Unit 1: Periodic Table, Periodic Properties & Variations
A Comprehensive Lecture by Prof. Anil Tyagi, Ph.D.
My dear students, future custodians of health and scientific inquiry, welcome. Before we can comprehend the complex biochemistry of the human body, the pharmacology of drugs, or the diagnosis of diseases, we must first become fluent in the language of matter itself. That language is written in the elements, and its grammar is defined by the Periodic Table. This is not merely a chart to be memorized; it is a profound reflection of the inherent order in the universe, a map that reveals the deepest relationships between the building blocks of everything we see, touch, and are. Consider this unit the most important chapter in your chemical education. Let’s begin.
1. The Early Classification: A Quest for Order
The 19th century was a time of rapid chemical discovery. New elements were being isolated regularly, and scientists, like detectives, were desperately looking for patterns to make sense of the chaos. They were trying to group elements based on their properties, much like a librarian organizes books.
a) Dobereiner’s Triads (1829): The First Hint of a Pattern
Johann Wolfgang Dobereiner, a German chemist, was the first to notice a tangible numerical relationship between elements and their properties.
- The Concept of Triads: He grouped elements into sets of three, called triads, which had strikingly similar chemical properties.
- The Law of Triads: Dobereiner’s crucial observation was that the atomic mass of the middle element in each triad was almost exactly the arithmetic mean (average) of the atomic masses of the other two elements.
- Classic Examples:
- Chlorine (35.5), Bromine (80), Iodine (127):
- Check: (35.5 + 127) / 2 = 162.5 / 2 = 81.25 (very close to Bromine’s atomic mass of 80).
- Calcium (40), Strontium (88), Barium (137):
- Check: (40 + 137) / 2 = 177 / 2 = 88.5 (very close to Strontium’s 88).
- Lithium (7), Sodium (23), Potassium (39):
- Check: (7 + 39) / 2 = 46 / 2 = 23 (a perfect match for Sodium).
- Chlorine (35.5), Bromine (80), Iodine (127):
- Why It Was Significant: This was the first time a numerical value (atomic mass) was linked to chemical behaviour. It was a revolutionary idea.
- The Limitation: The fatal flaw was that this pattern could only be applied to a limited number of elements. Only a handful of such triads could be identified. As more elements were discovered, they didn’t fit neatly into this model. It was a good start, but the puzzle was far more complex.
b) Newlands’ Law of Octaves (1864): The Musical Analogy
Inspired by Dobereiner’s work, the English chemist John Newlands took a more ambitious approach. He arranged all the known elements (about 56 at the time) in the strict order of increasing atomic masses.
- The ‘Octave’ Law: Newlands observed that every eighth element, starting from a given one, had properties resembling those of the first. This reminded him of the musical notes in an octave, where the eighth note is a repetition of the first at a higher scale. Hence, he called this the Law of Octaves.
- An Example: If we start with Lithium (Li) as the first element:
- Li (1), Be (2), B (3), C (4), N (5), O (6), F (7)
- The eighth element is Sodium (Na), which has properties similar to Lithium.
- The next eighth is Potassium (K), again similar.
- Why It Was Significant: It was the first attempt to assign periodic numbers to elements and suggest a repeating pattern.
- The Limitations: Newlands’ law was met with ridicule.
- It Worked Like a Bad Melody: The analogy held true only for the lighter elements. When he reached heavier elements like transition metals, the pattern completely broke down. Calcium (40) and Iron (56), for instance, had no similarity, yet they were placed in the same column under his scheme.
- The Assumption of Finality: Newlands assumed no new elements would be discovered. He forced the existing elements into his table, much like trying to force a square peg into a round hole. When noble gases like Helium and Argon were later discovered, they would have shattered his model completely.
- Lack of Gaps: He did not leave any gaps for undiscovered elements, making his table rigid and inflexible.
c) Mendeleev’s Periodic Table (1869): The Masterstroke
Then came Dmitri Ivanovich Mendeleev, a Russian chemist of unparalleled vision. His approach was both brilliant and bold.
- The Periodic Law: Mendeleev stated that “the properties of elements are a periodic function of their atomic masses.” This was a more refined version of the ideas before him.
- The Methodology: He formulated his table based on two fundamental principles:
- Increasing Atomic Masses: He arranged elements horizontally in rows (called periods) in the order of their increasing atomic masses.
- Similar Properties: He placed elements with similar chemical properties vertically in columns (called groups). This was his masterstroke.
