascending part of the curve. 4. The atomic volume generally increases for the elements of the same family with increasing atomic number but this increase in volume becomes less as we approach halogens and for the elements of VIII group, atomic volume was nearly constant. Lothar Meyer plotted other physical properties against atomic weight and showed that there is a periodicity in the other physical properties of elements also. Melting point, boiling point, density, hardness, malleability, ductility, compressibility, conductivity, refractive index etc. were all found to show periodicity with respect to the atomic weights of the elements. Lothar Meyer was always very cautious about his statements. He was more of a critic of his own work rather than an advocate. So, for the same work of classification of elements, Mendeleev received more credit than Lothar Meyer. The modified form of Lothar Meyer's atomic volume curve is shown here. The modification is that atomic number has been used in place of atomic weight.

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Lothar Meyer's atomic volume curve

1.4 MENDELEEV’S PERIODIC LAW AND PERIODIC TABLE Mendeleev’s periodic law Dmitry Mendeleev** a Russian chemist while trying to classify elements discovered that on arranging in the increasing order of atomic mass*, elements with similar chemical properties occurred periodically. In1869, he stated this observation in the following form which is known as Mendeleev’s Periodic Law. A periodic function is the one which repeats itself after a certain interval. Thus, according to the periodic law the chemical and physical properties of elements repeat themselves after certain intervals when they are arranged in the increasing order of their atomic mass. Now we shall learn about the arrangement of elements on the basis of the periodic law. The chemical and physical properties of elements are a periodic function of their atomic EDUDIGM 1B Panditya Road, Kolkata 29

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masses.

A tabular arrangement of the elements based on the periodic law is called periodic table. Mendeleev believed that atomic mass of elements was the most fundamental property and arranged them in its increasing order in horizontal rows till he encountered an element which had properties similar to the first element. He placed this element below the first element and thus started the second row of elements. Proceeding in this manner he could arrange all the known elements according to their properties and thus created the first periodic table.

Mendeleev’s periodic table Main features of Mendeleev’s periodic table Look at the Mendeleev’s periodic table shown in fig.4.2 carefully. What do you observe? Here, elements are arranged in tabular form in rows and columns. Now let us learn more about these rows and columns and the elements present in them. 1. The horizontal rows present in the periodic table are called periods. You can see that there are seven periods in the periodic table. These are numbered from 1 to 7 (Arabic numerals). 2. Properties of elements in a particular period show regular gradation (i.e. increase or decrease) from left to right.

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3. The vertical columns present in it are called groups. You must have noticed that these are nine in number and are numbered from I to VIII and Zero (Roman numerals). 4. Groups I to VII are subdivided into A and B subgroups. Groups Zero and VIII don’t have any subgroups. 5. All the elements in a particular group are chemically similar in nature. They show regular gradation in their physical properties and chemical reactivities. After learning about the main features we shall now learn about the main merits of Mendeleev’s periodic table. Merits of Mendeleev’s periodic classification 1. Classification of all elements Mendeleev’s was the first classification which successfully included all the elements. 2. Prediction of new elements Mendeleev’s periodic table had some blank spaces in it. These vacant spaces were for elements that were yet to be discovered. For example, he proposed the existence of an unknown element that he called eka-aluminium. The element gallium was discovered four years later and its properties matched very closely with the predicted properties of

ekaaluminium. In this section we have learnt about the success of Mendeleev’s periodic classification and also about its merits. Does it mean that this periodic table was perfect? No. Although it was a very successful attempt but it also had some defects in it. Now we shall discuss the defects in this classification. Defects in Mendeleev’s periodic classification In spite of being a historic achievement Mendeleev’s periodic table had some defects in it. The following were the main defects in it: 1. Position of hydrogen Hydrogen resembles alkali metals (forms H+ ion just like Na+ ions) as well as halogens (forms H- ion similar to Cl- ion).Therefore, it could neither be placed with alkali metals (group I ) nor with halogens (group VII ).

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2. Position of isotopes Different isotopes of same elements have different atomic masses, therefore, each one of them should be given a different position in the periodic table. On the other hand, because they are chemically similar, they had to be given same position. 3. Anomalous pairs of elements At certain places, an element of higher atomic mass has been placed before an element of lower atomic mass. For example, Argon (39.91) is placed before potassium (39.1).

1.5 MODERN CLASSIFICATION Henry Moseley, an English physicist discovered in the year 1913 that atomic number, is the most fundamental property of an element and not its atomic mass. Atomic number, (Z), of an element is the number of protons in the nucleus of its atom. The number of electrons in the neutral atom is also equal to its atomic number. This discovery changed the whole perspective about elements and their properties to such an extent that a need was felt to change the periodic law also. Now we shall learn about the changes made in the periodic law. Modern periodic law After discovery of atomic number the periodic law was modified and the new law was based upon atomic numbers in place of atomic masses of elements. The Modern Periodic Law states “The chemical and physical properties of elements are a periodic function of their atomic numbers” After the change in the periodic law many changes were suggested in the periodic table. Now we shall learn about the modern periodic table which finally emerged. Modern periodic table The periodic table based on the modern periodic law is called the Modern Periodic Table. Many versions of this periodic table are in use but the one which is most commonly used is the Long Form of Modern Periodic Table. It is shown in figure.

