NCERT Class 10 Science Chapter 5: Periodic Classification of Elements - A Complete Guide

Introduction to Periodic Classification of Elements

Welcome, students! Imagine walking into a massive library where books are scattered randomly. Finding the one you need would be a nightmare, right? Now, picture the same library with books neatly arranged by genre, author, and title. This organization makes finding any book effortless. In the early 19th century, the world of chemistry was like that chaotic library. Scientists had discovered many elements, each with unique properties, but there was no system to organize them. Chapter 5 of your Class 10 Science syllabus, 'Periodic Classification of Elements', takes you on a fascinating journey through the history of how scientists brought order to this chaos. We will explore the brilliant minds and groundbreaking ideas that led to the creation of the modern periodic table – one of the most powerful tools in chemistry. Understanding this classification is fundamental as it helps us predict the properties of elements and understand their relationships with one another.

Making Order Out of Chaos – Early Attempts at the Classification of Elements

Before the well-organized table we use today, several chemists made significant attempts to classify the elements. These early models, while not perfect, were crucial stepping stones that paved the way for future discoveries.

Johann Wolfgang Döbereiner’s Triads

In 1817, a German chemist named Johann Wolfgang Döbereiner made one of the first notable attempts to classify elements. He observed that certain elements could be grouped into sets of three, which he called 'triads'. The elements within a triad shared similar chemical properties.

The Law of Triads states: When the three elements in a triad are written in order of increasing atomic masses, the atomic mass of the middle element is roughly the average of the atomic masses of the other two elements.

Let's look at some examples Döbereiner identified:

• The Alkali Metal Triad: Lithium (Li), Sodium (Na), and Potassium (K). These are all reactive metals that form +1 ions.
  • Atomic mass of Li = 6.9
  • Atomic mass of K = 39.0
  • Average mass = (6.9 + 39.0) / 2 = 22.95
  • The actual atomic mass of the middle element, Sodium (Na), is 23.0. This is remarkably close!
• The Alkaline Earth Metal Triad: Calcium (Ca), Strontium (Sr), and Barium (Ba).
  • Atomic mass of Ca = 40.1
  • Atomic mass of Ba = 137.3
  • Average mass = (40.1 + 137.3) / 2 = 88.7
  • The actual atomic mass of Strontium (Sr) is 87.6, which is again very close.

Limitations of Döbereiner’s Triads:

While a brilliant insight, Döbereiner's classification had a major drawback. He could only identify a few such triads from the elements known at the time. The law was not applicable to all known elements, so it was dismissed as a mere coincidence and not a universal principle.

John Newlands’ Law of Octaves

In 1866, an English chemist named John Newlands took the next step. He arranged the then-known elements in order of increasing atomic masses. He was a music lover and found a pattern reminiscent of musical octaves.

Newlands' Law of Octaves states: When elements are arranged in order of increasing atomic masses, the properties of every eighth element are a repetition of the properties of the first element, just like the eighth note in a musical octave is a repetition of the first.

He compared this to the Indian musical notes (Sa, Re, Ga, Ma, Pa, Dha, Ni) and the Western ones (Do, Re, Mi, Fa, So, La, Ti). For example, the properties of Sodium (the eighth element after Lithium) were found to be similar to those of Lithium.

Limitations of Newlands’ Law of Octaves:

• Applicability: The law was found to be applicable only up to Calcium (Ca). After Calcium, every eighth element did not possess properties similar to that of the first.
• Assumption about Future Elements: Newlands assumed that only 56 elements existed in nature and no more elements would be discovered in the future. This was proven wrong.
• Placement of Elements: To fit elements into his table, Newlands sometimes placed two elements in the same slot (e.g., Cobalt and Nickel). He also placed some unlike elements under the same note. For instance, Cobalt (Co) and Nickel (Ni) were placed in the same slot as Fluorine (F), Chlorine (Cl), and Bromine (Br), which have very different properties.
• Inert Gases: The discovery of noble gases (like Neon and Argon) later completely disrupted the octave pattern.

