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NEET Important Chapter: Periodic Table and Periodicity in Properties

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Introduction

Introduction

Need of Classification of Elements

Generally, Elements are the basic constituents of all types of matter. Since ancient times, discovery of chemical elements has been an ongoing process. At the beginning of the 18th century, only a few elements were known. But during the 18th and early 19th century, advances in chemistry made it easier to isolate elements from their compounds. So, the number of known elements had more than doubled from 31 in 1800 to 63 by 1865. At present, 118 elements are known. 


As the number of known elements increased, it became rather impossible to individually study these elements and their innumerable compounds, so scientists began to investigate the possibilities of classifying them in useful ways. Hence, the basic object and need of classification can be stated in the words of Huxley, “The ideal or actual  arrangement together of those which are alike and separation of those which are unlike; the objective of this arrangement being primarily to disclose the correlations, or laws, or properties, or circumstances and secondarily to facilitate the operation of the mind in receiving and then retaining in the memory the characteristics of the objects in question.”


Classification of Elements and Periodicity in Properties is an important topic of the NEET Exam. Electronegativity, ionization enthalpy, atomic radii, ionic radii, electron gain enthalpy, valency and so on are all significant topics. This article provides an example of the types of questions that might be asked about this subject.


Important Topics of Chemical Thermodynamics

  • Dobereiner's Law of Triads

  • Newland's Law of Octaves

  • Mendeleev's Periodic Law

  • Modern Periodic Law

  • Atomic radii

  • Ionic radii

  • Electronegativity

  • Ionization Enthalpy

  • Electron gain enthalpy

  • Valency

  • Metallic character

  • Non-metallic character


Earlier Attempts of Classification of Elements

1. Dobereiner's Law of Triads: (In between 1815 and 1829)

  • This law states that “In certain triads (groups of three elements which possessed similar chemical properties) the atomic mass of the middle element was merely the arithmetic mean of the atomic masses of the first and third elements and the properties of the middle element were in between those of the end members.” These families of three elements became known as Dobereirier's Triads.

  • Drawback of this law was that the concept was not applicable to all the known elements but only to a limited number of elements. 

  • Example:

Triads

The Mean of atomic masses of 1st and 3rd elements

Li          Na           K

7           23           39

(7 + 39)/2 = 23

2. Newland's Law of Octaves: (In 1864)

  • Newland’s law states that, “When the elements arranged in order of their increasing atomic masses, the eighth succeeding element was the repetition of the first one like 8th note of the musical scale.

  • Newland’s law worked quite well for the lighter elements. For example, Elements H, F and Cl show similar properties and similarly Li, Na and K exhibit the same characteristics. This law, however, fails in the case of heavier elements as manganese has been placed along with nitrogen and phosphorus or iron has been placed along with oxygen and sulfur, i.e., dissimilar elements have been grouped which is against the law of this classification.


Mendeleev's Periodic Law And Original Periodic Table 

  • Mendeleev’s law states that, "If the elements are arranged in order of increasing atomic masses, the properties of the elements are a periodic function of their atomic masses.” 

  • A tabular arrangement of elements in rows and columns, highlighting the regular repetition of properties of elements is called a periodic table. Elements with similar properties were present in vertical columns called groups. The horizontal rows were called periods. 

  • When Mendeleev presented his periodic table, only 63 elements were known. Hence, a number of gaps for unknown elements were left in the table and the properties of these unknown elements were predicted on the basis of periodic law. These predicted properties helped the future scientists in the discovery of unknown elements. For example, gallium and germanium were not known when he presented his periodic table. These elements were named eka-aluminum and eka-silicon.

  • This table consisted of a number of defects such as anomalous pairs, position of isotopes, position of lanthanides and actinides, etc. Some of these defects do not exist in the modern Mendeleev periodic table.


Modern Periodic Law and the Long Form of Periodic Table

  • This law states that, "When the elements are arranged in order of their increasing atomic numbers, the physical and chemical properties of the elements are periodic functions of their atomic numbers.” 

  • With the replacement of atomic mass by atomic number as the basis of classification, two main defects of the original table based on atomic masses disappear. These defects are: (i) Anomalous pairs (ii) Position of isotopes.

  • In the Long Form of Periodic Table, there are 7 periods and 18 groups. 

  • In this, elements of the same group have a similar outer electronic configuration.

  • In this, the period number corresponds to the highest principal quantum number (n) of the elements in the period.

  • In the modern periodic table, 14 elements of both the 6th and 7th period are placed in the separate panels at the bottom. These elements are known as lanthanides and actinides respectively. 


Electronic Configurations of the Elements and Periodic Table

  • The period number indicates the value of (n) for the outermost shell or valence shell. An element placed in 2nd period will have its outermost electrons in 2s or 2p orbitals.

  • Elements of the same group have similar valence shell electronic configurations. Because they have the same number of electrons in the outer orbitals.


