notes -periodic properties

Long Form of the Periodic Table or Modern Periodic Table ;

The modern periodic table is based on modern periodic law given by Moseley in 1942. The modern periodic law states that “the physical and chemical properties of the elements are periodic function of their atomic numbers“. Thus, when the elements were arranged in the order of their increasing atomic numbers, the elements of similar properties recur at regular intervals.

This table consists of horizontal rows called as ‘periods’ and vertical columns called as ‘groups’.

Periods ;  There are seven periods in the periodic table and each period starts with a different principal quantum number.

The first period corresponding to ‘n’ = 1 consists of only two elements hydrogen (1s1) and helium (1s2). This is because the first energy shell has only one orbital (1s), which can accommodate only two electrons.

In the second period corresponding to ‘n’ = 2, there are four orbitals (one ’2s’ and three ’2p’) having a capacity of eight electrons and so contains eight elements. This period starts with lithium (Z = 3) with electron entering the ’2s’ orbital and ends with neon (Z = 10) where the second shell is complete (2s2 2p6).

In the third period corresponding to ‘n’ = 3, there are nine orbitals (one ’2s’ three ’2p’ and five ’3d’). As ’3d’ orbitals are higher in energy, they are filled after the ’4s’ orbitals.

The fourth period corresponding to ‘n’ = 4, consists of filling of one ’4s’ and three ’4p’ orbitals.. there are eighteen elements in this period starting from potassium with electron entering the ’4s’ orbital (Z = 19) to krypton (Z = 36) where the third shell gets completed (4s2 3d104p6).

In the fifth period there are 18 elements like the fourth period. It begins with rubidium (Z = 37) with the filling of ’5s’ orbital and ends with xenon (Z = 54) with the filling of ’5p’ orbital.

The sixth period contains 32 elements (Z=55 to 86) and the successive electrons enter into ’6s’, ’4f’, ’5d’ and ’6p’ orbitals in that order. It starts with caesium and ends with radon.

The seventh period, though expected to have 32 elements is incomplete and contains only 19 elements at present.   Number of elements in each period ;  There is a periodicity occurring at regular intervals of 2, 8, 8, 18, 18 and 32 and so the numbers 2, 8, 18 and 32 are called magic numbers. The first three periods are called short periods while the other three periods are called long periods.

Groups; The vertical column in the periodic table is called as group. There are 18 groups in the long form of the periodic table and they are numbered from 1 to 18 in the IUPAC system.

Merits of the long form of the periodic table

1. based on the most fundamental property of the elements – the atomic number, so it is more accurate.

2. the position of isotopes in one place is justified.

3. four blocks of ‘s’, ‘p’, ‘d’ and ‘f ‘ has made the study of elements simpler.

4. The position of the elements, which were misfit on the basis of atomic mass is now justified on the basis of atomic number

Grouping of Elements In the long form of the periodic table,; elements are grouped into four main blocks, purely on the basis of electronic configurations. Elements are grouped in blocks ‘s’, ‘p’, ‘d’ and ‘f’ depending on the nature of orbital(s) into which the last electron of the atom enters.

The ‘s’ block elements ;  ‘s’ block elements, (except helium) are those in which the valence-shell is the ‘ns’ orbital and where the last electron enters into the ‘s’ orbital of the outer element. The general outer electronic configuration of such elements is either ns1or ns2. The ‘s’ block elements belong to group 1 and 2 of the long form periodic table. They are situated on the left side of the table.

The ‘p’ Block Elements ‘p’ block elements are those in which the outer electronic configuration is of the type

ns2 np1to6 and where the last electron enters into any of the outermost ‘p’ orbitals. These elements belong to groups 13 to 18 of the long form of periodic table and are situated on the right hand side of the table.

The ‘d’ Block Elements ‘d’ block elements are those in which the added electron goes into one of the ‘d’ orbitals. These elements have valence electrons in both their outermost and penultimate shells (second outermost) and have a general outer electronic configuration of (n – 1) d1-10 ns1-2.   The penultimate shell in these elements is expanded from 8 to 18 by the inclusion of ten ‘d’ electrons. It is for this reason that these elements are called ‘d’ block elements. Thus, ‘d’ block elements are those, which in their elemental or combined forms have partially filled ‘d’ orbitals. They are also referred to as transition elements.   ‘d’ block elements belong to groups 3 to 12 of the long form of periodic table and are situated in the middle of the table between ‘s’ block and ‘p’ block elements.   All ‘d’ block elements are classified into four transition series, namely:   3d series in the 4th period having 10 elements   4d series in the 5th period having 10 elements   5d series in the 6th period having 10 elements   6d series in the 7th period having incomplete elements.

