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

Introduction to Periodic Classification of Elements

Welcome, students! In the world of chemistry, elements are the fundamental building blocks of all matter. Imagine a library with millions of books scattered randomly without any system. Finding a specific book would be a chaotic and time-consuming task. Similarly, as scientists started discovering more and more elements, they faced a similar challenge. By the early 19th century, dozens of elements were known, each with its unique set of physical and chemical properties. To study them effectively and understand their relationships, a systematic classification was desperately needed. This need gave birth to the concept of the periodic table.

Chapter 5, 'Periodic Classification of Elements,' takes us on a historical and scientific journey. We will explore the brilliant minds and the meticulous efforts that led to the organization of elements. We'll start with the earliest, simplest attempts and gradually move to the sophisticated and highly organized Modern Periodic Table that we use today. This chapter is not just about memorizing a table; it's about understanding the underlying principles, recognizing patterns (periodicity), and predicting the properties of elements. A solid grasp of this chapter will form a strong foundation for your future studies in chemistry.

Early Attempts at the Classification of Elements

Before the final, accepted version of the periodic table was developed, several scientists proposed their own methods of classifying elements. These early attempts, while not entirely successful, were crucial stepping stones that provided valuable insights and laid the groundwork for future discoveries.

Döbereiner's Triads

One of the first notable attempts was made by a German chemist, Johann Wolfgang Döbereiner, in 1817. He observed that certain elements could be grouped into sets of three, which he called 'triads'. The unique characteristic of these triads was the relationship between their atomic masses.

What is a Döbereiner's Triad?

Döbereiner arranged three chemically similar elements in a group in the increasing order of their atomic masses. He discovered that the atomic mass of the middle element was approximately the arithmetic mean (average) of the atomic masses of the other two elements. Furthermore, the properties of the middle element were also intermediate between the other two.

Let's look at some examples to understand this better:

• The Alkali Metal Triad: Lithium (Li), Sodium (Na), and Potassium (K).
  • Atomic mass of Li = 6.9 u
  • Atomic mass of K = 39.0 u
  • Average atomic mass = (6.9 + 39.0) / 2 = 22.95 u
  • Actual atomic mass of Sodium (Na), the middle element, is 23.0 u. This is remarkably close! All three are reactive metals with a valency of 1.
• The Alkaline Earth Metal Triad: Calcium (Ca), Strontium (Sr), and Barium (Ba).
  • Atomic mass of Ca = 40.1 u
  • Atomic mass of Ba = 137.3 u
  • Average atomic mass = (40.1 + 137.3) / 2 = 88.7 u
  • Actual atomic mass of Strontium (Sr) is 87.6 u. Again, a very close match.
• The Halogen Triad: Chlorine (Cl), Bromine (Br), and Iodine (I).
  • Atomic mass of Cl = 35.5 u
  • Atomic mass of I = 126.9 u
  • Average atomic mass = (35.5 + 126.9) / 2 = 81.2 u
  • Actual atomic mass of Bromine (Br) is 79.9 u. This is also quite close.

Limitations of Döbereiner's Triads:

Despite its cleverness, Döbereiner's system had significant drawbacks. He could only identify a few such triads from the elements known at that time. The law was not applicable to all elements, and as new elements were discovered, it became clear that this system of classification was not comprehensive enough. For instance, Fluorine (F), Chlorine (Cl), and Bromine (Br) are chemically similar, but they do not form a triad as the average mass of F and Br is not close to the mass of Cl. Therefore, his classification was dismissed as a mere coincidence rather than a fundamental law.

Newlands' Law of Octaves

In 1866, an English scientist named John Newlands arranged the then-known elements in the order of increasing atomic mass. He was a music lover and found a fascinating pattern connecting chemistry and music.

What is the Law of Octaves?

Newlands stated that when elements are arranged in increasing order of their atomic masses, the properties of every eighth element are a repetition of the properties of the first element, similar to the eighth note in an octave of Western music (Do, Re, Mi, Fa, So, La, Ti, Do). This is why his classification is called the 'Law of Octaves'.