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Periodic Properties 1 1 1 H 2 3 Li 3 11 Na 4 19 K 5 37 Rb 6 55 Cs 7 87 Fr

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2 4 Be 12 Mg 20 Ca 38 Sr 56 Ba 88 Ra

Lanthanides Actinides

3

4

5

6

7

8

9

10

21 Sc 39 Y 57 La 89 Ac

22 Ti 40 Zr 72 Hf 104 Rf

23 Y 41 Nb 73 Ta 105 Ha

24 Cr 42 Mo 74 W 106 Un

25 Mn 43 Tc 75 Re 107 Un

26 Fe 44 Ru 76 Os 108 Un

27 Co 45 Rh 77 Ir 109 Un

28 Ni 46 Pd 78 Pt 110 Un

59 Pr 91 Pa

60 Nd 92 U

58 Ce 90 Th

61 Pm 93 Np

62 Sm 94 Pu

63 Eu 95 Am

64 Gd 96 Cm

65 Tb 97 Bk

66 Dy 98 Cf

18 13 14 15 16 17 2 He 5 6 7 8 9 10 B C N O F Ne 11 12 13 14 15 16 17 18 Al Si P S Cl Ar 29 30 31 32 33 34 35 36 Cu Zn Ga Ge As Se Br Kr 47 48 49 50 51 52 53 54 Ag Cd In Sn Sb Te I Xe 79 80 81 82 83 84 85 86 Au Hg Ti Pb Bi Po At Rn

67 Ho 99 Ex

68 Er 100 Fm

69 Tm 101 Md

70 Yb 102 No

71 Lu 103 Lr

If you look at the modern periodic table shown in the fig.4.3 you will observe that it is not much different from Mendeleev’s periodic table. Now let us learn the main features of this periodic table. Groups There are 18 vertical columns in the periodic table. Each column is called a group. The groups have been numbered from 1 to 18 (in Arabic numerals) from left to right. Group 1 on extreme left position contains alkali metals (Li, Na, K, Rb, Cs and Fr) and group 18 on extreme right side position contains noble gases (He, Ne, Ar, Kr, Xe and Rn). All elements present in a group have similar electronic configurations and have same number of valence electrons. You can see in case of group 1 (alkali metals) and group 17 elements (halogens) that as one moves down a group, more and more shells are added.

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All elements of group 1 have only one valence electron. Li has electrons in two shells, Na in three, K in four while Rb has electrons in five shells. Similarly all the elements of group 17 have seven valence electrons however the number of shells is increasing from two in F to five in I. Elements present in groups 1 and 2 on left side and groups 13 to 17 on the right side of the periodic table are called normal elements or representative elements. Their outermost shells are incomplete. They are also called typical or main group elements. Elements present in groups 3 to 12 in the middle of the periodic table are called transition elements. (Although groups 11 and 12 elements are, strictly speaking, not transition elements). Their two outermost shells are incomplete. However, it should be noted here that more and more electrons are added to valence shell only in case of normal elements. In transitions elements, the electrons are added to incomplete inner shells. Elements 113, 115 and 117 are not known but included at their expected positions. Group 18 on extreme right side of the periodic table contains noble gases. Their outermost shells contain 8 electrons. Inner transition elements:14 elements with atomic numbers 58 to 71 (Ce to Lu) are called lanthanides# and they are placed along with the element lanthanum (La), atomic number 57 in the same position (group 3 in period 6) because of very close resemblance between them. However, for convenience sake they are shown separately below the main periodic table 14 elements with atomic numbers 90 to103 (Th to Lr) are called actinides* and they are placed along with the element actinium (Ac), atomic number 89 in the same position (group 3 in period 7) because of very close resemblance between them. They are shown also separately below the main periodic table along with lanthanides. Periods There are seven rows in the periodic table. Each row is called a period. The periods have been numbered from 1 to 7 (Arabic numerals). In each period a new shell starts filling up. The period number is also the number of shell which starts filling up in it. For example, in elements of 3rd period, the third shell (M shell) starts filling up as we move from left to right@ . The first element of this

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period sodium Na (2,8,1) has only one electron in its valence shell (third shell) while the last element of this period, argon Ar (2,8,8) has eight electrons in its valence shell. The gradual filing of the third shell can be seen below.

The first period is the shortest period of all and contains only 2 elements, H and He. The second and third periods are called short periods and contain 8 elements each. Fourth and fifth periods are long periods and contain 18 elements each. Sixth and seventh periods are very long periods containing 32 elements* * each. Merits of modern periodic table over Mendeleev’s periodic table The modern periodic table is based on atomic number which is more fundamental property of an atom than atomic mass. The long form of modern periodic table is therefore free of main defects of Mendeleev’s periodic table. 1. Position of isotopes All isotopes of the same elements have different atomic masses but same atomic number. Therefore, they occupy the same position in the modern periodic table which they should have because all of them are chemically similar. 2. Anomalous pairs of elements When elements are arranged in the periodic table according to their atomic numbers the anomaly regarding certain pairs of elements in Mendeleev’s periodic table disappears. For example, atomic numbers of argon and potassium are 18 and 19 respectively. Therefore, argon with smaller atomic number comes before potassium although its atomic mass is greater and properties of both the elements match with other elements of their respective groups.