- The Merits & Genius of Mendeleev:
- Prediction of New Elements: This is his greatest contribution. Where existing elements did not fit the pattern, Mendeleev had the courage to leave gaps. He didn’t see these gaps as failures but as predictions. He even named these unknown elements using the Sanskrit prefix ‘Eka’ (meaning one) and predicted their properties in astonishing detail.
- Eka-boron (predicted) → Scandium (discovered in 1879)
- Eka-aluminium (predicted) → Gallium (discovered in 1875)
- Eka-silicon (predicted) → Germanium (discovered in 1886)
The accuracy of his predictions, like the density and atomic mass of Germanium, stunned the scientific world and cemented the validity of his table.
- Adjustment for Anomalies: When the strict order of atomic mass contradicted the order of chemical properties, Mendeleev prioritized properties. For example:
- Cobalt (58.9) was placed before Nickel (58.7) because Co resembled other group members more than Ni did.
- Similarly, Tellurium (127.6) was placed before Iodine (126.9) for the same reason. This showed his deep understanding that atomic mass alone wasn’t the final word.
- Prediction of New Elements: This is his greatest contribution. Where existing elements did not fit the pattern, Mendeleev had the courage to leave gaps. He didn’t see these gaps as failures but as predictions. He even named these unknown elements using the Sanskrit prefix ‘Eka’ (meaning one) and predicted their properties in astonishing detail.
- The Limitations: Even a masterpiece can have flaws, which later paved the way for improvement.
- Position of Hydrogen: Hydrogen, with its unique ability to gain one electron (like halogens) or lose one electron (like alkali metals), could not be assigned a fixed position. It was placed in Group 1, but this was never a perfect fit.
- Isotopes: The discovery of isotopes posed a serious problem. Isotopes of the same element have different atomic masses but identical chemical properties. According to Mendeleev’s law, they should be placed in different positions, which is illogical. For example, Chlorine-35 and Chlorine-37 would need separate slots.
- Anomalous Pairs: A few pairs of elements still had to be placed against the order of atomic masses to maintain periodicity. For instance, Argon (39.9), a noble gas, had to be placed before Potassium (39.1), a highly reactive metal, because Argon’s properties matched other noble gases. Mendeleev’s table could not explain why this was necessary.
2. The Modern Periodic Table: The Final Key
The year 1913 was a turning point. The English physicist Henry Moseley, while working with X-rays, discovered a fundamental property of the atom: its atomic number (Z).
- Moseley’s Discovery: He found that the square root of the frequency of an element’s characteristic X-rays was directly proportional to its atomic number (the number of protons in its nucleus). This proved that atomic number, not atomic mass, was the true fundamental property of an element.
- The Modern Periodic Law: This discovery resolved all the limitations of Mendeleev’s table. The law was revised to: “The physical and chemical properties of elements are a periodic function of their atomic numbers.”
- Isotopes Solved: All isotopes of an element have the same atomic number (same protons), so they occupy the same position in the table.
- Anomalous Pairs Solved: When elements are arranged by atomic number, all anomalies vanish. Argon (Z=18) correctly comes before Potassium (Z=19).
- Structure of the Modern Table:
- Periods: There are 7 periods (horizontal rows).
- The period number (n) denotes the number of electron shells (energy levels) in an atom.
- Example: An element in the 4th period has electrons filling the K, L, M, and N shells.
- Groups: There are 18 groups (vertical columns), numbered 1 to 18 as per IUPAC convention.
- Elements in the same group have the same number of valence electrons (electrons in the outermost shell), which is why they exhibit similar chemical properties.
- Example: All elements in Group 1 (Alkali Metals) have 1 valence electron. All in Group 17 (Halogens) have 7 valence electrons.
- Blocks: The table is also divided into blocks (s, p, d, f) based on the subshell in which the differentiating electron (last electron) enters.
- Periods: There are 7 periods (horizontal rows).
3. Periodic Properties: The Heart of the Matter
Why do elements behave the way they do? The answer lies in periodic trends. These trends are logical and can be understood by considering three key factors:
- Nuclear Charge (Z): The positive charge of the nucleus. It increases across a period.
- Atomic Size (Radius): The distance from the nucleus to the outermost shell. It decreases across a period but increases down a group.
- Screening Effect/Shielding Effect: The repulsion between inner shell electrons and the valence electrons, which reduces the effective pull of the nucleus. It remains constant across a period but increases down a group.
