Periodic Classification of Elements — Maharashtra (SSC) Class 10 Science Solutions (Free)

The Modern Periodic Table arranges all known elements in order of increasing atomic number, into 18 vertical columns (called groups) and 7 horizontal rows (called periods). Its basis — the Modern Periodic Law: the physical and chemical properties of elements are periodic functions of their atomic numbers. Elements placed in the same group share similar chemical properties, because they have the same number of electrons in their outermost (valence) shell. How this differs from earlier attempts: Mendeleev's original periodic table (1869) arranged elements by increasing ATOMIC MASS rather than atomic number. This mostly worked, but caused a few genuine anomalies — for instance, tellurium (atomic mass ≈128) had to be placed before iodine (atomic mass ≈127) to keep it in the correct chemical family, which broke the mass-ordering rule. Once the atomic number (the actual number of protons) was discovered by Moseley in 1913, switching the ordering principle from atomic mass to atomic number resolved every such anomaly — which is exactly why the modern table uses atomic number, not mass, as its organizing basis.

A periodic property is a physical or chemical property that shows a predictable, repeating pattern of increase or decrease as elements are arranged in order of increasing atomic number. The four major periodic trends: 1. Atomic radius — DECREASES across a period (left to right), INCREASES down a group. 2. Ionisation energy — INCREASES across a period, DECREASES down a group. 3. Electronegativity — INCREASES across a period, DECREASES down a group. 4. Metallic character — DECREASES across a period, INCREASES down a group. Notice the pattern: atomic radius and metallic character move together (both decrease across a period), while ionisation energy and electronegativity move together, in the OPPOSITE direction to atomic radius. This isn't a coincidence — as explained in later answers, all four trends trace back to the same underlying cause: how strongly the nucleus attracts the outermost electrons, which depends on nuclear charge and atomic size together.

All three elements — Na, Mg, and Al — belong to Period 3 of the periodic table, meaning they all have electrons filling the same number of shells (three). Order of atomic radius: Na > Mg > Al (radius decreases left to right) Why: moving from Na (atomic number 11) to Mg (12) to Al (13), each successive element has one more proton in its nucleus, increasing the nuclear charge, while the added electron goes into the SAME third shell (the number of shells stays constant across a period). A stronger nuclear charge pulls all the electrons — including those in the outer shell — inward more tightly, without adding any new shell to compensate. The net effect is that the atom's electron cloud is pulled in, so the atomic radius shrinks as you move across a period, even though more electrons are being added.

Ionisation energy is the minimum energy required to remove the most loosely held electron from a neutral, gaseous atom. Across a period (left to right), two things happen together: 1. The nuclear charge increases (one more proton is added to each successive element). 2. The atomic radius decreases (as explained above), because the number of electron shells stays the same while nuclear charge grows. Together, this means the outermost electron of an atom further right in a period is BOTH closer to the nucleus AND experiencing a stronger positive pull from it. Since removing an electron means overcoming the nucleus's attraction, an electron that is held more tightly requires more energy to remove — so ionisation energy rises steadily as you move across a period. The shielding effect (how much inner-shell electrons block the nucleus's pull on outer electrons) stays roughly constant across a period, since no new inner shell is added — this is why the ionisation-energy increase across a period is fairly smooth, unlike the sharper drop seen down a group.

Order of increasing ionisation energy: Cs < K < Na All three belong to Group 1 (the alkali metals), meaning they all have exactly one electron in their outermost shell, but that shell is different for each: Na's valence electron is in the 3rd shell, K's in the 4th, and Cs's in the 6th shell. Why ionisation energy DECREASES down this group: 1. Atomic radius increases going down the group, since each successive element adds an entirely new electron shell. 2. The valence electron therefore sits progressively FARTHER from the nucleus. 3. At the same time, the number of inner electron shells between the nucleus and the valence electron increases, increasing the shielding effect — the inner electrons partly block the nucleus's pull on the outer electron. 4. Both effects combine to weaken the nucleus's grip on the valence electron the further down the group you go, so LESS energy is needed to remove it. This is exactly why caesium is one of the most reactive metals known — its single valence electron is held so weakly that it reacts violently even with the moisture in air.