1.6 Classification of Elements (s, p, d, f) s, p, d and f-BLOCK ELEMENTS: We divide the whole periodic table in parts based on similar properties. Elements in which the last electron enters the: EDUDIGM 1B Panditya Road, Kolkata 29

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>> s-orbital of the outermost energy level are s-block elements. >> p-orbital of the outermost energy level are p-block elements. >> d-orbital of the outermost energy level are d-block elements. >> f-orbital of the outermost energy level are f-block elements. s s block

s p - block d - block

f – block S-Block The S block contains group-1 and group-2 elements. Group-1 elements also called alkali metals has the electronic configuration ns1, while group-2(alkaline earth metals) has the EC of ns2. All the elements in this group are reactive and lose electrons readily. So they are metallic. The metallic character and reactivity increases down the group P-Block P-block contains Group-13 to group-18 elements. The electronic configuration is ns2np1 to ns2np6. The p-block elements are mostly non-metals. D-Block They are called as transition elements and include the groups from group3-group12. Their valence electronic configuration is (n-1)d1-10ns1-2. They are metals and are less reactive than s-block elements but more reactive than the p-block elements. F-Block These are groups are put separately and are called as inner transition elements. Their valence electronic configuration is (n-2)f1-14(n-1)d0-1ns2. The elements from atomic number 58-71 are called lanthanides which that of 90-103 are called actinides.

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Section 2 Peridicity 2.1 2.2 2.3 2.4 2.5 2.5

Valency ................................................................................................................. 16 Atomic Radii ........................................................................................................ 17 Ionization Energy ............................................................................................... 18 Electron Gain Enthaly ........................................................................................ 19 Electronegativity ................................................................................................ 21 Metallic Character .............................................................................................. 21

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2 PERIODIC PROPERTIES IN GENERAL In the previous section we have learnt about the main features of the Modern Periodic Table. We have also learnt that in a period the number of valence electrons and the nuclear charge increases from left to right. It increases the force of attraction between them. In a group the number of filled shells increases and valence electrons are present in higher shells. This decreases the force of attraction between them and the nucleus of the atom. These changes affect various properties of elements and they show gradual variation in a group and in a period and they repeat themselves after a certain interval of atomic number. Such properties are called periodic properties. In this section we shall learn about some periodic properties and their variation in the periodic table.

2.1 VALENCY (a) Valency in a period : You have already learnt in the previous section that the number of valence electrons increases in a period. In normal elements it increases from 1 to 8 in a period from left to right. It reaches 8 in group 18 elements (noble gases) which show practically no chemical activity under ordinary conditions and their valency is taken as zero. Carefully look at the table given below. What do you observe? Valency of normal

elements with respect oxygen increases from 1 to 7 as shown below for elements of third period. This valency is equal to the number of valence electrons or group number for groups 1 and 2, or (group number-10) for groups 13 to 17.

Group Element No. of valence electrons Valency with respect to oxygen Formula of oxide

1 2 13 14 Na Mg Al Si 1 2 3 4 1 2 3 4 Na2O MgO Al2O3 SiO2

15 P 5 5 P4O10

16 17 S Cl 6 7 6 7 SO3 Cl2O7

In the following table for elements of second period you will observe that valency of

elements of with respect to hydrogen and chlorine increases from 1 to 4 and then decreases to 1 again. Group Element

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1 Li

2 Be

13 B

14 C

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16 O

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Periodic Properties No. of valence electrons Valency with respect to hydrogen and chlorine Formula of hydride Formula of chloride

Page 17 1 2 3 4 5 6 7 1 2 3 4 3 2 1 LiH BeH2 BH3 CH4 NH3 H2O HF LiCl BeCl2 BCl3 CCl4 NCl3 Cl2O ClF

(b) Valency in a group : All the elements of a group have the same number of valence

electrons. Therefore, they all have the same valency. Thus valency of all group 1 elements, alkali metals, is 1. Similarly valency of all group 17 elements, halogens, is 1 with respect to hydrogen and 7 with respect to oxygen.

2.2 Atomic radii In simple terms it means, the distance from the center of the nucleus to the outermost shell of electrons. Generally, atomic radii decreases along a period and increases down a group. With the increase in the atomic number (increased number of protons, electrons and neutrons) in the 3rd period, the net positive charge of the nucleus gradually increases. This increased positive charge exerts a greater attraction on the shells and attract the electrons in the shells a little closer to the nucleus. Hence, sodium has the largest atom and chlorine the smallest. This is true of other periods as well. The radius of atom of elements in the same group increases downward. For e.g., in group 1, starting from lithium to caesium, the atomic size increases because there is a gradual increase in the number of shells. Positive ion is always smaller than the neutral atom, owing to the diminished electronelectron repulsion. If a second electron is lost, the ion gets even smaller. Negative ions are always larger than the parent ion; the addition of one or more electrons to an existing shell increases electron-electron repulsion which results in a general expansion of the atom. An isoelectronic series is a sequence of species all having the same number of electrons (and thus the same amount of electron-electron repulsion) but differing in nuclear charge. Of course, only one member of such a sequence can be a neutral atom (neon in the series shown below.) The effect of increasing nuclear charge on the radius is clearly seen.

