Periodic Classification of Elements Chapter 5 Notes PDF Download
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Introduction:
In the beginning, scientists had classified elements into two broad categories as metals and non-metals. Some elements exhibited properties because of which they could neither be classified as metals nor non-metals and hence were called metalloids. This classification was not sufficient for scientific study. Over the years, many chemists attempted to make a rational and systematic classification. It was based on the physical and chemical properties of each element. These results were then tabulated in the form of a table.
Periodic table – The table giving the arrangement of the known elements according to their properties so that similar elements fall within the same vertical column and dissimilar elements are separated.
Dobereiner's Triads:
In 1817, a German chemist named Johann Wolfgang Dobereiner arranged the elements with similar properties into groups.
He identified many such groups which had three elements in them. Hence, these groups were termed triads.
In a triad, the elements were arranged in increasing order of their atomic masses. When the arithmetic means of the atomic mass of the first and the third element was taken, it came out to be approximately the same as that of the second or the middle element. This is a distinctive feature that can be used to identify if elements form a triad.
This classification was also not sufficient as not many triads could be identified. Only three were identified as:
Newlands' Law of Octaves:
In 1866, an English scientist, John Newlands arranged many of the then-known elements in the increasing order of their atomic masses. So, he started with the element having the lowest atomic mass (hydrogen) and ended at Thorium which was the ${{56}^{th}}$ element. He noticed that the eighth element was similar in properties to the first element.
It was concluded that there exists some systematic relationship between the order of atomic masses and the repetition of properties of elements. This relationship, when represented in a tabular form, presented a periodic repetition of the properties of the elements. Hence, the term periodicity was introduced.
This had a resemblance to the eight musical notes in both Western as well as Indian music.
It was found that the Law of Octaves was applicable only up to Calcium because, after Calcium, every eighth element did not possess properties similar to that of the first. John Newlands had assumed that only 56 elements existed in nature and no more elements would be discovered in the future. But several new elements were discovered later on, whose properties did not fit into the Law of Octaves. So, this was not sufficient either.
Mendeleev's Periodic Table:
In 1869, a Russian chemist, Dmitri Ivanovich Mendeleev classified the then known 63 elements based on their physical and chemical properties in the increasing order of the atomic masses in the form of a table.
He had observed that properties of the elements recur cyclically when they were arranged in the order of their increasing atomic masses. This observation led to the conclusion that the physical and chemical properties of the elements are periodic functions of their atomic masses. This came to be known as the law of chemical periodicity.
A periodic Table is the tabulation of all the known elements in a tabular format based on this law. It contains eight vertical columns called 'groups' and seven horizontal rows called 'periods’. Each of the eight groups has two sub-groups A and B. The properties of elements of a sub-group resemble each other more markedly than the properties of those between the elements of the two sub-groups.
Achievements of Mendeleev's Periodic Table:
Some of the important contributions of Mendeleev's periodic table are as follows:
Systematic Study of Elements – The table provided the arrangements of elements showing similar properties into groups. This was very useful in studying and remembering the properties of a large number of elements in a systematic way.
Prediction of New Elements – Mendeleev had predicted new elements and had left three blanks for these undiscovered elements. He was able to predict their properties more or less accurately. He named them eka-boron, eka-aluminium and eka-silicon.
Correction of Atomic Masses - Based on the elements’ positions in the periodic table, Mendeleev was able to correct their atomic masses. The atomic mass of beryllium was corrected from 13.5 to 9.0.
Limitations of Mendeleev’s Classification:
Although Mendeleev's periodic table has many advantages, it could not explain certain things, which are considered as its limitations. They are as follows:
Assigning a position to Hydrogen: Hydrogen has an electronic configuration as that of alkali metals and combines with halogens, oxygen, sulphur to form compounds, like $HCL,{{H}_{2}}O,{{H}_{2}}S$ and at the same time it exists as a diatomic molecule like halogens. So, Mendeleev was not able to assign a proper position for hydrogen.
Assigning position to isotopes: The isotopes have similar properties but differ in their atomic masses. Mendeleev’s classification would place them in different groups due to their different atomic masses, but isotopes were not placed so as their properties were similar.
