Periodic Properties of Elements

Periodic Properties

EDUDIGM 1B Panditya Road, Kolkata 29 www.edudigm.in 40034819

Periodic Table

Periodic Properties


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

Periodic Properties


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.

Periodic Properties


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.

Periodic Properties


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 Beryllium Boron Carbon Nitrogen Oxygen Fluorine

Li Be B C N O F

Sodium Magnesium Aluminium Silicon Phosphorus Sulphur Chlorine

Na Mg Al Si P S Cl

Potassium Calcium Titanium Chromium Manganese Iron Cobalt,Nickle

K Ca Ti Cr Mn Fe 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.

Periodic Properties


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 non-

metallic 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.

Periodic Properties


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

Periodic Properties


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.

Periodic Properties


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 ).

Periodic Properties


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.

Periodic Properties


1 18 1 1

H 2 13 14 15 16 17 2

He

2 3 Li

4 Be

5 B

6 C

7 N

8 O

9 F

10 Ne

3 11 Na

12 Mg

3 4 5 6 7 8 9 10 11 12 13 Al

14 Si

15 P

16 S

17 Cl

18 Ar

4 19 K

20 Ca

21 Sc

22 Ti

23 Y

24 Cr

25 Mn

26 Fe

27 Co

28 Ni

29 Cu

30 Zn

31 Ga

32 Ge

33 As

34 Se

35 Br

36 Kr

5 37 Rb

38 Sr

39 Y

40 Zr

41 Nb

42 Mo

43 Tc

44 Ru

45 Rh

46 Pd

47 Ag

48 Cd

49 In

50 Sn

51 Sb

52 Te

53 I

54 Xe

6 55 Cs

56 Ba

57 La

72 Hf

73 Ta

74 W

75 Re

76 Os

77 Ir

78 Pt

79 Au

80 Hg

81 Ti

82 Pb

83 Bi

84 Po

85 At

86 Rn

7 87 Fr

88 Ra

89 Ac

104 Rf

105 Ha

106 Un

107 Un

108 Un

109 Un

110 Un

Lanthanides 58 Ce

59 Pr

60 Nd

61 Pm

62 Sm

63 Eu

64 Gd

65 Tb

66 Dy

67 Ho

68 Er

69 Tm

70 Yb

71 Lu

Actinides 90 Th

91 Pa

92 U

93 Np

94 Pu

95 Am

96 Cm

97 Bk

98 Cf

99 Ex

100 Fm

101 Md

102 No

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.

Periodic Properties


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

Periodic Properties


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:

Periodic Properties


>> 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

s b

lock

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.

Periodic Properties


Section 2

Peridicity

2.1 Valency ................................................................................................................. 16

2.2 Atomic Radii ........................................................................................................ 17

2.3 Ionization Energy ............................................................................................... 18

2.4 Electron Gain Enthaly ........................................................................................ 19

2.5 Electronegativity ................................................................................................ 21

2.5 Metallic Character .............................................................................................. 21

Periodic Properties


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 1 2 13 14 15 16 17

Element Na Mg Al Si P S Cl

No. of valence electrons 1 2 3 4 5 6 7

Valency with respect to oxygen 1 2 3 4 5 6 7

Formula of oxide Na2O MgO Al2O3 SiO2 P4O10 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 1 2 13 14 15 16 17

Element Li Be B C N O F

Periodic Properties


No. of valence electrons 1 2 3 4 5 6 7

Valency with respect to hydrogen and chlorine 1 2 3 4 3 2 1

Formula of hydride LiH BeH2 BH3 CH4 NH3 H2O HF

Formula of chloride 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 electron-

electron 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.

Periodic Properties


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:

Periodic Properties


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 d-

block 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. …

Periodic Properties


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.

Periodic Properties


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.

Periodic Properties


Across

1 The lightest inert gas. (6)

4 The element we need to breathe. (6)

5 Alkaline metal in table salt. (6)

7 Inert gas used to make bright city lights. (4)

8 2nd place in the Olympics. (6)

10 An important element in bones. (7)

13 A radioactive element often used in nuclear power stations. (7)

14 Poison gas in WWI. (8)

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

17 A metal sought after during the Klondike. (4)

18 The element diamonds are made from. (6)

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

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)

1 2

3

4

5 6

7 8

9 10 11 12

13

14

15 16

17

18

19

Periodic Properties


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 . But

in case of bond the difference of electronegativity of and 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.

Periodic Properties


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

Example :

atoms of X(g) are converted into ( ) by energy .

atoms of X(g)

are converted into ( ) and

atoms of X(g) are converted into ( ) 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 atoms have the electronic configuration

[ ]3 3 [ ]3 3 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 2762

[ ]3 3 4 692

Completely 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 which will have stable

half filled electronic configuration ( ) with respect to configuration ( )

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

sulphur

Periodic Properties


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.