Mendeleev's Periodic Table: A Monumental Leap

The real breakthrough in classification came from the Russian chemist, Dmitri Ivanovich Mendeleev. In 1869, he proposed a periodic table that became the foundation for the one we use today. He not only organized the known elements but also used his table to predict the existence and properties of elements yet to be discovered.

The Foundation: Atomic Mass and Chemical Properties

Mendeleev's classification was based on a fundamental relationship he observed between the atomic mass of elements and their physical and chemical properties. He concentrated on the chemical properties, specifically the compounds elements formed with oxygen and hydrogen, as they are very reactive and form compounds with most elements. He wrote down the properties of each of the 63 known elements on a separate card and then arranged the cards, looking for a repeating pattern.

This led him to formulate the Mendeleev's Periodic Law, which states:

'The properties of elements are the periodic function of their atomic masses.'

This means that if you arrange elements by increasing atomic mass, their properties will repeat at regular intervals or 'periods'. His table consisted of vertical columns called 'groups' and horizontal rows called 'periods'.

Achievements of Mendeleev's Periodic Table

Mendeleev's table was widely accepted because of its remarkable accuracy and predictive power.

• Systematic Study of Elements: For the first time, elements were arranged in a way that grouped similar elements together (e.g., Li, Na, K in Group I). This made the study of chemistry much more organized and systematic.
• Prediction of New Elements: This was Mendeleev's most spectacular achievement. He left some gaps in his periodic table, boldly claiming that these gaps represented undiscovered elements. He even predicted their properties based on their position in the table. He named them using Sanskrit numerals: Eka (one).
  • Eka-aluminium: Predicted element below Aluminium. This was later discovered as Gallium (Ga), and its properties closely matched Mendeleev's predictions.
  • Eka-silicon: Predicted element below Silicon. This was later discovered as Germanium (Ge), which also had properties almost identical to his predictions.
  • Eka-boron: Predicted element below Boron. This was later discovered as Scandium (Sc).
• Correction of Doubtful Atomic Masses: Mendeleev's table helped in correcting the atomic masses of some elements. For example, the atomic mass of Beryllium (Be) was previously determined to be 13.5. Based on this, it would not fit correctly in the table. Mendeleev suggested its mass should be 9, placing it correctly before Boron. This was later verified experimentally.
• Placement of Noble Gases: When noble gases like Helium (He), Neon (Ne), and Argon (Ar) were discovered, they could be placed in a new group (Group 0) without disturbing the existing order of Mendeleev's table. This showed the robustness of his model.

Limitations of Mendeleev's Classification

Despite its immense success, Mendeleev's table had a few anomalies that could not be explained:

• Anomalous Pairs of Elements: In a few cases, Mendeleev had to place an element with a slightly higher atomic mass before an element with a slightly lower atomic mass to maintain the grouping of similar properties. For example, Cobalt (Co, atomic mass 58.9) was placed before Nickel (Ni, atomic mass 58.7).
• Position of Isotopes: Isotopes are atoms of the same element with the same chemical properties but different atomic masses (e.g., Chlorine-35 and Chlorine-37). According to Mendeleev's law, they should be placed in different positions due to their different masses. However, they were not given separate places, which violated his own law.
• Position of Hydrogen: The position of Hydrogen was not fixed. It resembles alkali metals (Group 1) as it forms a +1 ion, and it also resembles halogens (Group 17) as it exists as a diatomic molecule (H₂) and can form a -1 ion. Mendeleev could not assign a correct position to hydrogen.

The Modern Periodic Table

The limitations of Mendeleev's table pointed towards a need for a more fundamental property than atomic mass for classification. The solution came in 1913 from the work of an English physicist, Henry Moseley.

Henry Moseley's Contribution: Atomic Number

Henry Moseley, through his experiments with X-rays, demonstrated that the atomic number (Z) of an element is a more fundamental property than its atomic mass. The atomic number of an element is the number of protons in the nucleus of an atom. Since the number of protons is unique for each element and is always a whole number, it provided a much more consistent basis for classification.

The Modern Periodic Law

With atomic number as the basis, the Modern Periodic Law was proposed. It states:

'The properties of elements are a periodic function of their atomic number.'