The anomalous pairing of some elements: Mendeleev did not follow the increasing atomic masses but grouped some elements based on similar properties. Argon with an atomic mass of 39.9 was placed before potassium with an atomic mass of 39.1. Also, some elements with similar properties like copper (Cu) and mercury (Hg) were placed separately, and some very dissimilar elements were placed in one group. Copper was placed in group I, the elements of which had no similarities with copper.
The Modern Periodic Table:
In the year 1913, an English physicist named Henry Mosely found that the atomic number of an element, which was denoted by the symbol ‘Z’ was a more basic property to group them instead of their atomic masses. Thus Mendeleev's periodic table was modified for the same. The elements were now grouped based on the increasing atomic number.
This came to be known as the Modern Periodic Law and it states, ‘properties of the elements are a periodic function of their atomic number’. Hence the new classification of the elements based on this came into existence and was termed as ‘Modern Periodic Table.
With this system of grouping, it was easy to predict the properties of the elements when they were arranged in the order of increasing atomic numbers. It is to be noted that the periodicity of the elements is based on the electronic configuration or the number of protons in the nucleus.
Salient Features of the Modern Periodic Table:
The table has 18 vertical columns that are known by the name of groups and 7 horizontal rows that are named as periods.
Periods:
There are 7 periods in this table. The periods have the same elements that have the same valence shell or the energy shell. Example - Na, Mg, Al, Si, P, S, Cl are placed in the same shell as they have the electronic shells as K, L and M.
In a period, the number of electrons present in the energy shells increases by 1 on moving from left to right within a period. Example - Na - 1, Mg - 2, Al - 3, and so on.
The number of elements present in a period can be determined by the formula $2{{n}^{2}}$, where n is the number of the shell from the nucleus. Example:
The first period consists of two elements only namely, hydrogen and helium as they have only 1 valence shell. Example - hydrogen (Z = 1 or shell as K = 1), helium (Z = 2 or shell as K = 2)
The second period has 8 elements with 2 shells and it starts with lithium (Z = 3 or shells as K = 2, L = 1) and ends with neon (Z = 10 or shells as K = 2, L = 8).
The third period has 8 elements with 3 shells and it starts with sodium (Z = 11 or shells as K = 2, L = 8, M = 1) and ends with argon (Z = 18 or shells as K = 2, L = 8, M = 8).
Similarly, the fourth period has 18 elements with 4 shells and starts with potassium (Z = 19) and ends with krypton (Z = 36).
The fifth period having 18 elements with 5 shells starts with rubidium (Z = 37) and ends with xenon (Z = 54).
The sixth period with 32 elements has 6 shells and it starts with caesium (Z = 55) ending with radon (Z = 86).
The seventh and last period is incomplete with 19 elements starts francium (Z = 87) and going on till oganesson (Z = 118).
Groups:
There are 18 groups in the periodic table. The group consists of elements that have the same number of electrons in the valence shell or outermost shell of the atom.
The valence shell predominantly decides the physical or chemical properties of the elements, so the elements of the same group have the same properties due to the same number of valence electrons. For example - fluorine and chlorine have valence electrons of 7 and they belong to group 17.
The number of shells increases by one unit as the elements move down in the periodic table in a group.
Blocks:
The periodic table is also divided into 4 blocks that are based on the subshell of the valence electrons. They are:
s-Block elements: All the elements of group 1 and 2 are included in this block and their general electronic configuration is $n{{s}^{1-2}}$ Example - Hydrogen (H), Sodium (Na), etc from group 1 and Magnesium (Mg), Calcium (Ca), etc from group 2.
p-Block elements: This includes the elements from group 13 to 18. They have an electronic configuration as $n{{s}^{2}}n{{p}^{1-6}}$.
d-block elements: This includes group 3 to 12 elements. They have a general electronic configuration as $\left( n-1 \right){{d}^{1-10}}n{{s}^{1-2}}$.
f-block elements: This block has sets of elements, lanthanides and the actinides. They have the electronic configuration of $\left( n-2 \right){{f}^{1-14}}\left( n-1 \right){{d}^{0-1}}n{{s}^{2}}$. The lanthanides starts from Lanthanum (La) - Lutetium (Lu) and the actinides starts from Actinium (Ac) - Lawrencium (Lr).
