Periodic Classification of Elements: Class 10 Science NCERT Guide

The Journey to Organisation: Early Classification Attempts

In the early 19th century, as more and more elements were discovered, scientists felt an urgent need to organize them. The first attempts were based on simple properties, but they laid the groundwork for the future.

Döbereiner's Triads (1817): German chemist Johann Döbereiner noticed that some elements could be grouped into threes, which he called 'triads'. When arranged in order of increasing atomic mass, the atomic mass of the middle element in a triad was roughly the average of the other two. For example, in the triad of Lithium (Li, mass ~7), Sodium (Na, mass ~23), and Potassium (K, mass ~39), the average mass of Li and K is (7+39)/2 = 23, which is the atomic mass of Sodium! However, Döbereiner could only identify a few such triads, so this system was not widely accepted.

Newlands' Law of Octaves (1866): British chemist John Newlands arranged the known elements in order of increasing atomic mass. He found that every eighth element had properties similar to the first, much like the eighth note in a musical octave. For example, the properties of sodium were similar to lithium (the eighth element after it). While clever, this law worked only up to calcium. With the discovery of noble gases, the 'eighth element' pattern broke down completely, and the idea was initially ridiculed.

Mendeleev's Periodic Table: A Revolutionary Breakthrough

1. The Guiding Principle: Atomic Mass and Chemical Properties — Russian chemist Dmitri Mendeleev created the first truly functional periodic table. His periodic law stated: 'The properties of elements are the periodic function of their atomic masses.' He focused on arranging elements with similar chemical properties (especially the formulae of their hydrides and oxides) under each other in vertical columns (groups), while arranging them by increasing atomic mass in horizontal rows (periods).
2. Major Achievements of Mendeleev's Table — Mendeleev's table was a huge success for three main reasons. 1. Systematic Study: It organized all 63 known elements systematically. 2. Prediction of New Elements: He boldly left gaps in his table for elements that he predicted would be discovered later. He even predicted their properties! For instance, he predicted 'Eka-aluminium' and 'Eka-silicon', which were later discovered as Gallium and Germanium, and their properties matched his predictions almost exactly. 3. Correction of Atomic Masses: He corrected the atomic masses of several elements, like Beryllium, based on their expected position in the table.
3. Limitations and Anomalies — Despite its success, Mendeleev's table had some flaws. 1. Position of Hydrogen: Hydrogen resembled both alkali metals (Group 1) and halogens (Group 17), so it couldn't be given a fixed position. 2. Anomalous Pairs: To group elements with similar properties, he had to place some elements with slightly higher atomic mass before those with lower mass (e.g., Cobalt before Nickel). 3. Position of Isotopes: Isotopes are atoms of the same element with different atomic masses. Since his table was based on atomic mass, isotopes would need separate positions, which would disrupt the table's order. These limitations hinted that atomic mass might not be the most fundamental property.

The Modern Periodic Table: Based on Atomic Number

The solution to Mendeleev's problems came in 1913 with the work of Henry Moseley. He showed that an element's atomic number (Z), which is the number of protons in its nucleus, is a more fundamental property than its atomic mass. This led to the Modern Periodic Law: 'The properties of elements are a periodic function of their atomic number.'

When elements are arranged this way, the anomalies of Mendeleev's table disappear. For instance, Cobalt (Z=27) naturally comes before Nickel (Z=28). Since all isotopes of an element have the same atomic number, they get a single, unified position. The modern table has 18 vertical columns called groups and 7 horizontal rows called periods. An element's position is determined by its electronic configuration, making the table a powerful tool for predicting chemical behaviour.

Understanding Trends in the Modern Periodic Table

• Valency: Valency is the combining capacity of an element. For main group elements, it's determined by the number of valence electrons (electrons in the outermost shell). Down a group, the number of valence electrons is the same, so valency remains constant (e.g., Li, Na, K all have valency 1). Across a period, valency first increases from 1 to 4 and then decreases to 0 (for noble gases). For Period 3: Na(1), Mg(2), Al(3), Si(4), P(3), S(2), Cl(1), Ar(0).
• Atomic Size (Radius): Atomic size refers to the radius of an atom. It decreases as you move from left to right across a period. This is because the atomic number increases, leading to a stronger pull (increased nuclear charge) on the electrons, shrinking the atom. Atomic size increases as you move down a group. This is because a new electron shell is added for each element down the group, making the atom larger.
• Metallic and Non-metallic Character: Metallic character refers to the tendency of an atom to lose electrons. This tendency increases down a group (as electrons are further from the nucleus and easier to lose) and decreases across a period (as the increasing nuclear charge holds electrons more tightly). Consequently, non-metallic character (tendency to gain electrons) increases across a period and decreases down a group.