Position of the Element on the Basis of Electronic Configuration

  • The block, period and group of an element can be easily decided with the help of its electronic configuration of the element.

  •  Classification of Elements:

Elements

Electronic configuration

Position in Modern periodic table

Type of elements

Nature

s-Block Elements

ns1-2 (n = 1 to 7)

Group 1 and 2

Alkali metals and Alkaline earth metals

Metals

p-Block Elements

ns2 np1-6 (n = 2 to 7)

Group 13 - 18

Representative or main group elements

Metalloids & non-metals but some are metals also

d-Block Elements

(n – 1)d1-10 ns1-2 (n = 4 - 7)

Group 3 - 12 (3d, 4d and 5d series)

Transition elements

Metals

f-Block Elements

(n-2)f1-14 (n-1)s2 (n-1)p6 (n-1)d0-1 ns2 (n = 6 & 7)

Group 3 (4f series - Lanthanides and 5f series - Actinides)

Inner transition elements

Rare earth metals and actinides are radioactive.


Periodic Trends in Properties of Elements 

Atomic Radius

  • When we move from left to right across a period, atomic radius gradually decreases. Because as we move left to right in a period, the atomic number of the elements increases so nuclear charge increases while the number of shells in elements remain the same. 

  • Noble gases show exceptional behavior. The atomic radii of inert gases suddenly increase as compared to its predecessor halogen atom. And the reason for this  exceptional behavior is that atomic radius refers to van der Waals radius in case of noble gases while in case of other elements it refers to covalent radius. 

  • When we move from top to bottom in a group, atomic radii gradually increase as nuclear charge and number of shells also increase. 


Ionization Enthalpy

  • When we move from left to right across a period, ionization energy gradually increases. Because as we move left to right in a period atomic size or atomic radius decreases while nuclear charge increases. 

  • Beryllium has more first ionization energy than that of Boron. Because beryllium has a full - filled s -orbital and more energy is required to remove an electron from half or completely filled orbitals. Due to this reason noble gases also show exceptionally high ionization energies. 

  • When we move from top to bottom in a group, ionization energy gradually decreases as atomic radius increases.


Electronegativity 

  • When we move from left to right across a period, electronegativity increases in the periodic table. Fluorine is the most electronegative element. Because the nuclear charge increases in an atom, its electron loving character also increases. 

  • When we move from top to bottom in a group, electronegativity decreases. 


Valency

  • Valency is defined as “the combining capacity of an atom”.

  • When we move from left to right across a period in the periodic table, first valency increases then decreases. 

  • When we move top to bottom in a group, valency remains the same. Elements of the same groups show the same valency as they have same valence electrons.


Metallic Character of the Elements  

  • When we move from left to right across a period in the periodic table, the metallic character of elements decreases. 

  • When we move top to bottom in a group of the periodic table, the metallic character of elements increases. 


Non - Metallic Character of the Elements

  • When we move from left to right across a period in the periodic table, the nonmetallic character of elements increases.

  • When we move from top to bottom in a group of periodic tables, non metallic character decreases.


Ionic Radii

  • In any particular group, the ions (cations or anions) increase in size on moving from top to bottom. Because the number of shells increases. 

  • The size of the cations of the same element decreases as the positive charge increases. For example, Pb4+ is smaller in size than Pb2+

  • In a set of species having the same number of electrons (isoelectronic species), the size decreases as the charge on the nucleus increases. 

  • The radius of cation (positive ion) is always smaller than that of its parent atom. But the radius of an anion (negative ion) is always larger than that of the parent atom. 


Electron Gain Enthalpy

  • In general, the electron gain enthalpy becomes less negative in going from top to bottom in a group and more negative in going from left to right in a period.

  • Electron gain enthalpies of some of the members of alkaline earth metals, noble gases and nitrogen are positive because they possess stable configurations.

  • Halogens have highest negative electron gain enthalpies because halogens have a very strong tendency to accept an additional electron.


Solved Examples from the Chapter

Example 1: Alkali metals have very high second ionization enthalpy values. Why? 

Solution: 

General electronic configuration of alkali metals is ns1. By losing this outermost shell electron, alkali metals acquire the configuration of the nearest noble gas element. Hence, it is a stable configuration. And removing the electron from this configuration is difficult, as it requires a large amount of energy. Thus, alkali metals have very high second ionization enthalpy values. 

Key point to remember: Ionization energy of an atom is defined as the amount of energy required to remove an electron from the gas state to form a provided atom or ion. And removing the electron from half filled or full filled configuration (stable configuration) is difficult, as it requires a large amount of energy.

Example 2: Among the elements with atomic numbers 9, 12 and 36, identify by atomic number of an element which is highly electronegative?

Solution: Electronic configuration of the elements with given atomic numbers 9, 12 and 36 are: -

Atomic number 9 = 2, 7

Atomic number 12 = 2, 8, 2

And atomic number 36 = 2, 8, 18, 8

The element with given atomic number 9 can accept one more electron to have 8 electrons in the outermost orbit i.e. to attain stable electronic configuration. Thus an element with atomic number 9 is highly electronegative. 