The ‘f’ Block Elements The ‘f’ block elements are those, which in their elemental or ionic forms have partially filled ‘f’ orbitals. The differentiating (last) electron enters ‘f’ orbitals, which lie inner to the second outermost (penultimate) shell. Thus these elements are also known as inner-transition elements.   There are two series of ‘f’ block elements, each having 14 elements. Lanthanides (atomic number 58 – 71) are those inner-transition elements in which 4 ‘f’ orbitals are progressively filled. Actinides (atomic number 90 – 103) are those elements in which 5 ‘f’ orbitals are progressively filled.   The ‘f’ block elements are placed at the bottom of the long form of periodic table in the form of two rows. Most actinides are radioactive.   The general outer electronic configuration of ‘f’ block elements is, (n – 2) f1-14 (n – 1) d0-1 ns2

Representative or Main Group Elements; These consist of all ‘s’ and ‘p’ block elements excluding the noble gases (group 18 elements). The chemical properties of the representative elements are determined by the number of valence electrons in their atoms. The number of valence electrons belonging to the above two blocks are:   ‘s’ block elements: 1 or 2: Same as the group number   ‘p’ block elements: 3 to 8: Group number minus ten   The representative elements are also known as normal or typical elements. Thus on the basis of their location in the periodic table, these are elements, which belong to groups 1, 2 and 13 to 17.

Noble gases or Aerogens These consist of gaseous elements of group 18, the last group in ‘p’ block. All noble gases have a general outermost electronic configuration of ‘ns’ and ‘np’. The only exception is helium having an outermost configuration of 1s2. Because they display these stable characteristics (having completely filled outermost shells), they do not show any chemical reactivity. Only Xenon shows limited chemical reactivity under severe conditions.

Transition elements ‘d’ block elements are called transition elements. These elements occupy groups 3 to 12 on the periodic table. There are four series of transitional elements, ’3d’, ’4d’, ’5d’ and ’6d’ depending on the energy levels of ‘d’ orbitals. They have a general outermost electronic configuration of (n – 1) d1-10 ns1-2.

Inner transition elements The ‘f’ block elements are called inner transition elements. There are two series of ‘f’ block elements. These are the ’4f’ and ’5f’ series called Lanthanides and Actinides. The general outer electronic configuration of ‘f’ block elements is, (n – 2) f1-14 (n – 1) d0-1 ns2.

Periodic Properties;

Atomic radii ; defined as one half of the inter-nuclear distance between the two centres of the neighbouring atoms in a covalently bonded molecule.

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Variation of Atomic Radii

Variation in a period Atomic radii in general, decrease with increase in atomic number, going from left to right in a period. This is explained on the basis of increasing nuclear charge along a period. The nuclear charge increases progressively by one unit while the corresponding addition of one electron takes place in the same principal shell. As the electrons in the same shell do not screen each other from the nucleus, the nuclear charge is not neutralized by the extra valence electron. Consequently the electrons are pulled closer to the nucleus by the increased effective nuclear charge resulting in the decrease in the size of the atom. In this way the atomic size goes on decreasing across the period.

element

Li

Be

B

C

N

O

F

Atomic radius-pm

135

90

82

70

70

66

64

The atomic radius abruptly increases in the case of noble gas element Neon as it does not form covalent bonds. So the value of Neon radius is Van der Waals radius which is considerably higher than the value of other covalent radii.

Variation in a group The atomic radii of elements increases from top to bottom in a group because the nuclear charge increases with increasing atomic number. Although, there is an increase in the principal quantum number from one atom to another, the number of electrons in the valence shell remain the same. The effect of increase in the size of the electron cloud out weighs the effect of increased nuclear charge and so the distance of the valence electron from the nucleus increases down the group. Thus the size of the atom goes on increasing down the group in spite of increasing nuclear charge.

element

Li

Na

K

Rb

Cs

Atomic radius

135

154

196

211

225

Ionic Radii These are radii of ions in ionic crystals. Ionic radius may be defined as the effective distance from the center of nucleus of an ion up to which it has an influence on its electron cloud. In ionic compounds the inter nuclear distance may be taken as equal to the sum of the ionic radii of the two ions. The inter nuclear distance in ionic crystals are obtained from X-ray studies.

Radius of the cation A cation is formed by the loss of one or more electrons from the gaseous atom. Thus, the whole of the outer most shell of electrons is removed resulting in the smaller size in the cation.   For example, in lithium atom, there is only one electron in the outermost ’2s’ shell. As the lithium atom changes to Li+ ion the outer most ’2s’ shell disappears completely. This disappearance results in the decrease in size.

With the removal of electrons from an atom the magnitude of the nuclear charge remains the same while the number of electrons decreases. As a result the nuclear charge acts on less number of electrons. The effective nuclear charge per electron increases and the electrons are more strongly attracted and pulled towards the nucleus. This causes a decrease in the size of the ion.

Radius of the anion The negative atom is formed by the gain of one or more electrons in the neutral atom. The number of electrons increases while the magnitude of nuclear charge remains the same. The same nuclear charge acts on larger number of electrons than were present in the neutral atom. The effective nuclear charge per electron is reduced and the electron cloud is held less tightly by the nucleus. This causes an increase in the size of the ion. Thus anions are larger in size than the corresponding atom.