For example, he started with Hydrogen (H) and ended at Thorium (Th), which was the 56th element. He found that the properties of Lithium (Li), the second element in his table, were similar to Sodium (Na), the ninth element. Similarly, Beryllium (Be) showed properties similar to Magnesium (Mg).

A part of Newlands' Octaves table looked like this:

Limitations of Newlands' Law of Octaves:

Newlands' idea was initially ridiculed by the scientific community, but it did contain the seed of the concept of 'periodicity'. However, it suffered from several major flaws:

• Limited 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 Elements: Newlands assumed that only 56 elements existed in nature and no more elements would be discovered in the future. This assumption was proven wrong with the discovery of new elements whose properties did not fit into his Law of Octaves.
• Misfit Elements: To fit elements into his table, Newlands put two elements in the same slot. For example, Cobalt (Co) and Nickel (Ni) were placed in the same slot as Fluorine (F), Chlorine (Cl), and Bromine (Br). However, Co and Ni have very different properties compared to the halogens. He also placed Iron (Fe), which resembles Co and Ni in properties, far away from them.
• Inapplicability to Noble Gases: The law completely failed when the noble gases (like Helium, Neon, Argon) were discovered. Their inclusion disrupted the entire octave pattern.

Mendeleev's Periodic Table: A Major Breakthrough

The real breakthrough in the classification of elements came from a Russian chemist, Dmitri Ivanovich Mendeleev. His work transformed the scattered knowledge of elements into a structured, logical system. He is often regarded as the 'father of the periodic table'.

The Architect of the Periodic Table: Dmitri Mendeleev

Mendeleev meticulously studied the relationship between the atomic masses of elements and their physical and chemical properties. He concentrated on the chemical properties, particularly the compounds formed by elements with oxygen and hydrogen, as they are very reactive and formed compounds with most elements. He wrote down the properties of each of the 63 known elements on a separate card and then sorted these cards. He found that when elements were arranged in order of increasing atomic mass, a periodic recurrence of properties was observed.

Based on these observations, Mendeleev formulated his famous Periodic Law, which states:

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

Features of Mendeleev's Periodic Table

Mendeleev's Periodic Table, published in 1869, was a systematic arrangement of elements into vertical columns and horizontal rows.

• Groups: The vertical columns were called 'groups'. There were eight groups in his table, numbered I to VIII in Roman numerals. Elements within the same group exhibited similar chemical properties. For instance, elements in Group I (like Li, Na, K) were all reactive metals that formed oxides with the formula R₂O and hydrides with the formula RH.
• Periods: The horizontal rows were called 'periods'. There were seven periods. The properties of elements changed gradually as one moved from left to right across a period.
• Sub-groups: Each group from I to VII was further divided into two sub-groups, A and B. Elements within a sub-group showed more resemblance to each other than to the elements of the other sub-group in the same group.

Achievements of Mendeleev's Periodic Table

Mendeleev's table was not just a classification; it was a powerful predictive tool. Its success was due to several remarkable achievements:

1. Systematic Study of Elements

By arranging elements into groups and periods, Mendeleev's table made the study of chemistry much simpler and more organized. If you knew the properties of one element in a group, you could predict the properties of the others with reasonable accuracy.

2. Correction of Doubtful Atomic Masses

Mendeleev's bold belief in his periodic law allowed him to correct the atomic masses of some elements. For example, the atomic mass of Beryllium (Be) was previously determined to be 13.5. This would have placed it between Carbon and Nitrogen, which didn't fit its properties. Mendeleev calculated that its valency should be 2, not 3, which gave it an atomic mass of 9. This placed Beryllium in the correct position before Boron, fitting perfectly with its properties resembling Magnesium and Calcium.

3. Prediction of New Elements

This is considered the most significant achievement of Mendeleev's table. Instead of seeing the gaps in his table as defects, he boldly declared that these were undiscovered elements waiting to be found. He didn't just predict their existence; he also predicted their properties with astonishing accuracy based on their positions in the table. He named them using Sanskrit numerals 'Eka' (meaning one) as a prefix to the name of the preceding element in the same group.