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O2->F->Ne>Na+>Mg2+ Example The radii of Ar is greater than the radii of chlorine Solution: In chlorine the radii means the atomic or covalent radii which are actually half the inter – nuclear distance between 2 atoms whereas in Argon the radii means the van der Waals radii as Argon is not a diatomic molecule. Van der Walls radii are actually half the distance between adjacent molecules. So ven der Waals radii being larger than atomic radii, Argon, has got a larger radii then chlorine

2.3 Ionization enthalpy (I.E.) Energy required to remove the most loosely held electron from the gaseous isolated state of an atom. Remember: Energy

required to bring about the change

A

e is called

1st Ionization energy (IE) i.e. the energy required to remove the 1st electron from an isolated atom. Similarly, 2nd I.E. is the energy required to remove the 2nd electron from a isolated atom A

A

e

So, in general, the I.E. increases along a period and decreases along a group. [Why? Hint: - The closer the e- to the nucleus, the more tightly held it is and hence higher the I.E.] Also, 2nd ionization energy is always greater than the first ionization energy. (Because it is easy to separate an electron from a neutral atom than to separate it from a +vely charged atom). Ionization potential increases across the period because of increase in nuclear charge due to which the atomic size decreases. Thus, more energy is required to pull away the electron from the outermost shell of the atom of smaller size. Ionization potential decreases down the group because of increase in the number of shells. The effective nuclear charge decreases as atomic size increases. Thus it is easier to pull one electron from the outermost shell of the atom. Group 18 elements have the highest Ionization enthalpy because of their full-filled electronic configuration. Similarly it is more for half filled electronic configurations (e.g. N) Some points to note:

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The noble gases have the highest IE's of any element in the period. This is because of very high nuclear charge and fulfilled electronic configuration.



IE's (as well as many other properties) tend not to vary greatly amongst the dblock elements. This reflects the fact that as the more-compact d orbitals are being filled, they exert a screening effect that partly offsets that increasing nuclear charge on the outermost s orbitals of higher principal quantum number.



Each of the Group 13 elements has a lower first-IE than that of the element preceding it. The reversal of the IE trend in this group is often attributed to the more easy removal of the single outer-shell p electron compared to that of electrons contained in filled (and thus spin-paired) s- and d-orbitals in the preceding elements.

Example The first I.P. of nitrogen is greater than oxygen white the reverse is true for their second I.P. values. Solution: The first I.P corresponds to the removal if first electron. Since nitrogen is already half filled. So more energy is required to remove the electron. But once the electron is removed from oxygen it gains half filled stability and therefore the 2nd I.P. becomes high. Example The ionization energy of the coinage metals fall in the order

.

Solution: In all the 3 cases as s – electron in the unpaired state is to be removed. In the case of Cu a 4s electron is to be removed which is closer to the nucleus than the 5s electron of Ag. So I.P. decreases from Cu to Ag. However form Ag to Au the 14 f electrons are added which provide very poor shielding effect. The nuclear charge is thus enhanced and therefore the outer electron of Au is more tightly held and so the IP is high.

2.4 Electron gain enthalpy (E.G.E) Energy released when an electron is added to the gaseous isolated state of an atom. Again sounding tough? Well, take the reaction: The energy

released in this reaction is E.G.E.

More the E.G.E., easier the addition of electrons. So, in general, the E.G.E. … EDUDIGM 1B Panditya Road, Kolkata 29

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Becomes more negative along a period and less negative down the group. [Why?.... Hint: - smaller the atomic radius, higher the “effective nuclear charge” and hence, gaining of e- becomes easier. Thus E.G.E. is more negative.] Electron affinity is the ability of an atom to hold an additional electron. If an atom has greater tendency to accept an electron then the energy released will be large and consequently the electron affinity will be high. Electron Affinity can be positive or negative. Electron affinity increases from left to right across the period because of increase in nuclear charge and decrease in atomic size. This causes the incoming electron to experience a greater pull of the nucleus. Electron affinity decreases down the group because the number of shells increases i.e., the atomic size increases and the effective nuclear charge decreases. The electron affinity of completely filled atoms is almost zero. An atom does not accept an electron in its outermost shell if it already has stable configuration e.g. inert gases So it depends mainly of effective nuclear charge, atomic size and electronic configuration of the element. Chlorine has the highest electron affinity. Some irregularities are 

In the Group 2 elements, the filled 2s orbital apparently shields the nucleus so effectively that the electron affinities are slightly endothermic.



