In the year 1869 Lother Mayer plotted various properties of the elements like atomic volume, density, melting and boiling points etc. against atomic weights and discovered that the variation in propertiess of the elements is a periodic function of the atomic weight. The most noteworthy attempt to classify the elements is made by Russian Scientist Dimitri Mendeleev.
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| Periodic table |
According to Mandeleev 'the Physical and Chemical properties of the elements are the periodic functions of their atomic weights.' On the basis of this law, he arranged elements. This arrangement of elements is called Mendeleev's Periodic Table. Simultaneously working independently, a German chemist J, Lother Mayer reported identical observations.
Later on knowledge of electronic model of atomic structure, Mosley (1913) arranged the elements in increasing order of atomic numbers and prepared a table which is called as Modern Periodic Table.
Modern Periodic Table:
Modern periodic table or long form of the periodic table is based on 'Modern Periodic Table Law'.
Modern periodic law states that "The physical and chemical properties of the elements are periodic functions of their atomic numbers".
According to International Union of Pune and Applied Chemistry (IUPAC) recommendations, for the elements having atomic number greater than 100, a systematic nomenclature is derived directly from its atomic number using the numerical roots for zero and digits 1 to 9.
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| IUPAC name of elements |
Locations of different types of elements in long form of the periodic table are as follows:
s block elements, i.e. alkali metals and alkaline earth metals are placed at the left hand side, p block elements at the right hand side, d block elements occupy central position and f block elements are placed at the bottom of table.
Classification of elements:
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| Classification of Elements |
Classification of elements in the periodic table is based on the type of orbital in the incomplete shell where the last electron enters.
On this basis, the elements are classified into four blocks:
1) s-block elements.
2) p-block elements.
3) d-block elements.
4) f-block elements.
s-block elements:
In the atoms of these elements, the last electron enters s-orbital of the valence shell. General valence shell electronic configuration of these elements is ns¹ or ns². Elements of group 1 (alkali metals) and group 2 (alkaline earth metal) belong to s-block. These elements are placed on extreme left of the periodic table.
p-block elements:
In the atoms of these element, the last electron enters in the p-orbital of valence shell. Valence shell electronic configuration of these elements varies from ns²np¹ to ns² np6. Elements of groups 13,14,15,16,17 and 18 except helium belong to p-block element. These elements are placed on extreme right of periodic table. Most of these elements are non-metals, some are metalloids while a few are metals. Eighteen group consists of noble gases.
d-block elements:
In the atoms of these elements the last electron enters into s-orbital of penultimate shell (inner to ultimate) i.e. (n-1)d orbital. Electronic configuration of these elements varies from (n-1)d¹ns^1-2 to (n-1)d^10 ns^1-2. Elements of group 3 to 12 are included in d-block. These elements are placed in middle portion of periodic table, between the s and p block elements. Since the properties of d-block elements are in intermediate between s and p block elements and are called transition elements.
The d-block elements are further classified into 4 series, depending upon filling of 3d, 4d, 5d, 6d, orbital.
3d-series:
This series has 10 elements from Sc (Z=21) to Zn (Z=30).
4d-series:
This series has 10 elements from Y (Z=39) to Cd (Z=48).
5d-series:
This series has 10 elements from La(Z=57) and from Hf (Z=58) to Hg (Z=80).
6d-series:
Ac (Z=89) and from Rf (Z=104) to Uub (Z=112).
f-block elements:
In the atoms of these elements the last electron enters into f-orbital of pre-penultimate shell. Electronic configuration of these elements is (n-2)f^1-14 (n-1)d^0-1 ns², i.e. last three shells are incomplete. These elements are included in 3rd group of sixth and seventh period and are all metals. This block consist of series of lanthanoids and actinoids, placed at the bottom of the periodic table. As addition of elections takes place in the f-orbital of pre-penultimate shell these elements are called inner transition elements.
General characteristics of s, p, d, and f block elements.