Variation of ionic radii in a group The ionic radii in a particular group increases in moving from top to bottom because of the increase in the principal quantum number though the number of electrons in the valence shell remains the same.

Isoelectronic ions are ions with same number of electrons.

In isoelectronic series of ions, as the nuclear charge increases the electrons are pulled more and more strongly and the size decreases.

ion

Na +

Mg 2+

Al 3+

P 3- S 2-

Cl -

Ionic radius

95

65

50

212 184

181

Ionization Energy ;The amount of energy required to remove the most loosely bound electron from an isolated gaseous atom is called ionization energy (IE).

Ionization energy is also called as ionization potential because it is measured as the minimum potential required to remove the most loosely held electron from the rest of the atom. It is measured in the units of electron volts (eV) per atom or kilo joules per mole of atoms (kJ mol-1).   1 eV per atom = 96.64 kJ mol-1 = 23.05 k cal mol-1 Thus, the ionization energy gives the ease with which the electron can be removed from an atom. The smaller the value of the ionization energy, the easier it is to remove the electron from the atom.

.

Factors affecting Ionization Energy; 1. Size of the atom   2.Charge on the nucleus   3.Screening effect  . 4.Electronic arrangement  

Variation along a period ; The ionization energy increases with increasing atomic number in a period. This is because  1) The nuclear charge increases on moving across a period from left to right.  2) The atomic size decreases along a period though the main energy level remains the same.   Due to the increased nuclear charge and simultaneous decrease in atomic size, the valence electrons are more tightly held by the nucleus. Therefore more energy is needed to remove the electron and hence ionization energy keeps increasing.

However some irregularities have been noticed due to the extra stability of the half filled and completely filled configurations.   For example, the nuclear charge on Boron is more than Beryllium, yet there is slight decrease in ionization energy from Be to B. This is because, in boron the last electron goes to ’2p’ orbital which is at a slightly higher energy than ’2s’ orbital. Also, the electronic configuration of B is less stable than that of Be (has completely filled orbitals). Hence the ionization energy is less than that of Be. Similarly, nitrogen, which has half filled ’2p’ orbitals, is more stable than oxygen. Therefore the ionization energy of nitrogen is more than that of oxygen.

element

Li

Be

B

C

N

O

F Ne

Ionization energy

520

899

801

1056

1402

1314

1651 2050

Variation down a group ; The ionization energy gradually decreases in moving from top to bottom in a group. This is due to the fact that:   1)The nuclear charge increases in going from top to bottom in a group.   2)An increase in the atomic size due to an additional energy shell (level) ‘n’.   3)Due to the increase in the number of inner electrons there is an increase in the shielding effect on the outer most electron. The effect of increase in atomic size and the shielding effect is much more than the effect of increased nuclear charge.   As a result , the electron becomes less firmly held to the nucleus and so the ionization energy decreases as we move down the group.

element

Li

Na

K

Rb

Cs

Ionization energy

520

496

419

403

376

Electron Affinity; Electron affinity is the amount of energy released when an electron is added to an isolated gaseous atom.

Electron affinity is the ability of an atom to hold an additional electron. If the atom has more tendency to accept an electron then the energy released will be large and consequently the electron affinity will be high. Electron affinities can be positive or negative. It is taken as positive when an electron is added to an atom. It is expressed as electron volts per atom (eV per atom) or kilo joules per mole.

Variation along a period -  The size of an atom decreases and the nuclear charge increases on moving across a period. This results in greater attraction for the incoming electron. Hence the electron affinity increases in a period from left to right.

Variation down a group – As we move down a group the atomic size and nuclear size increases. As the effect of increase in atomic size is more pronounced the additional electron feels less attracted by the large atom. Consequently the electron affinity decreases.

Electronegativity The relative tendency of an atom in a molecule to attract a shared pair of electrons towards itself is termed as Electronegativity.

The value of electronegativity of an element describes the ability of its atom to compete for electrons with the other atom to which it is bonded. Electronegativity is however not the property of an isolated atom. Electronegativity is measured on a number of scale levels, the most commonly used are of Pauling or Mulliken.

Factors Affecting Electronegativity

1)Atomic size   2)Ionisation energy and electron affinity   3)Number and nature of atoms   4)Type of hybridization   5)Charge on the ion

Variation along a period -  As the nuclear charge increases from going left to right in a period because the electrons enter the same shell, the shielding is less effective. Thus the increased nuclear charge attract the shared pair of electrons more strongly resulting in higher electronegativity from going left to right in a period.

element

Li

Be

B

C

N

O

F

electronegetivity

1.0

1.5

2

2.5

3

3.5

4

Variation down the group -  Electronegativity decreases down the group because the atomic size increases. The larger the size of the atom the lesser the tendency to attract the shared pair of electrons.

element

Li

Na

K

Rb

Cs

electronegetivity

1

0.9

0.8

0.8

0.7

Comments

  • sandeep  On March 22, 2009 at 3:39 am

    ………….

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