The Group 15 elements have rather low values, due possibly to the need to place the added electron in a half-filled p orbital; why the electron affinity of nitrogen should be endothermic is not clear. The vertical trend is for electron affinity to become less exothermic in successive periods owing to better shielding of the nucleus by more inner shells and the greater size of the atom, but here also there are some apparent anomalies.

Example The electron affinity of sulfur is greater than oxygen. Why? Solution: This is because of smaller size of oxygen due to to which it has got higher change density and thus electronic repulsion increases as it takes electron. So its E.A. is less than sulphur.

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2.5 Electronegativity It is the property of an atom which is bonded to another atom. The atom which is more electronegative tries to keep more share of electron. It increases along the period, as effective nuclear charge increases and atom will try to pull the shared electron because of more nuclear charge. It decreases down the group. Fluorine is the most electronegative element.

2.6 Metallic Character The metallic character is actually tendency of atom to lose electrons and form positive ions. It has similar trend as ionization energy. The metallic character decreases along the period because of increase in effective nuclear charge which holds the valence electron with greater force. The metallic character increases down the group, as down the group the nuclear force of attraction decreases. Some important facts! 1. Size of anion > size of atom > size of cation. [Why?] 2. The ions having same no. of electrons are called isoelectronic ions. The size in an isoelectronic series decreases with increase in nuclear charge. E.g.

[Why?]

3. Be, Mg, N, P and noble gases have exceptionally high values of I.E. due to their ‘stable’ half-filled and completely-filled electronic configuration. For the same reason they have very low E.G.E. also. 4. Chlorine has highest

. . .

among all the elements.

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Periodic Properties

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1

Across

2

1 The lightest inert gas. (6)

3 4 5

7

6

10

5 Alkaline metal in table salt. (6) 7 Inert gas used to make bright city lights. (4)

8

9

4 The element we need to breathe. (6)

11

8 2nd place Olympics. (6)

12

13

in

the

10 An important element in bones. (7) 13 A radioactive element often used in nuclear power stations. (7)

14

14 Poison gas in WWI. (8) 15

15 A famous poison that turns your tongue black. (7)

16

17

17 A metal sought after during the Klondike. (4) 18 The element diamonds are made from. (6)

18

19 This metal is used along with carbon to make steel. (4)

19

Down 2 Heavy metal used in paints, batteries, and radiation shields. (4) 3 The most common element in the universe. (8) 6 A liquid metal that was used in thermometers. (7) 7 Most common element in the earth's atmosphere. (8) 9 A component of gunpowder that smells like rotten eggs. (7) 11 A metal used in foil. (World Spelling) (9) 12 A metal used in wires. (6) 15 The most common inert gas in the atmosphere. (5) 16 Element used to make semi-conductors (computer chips). (7)

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Solved Examples Example Berilium and Al are placed in different periods and groups by they show the similar properties. Solution: On moving across a period the charge on the ions increases and the size decreases, causing the polarizing power to increases. On moving down a group the size increases and polarizing power decreases. On moving diagonally i.e., form Be to Al these two effects partly cancel each other and so there is no marked change in properties. Example (i) NaOH behaves as a base while

(

) is amphoteric why?

(ii) Among fluorine – fluorine bond and chlorine – chlorine bond. Which is more stronger and why? Solution: (i) In NaOH the bond electronegativity difference between Na and oxygen is greater than between H and O and therefore it is the Na- O bond that breaks releasing in case of

bond the difference of electronegativity of

and

. But are

almost same. So there is equal probability that the bond breaks in both ways leading to an amphoteric behaviour (ii) In Cl – Cl bond, a filled p – orbital of chlorine can overalap with a suitable vacant d – orbital of adjacent chlorine thereby introducing some double bond character. Thus the bond strength increases. This is not possible in fluorine as it has got no vacant d – orbital Example (i) In alkali metal group which is the strength reducing agent and why? (ii) Although aluminium is above hydrogen in the electrochemical series, it is stable in air and water. Explain. Solution: (i) Li is the strongest reducing agent. Since I.P. decreases down the group we would expect that Li will have the lowest reducing power in the group. But since it’s hydration energy is very high and which in fact decreases down the group, Li will have highest reducing power.

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(ii) Due to the formation of protective oxide layer on its surface. Example :

atoms of X(g) are converted into

are converted into

( ) and

( ) by energy

.

atoms of X(g) are converted into

atoms of X(g)

( ) by energy

.