Position of Elements in The Periodic Table:
The position of the various elements is decided on the basis of their valence shells and the number of electrons present in their valence shells. Example - Sodium (Z - 11, 2,8,1), so it has 3 shells, so it is placed in period 3 and since it has 1 valence electron in the outermost shell, it is placed in group 1.
The position of an element in the periodic table determines its chemical nature. Based on the position of the elements in the periodic table, they can be classified as follows:
Noble gases: These are a group of elements placed in group 18, which are tasteless, odourless monoatomic gases that have very low chemical reactivity. There are 6 such gases, namely, Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn). They are also referred to as inert gases and due to their inertness, they are suitable to be used where reactions are not required. Example - He is used by deep-sea divers in the breathing gas to prevent toxicity of oxygen, nitrogen and carbon dioxide.
Normal elements: All the elements that are placed in groups 1 to 7 are included in this.
Alkali metals: The elements in group 1, namely Lithium (Li) - Francium (Fr), except Hydrogen (H) are termed alkali metals as they tend to form hydroxide with water which is strong alkalis. Thus alkali metals are very reactive and react quickly with water or air. Example - Sodium (Na) reacts violently with oxygen in the air, so it is stored in mineral oil.
Alkaline earth metals: These include group 2 elements starting from Beryllium (Be) - Radium (Ra). They are less reactive than the alkali metals that are found as a compound.
Transition elements: These include the elements from groups 3 to 11. These are so named as they exhibit a transition in their properties from the left to the right, including an increase in atomic size, ionization energy, electronegativity.
Inner transition elements: These are elements with similar properties, placed at the end of group 3 in periods 7 and 8. These are called the lanthanide series starting from Lanthanum (La) - Lutetium (Lu) in period 6 with 14 elements. Period 7 contains 14 elements starting from Actinium (Ac) - Lawrencium (Lr).
Halogens: These include a group of elements in group 17 that are generally non-metals that can exist in solid, liquid and gas form. They react with the metals to form salts. They are Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At) and Tennessine (Ts).
Properties of the Periodic Table:
Valency:
The term valency denotes the number of electrons that are gained or lost by an atom in order to complete its outermost shell to have a stable electronic configuration. This valency is the number of electrons present in the valence shell.
It can be noted that the valency increases from left to right in a period, and then decreases.
The valency in the group remains the same through the group, going downward.
Example - Period 2 elements have atomic numbers from 3 to 10, so they have 2 shells with an increasing number of valence till C and then it decreases. But all the group 2 elements have 2 electrons in their outermost shell, so their valency is 2.
Atomic Size:
The atomic size is determined by the atomic radius of the atom and it can be termed as the distance from the centre of the nucleus of the atom to its outermost shell.
It is seen that the elements are placed across a period from left to right, the atomic radius decreases. This is because an increased nuclear charge has the tendency to pull the electrons towards the nucleus, thereby decreasing the radius of the atom, and thus the atomic size decreases.
It is not the same in a group. As the elements move down a group, there is an addition of an extra shell, hence their atomic radius and thus atomic size increase. Example -
Ionization Energy:
It is termed as the amount of energy that is required to remove the electrons that are present in the outermost shell of the atom.
Across the period, since the atomic radius decreases, these electrons in the outermost shell are much closer to the nucleus and therefore it requires more energy to remove them. So, ionisation energy increases across a period.
In the groups, since a shell is being added down the group, so the atomic radius increases which lead to the electrons in the outermost shell being far from the nucleus and therefore it is easy to remove them. Hence in a group, the ionisation energy decreases going from top to bottom.
Electron Affinity:
It is termed as the amount of energy change as a result of an addition of an electron to the atom or the ability of an electron to accept electrons.
In a period, from left to right as the nuclear charge increases and atomic size decreases, it is easy for the addition of electrons, which leads to the generation of more energy. Hence electron affinity increases across a period.
In the group, as atomic size increases, nuclear charge decreases, so a lesser number of electrons can be added which leads to lesser energy generation. Hence electron affinity decreases from top to bottom in a group.