Let’s analyze the trends in detail:
| Property & Definition | Trend Across a Period (Left to Right) | Trend Down a Group (Top to Bottom) | Underlying Reason |
|---|---|---|---|
| Atomic Radius Distance from nucleus to outermost electron | Decreases | Increases | Across: Increased nuclear charge pulls the electron cloud closer with greater force, shrinking the atom. Down: A new electron shell is added with each period, increasing the size despite the increased nuclear charge. |
| Metallic Character Tendency to lose electrons and form positive ions | Decreases | Increases | Across: Increased nuclear charge holds onto valence electrons more tightly, making them harder to lose. Down: Increased atomic size and shielding effect mean valence electrons are farther from the nucleus and are held less tightly, making them easier to lose. |
| Non-Metallic Character Tendency to gain electrons and form negative ions | Increases | Decreases | This is the exact opposite of metallic character. It increases as the effective nuclear charge increases and the atomic size decreases, making it easier to attract electrons. |
| Ionisation Energy (IE) Minimum energy required to remove the most loosely bound electron from an isolated gaseous atom | Increases | Decreases | Across: Increasing nuclear charge and decreasing atomic size make it harder to remove an electron (more energy needed). Down: Increased atomic size and shielding effect make it easier to remove an electron (less energy needed). |
| Electron Affinity (EA) Energy change when an electron is added to a neutral gaseous atom | Increases | Decreases | Across: Smaller atomic size and high nuclear charge facilitate easier addition of an electron, releasing more energy. Down: Larger atomic size means the incoming electron is less attracted by the nucleus, so less energy is released. |
| Electronegativity (EN) The relative tendency of an atom to attract the shared pair of electrons towards itself in a covalent bond | Increases | Decreases | Across: High nuclear charge and small size allow an atom to attract shared electrons strongly. Down: Larger size and increased shielding reduce the atom’s ability to attract shared electrons. |
Note on Noble Gases: They often are exceptions to trends like Electron Affinity and Electronegativity as they have a stable octet and little tendency to gain or attract electrons.
4. Core Concepts: Valency, Atomic Number, Mass Number
These are the fundamental identifiers and descriptors of an element.
- Atomic Number (Z): This is the fingerprint of an element. It is defined as the number of protons in the nucleus of an atom. It is represented by the symbol Z.
Atomic Number (Z) = Number of Protons- In a neutral atom,
Number of Protons = Number of Electrons. - It is the atomic number that defines the identity of an element. Change the proton number, and you change the element itself.
- Mass Number (A): This is the total mass of the atom, concentrated in the nucleus. It is the sum of the number of protons and neutrons. It is represented by the symbol A.
Mass Number (A) = Number of Protons (Z) + Number of Neutrons (N)- Electrons have negligible mass and are not counted.
- Atoms of the same element (same Z) can have different numbers of neutrons, leading to different mass numbers. These are called isotopes (e.g., C-12, C-13, C-14).
- Valency: This is a crucial concept for understanding how elements bond. It is the combining capacity of an atom.
- It is determined by the number of electrons an atom can lose, gain, or share to achieve a stable, filled outer shell (octet or duplet).
- For Metals: Valency = Number of Valence Electrons (e.g., Magnesium (2,8,2) has valency 2).
- For Non-Metals: Valency = 8 – Number of Valence Electrons (e.g., Oxygen (2,6) has valency 8-6=2).
- Trend in a Period: Valency first increases from 1 to 4 and then decreases to 0 (for noble gases). For example, in Period 3: Na(1), Mg(2), Al(3), Si(4), P(3), S(2), Cl(1), Ar(0).
Professor Tyagi’s Final Words for NEET Aspirants:
My young scholars, internalize this unit. Do not rote-learn it.
- Think like Mendeleev: His story teaches us about the power of prediction and the scientific method.
- Understand Moseley’s Contribution: The shift from mass to atomic number is a paradigm shift in scientific thought.
- Master the Trends: For NEET, you will be questioned on the why behind the trends. Always link them to nuclear charge, atomic size, and shielding effect. Draw the curves in your mind.
- Valency is Key: You cannot write a single chemical formula or balance a reaction without understanding valency.
This chapter is the alphabet of chemistry. Master it, and you will be able to read the entire language of science with ease. Your journey to AIIMS and other top medical colleges begins with a strong command of such fundamental concepts.
Work hard, stay curious, and never stop questioning. I am here to guide you every step of the way.