Calculate the ionization energy and electron affinity of X Solution: let the ionization energy of X beI (which is always in kJ/ mol) and electron affinity of X be EA kJ/mol ( )

( )

(

2

)

(

)

(where

is the Avogadro’s number)

Similarly ( )

( )

( )

( )+ (

2

) )

2( Example : [

two

]3 3

[

atoms ]3 3

have

the

electronic

configuration

4 . the first ionization energy of 2762 kJ/mole and that of

the other is 692 Kj/mole. Match each ionization energy with one of the electronic configuration. Justify your choice Solution: [ [

]3 3

Completely

]3 3 4

2762

692 filled

configurations

have

higher

ionization

energy

than

other

configurations due to their extra stability Example : explain why the 2nd ionization energy of Cr is higher than that of Mn Solution: after losing one electron Cr will get converted into half filled electronic configuration (

which will have stable

) with respect to configuration

(

)

Example : why the 1st Ionisation energy of the phosphorus is greater than that of sulphur

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Solution: phosphorus has got a stable half filled electronic configuration hence the 1st I.E. of P is higher than that of Example : the 1st and 2nd I.E. of K are 419 kJ/m and 3052 kJ/mol and Ca are 590 kJ/mol and 1145 kJ/mol respectively. Compare their values and explain the differences Solution: Removal of 2nd electron from K is extremely difficult because K acquires stable noble gas configuration after removing one electron while removal of both the electrons for Ca is comparatively easy, as it acquires stable configuration after removal of both the electons Example : third ionization energy of C is higher than that of N, explain. Solution: after ejecting two electrons from 2p orbitals, the third electron is being ejected from filled (stable) 2s orbital of carbon, hence a greater amount of energy is required. But in nitrogen the third electron is to be ejected from 2

orbital, thus lesser energy is

required to remove that electron. Example :

Among the elements with atomic number 9, 12 and 36, identify by atomic number of an element which is (a) highly electronegative (b) an inert gas (c) highly electropositive and give reasons for your choice

Solution:

The electronic configuration of the elements with atomic number 9, 12

and 36 are: Atomic number 9 : 2, 7 Atomic number 12 : 2, 8, 2 Atomic number 36 : 2, 8, 18, 8 (a) The element with atomic number 9 can accept one more electron to have 8 electrons in the outermost orbit, thus it is an electronegative element. (b) The element with atomic number 12 looses two electrons to acquire inert gas configuration thus it is electropositive in nature.

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(c) The element with atomic number 36 has 8 electrons in the outermost orbit, hence it has no tendency either to lose or accept electrons. Thus, it is an inert gas. Example :

The first ionization energy of carbon atom is greater than that of boron atom, whereas reverse is true for the second ionization energy. Explain.

Solution:

The electronic configurations of carbon and boron are as follows: 2

2

2

,2 ,2 Due to higher nuclear charge in carbon, the force of attraction of valency electron is more in carbon atom and hence the first ionization energy is greater than boron atom. After loss of one electron, the monovalent cations have the configurations as follows: ,2 ,2 ,2 The

configuration is stable one and hence the removal of electron is

difficult in comparison to

. Hence, second ionization potential of boron

is higher than carbon. Example :

The formation of

(g) from F (g) is exothermic whereas that of

(g)

from O (g) is endothermic, explain Solution:

The addition of an electron to a neutral atom is an exothermic process. …… ( ) The addition of second electron to a monovalent anion,

, as to make it

is difficult because both have the same charge and experience a lot of repulsion. Thus, the addition of an electron to

requires energy to

overcome the force of repulsion. …….. (2) The energy absorbed in step 2 is more than the energy released in the step -1. Hence, the formation of

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from O is endothermic in nature.

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The diagram given below is a part of the periodic table. Study the table and answer the questions that follow – (i) Name the elements in the same group of the periodic table (ii) Name a transition metal (iii) Give the atomic number of an element which is inert

1 3

4 Be 11 12

5

6

7

8

9

2 He 10

13 14 15 16 17 18 Si S 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 Ca Cr Kr Solution: (i) The element placed vertically one below another belong to the same subgroup. Thus Be and Ca belong to the same group. (ii) First transition series starts from element – 21 and ends at element – 30. Thus chromium – 24 is a transition metal. (iii) The atomic number of an inert gas is 36 (Kr.)

Assignments Level-I 1.

From the following point of elements decide which element is going to have higher I.E? (a) He,

Li

(b) Be,

B

(c) N, O

2.

What are the diagonal relationship in periodic table and why do these occur?

3.

Write down the increasing order of electron affinity in VII group. VII group elements are – F, CI, Br, I, At,

4.

The order of electronegativity in VII group would be ………………………..

5.

HClO3 behaves as a stronger acid than HClO. How can you explain this fact on the basis of electronegativity?

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Periodic Properties 6.

Page 28

is more acidic than

which in turn is more acidic

. Explain this on

the basis of hybridization electronegativity. 7.

Compound

Electronegativity of central atom

Bond angle

3.0

2

2.1

97

2.0

96

1.9

88

Explain this table by trying to correlate the electronegativity of the central atom and bond angle. 8.

Arrange the following isoelectronic ions in increasing order of ionic radii:,

9.

,

,

,

?

How many elements are there in the periodic table? If an element with Z = 107 is discovered, where would you place it in the periodic table?

10.

Arrange the following species in the order of increasing size: (i)

11.

,

, ,

(b)

,

,

,

Electronic configurations

Type of element

(a)

(1)

alkali metal

(2)

Halogen

62

(C)

(3)

(d) (e)

(

)

(f)

,

,

alkaline earth metal

(4)

Transition metal

(5)

Inert gas (6)

Non metal

Which of the following elements are going to have similar properties, and why? (A)

13.