• Eka-aluminium: He predicted an element that would be below Aluminium. Later, Gallium (Ga) was discovered, and its properties matched Mendeleev's predictions almost perfectly.
• Eka-silicon: He predicted an element below Silicon. This element, Germanium (Ge), was discovered later, and its properties were remarkably similar to those predicted by Mendeleev.
• Eka-boron: This predicted element, later discovered as Scandium (Sc), also matched his predictions.

Let's compare the predicted properties of Eka-silicon with the actual properties of Germanium:

The stunning accuracy of these predictions provided strong evidence for the correctness of Mendeleev's Periodic Table and made the scientific community accept it widely.

4. Position of Noble Gases

Noble gases like Helium (He), Neon (Ne), and Argon (Ar) were discovered much later. Since they are chemically inert, they could be placed in a new, separate group (called the 'zero group') without disturbing the existing order of Mendeleev's table. This showed the flexibility and robustness of his classification.

Limitations of Mendeleev's Classification

Despite its monumental success, Mendeleev's Periodic Table had a few anomalies that could not be explained:

1. Anomalous Position of Hydrogen

Hydrogen was a puzzle. It showed similarities to both Group I elements (alkali metals) and Group VII elements (halogens).

• Similarity with Alkali Metals (Group I): Like alkali metals, Hydrogen has one electron in its outermost shell and forms positive ions (H+). It combines with halogens, oxygen, and sulphur to form compounds with similar formulae (e.g., HCl, H₂O, H₂S vs. NaCl, Na₂O, Na₂S).
• Similarity with Halogens (Group VII): Like halogens, Hydrogen exists as a diatomic molecule (H₂ vs. F₂, Cl₂). It can also gain an electron to form a negative ion (hydride ion, H-), similar to halide ions (F-, Cl-).

2. Position of Isotopes

Isotopes are atoms of the same element that have the same chemical properties but different atomic masses (e.g., Chlorine has isotopes with masses 35 u and 37 u). Since Mendeleev's table was based on increasing atomic mass, isotopes of an element would have to be given different positions. This would disrupt the symmetry of the table and would be chemically illogical, as they are the same element. This was a major challenge to Mendeleev's law.

3. Anomalous Pairs of Elements

In a few instances, Mendeleev had to place an element with a slightly higher atomic mass before an element with a slightly lower atomic mass to maintain the similarity in properties. This violated his own law. For example:

• Cobalt (Co) and Nickel (Ni): The atomic mass of Co is 58.9, and Ni is 58.7. Yet, Co was placed before Ni to group it with rhodium and iridium, which it resembles.
• Argon (Ar) and Potassium (K): The atomic mass of Ar is 39.9, and K is 39.1. However, Ar (a noble gas) was placed before K (an alkali metal) to fit the property patterns.

The Modern Periodic Table: The Final Form

The limitations of Mendeleev's table hinted that a more fundamental property was at play. The discovery that would revolutionize the periodic classification came in the early 20th century.

Henry Moseley's Contribution

In 1913, a young English physicist named Henry Moseley conducted experiments where he bombarded different elements with high-speed electrons and studied the X-rays emitted. He discovered a precise mathematical relationship between the frequency of the emitted X-rays and a property of the atom he called the atomic number (Z). He concluded that the atomic number of an element was a more fundamental property than its atomic mass. The atomic number corresponds to the number of protons in the nucleus of an atom, and it is a unique, whole number for each element.

The Modern Periodic Law

Based on Moseley's discovery, the periodic law was modified. The Modern Periodic Law states:

"Properties of elements are a periodic function of their atomic number."

When elements are arranged in order of increasing atomic number, the periodicity in their properties becomes even more pronounced, and most of the anomalies of Mendeleev's table are resolved.

Structure of the Modern Periodic Table (Long Form)

The Modern Periodic Table is an arrangement of elements in order of increasing atomic number. Its structure is directly linked to the electronic configuration of the atoms.