Characteristics of s-block elements:
a) s-block includes Alkali metals and Alkaline earth metals.
b) Valence shell electronic configuration of alkali and alkaline earth metals is ns¹ and ns² respectively.
c) By losing one and two electrons respectively from their valence shell, these elements acquire stable inert gas configuration and hence show following properties:
1) These elements are elecropositive in nature.
2) These elements show 1+, 2+ oxidation states which are very stable.
3) These elements are powerful reducing agent.
4) These elements are good conductor of heat and electricity.
5) These elements form ionic compounds.
6) Form strongly basic oxides of these elements.
7) These elements have low ionization energy, electron affinity and electro negativity, melting point, boiling point.
8) Ions of these elements impart characteristic colours to Bunsen flame.
9) These elements are soft and light.
10) These elements are highly reactive and hence are never found in free state in nature.
Characteristics of p-block elements:
1) Valence shell electronic configuration of p-block elements is ns²np¹ to ns² np6.
2) The p-block consists of metals, non-metals, metalloids and noble gases.
3) The p-block elements can form covalent compounds and ionic compounds.
4) Many of the p-block elements exhibit allotropy.
5) Metals are good conductor of heat and electricity, non-metals are insulators, silicon, germanium are semi-conductors.
6) Most of these elements form acidic oxides.
7) These elements form colourless compounds.
8) These elements have high ionization energies.
9) preceding the noble gases family, there are two chemically important groups of non-metals, viz. Halogens (group 17) and chalcogens (group 16). Elements of group 16 and group 17 have high electron affinity.
Characteristics of d-block elements:
1) The electronic configuration of d-block element is (n-1)d^ 1-10 ns^ 1-2.
2) All these elements are heavy metals.
3) These elements are good conductor of heat and electricity.
4) These elements are malleable and ductile.
5) Due to the presence of incompletely filled (n-1) d orbital, d-block elements from coloured compounds such as e.g. CuSO4, NiCl2, They form complex compounds and are used as catalysts. These elements form paramagnetic compounds and show variable oxidation state.
6) d-block metals from complex compounds and metals are used as catalysts.
7) d-block elements form industrial compounds.
Characteristics of f-block elements:
1) Electronic configuration of f-block elements is (n-2)f^1-14 (n-1)d^0-1 ns².
2) Most of the elements are metallic, having high melting and boiling point.
3) f-block elements form coloured compounds.
4) Some of these elements are radioactive.
5) As f-block elements have unpaired electrons in their f-orbital, they show paramagnetic properties.
Characteristics of noble gases or 18 group elements:
These elements are known as rare gases or inert gases or noble gases.
1) The general electronic configuration of these elements is ns² np6. The first member of 18th group element. However has electronic configuration 1s².
2) Due to the presence of completely filled s and p outermost orbitals octet is completed. Hence these elements are chemically inert.
3) These elements are zero valent.
4) These elements form mono-atomic molecules.
5) These elements have large atomic size.
6) These elements have high ionization energy, zero electro negativity and zero electron affinity.
7) Noble gases are diamagnetic.
8) These elements have low melting and boiling point.
Periodic trends in properties:
Periodic properties:
The modern periodic law states that "The physical and chemical properties of the elements are periodic functions of their atomic numbers".
Certain characteristics properties of elements that reappear at definite interval are known as periodic properties.
There is a gradual increase or decrease in a particular property of an element, in the given period or group, with increase in atomic number of an element. These properties depend on valence shell electronic configuration of atom of the elements. Some of the properties like atomic size, ionic radii, ionization energy, electron affinity, alectro negativity are discussed below:
Atomic size (atomic radius) :
Definition:
Atomic size or atomic radius is defined as the distance of valence shell of electrons from the centre of the nucleus of an atom.
Factors affecting atomic size:
1) Number of shells:
The atomic size increases with increase in the number of electronic shells. Atomic size is proportional to numbers of electronic shells.
2) Nuclear charge:
Atomic size decreases with increase in the nuclear charge. If nuclear charges is more, then nucleus attracts the electrons towards itself and atomic size decreases. i.e. atomic size is inversely proportional to the nuclear charge.