When elements are arranged in order of increasing atomic number, the anomalies of Mendeleev's table disappear. For instance, the atomic number of Cobalt (27) is lower than that of Nickel (28), so their placement is justified. All isotopes of an element have the same atomic number, so they can be placed in the same slot.

Position of Elements in the Modern Periodic Table

The modern periodic table, also known as the long form of the periodic table, is based on the modern periodic law. Let's understand its structure:

• Groups: The vertical columns in the periodic table are called groups. There are 18 groups numbered from 1 to 18.
  • Elements in the same group have the same number of valence electrons (electrons in the outermost shell).
  • Because they have the same number of valence electrons, elements in the same group have similar chemical properties. For example, Group 1 elements (alkali metals) all have 1 valence electron and are highly reactive. Group 17 elements (halogens) all have 7 valence electrons and are also highly reactive non-metals. Group 18 elements (noble gases) have a full outermost shell, making them chemically inert.
• Periods: The horizontal rows in the periodic table are called periods. There are 7 periods.
  • The period number of an element signifies the number of electron shells in its atoms. For example, elements in Period 2 (like Li, Be, B, C) have 2 electron shells (K and L). Elements in Period 3 (like Na, Mg, Al) have 3 electron shells (K, L, and M).
  • The number of elements in a period is determined by the maximum number of electrons that can be accommodated in the outermost shell. The first period has 2 elements, the second and third have 8, the fourth and fifth have 18, and so on.

Trends in the Modern Periodic Table

The arrangement of elements in the modern periodic table reveals fascinating patterns or 'trends' in their properties. Understanding these trends helps us predict the behavior of elements without memorizing individual properties.

Valency

Valency is the combining capacity of an element, determined by the number of valence electrons.

• Across a Period (left to right): The valency first increases and then decreases. For example, in the 3rd period:
  • Na (1 valence electron) -> Valency = 1
  • Mg (2 valence electrons) -> Valency = 2
  • Al (3 valence electrons) -> Valency = 3
  • Si (4 valence electrons) -> Valency = 4
  • P (5 valence electrons) -> Valency = 3 (8-5)
  • S (6 valence electrons) -> Valency = 2 (8-6)
  • Cl (7 valence electrons) -> Valency = 1 (8-7)
  • Ar (8 valence electrons) -> Valency = 0
• Down a Group: The number of valence electrons remains the same. Therefore, the valency remains the same for all elements in a group. For example, all elements in Group 1 have a valency of 1.

Atomic Size (Atomic Radius)

Atomic size refers to the radius of an atom, which is the distance from the center of the nucleus to the outermost electron shell.

• Across a Period (left to right): The atomic size decreases. Why? As we move across a period, the atomic number increases by one, meaning one proton is added to the nucleus and one electron is added to the same outermost shell. The increased positive charge in the nucleus (increased nuclear charge) pulls the electrons closer to the nucleus, causing the atom to shrink.
• Down a Group: The atomic size increases. Why? As we move down a group, a new electron shell is added for each successive element. Even though the nuclear charge also increases, the effect of adding a new shell is much more significant. The new shell is farther from the nucleus, and the inner electrons shield the outer electrons from the full pull of the nucleus. This causes the atomic radius to increase.

Metallic and Non-metallic Properties

Metallic character refers to the tendency of an atom to lose electrons and form positive ions (cations). Non-metallic character refers to the tendency to gain electrons and form negative ions (anions).

• Metallic Character:
  • Across a Period (left to right): Metallic character decreases. This is because the increasing nuclear charge holds the valence electrons more tightly, making it harder for the atom to lose them.
  • Down a Group: Metallic character increases. As the atomic size increases, the valence electrons are farther from the nucleus and experience less of its pull. They can be lost more easily.
• Non-metallic Character:
  • Across a Period (left to right): Non-metallic character increases. The tendency to gain electrons increases because of the increasing nuclear charge and decreasing atomic size.
  • Down a Group: Non-metallic character decreases. As the atomic size increases, the nucleus's ability to attract an incoming electron weakens.

In the periodic table, metals are on the left, non-metals are on the right, and the elements in between, called metalloids or semi-metals (like Boron, Silicon, Germanium), exhibit properties of both metals and non-metals. They form a zigzag line separating metals from non-metals.