Exam Tips: Avoid These Common Errors

A very common question in exams asks you to explain why a trend occurs. Don't just state the trend!

• For trends across a period (left to right): The key reason is the increase in effective nuclear charge. The number of shells remains the same, but more protons are added to the nucleus, pulling the electrons in more tightly. This explains why atomic size decreases and it becomes harder to lose electrons (metallic character decreases).
• For trends down a group (top to bottom): The key reason is the addition of new electron shells. Even though the nuclear charge increases, the outermost electrons are in a new shell, further away from the nucleus. This new shell effect is dominant, causing the atomic size to increase and making it easier to lose the valence electrons (metallic character increases).
• Mendeleev vs. Modern: Always be clear: Mendeleev used atomic mass. The Modern Periodic Table uses atomic number. This is a fundamental difference and a frequent source of questions.

Practice Questions with Solutions

• Q: An element 'X' has atomic number 17. (a) What is its electronic configuration? (b) To which group and period does it belong? (c) Is it a metal or a non-metal? Justify your answer. A: Step 1: Write the electronic configuration. The atomic number is 17, so there are 17 electrons. The distribution will be 2 electrons in the K shell, 8 in the L shell, and 7 in the M shell. So, the electronic configuration is 2, 8, 7. Step 2: Determine the group and period. The number of shells (K, L, M) is 3, so it belongs to Period 3. The number of valence electrons (electrons in the outermost M shell) is 7. For valence electrons > 2, Group number = 10 + valence electrons = 10 + 7 = 17. So, it belongs to Group 17. Step 3: Classify as metal or non-metal. An atom with 7 valence electrons will tend to gain one electron to complete its octet, rather than lose 7 electrons. The tendency to gain electrons is a characteristic of non-metals. Final answer: (a) Electronic configuration is 2, 8, 7. (b) It belongs to Period 3 and Group 17. (c) It is a non-metal because it needs to gain only one electron to achieve a stable configuration, which is easier than losing seven electrons.
• Q: How did Mendeleev's prediction of 'Eka-aluminium' and its subsequent discovery as Gallium support his periodic table? A: Step 1: Explain what Mendeleev did. Mendeleev observed a gap in his periodic table below Aluminium. He was so confident in his periodic law that he proposed the existence of an undiscovered element that should fit in this gap. He named it 'Eka-aluminium' (meaning 'one place below aluminium'). Step 2: Describe his predictions. Based on its position, Mendeleev predicted the properties of Eka-aluminium, including its atomic mass (around 68), its low melting point, and the formula of its oxide (E2O3) and chloride (ECl3). Step 3: Explain the discovery. Years later, the element Gallium (Ga) was discovered. Its properties were found to be remarkably similar to Mendeleev's predictions. Gallium's atomic mass is 69.7, its oxide is Ga2O3, and its chloride is GaCl3. Final answer: The accurate prediction of Gallium's properties before its discovery provided powerful evidence for the correctness and usefulness of Mendeleev's periodic table, leading to its wide acceptance by the scientific community.
• Q: Why does atomic radius decrease on moving from left to right in a period? A: Step 1: Define the movement. Moving from left to right in a period means the atomic number is increasing by one for each subsequent element (e.g., from Na (11) to Mg (12) to Al (13) in Period 3). Step 2: Identify what changes and what stays the same. As the atomic number increases, the number of protons in the nucleus increases. However, the electrons are being added to the same outermost shell for all elements within that period. Step 3: Explain the effect. The increased number of protons leads to a greater force of attraction from the nucleus (increased effective nuclear charge) on the electrons. This stronger pull draws the electron shells closer to the nucleus. Final answer: Atomic radius decreases across a period because the effective nuclear charge increases. While electrons are added, they go into the same principal shell and do not provide enough shielding to counteract the stronger pull from the increasing number of protons in the nucleus, thus shrinking the atom.
• Q: Compare and contrast the arrangement of elements in Mendeleev's and the Modern Periodic Table with respect to: (a) The basis of arrangement (b) The problem of isotopes A: Step 1: Compare the basis of arrangement. Mendeleev's Periodic Table arranged elements in order of increasing atomic mass. The Modern Periodic Table arranges elements in order of increasing atomic number. Step 2: Contrast the handling of isotopes. In Mendeleev's table, isotopes posed a problem. Since isotopes of an element have different atomic masses but the same chemical properties, they would need different positions according to his law, which would disrupt the table's structure. In the Modern Periodic Table, this problem is solved. All isotopes of an element have the same atomic number (same number of protons), so they all occupy the same single position in the table. Final answer: (a) Mendeleev's basis was atomic mass; the Modern table's basis is atomic number. (b) Isotopes would require separate spots and disrupt Mendeleev's table, while in the Modern table, all isotopes of an element correctly occupy a single spot as they share the same atomic number.

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