(c)

Match the correct pairs –

(b)

12.

,

2 2

(B)

2 2

3 3

3

4

(c)

2

Arrange the following according to given instructions: (i)

,

,

,

(ii)

,

, ,

(increasing radius) (Increasing radius)

(iii)

,

,

,

(increasing electronegativity)

(iv)

,

,

,

(increasing electron affinity)

(v)

, , ,

(vi)

,

,

(increasing electronegativity) ,

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(increasing 1st ionization potential)

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Periodic Properties (vii)

,

(viii)

,

Page 29 ,

, ,

(ix)

,

(x)

,

(increasing peramagnetism) ,

(increasing electropositive character)

, ,

,

(increasing atomic volume)

,

(increasing IInd ionization potential)

Level-II 14.

Explain the following

15.

(A)

Zinc salt do not have any colour but copper salts are blue.

(B)

halogens have high electron affinity.

(C)

s – block elements are good conductor of electricity

(D)

d-block elements are called transition metals.

(E)

Nitrogen has very high 1st ionization potential

(F)

CsF is the most ionic compound which one can expect.

(G)

The electron affinity of noble gases are zero.

(H)

Gallium is smaller in size than aluminium.

The heat of formation of the oxides of third period are given in KJ mol–1.

–416

–602 –1676

–911

–2984

–395

+250 Arrange these oxides in increasing order of stability. 16.

Arrange the following in the order of reducing character: (a) Na, K, Rb (b) Na, Mg, Al ,

17.

,

,

,

(d)

, Indicate whether the following process is exothermic or endothermic: ( )

2

( )

( )

2

( . .)

( )

737.7

( . .)

( )

45

( . .) 18.

(c)

( )

328

( )

Arrange the following ions as indicated; ,

,

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Periodic Properties

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In increasing order of (i) degree of hydration

(ii) hydration energy

(iii) size of

hydrated ions (iv) ionic mobility 19.

(v) standard reduction potential

Calculate the electron affinity of the hydrogen atom using the following enthalpy

data: 2

( ) ( )

2

20.

2

( )

( )

( )

( )

( )

( )

436

( )

8

( )

83 4 3 742

( )

( )

Arrange the following species in decreasing order of electropositive character Fe, Sc, Rb, Br, Te,F, Ca

21.

Select paramagnetic and diamagnetic species among the following ,

22. , 23.

,

Arrange ,

,

,

,

following

,

,

species

. in

decreasing

order

of

atomic

size

,

Calculate the electronegativity of fluorine from following data 4.2 36.3 34.6 Electronegativity of H is 2.05.

24.

First and second ionization energies of Calculate percentage of

25.

( ),

( )

74

( )

if 1 g of

( )

45

.

absorbs 50 KJ of energy.

The first four ionization energies of an element are approximately 738 kj/mole, 1450 kj/mol, 7700 kj/mol and 1000 kj/mol. Identify the periodic group to which this element belongs.

26. 27.

Arrange the following isoelectronic species in order of (a) increasing ionic radius and (b) increasing ionization energy:

,

Which oxide is more basic,

Why?

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or

,

,

.

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Periodic Properties

Page 31

Level-III 28.

Write balanced equation for the reactions between each of the following oxides and water.

29.

(a)

(b)

(c)

Arrange the elements in each of the following groups in increasing order of the most positive electron affinity. (a) Li, Na, K

30.

(b) ,

,

,

Among the elements of the third period (Na to Ar), pick out the element (i) With the highest first ionization enthalpy (ii) With the largest atomic radius (iii) That is the most reactive nonmetal (iv) That is the most reactive metal

31.

Arrange the following elements in the increasing order of nonmetallic character: B, C, Si, N and F.

32.

Account for the large decrease in electron affinity between lithium and beryllium despite the increase in nuclear charge.

33.

In general, ionization energy across a period from left to right. Explain why the second ionization energy of chromium is higher than that of manganese.

34.

The ionization energies of Li and K are 5.4 and 4.3 eV, respectively. What do we predict for the ionization energy of

35.

?

The ionization energies of Li, Be and C are 5.4, 9.3 and 11.3 eV. What do we predict for the ionization energies of B and N?

36.

Explain in terms of their electronic configurations, why oxidized to

37.

than

to

.

The electron affinity of chlorine is 3.7 eV. How much energy in kcal is released when 2 g of chlorine is completely converted to

38.

is more easily

ion in a gaseous state?

The first ionization potential of Li is 5.4 eV and the electron affinity of Cl is 3.6eV. Calculate ( )

in kcal

for the reaction

( )

Carried out at such low pressures that resulting ions do not combine with each other. 39.

The ionization potentials of atoms A and B are 400 and 300 kcal respectively. The electron affinities of these atoms are 80.0 and 85.0 kcal respectively. Prove that which of the atoms has higher electronegativity.

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Periodic Properties 40.

Arrange the following species in decreasing order of atomic size. ,

41.

Page 32

,

,

and

Indicate whether the following process is exothermic or endothermic. ( )

2 ( )

Given: ( . . ) of

( ) ( )

( . . ) of ( ) 42.