Metallic and non-metallic properties:
The metals are those elements that have a tendency to lose electrons and attain a positive charge. So, the metallic nature of elements means that they are electropositive and have low ionization energy. This metallic character decreases along a period.
The non-metallic nature of elements indicates that they have the ability to gain electrons and attain a negative charge. So, this non-metallic nature indicates that they are electronegative and have high ionization energies.
Hence this non-metallic nature increases along a period. Hence, moving along a period from left to right, the metallic character decreases and non-metallic character increases and there is a semi-metallic nature in between them.
The metallic nature increases down a group as the atomic size increases and they have an increased tendency of losing electrons. Example -
These trends can be summarised with the chart for easy reference:
Class 10 Science Chapter 5 Notes PDF Download
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Chapter 5 - Periodic Classification of Elements Revision Notes
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Class 10 Chemistry Chapter 5 Notes Summary
Periodic Classification of Elements Class 10 Notes comprise of the following crucial topics:
Dobereiner's Triads
Examples of Dobereiner's Triads
Defects
The first section in class 10 Science Ch 5 Notes introduces the topic of Dobereiner's Triads. You will also get familiar with the examples and defects of the same.
Newland's Law of Octaves
Groups of the 8th element
Law of Octaves achievements
Defects
Next, comes the Newland's Law of Octaves. This is another way of arranging chemical elements, and you must thoroughly go through class 10 Chapter 5 Science Notes to know the differences between each arrangement.
Lethar Meyer's Curves
This is again another way of showing the periodic arrangement of elements plotted in a graph by a German chemist described in Class 10 Science Chapter 5 Notes.
Mendeleev's Periodic Table
Achievements
Systematic study of elements
Prediction of new elements
Atomic mass correction
Limitations
Element grouping
Anomalous pair
Isotopes
Lanthanides and actinides
This section in notes of Chemistry Class 10 Chapter 5 contains a detailed explanation of all related aspects mentioned above of Mendeleev's periodic table. Plus, a table is also provided that represents the late discovered elements.
Modern Periodic Table
Periods
Groups
Position of elements
Merits and demerits
Next comes the modern periodic table which is also termed as the long form of the periodic table. This is an important topic, and questions usually get repeated. All necessary details associated with the modern periodic table are laid out concisely in the Class 10 Science Chapter 5 Notes, and you will also get to know how electrons are accommodated in a shell.
Furthermore, classification of elements in the periodic table is also listed in class 10 Science Notes Chapter 5. The elements are classified as:
Transition elements
Inner-transition elements
Alkali metals
Noble gas
Normal elements
Halogens
Periodic Properties
Valency
Atomic volume
Atomic radius
Ionisation energy
Electron affinity
Comprehensive elucidation of the mentioned topics is provided in the Class 10 Science Chapter 5 Notes, making it easier for you to prepare the periodic properties of elements.
Electronegativity
Types of electronegativity
Metallic and non-metallic character
In the last portion of class 10th Science Chapter 5 Notes, students will get well-versed with the concepts of electronegativity, its types and an element's metallic and non-metallic character.
We recommend you to download notes of Science Class 10 Chapter 5 prepared by Vedantu for better comprehension, and this material also acts as a last-minute guide before exams.
FAQs on Periodic Classification of Elements Class 10 Notes CBSE Science Chapter 5 (Free PDF Download)
1. State two differences between the modern periodic table and Mendeleev’s periodic table.
The two differences between the modern periodic table and Mendeleev’s periodic table are as follows.
- The elements are arranged on the basis of atomic number in the modern periodic table, whereas the elements in Mendeleev’s periodic table are arranged on the basis of atomic mass.
- The modern periodic table consists of 118 elements, whereas Mendeleev’s periodic table had 63 elements.
2. What are halogens in the periodic table?
The elements in Group 17 of the modern periodic table are called halogens. Chlorine, Fluorine, Bromine and Iodine are some of the elements found in Group 17. These non-metals are highly reactive. On reaction with metals, they commonly form ionic bonds, and with non-metals, they form covalent bonds. The halogens consist of elements in all three states of matter: solid, liquid and gas.