3) Screening effect or Shielding effect:
The electrons are present in the shells, between the nucleus and valence shell. The inner electrons screen or shield the nucleus from valence electrons. Hence force of attraction between the nucleus and valence electrons decreases. This effect is called as screening effect or shielding effect. With increase in screening or shielding effect, atomic size increases.
Atomic size is directly proportional to shielding effect.
Trends in atomic size:
Period:
In any given period, the atomic size of an element, decreases from left to right with increase in the atomic number.
In a given period, as the atomic number increases from left to right, nuclear charge also increases gradually by addition of electrons in the same shell. The attraction between nucleus and valence electrons increases and atomic size decreases.
Group:
In a group, atomic size increases from top to bottom as we move down the group in the periodic table.
In a group, as we move down the group with increase in the atomic number, nuclear charge increases along with increase in the number of shells. Due to this magnitude of effective nuclear charge decreases and attraction between nucleus and valence electron becomes less and Atomic size increases in a group from top to bottom.
Ionic radius:
Definition:
Ionic radius is defined as the distance of valence shell electrons from the centre of nucleus in an ion.
When an atom loses one or more electrons, cation is formed. As the no of electrons in a cation is less than the parent atom, attraction between the nucleus and remaining electrons in cation increases. As a result size, of cation is less than the atom.
Trends in ionic radius:
Period:
In a given period, ionic radius decreases from left to right with increase in the atomic number.
Group:
In a given group, ionic radius increases down the group with increase in the atomic number.
Ionization energy:
Ionization energy of an element is defined as the amount of energy required to remove electron, from an isolated gaseous atom in its ground state.
The energy required to remove first electron valence shell is called as first ionization energy. The energy required to remove second electron is called as second ionization energy and so on.
Factors affecting ionization energy:
1) Atomic size or atomic radius:
If atomic size increases, attraction between nucleus and valence shell electrons decreases and ionization energy decreases.
2) Nuclear charge:
As the magnitude of nuclear charge increases in an atom, attraction between the nucleus and electrons in the valence shell increases. Hence higher energy is required to remove the electron from valence shell.
Ionization energy is directly proportional to nuclear charge.
3) Shielding or Screening effect:
More the number of inner filled shells more will be the shielding effect and effective nuclear force of attraction becomes less and hence ionization energy decreases.
Ionization energy is inversely proportional to Shielding effect or Screening effect.
4) Type of electrons to be removed:
The electrons in s orbital are closer to the nucleus than electrons in p, d and f orbitals of same principal shell. Therefore the attraction between nucleus and electron of s orbital is more. Hence maximum energy is required to remove electrons from s orbital. The amount of energy goes on decreasing for the electrons of p, d, f orbitals in the order.
s> p > d> f
5) Electronic configuration:
f-block elements:
In the atoms of these elements the last electron enters into f-orbital of pre-penultimate shell. Electronic configuration of these elements is (n-2)f^1-14 (n-1)d^0-1 ns², i.e. last three shells are incomplete. These elements are included in 3rd group of sixth and seventh period and are all metals. This block consist of series of lanthanoids and actinoids, placed at the bottom of the periodic table. As addition of elections takes place in the f-orbital of pre-penultimate shell these elements are called inner transition elements.
General characteristics of s, p, d, and f block elements.
Characteristics of s-block elements:
a) s-block includes Alkali metals and Alkaline earth metals.
b) Valence shell electronic configuration of alkali and alkaline earth metals is ns¹ and ns² respectively.
c) By losing one and two electrons respectively from their valence shell, these elements acquire stable inert gas configuration and hence show following properties:
1) These elements are elecropositive in nature.
2) These elements show 1+, 2+ oxidation states which are very stable.
3) These elements are powerful reducing agent.
4) These elements are good conductor of heat and electricity.
5) These elements form ionic compounds.
6) Form strongly basic oxides of these elements.
7) These elements have low ionization energy, electron affinity and electro negativity, melting point, boiling point.
8) Ions of these elements impart characteristic colours to Bunsen flame.
9) These elements are soft and light.
10) These elements are highly reactive and hence are never found in free state in nature.