2

( ) ; ( . .)

737.7

( )

of

45

&

328

Arrange the following atoms/species in the order of reducing character. (a)

, ,

(b)

,

,

(c)

,

,

,

Hints & Answers 1.

(a) He (inert Gas)

(b) Be (Fully filled shell)

(c) N (Half filled

configuration) 2.

In the periodic table the elements show similar properties across diagonal. This behaviour is because of identical size of the elements. E. g.: GP I

GP II

GP III

GP IV

Li

Be

B

C

Na

Mg

Al

Si

3.

The order of electron affinity in VII gp is as F < Cl > Br > I > At

4.

F > Cl > Br > I > At

5.

O H – Cl = O

+ HClO 1

O O Chlorine in (+5) oxidation state is less stable than (+1) oxidation sate and has greatest tendency to attract electrons than in (+1) oxidation sate. Therefore HClO as a stronger acid. H H3 behaves H H H 6.

H H H –HC –HC – H < H – C = C – H < H – C

C–H

H In acetylene carbon in ‘sp’ hybridized and thus the percentage of ‘’s’’ character is maximum in acetylene. With the increase of a character shielding of electrons

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Page 33

around the carbon nucleus is maximum. Thus and

can be lost easily. But in

the percentage of s character decreases so the acidic character also

decreases. 7.

The lone pair on nitrogen in

is strongly attracted so there is less pair – bond

pair repulsion and hence the bond angle will be maximum. As the electronegativity of the central atom decreases the lone pair – bond pair repulsion increases and hence bond angle decreases. 8.

(increasing order of size)

9.

107 elements; 7th period & VII B Group.

10.

(i) Cl < S < P < Si

11.

(a) 2.

12.

B & C; Because of bipositive ions formation

13.

(i)

(ii)

(iii)

(iv)

(v)

(vi)

(vii)

(viii)

(ix)

(x)

14.

(a)

(b)

(c)

(b) 1

(c) 5

(d) 3

(e) 4

(g) 6

Zinc do not have partly filled d – orbitals so cannot undergo excitation of

electrons to higher energy shells (no d – d transition). Therefore they are colourless. (b)

They have a stron tendency to gain an electron to acquire noble gas

configuration, viz,

.

(c)

Are strongly metallic in nature.

(d)

Because their properties are transitional between s and p block elements.

(e)

Due to extra stability of half filled orbital.

(f)

Due to largest difference in the size and their electronegativity values.

(g)

Noble gases are stable.

(h)

Poor screening effect.

15. Hint:

. Oxide

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MgO

– 416

– 416

– 602

– 602

– 1676

– 558.7

– 911

– 455.5

– 2984

– 298.4

– 395

– 131.7

+ 250

+ 35.7

Greater is the negative value of heat of formation per oxygen atom more is the stability of oxide. 16.

(a) Na < K < Rb

(b) Na > Mg > Al

(c)

(d)

Mg < Ca < Sr 17.

The process is endothermic Hint:

( )

( )

[( . ) 2 . . 2

( . )

( )

2

328

2

19.

]

2 88.7 2

( )

2 2 88.7

18.

2

328

2

532.7

(i)

(ii)

(iii)

(iv)

37 Hint: it can be solved by forming the born haber’s cycle.

20. 21.

Paramagnetic species:

,

Diamagnetic species:

,

,

,

,

22. Hint: in isoelectronic species, size decreases with increase in number of protons in the nucleus.

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Periodic Properties 23.

Page 35

3.6034 Hint: On pauling scale . 82√ (

…………….. ( ) )

. .

34.6

√ ()

. 82√72.84 .5534

24.

2. 5

68.35

4.2

√

36.6

72.84

.5534

.5534

3.6 34

3 .65

Hint: Number of moles of 1 g of ( )

Energy required to convert Remaining energy

5

. 4 7 ( )

3 .83

74

3 .83

9. 7 .

Number of moles of Thus, remaining

. 4 7

. 4 7

. .

32

32

. 285 .

68.35

.

68.35

3 .65

25. II A 26. (a)

(b)

27. BaO, because basic character of oxides increases in group 28. (a)

2

(

(b)

)

(c) 29. (a)

(b)

30. (i) Ar (II) Na (III) Cl (IV) Na 31. Si < B< C< N < F 32. Be has a stable filled 2s configuration. 33.

will have stable half filled configuration ( (

) with respective configuration

).

34. Ionization energy of Na is intermediate between that of Li & K i.e. 4.9 eV. 35. I.E. (B) < I.E (Be) 36.

I.E. (N)>I.E (C)

has outer electronic configuration of 3

gets oxidized to

, as

while

has 3

has a stable half filled configuration (3

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easily will

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Periodic Properties

Page 36

not easily get oxidized to (3

as it goes to a stable having less stable configuration

).

37. 4.8 kcal 38. 41.508 kcal 39. Electronegativity of A = 3.84; electronegativity of B = 3.08. Therefore A has higher electronegativity. 40. 41. The process is endothermic ( 42. (a)

(b)

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532.7

)

(c)

(d)

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