Important Questions and Answers (from NCERT Exercise)

Question 1: Use Mendeleev’s Periodic Table to predict the formulae for the oxides of the following elements: K, C, Al, Si, Ba.

Answer: To predict the formulae, we first need to identify the group in which these elements are placed in Mendeleev's table. The formula of the oxide is determined by the valency of the element, which is typically the same as its group number.

• K (Potassium): It is in Group I. The valency is 1. Oxygen has a valency of 2. So, the formula of the oxide will be K₂O.
• C (Carbon): It is in Group IV. The valency is 4. The formula of the oxide will be C₂O₄, which simplifies to CO₂.
• Al (Aluminium): It is in Group III. The valency is 3. The formula of the oxide will be Al₂O₃.
• Si (Silicon): It is in Group IV. The valency is 4. The formula of the oxide will be Si₂O₄, which simplifies to SiO₂.
• Ba (Barium): It is in Group II. The valency is 2. The formula of the oxide will be Ba₂O₂, which simplifies to BaO.

Question 2: An element X has mass number 35 and the number of neutrons is 18. What is the atomic number of X? Write the name, group number, and period of X.

Answer:

1. Find the Atomic Number (Z):
   • Mass Number (A) = Number of Protons + Number of Neutrons
   • Atomic Number (Z) = Number of Protons
   • Given: Mass Number = 35, Number of Neutrons = 18
   • Number of Protons = Mass Number - Number of Neutrons = 35 - 18 = 17
   • So, the Atomic Number (Z) is 17.
2. Identify the Element:
   • The element with atomic number 17 is Chlorine (Cl).
3. Determine Group and Period:
   • The electronic configuration of Chlorine (Z=17) is 2, 8, 7.
   • Period: The number of shells is 3 (K, L, M). Therefore, it belongs to Period 3.
   • Group: The number of valence electrons is 7. For p-block elements, Group Number = 10 + (number of valence electrons). So, Group Number = 10 + 7 = 17.

Question 3: How does the electronic configuration of an atom relate to its position in the Modern Periodic Table?

Answer: The electronic configuration of an atom is directly linked to its position in the Modern Periodic Table in two key ways:

• Period Number: The number of the outermost shell (or principal energy level) of an atom's electronic configuration determines its period. For example, Sodium (Na) has the configuration 2, 8, 1. Its outermost electron is in the 3rd shell (M-shell), so Sodium is in the 3rd Period.
• Group Number: The number of valence electrons (electrons in the outermost shell) determines its group. For elements in groups 1 and 2, the group number is equal to the number of valence electrons. For elements in groups 13 to 18, the group number is equal to 10 + the number of valence electrons. For example, Chlorine (Cl) has the configuration 2, 8, 7. It has 7 valence electrons, so it belongs to Group 10 + 7 = 17.

Question 4: Compare and contrast the arrangement of elements in Mendeleev’s Periodic Table and the Modern Periodic Table.

Answer:

Chapter Summary

Here are the key takeaways from our deep dive into the Periodic Classification of Elements:

• Need for Classification: Organizing elements helps in their systematic study and in predicting their properties.
• Döbereiner’s Triads: An early attempt grouping three elements with similar properties, where the atomic mass of the middle element was the average of the other two. It was limited in application.
• Newlands’ Law of Octaves: Arranged elements by atomic mass, where every eighth element had properties similar to the first. This worked only up to Calcium.
• Mendeleev's Periodic Law: 'The properties of elements are the periodic function of their atomic masses.' His table successfully organized elements and predicted new ones but had anomalies regarding isotopes and some element pairs.
• Modern Periodic Law: 'The properties of elements are a periodic function of their atomic number.' This law, based on Henry Moseley's work, forms the basis of the modern table.
• Modern Periodic Table Structure: It has 18 vertical Groups and 7 horizontal Periods. Group number is related to valence electrons, and Period number is related to the number of electron shells.
• Periodic Trends:
  • Valency: Remains the same down a group; first increases then decreases across a period.
  • Atomic Size: Increases down a group; decreases across a period.
  • Metallic Character: Increases down a group; decreases across a period.
  • Non-metallic Character: Decreases down a group; increases across a period.