Characteristics of p-block elements:
1) Valence shell electronic configuration of p-block elements is ns²np¹ to ns² np6.
2) The p-block consists of metals, non-metals, metalloids and noble gases.
3) The p-block elements can form covalent compounds and ionic compounds.
4) Many of the p-block elements exhibit allotropy.
5) Metals are good conductor of heat and electricity, non-metals are insulators, silicon, germanium are semi-conductors.
6) Most of these elements form acidic oxides.
7) These elements form colourless compounds.
8) These elements have high ionization energies.
9) preceding the noble gases family, there are two chemically important groups of non-metals, viz. Halogens (group 17) and chalcogens (group 16). Elements of group 16 and group 17 have high electron affinity.
Characteristics of d-block elements:
1) The electronic configuration of d-block element is (n-1)d^ 1-10 ns^ 1-2.
2) All these elements are heavy metals.
3) These elements are good conductor of heat and electricity.
4) These elements are malleable and ductile.
5) Due to the presence of incompletely filled (n-1) d orbital, d-block elements from coloured compounds such as e.g. CuSO4, NiCl2, They form complex compounds and are used as catalysts. These elements form paramagnetic compounds and show variable oxidation state.
6) d-block metals from complex compounds and metals are used as catalysts.
7) d-block elements form industrial compounds.
Characteristics of f-block elements:
1) Electronic configuration of f-block elements is (n-2)f^1-14 (n-1)d^0-1 ns².
2) Most of the elements are metallic, having high melting and boiling point.
3) f-block elements form coloured compounds.
4) Some of these elements are radioactive.
5) As f-block elements have unpaired electrons in their f-orbital, they show paramagnetic properties.
Characteristics of noble gases or 18 group elements:
These elements are known as rare gases or inert gases or noble gases.
1) The general electronic configuration of these elements is ns² np6. The first member of 18th group element. However has electronic configuration 1s².
2) Due to the presence of completely filled s and p outermost orbitals octet is completed. Hence these elements are chemically inert.
3) These elements are zero valent.
4) These elements form mono-atomic molecules.
5) These elements have large atomic size.
6) These elements have high ionization energy, zero electro negativity and zero electron affinity.
7) Noble gases are diamagnetic.
8) These elements have low melting and boiling point.
Periodic trends in properties:
Periodic properties:
The modern periodic law states that "The physical and chemical properties of the elements are periodic functions of their atomic numbers".
Certain characteristics properties of elements that reappear at definite interval are known as periodic properties.
There is a gradual increase or decrease in a particular property of an element, in the given period or group, with increase in atomic number of an element. These properties depend on valence shell electronic configuration of atom of the elements. Some of the properties like atomic size, ionic radii, ionization energy, electron affinity, alectro negativity are discussed below:
Atomic size (atomic radius) :
Definition:
Atomic size or atomic radius is defined as the distance of valence shell of electrons from the centre of the nucleus of an atom.
Factors affecting atomic size:
1) Number of shells:
The atomic size increases with increase in the number of electronic shells. Atomic size is proportional to numbers of electronic shells.
2) Nuclear charge:
Atomic size decreases with increase in the nuclear charge. If nuclear charges is more, then nucleus attracts the electrons towards itself and atomic size decreases. i.e. atomic size is inversely proportional to the nuclear charge.
3) Screening effect or Shielding effect:
The electrons are present in the shells, between the nucleus and valence shell. The inner electrons screen or shield the nucleus from valence electrons. Hence force of attraction between the nucleus and valence electrons decreases. This effect is called as screening effect or shielding effect. With increase in screening or shielding effect, atomic size increases.
Atomic size is directly proportional to shielding effect.
Trends in atomic size:
Period:
In any given period, the atomic size of an element, decreases from left to right with increase in the atomic number.
In a given period, as the atomic number increases from left to right, nuclear charge also increases gradually by addition of electrons in the same shell. The attraction between nucleus and valence electrons increases and atomic size decreases.
Group:
In a group, atomic size increases from top to bottom as we move down the group in the periodic table.