An element with atomic number 12 will lose 2 electrons to attain the nearest noble gas configuration so it will be an electropositive element. While an element with atomic number 36 has a stable electronic configuration therefore it will neither gain nor lose the electron.

Key point to remember: Electronegativity is the relative tendency of an atom to attract the shared electron pair towards itself. 


Solved Questions From The Previous Year Question Papers

Question 1:  From the following pairs of ions which one is not an iso-electronic pair? Options:  

(a) Mn2+, Fe3+                                             (b) Fe2+, Mn2+                       

(c) O2–, F                                                   (d) Na+, Mg2+

Solution: Isoelectronic pairs are defined as which have the same number of electrons. 

The number of electrons in Mn2+ = 23, Fe3+ = 23 (iso-electronic pair)

The number of electrons in Fe2+ = 24, Mn2+ = 23 (these ions not an iso-electronic pair)

The number of electrons in O2– = 10, F = 10 (iso-electronic pair)

The number of electrons in Na + = 10, Mg2+ = 10 (iso-electronic pair)

Thus, the correct answer is option (b) Fe2+, Mn2+

Trick: Isoelectronic pairs have the same number of electrons. So, with the help of the total number of electrons in atoms/ions, isoelectronic pairs are found.

Question 2: For the given second period elements the correct increasing order of first ionization enthalpy is:

(a) Li < Be < B < C < N < O < F < Ne

(b) Li < B < Be < C < O < N < F < Ne

(c) Li < B < Be < C < N < O < F < Ne

(d) Li < Be < B < C < O < N < F < Ne

Solution: Ionization enthalpy generally increases from left to right across a particular period. But ‘Be’ and ‘N’ both have a full filled and half filled electronic configuration respectively that means both have comparatively more stable valence shells than ‘B’ and ‘O’. Thus, the correct answer is option (b) Li < B < Be < C < O < N < F < Ne.

Trick:  Ionization enthalpy is defined as the minimum amount of energy required to remove the most loosely bound electron from an isolated atom in the gaseous state of an element so as to convert it into gaseous monovalent positive ions. And generally it increases from left to right across a period.

Question 3: Magnesium (Mg) reacts with an element (X) to form an  ionic  compound.  If  the  ground  state electronic configuration of (X) is 1s2 2s2 2p3, then the simplest formula for this compound is:

(a) Mg2X                         (b) MgX2                          (c) Mg2X3                      (d) Mg3X2

Solution: Electronic configuration of element (X) = 1s2 2s2 2p3 

Element (X) will gain 3 electrons to attain the noble gas electronic configuration. So, the valency of element (X) will be = 3.

And Electronic configuration of Magnesium (Mg) = 1s2 2s2 2p6 3s2

Mg will lose 2 electrons to attain the noble gas electronic configuration. So, the valency of Mg will be = 2. Hence, the formula of the compound formed by Mg and X will be  Mg3X2. Thus, the correct answer is option (d) Mg3X2.

Trick:  Valency is the “combining capacity of the atom”. So, valency can be calculated with the help of the electronic configuration of the element.


Practice Questions

Question 1: Na+, Mg2+, Al3+ and Si4+ are isoelectronic ions. What is the order of their ionic size?

Answer: Na+ > Mg2+ > Al3+ > Si4+

Question 2: Which one of the following configurations represents a metallic character?

(a) 2,8,2                  (b) 2,8,7                     (c) 2,8,4                       (d) 2,8,8

Answer: (a) 2,8,2


Conclusion

Thus, the study of classification of elements with the help of periodic table predicts how chemicals behave. And with the help of classification of elements in the periodic table, we can easily determine the properties of all possible elements and their compounds. Also we can study the group, period, block and type of elements.


A variety of classification of elements are involved and their periodic properties help to understand the properties of elements and their compounds in detail. Hence, this is important not only for competitive exams like JEE or NEET 2022 but also in better understanding the study of every element and their compounds.

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FAQs on NEET Important Chapter: Periodic Table and Periodicity in Properties

FAQ

1. What difficulties arise if the classification of elements is not made?

There will be a lot of confusion in learning things without classification. Without classification, the properties of elements or features will overlap each other. Hence, classification is done to list down the properties of the elements and their compounds.

2. What is that tool used by scientists to determine  the properties of elements in an easier way?

In the periodic table, all the known elements are put into groups with similar properties. So, the periodic table makes an important tool for chemists, nanotechnologists and other scientists. If we get to understand the periodic table, and learn the use of it, we’ll be able to predict how these chemicals will behave.

3. Why is the atomic number known as the fingerprint of the elements?

If atomic number changes then the number of protons in the nucleus changes and so does the number of electrons. This procedure completely changes the element because the properties of the atom will change. Hence, the atomic number is unique for each element therefore it is considered as the "fingerprint" of elements”.