Metallic character is an element's tendency to lose electrons easily and form positively charged ions (cations) — the defining chemical behaviour of metals. How it varies: - Down a group: metallic character INCREASES. As explained above, ionisation energy falls down a group (electrons are held more loosely, further from the nucleus, with more shielding), so it becomes progressively easier for the atom to lose an electron — i.e., to behave more metallically. - Across a period (left to right): metallic character DECREASES. Ionisation energy rises across a period (electrons are held more tightly), so elements become progressively LESS willing to lose electrons, and instead tend toward gaining electrons — the defining behaviour of non-metals. This is why the alkali metals (Group 1, far left, bottom of the group especially) show the strongest metallic character of any elements, while the halogens and noble gases (Groups 17-18, far right) show the weakest — the periodic table is, in this sense, organised from most metallic (bottom-left) to least metallic (top-right).

Electronegativity measures how strongly an atom attracts a shared pair of electrons in a chemical bond. Fluorine has an electronegativity of about 4.0 (the highest of any element), compared to chlorine's roughly 3.0, on the Pauling scale. This might seem surprising, since chlorine (Period 3) is a larger atom than fluorine (Period 2), and larger atomic size usually reduces an atom's pull on its own electrons. The key is that electronegativity depends on how close the NUCLEUS is to the SHARED electron pair in a bond, not on the atom's overall size: 1. Fluorine's valence electrons are in its 2nd shell, much closer to its nucleus than chlorine's valence electrons, which sit in the 3rd shell. 2. Fluorine has fewer inner electron shells shielding its nucleus, so its nuclear charge is felt more strongly by any electrons nearby, including a shared bonding pair. 3. This closer, less-shielded nuclear attraction lets fluorine pull a shared electron pair toward itself more strongly than chlorine can, despite chlorine technically having a higher raw nuclear charge (17 protons vs fluorine's 9). This is why fluorine is the most electronegative element on the entire periodic table, and why it forms such strongly polar (and highly reactive) bonds.

When Mendeleev arranged the known elements of his time (1869) by atomic mass into families with similar properties, he found a few gaps where no known element fit the expected pattern. Rather than assume his table was wrong, he boldly left these gaps empty and predicted that undiscovered elements would eventually be found to fill them — going as far as predicting their approximate atomic mass and chemical properties from the trend of their neighbours in the table. One famous example: Mendeleev predicted an element he called "eka-silicon" (meaning "one place below silicon"), estimating its atomic mass at about 72 and describing several of its expected properties, based on the pattern of elements around the gap. In 1886 — 15 years after the prediction — the German chemist Clemens Winkler discovered a new element, germanium, with an atomic mass of about 72.6 and properties matching Mendeleev's prediction almost exactly. This successful prediction was one of the strongest pieces of evidence that Mendeleev's underlying principle — that element properties really do repeat periodically — was capturing something fundamentally true about atomic structure, well before atomic number or electron configuration were even understood.

Step 1: Write the electron configuration using the shell-filling rule (2, 8, 8...). Atomic number 16 means 16 electrons: K shell = 2, L shell = 8, remaining = 16 − 2 − 8 = 6 electrons in the M shell. Configuration: 2, 8, 6 Step 2: Determine the period. The period number equals the number of the highest (outermost) shell containing electrons. Here, the outermost shell is the 3rd (M) shell, so this element is in Period 3. Step 3: Determine the group. This element has 6 electrons in its valence shell, in the p-block. For p-block elements, the modern IUPAC group number is 10 + (number of valence electrons) = 10 + 6 = Group 16. This element (atomic number 16) is Sulfur, correctly placed in Period 3, Group 16 — the same group as oxygen, which is why sulfur shares oxygen's tendency to form 2− ions and similar compounds (like H₂S being chemically analogous to H₂O).