• Periods (Horizontal Rows): There are 7 periods in the Modern Periodic Table. The period number of an element signifies the principal energy level or the number of shells being filled in its atoms. For example, elements of Period 3 (Na, Mg, Al, etc.) have electrons in three shells (K, L, M). The number of elements in each period is determined by the maximum number of electrons that can be accommodated in the shells:
  • 1st Period: 2 elements (filling the K shell)
  • 2nd Period: 8 elements (filling the L shell)
  • 3rd Period: 8 elements (filling the M shell)
  • 4th Period: 18 elements
  • 5th Period: 18 elements
  • 6th Period: 32 elements
  • 7th Period: Incomplete, but also has a capacity for 32 elements.
• Groups (Vertical Columns): There are 18 groups in the Modern Periodic Table. Elements in the same group have the same number of valence electrons (electrons in the outermost shell). For example, all elements in Group 1 (Li, Na, K, etc.) have 1 valence electron, and all elements in Group 17 (F, Cl, Br, etc.) have 7 valence electrons. Since chemical properties are primarily determined by valence electrons, elements in the same group exhibit similar chemical behavior.

How the Modern Periodic Table Solves Mendeleev's Anomalies

Basing the table on atomic number beautifully resolved the earlier issues:

• Position of Isotopes: All isotopes of an element have the same number of protons, and therefore, the same atomic number. So, they all occupy the same single position in the table. For example, both Chlorine-35 and Chlorine-37 have an atomic number of 17, so they are placed in the same slot. Problem solved.
• Anomalous Pairs: The problem of 'wrong order' of atomic masses disappears. Cobalt (Z=27) rightfully comes before Nickel (Z=28), and Argon (Z=18) comes before Potassium (Z=19), following the strict sequence of increasing atomic number. Problem solved.
• Position of Hydrogen: While its position is still debated due to its unique properties, placing it in Group 1 is justified based on its atomic number (Z=1) and its electronic configuration (1s¹), which is similar to other Group 1 elements.

Trends in the Modern Periodic Table

The beauty of the Modern Periodic Table lies in the clear, predictable patterns or 'trends' in the properties of elements as we move across a period or down a group. Understanding these trends is key to understanding chemical behavior.

Valency

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

• Across a Period (Left to Right): The number of valence electrons increases from 1 to 8. Accordingly, the valency with respect to hydrogen or metals first increases from 1 to 4 and then decreases from 4 to 0. For example, in Period 3: Na(1), Mg(2), Al(3), Si(4), P(3), S(2), Cl(1), Ar(0).
• Down a Group: The number of valence electrons remains the same for all elements in a group. Therefore, all elements in a group have the same valency. For example, all elements in Group 1 (Li, Na, K, etc.) have a valency of 1.

Atomic Size (Atomic Radius)

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

• Across a Period (Left to Right): The atomic size decreases. This might seem counterintuitive as electrons are being added. However, the electrons are added to the same shell, while the number of protons in the nucleus (nuclear charge) increases. This increased positive charge pulls the electrons closer to the nucleus, causing the atom to shrink. So, Sodium (Na) is larger than Chlorine (Cl).
• Down a Group: The atomic size increases. As we move down a group, a new electron shell is added at each step. The addition of new shells increases the distance between the outermost electrons and the nucleus, overriding the effect of increased nuclear charge. Therefore, the atomic size increases. For instance, Lithium (Li) is smaller than Potassium (K).

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 is the tendency to gain electrons and form negative ions (anions).

• Across a Period (Left to Right): Metallic character decreases while non-metallic character increases. This is because the effective nuclear charge increases across a period, making it harder for an atom to lose its valence electrons. Conversely, the strong pull makes it easier to gain electrons. Thus, elements on the left side of the table (like Na, Mg) are metals, while those on the right (like S, Cl) are non-metals.
• Down a Group: Metallic character increases. As we go down a group, the atomic size increases. The outermost electrons are farther from the nucleus and experience a weaker pull. This makes it easier for them to be lost. Therefore, the tendency to lose electrons (metallic character) increases down a group. For example, in Group 14, Carbon is a non-metal, Silicon is a metalloid, while Tin and Lead are metals.

A zig-zag line separates the metals from the non-metals in the periodic table. Elements bordering this line, such as Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), and Tellurium (Te), exhibit properties of both metals and non-metals and are called metalloids or semi-metals.

Nature of Oxides

The chemical nature of the oxides formed by elements also shows a clear trend.