In a group, as we move down the group with increase in the atomic number, nuclear charge increases along with increase in the number of shells. Due to this magnitude of effective nuclear charge decreases and attraction between nucleus and valence electron becomes less and Atomic size increases in a group from top to bottom.
Ionic radius:
Definition:
Ionic radius is defined as the distance of valence shell electrons from the centre of nucleus in an ion.
When an atom loses one or more electrons, cation is formed. As the no of electrons in a cation is less than the parent atom, attraction between the nucleus and remaining electrons in cation increases. As a result size, of cation is less than the atom.
Trends in ionic radius:
Period:
In a given period, ionic radius decreases from left to right with increase in the atomic number.
Group:
In a given group, ionic radius increases down the group with increase in the atomic number.
Ionization energy:
Ionization energy of an element is defined as the amount of energy required to remove electron, from an isolated gaseous atom in its ground state.
The energy required to remove first electron valence shell is called as first ionization energy. The energy required to remove second electron is called as second ionization energy and so on.
Factors affecting ionization energy:
1) Atomic size or atomic radius:
If atomic size increases, attraction between nucleus and valence shell electrons decreases and ionization energy decreases.
2) Nuclear charge:
As the magnitude of nuclear charge increases in an atom, attraction between the nucleus and electrons in the valence shell increases. Hence higher energy is required to remove the electron from valence shell.
Ionization energy is directly proportional to nuclear charge.
3) Shielding or Screening effect:
More the number of inner filled shells more will be the shielding effect and effective nuclear force of attraction becomes less and hence ionization energy decreases.
Ionization energy is inversely proportional to Shielding effect or Screening effect.
4) Type of electrons to be removed:
The electrons in s orbital are closer to the nucleus than electrons in p, d and f orbitals of same principal shell. Therefore the attraction between nucleus and electron of s orbital is more. Hence maximum energy is required to remove electrons from s orbital. The amount of energy goes on decreasing for the electrons of p, d, f orbitals in the order.
s> p > d> f
5) Electronic configuration:
Atoms having completely filled or half filled orbitals are more stable. Due to this extra stability Ionization energy of such atoms is high.
Trends in ionization energy:
Trends in ionization energy:
Period:
In a period, ionization energy increases from left to right. It reaches to maximum for inert gas element.
In a period, as atomic number increases from left to right, nucleus charge also increases. Though there is increase in nuclear charge electrons get added to the same shell. Hence effective nuclear force of attraction for valence electrons increases and atomic size decreases. As atomic size decreases, it is more difficult to remove the electron from valence shell and hence ionization energy increases from left to right in a period.
Group:
In a group, ionization energy value, decreases from top to bottom, with increase in the atomic number.
In a group, as atomic number increases down the group, nuclear charge increases along with the number of shell. Hence effective nuclear force of attraction decreases for valence electrons and atomic size increases and therefore the ionization energy decreases down the group in the order H> Li > Na >K > Rb >Cs.
Electron Affinity:
The electron affinity is defined as the amount of energy released, when neutral gaseous atom, accepts an electron to form an anion.
Greater the electron affinity greater the non-metallic character. Unlike Ionization energies electron affinities may have negative values.
Factors affecting the electron affinity:
1) Atomic size or Atomic Radius:
Electron affinity is inversely proportional to the atomic size. Electron affinity increases with decrease in atomic size.
2) Nuclear charge:
Electron affinity increases with increase in nuclear charge. Electron affinity is directly proportional to nuclear charge.
3) Screening effect:
Electron affinity is inversely proportional to shielding effect or screening effect of electrons.
4) Electronic configuration:
The elements having half filled and completely filled valence shells have extra stability and show very little tendency to accept electrons and hence electron affinity is less.
Trends in electron affinity:
Period:
In a given period, electron affinity increases with increase in atomic number from left to right. As atomic number increases from left to right, atomic size decreases and electron affinity increases from left to right.
Group:
In a group, as we move down the group, electron affinity decreases with increase in the atomic number.
In a group, as atomic number increases, nuclear charge also increases along with atomic size. Since atomic size increases, additional electrons feel less attraction and therefore electron affinity decreases from top to bottom.