Periodic Properties


(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 from O is endothermic in nature.

Periodic Properties


Example : 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

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?

1

2

He

3

4 Be

5

6

7

8

9

10

11

12

13

14 Si

15

16 S

17

18

19

20 Ca

21

22

23 24 Cr

25

26

27

28

29

30

31

32

33

34

35

36 Kr

Periodic Properties


6. 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) , , , (b) , , , (c) , , ,

11. Match the correct pairs –

Electronic configurations Type of element

(a) 62 (1) alkali metal

(b) (2) Halogen

(C) (3) alkaline earth metal

(d) (4) Transition metal

(e) ( ) (5) Inert gas

(f) (6) Non metal

12. Which of the following elements are going to have similar properties, and why?

(A) 2 2 (B) 2 2 3 3 3 4 (c) 2

13. Arrange the following according to given instructions:

(i) , , , (increasing radius)

(ii) , , , (Increasing radius)

(iii) , , , (increasing electronegativity)

(iv) , , , (increasing electron affinity)

(v) , , , (increasing electronegativity)

(vi) , , , (increasing 1st ionization potential)

Periodic Properties


(vii) , , , (increasing peramagnetism)

(viii) , , , (increasing electropositive character)

(ix) , , , (increasing atomic volume)

(x) , , , (increasing IInd ionization potential)

Level-II

14. Explain the following

(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.

15. 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 (c) , , , (d)

, ,

17. Indicate whether the following process is exothermic or endothermic:

( ) 2 ( ) ( ) 2 ( )

( . . ) ( ) 737.7

( . . ) ( ) 45

( . . ) ( ) 328

18. Arrange the following ions as indicated;

, ,

Periodic Properties


In increasing order of

(i) degree of hydration (ii) hydration energy (iii) size of

hydrated ions

(iv) ionic mobility (v) standard reduction potential

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

data:

( ) 2 ( ) 436

( ) 2 ( ) 2 ( ) 8

( ) ( ) 83

( ) ( ) 4 3

( ) ( )

( ) 742

20. 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. Arrange following species in decreasing order of atomic size

, , , ,

23. 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 ( ) 74 45 .

Calculate percentage of ( ) ( )

, if 1 g of ( ) absorbs 50 KJ of energy.

25. 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. Arrange the following isoelectronic species in order of (a) increasing ionic radius

and (b) increasing ionization energy: , , , .

27. Which oxide is more basic, or Why?

Periodic Properties


Level-III

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

and water. (a) (b) (c)

29. Arrange the elements in each of the following groups in increasing order of the

most positive electron affinity. (a) Li, Na, K (b) , , ,

30. 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 is more easily

oxidized to than to .

37. 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 ion in a gaseous state?

38. 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.

Periodic Properties


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

, , , and

41. Indicate whether the following process is exothermic or endothermic.

( ) 2 ( ) ( ) 2 ( )

Given: ( . . ) of ( ) 737.7 ; ( . . ) of ( ) 45 &

( . . ) of ( ) 328

42. 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. H – Cl = O HClO

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

HClO3 behaves as a stronger acid.

6. H – C – C – H < H – C = C – H < H – C C – 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

H H

H

+

1

O

O O

H H H

H

H H H

H H

H H

H

Periodic Properties


around the carbon nucleus is maximum. Thus can be lost easily. But in

and 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 (b) (c)

11. (a) 2. (b) 1 (c) 5 (d) 3 (e) 4 (g) 6

12. B & C; Because of bipositive ions formation

13. (i) (ii)

(iii) (iv)

(v) (vi)

(vii) (viii)

(ix) (x)

14. (a) 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

Periodic Properties


– 416 – 416

MgO – 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 88.7 ]

2 ( ) 2 2 ( )

. . 328 2

2 2

2 88.7 328 2 532.7

18. (i) (ii)

(iii) (iv)

19. 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.

Periodic Properties


23. 3.6034

Hint: On pauling scale

. 82√ …………….. ( )

( . . )

√ 34.6 √ 4.2 36.6 72.84

( ) . 82√72.84 .5534

.5534 2. 5 .5534 3.6 34

24. 68.35 3 .65

Hint: Number of moles of 1 g of

. 4 7

Energy required to convert ( ) ( ) . 4 7 74 3 .83

Remaining energy 5 3 .83 9. 7

Number of moles of .

. 32

Thus, remaining . 4 7 . 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) I.E. (N)>I.E (C)

36. has outer electronic configuration of 3 while has 3 , so easily

gets oxidized to , as has a stable half filled configuration (3 ) but will

Periodic Properties


not easily get oxidized to as it goes to a stable having less stable configuration

(3 ).

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 ( 532.7 )

42. (a) (b) (c) (d)

Periodic Properties of Elements

Documents

Transcript of Periodic Properties of Elements