• Generally, metals form basic oxides (e.g., Na₂O, MgO), while non-metals form acidic oxides (e.g., SO₂, CO₂). Some metallic oxides, like Al₂O₃ and ZnO, show both acidic and basic behavior and are called amphoteric oxides.
• Across a Period (Left to Right): The basic nature of oxides decreases, and the acidic nature increases. For example, in Period 3: Na₂O (strongly basic) → MgO (basic) → Al₂O₃ (amphoteric) → SiO₂ (weakly acidic) → P₂O₅ (acidic) → SO₃ (strongly acidic).
• Down a Group: For metals, the basic nature of their oxides increases. This is because the metallic character increases down the group, making the element more likely to donate electrons to oxygen, resulting in a more ionic and basic oxide.

Important Questions and Answers

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 of the oxides, we first need to find the group number of each element in Mendeleev's table, which tells us its valency.

• K (Potassium): It is in Group I. Elements in Group I have a valency of 1. The formula of their oxide is R₂O. So, the formula for potassium oxide is K₂O.
• C (Carbon): It is in Group IV. Elements in Group IV have a valency of 4. The formula of their oxide is RO₂. So, the formula for carbon oxide is CO₂.
• Al (Aluminium): It is in Group III. Elements in Group III have a valency of 3. The formula of their oxide is R₂O₃. So, the formula for aluminium oxide is Al₂O₃.
• Si (Silicon): It is in Group IV. Elements in Group IV have a valency of 4. The formula of their oxide is RO₂. So, the formula for silicon oxide is SiO₂.
• Ba (Barium): It is in Group II. Elements in Group II have a valency of 2. The formula of their oxide is RO. So, the formula for barium oxide is BaO.

Question 2: An element X has atomic number 17. Predict its position in the periodic table, its valency, and its chemical nature (metal or non-metal).

Answer:

• Electronic Configuration: The atomic number of element X is 17. Its electronic configuration is 2, 8, 7 (K=2, L=8, M=7).
• Position in the Periodic Table:
  • Period: Since it has 3 electron shells (K, L, M), it belongs to the 3rd Period.
  • Group: It has 7 valence electrons. For p-block elements, the group number is 10 + (number of valence electrons). So, Group = 10 + 7 = 17.
• Valency: To achieve a stable octet, the atom needs to gain 1 electron. Therefore, its valency is 1.
• Chemical Nature: Since it has a strong tendency to gain an electron to complete its octet, it is a typical non-metal. (Element X is Chlorine, Cl).

Question 3: Compare and contrast the arrangement of elements in Mendeleev’s Periodic Table and the Modern Periodic Table.

Answer:

Question 4: Why does atomic radius decrease as we move from left to right in a period?

Answer: As we move from left to right across a period in the periodic table, the atomic number of the elements increases by one at each step. This means that one proton is added to the nucleus and one electron is added to the valence shell (the outermost shell) of the atom.

While an electron is added, it enters the same main energy level or shell. The addition of a proton in the nucleus increases the nuclear charge (the positive pull from the nucleus). This increased nuclear charge attracts the electrons, including the valence electrons, more strongly, pulling them closer to the nucleus. As a result, the size of the atom, or its atomic radius, contracts and decreases across the period.

Chapter Summary

Here are the key takeaways from our exploration of the Periodic Classification of Elements:

• The classification of elements is necessary for their systematic study.
• Döbereiner's Triads grouped three elements where the atomic mass of the middle element was the average of the other two. This was a limited classification.
• Newlands' Law of Octaves arranged elements by increasing atomic mass and noted that every eighth element had similar properties to the first. This was applicable only up to Calcium.
• Mendeleev's Periodic Law states that the properties of elements are a periodic function of their atomic masses. His table was a major success as it allowed for systematic study and the prediction of new elements.
• Limitations of Mendeleev's table included the position of hydrogen, the position of isotopes, and certain anomalous pairs of elements.
• The Modern Periodic Law, based on Henry Moseley's work, states that the properties of elements are a periodic function of their atomic number (Z).
• The Modern Periodic Table has 7 horizontal rows called periods and 18 vertical columns called groups.
• The position of an element in the table reveals its chemical properties. Elements in the same group have the same number of valence electrons and similar properties.
• Trends in the Modern Periodic Table:
  • 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.