Electronegativity:
In a period, as atomic number increases from left to right, nucleus charge also increases. Though there is increase in nuclear charge electrons get added to the same shell. Hence effective nuclear force of attraction for valence electrons increases and atomic size decreases. As atomic size decreases, it is more difficult to remove the electron from valence shell and hence ionization energy increases from left to right in a period.
Group:
In a group, ionization energy value, decreases from top to bottom, with increase in the atomic number.
In a group, as atomic number increases down the group, nuclear charge increases along with the number of shell. Hence effective nuclear force of attraction decreases for valence electrons and atomic size increases and therefore the ionization energy decreases down the group in the order H> Li > Na >K > Rb >Cs.
Electron Affinity:
The electron affinity is defined as the amount of energy released, when neutral gaseous atom, accepts an electron to form an anion.
Greater the electron affinity greater the non-metallic character. Unlike Ionization energies electron affinities may have negative values.
Factors affecting the electron affinity:
1) Atomic size or Atomic Radius:
Electron affinity is inversely proportional to the atomic size. Electron affinity increases with decrease in atomic size.
2) Nuclear charge:
Electron affinity increases with increase in nuclear charge. Electron affinity is directly proportional to nuclear charge.
3) Screening effect:
Electron affinity is inversely proportional to shielding effect or screening effect of electrons.
4) Electronic configuration:
The elements having half filled and completely filled valence shells have extra stability and show very little tendency to accept electrons and hence electron affinity is less.
Trends in electron affinity:
Period:
In a given period, electron affinity increases with increase in atomic number from left to right. As atomic number increases from left to right, atomic size decreases and electron affinity increases from left to right.
Group:
In a group, as we move down the group, electron affinity decreases with increase in the atomic number.
In a group, as atomic number increases, nuclear charge also increases along with atomic size. Since atomic size increases, additional electrons feel less attraction and therefore electron affinity decreases from top to bottom.
Electronegativity:
Electro-negatively of an atom in a molecule is defined as the tendency of an atom to attract towards itself the shared pair of electrons.
Electro-negativity values of combining atoms give idea about the nature of chemical bond. When combining atoms have nearly same values of electro-negativities covalent bond is formed between the two atoms e.g. Cl2. When combining atoms have large difference in electronegativity values a covalent bond with ionic character is formed between the two atoms. e.g. HCl, HF.
Factors affecting electro negativity are:
1) Atomic size:
Electro-negativity increases with decrease in atomic size. Electro negativity is inversely proportional to the atomic size.
2) Nuclear charge:
3) Screening effect:
Electro-negativity values of combining atoms give idea about the nature of chemical bond. When combining atoms have nearly same values of electro-negativities covalent bond is formed between the two atoms e.g. Cl2. When combining atoms have large difference in electronegativity values a covalent bond with ionic character is formed between the two atoms. e.g. HCl, HF.
Factors affecting electro negativity are:
1) Atomic size:
Electro-negativity increases with decrease in atomic size. Electro negativity is inversely proportional to the atomic size.
2) Nuclear charge:
Electro negativity is directly proportional to the nuclear charge.
3) Screening effect:
With increase in the s screening effect or shielding effect, electro negativity decreases.
Trend in electronegativity:
1. In a period:
In a period, as atomic number increases from left to right atomic size decreases and nuclear charge increases from left to right. Hence, tendency to attract shared pair of electron increases from left to right in a period.
2. In a group:
In a group, electronegativity decreases down the group.
In a group, with increase in the atomic number, atomic size increases and due to which tendency to attract shared pair of electron decreases and hence electronegativity decreases down the group.
Trend in electronegativity:
1. In a period:
In a period, as atomic number increases from left to right atomic size decreases and nuclear charge increases from left to right. Hence, tendency to attract shared pair of electron increases from left to right in a period.
2. In a group:
In a group, electronegativity decreases down the group.
In a group, with increase in the atomic number, atomic size increases and due to which tendency to attract shared pair of electron decreases and hence electronegativity decreases down the group.
