Periodic Classification of ElementsClass 10 Science Part 1 Chapter 2 Notes – Maharashtra Board

Summary: Periodic Classification of Elements

Around 1800 only about 30 elements were known; today we know 118. Studying each one separately would be impossible, so scientists looked for patterns. First they simply split elements into metals and non-metals, and later noticed a third class, the metalloids. Then came cleverer ideas based on atomic mass.

Dobereiner grouped similar elements in threes, called triads, where the middle element's atomic mass was close to the average of the other two. Newlands lined up elements by atomic mass and found every eighth element behaved like the first, just like notes in music. Both ideas worked only for a few elements.

Mendeleev made the first truly useful table. He arranged 63 elements by atomic mass and by their properties, and boldly left gaps for elements not yet discovered. When scandium, gallium and germanium were found with the properties he predicted, the world was convinced. Still, his table could not place isotopes, cobalt and nickel, or hydrogen neatly.

In 1913 Moseley showed that atomic number, not atomic mass, is the basic property of an element. The modern periodic table arranges elements by atomic number in 18 groups and 7 periods, divided into s, p, d and f blocks. An element's position comes straight from its electronic configuration.

The table also shows clear trends. Going across a period, atoms get smaller and become less metallic and more non-metallic. Going down a group, atoms get bigger and more metallic. Elements in the same group share the same valency. Knowing an element's place lets you predict how it will behave.

Key points to remember

• Early classification divided elements into metals and non-metals; metalloids were recognised later.
• Dobereiner (1817) made triads of three similar elements; the atomic mass of the middle one is about the mean of the other two.
• Newlands (1866) proposed the Law of Octaves: every eighth element has properties similar to the first. It worked only up to calcium.
• Mendeleev's periodic law: Properties of elements are a periodic function of their atomic masses.
• Mendeleev arranged 63 elements; vertical columns are groups and horizontal rows are periods.
• Mendeleev left gaps and predicted eka-boron, eka-aluminium and eka-silicon, later found as scandium, gallium and germanium.
• Demerits of Mendeleev's table: no proper place for isotopes, the Co–Ni order problem, uneven rise in atomic mass, and the uncertain position of hydrogen.
• Moseley (1913) showed that atomic number (number of protons) is more fundamental than atomic mass.
• Modern periodic law: Properties of elements are a periodic function of their atomic numbers.
• The modern periodic table has 7 periods, 18 groups, 118 places and four blocks: s (groups 1–2), p (13–18), d (3–12) and f (lanthanides and actinides).
• Elements in a group have the same number of valence electrons; elements in a period have the same number of shells.
• A shell can hold at most 2n² electrons: K = 2, L = 8, M = 18, N = 32.
• Across a period (left to right): atomic radius decreases, metallic character decreases, non-metallic character increases.
• Down a group: atomic radius increases, metallic character (electropositivity) increases, electronegativity decreases.

Periodic Classification of Elements – Notes

1. Why classify elements? and Dobereiner's triads

• As more elements were discovered and more facts about them were collected, it became hard to study each one separately. Arranging them in groups with similar properties makes study easy.
• The first simple classification was into metals and non-metals. Later, elements with properties in between, called metalloids (like silicon), were noticed.
• In 1817 the German scientist Dobereiner linked the properties of elements with their atomic masses.
• He formed groups of three chemically similar elements called triads, written in increasing order of atomic mass.
• In each triad, the atomic mass of the middle element is roughly the average (mean) of the atomic masses of the first and third elements.
• Examples of his triads: lithium–sodium–potassium, calcium–strontium–barium, chlorine–bromine–iodine.
• Limitation: only a few of the known elements could be fitted into triads, so the idea could not classify all elements.

2. Newlands' Law of Octaves

• In 1866 the English scientist John Newlands arranged the elements known then in increasing order of atomic mass, from hydrogen to thorium.
• He noticed that every eighth element showed properties like the first. For example, sodium is the eighth element from lithium and both are similar; magnesium resembles beryllium; chlorine resembles fluorine.
• He compared this with music. In Indian music the seven notes Sa, Re, Ga, Ma, Pa, Dha, Ni repeat, and the next Sa is the eighth note. In western music it is Do, Re, Mi, Fa, So, La, Ti.
• Law of Octaves: when elements are arranged in increasing order of atomic mass, the properties of every eighth element are similar to those of the first.
• Limitation 1: the law worked well only up to calcium.
• Limitation 2: to fit all elements into 56 boxes, he squeezed two elements into one box, e.g. Co and Ni, Ce and La.
• Limitation 3: some dissimilar elements came together. Metals Co and Ni were placed with halogens, while iron, which is like Co and Ni, was placed with non-metals O and S.
• Limitation 4: there was no room for newly discovered elements, and their properties did not fit the pattern.

3. Mendeleev's periodic table

• The Russian scientist Dmitri Mendeleev developed his periodic table between 1869 and 1872. It is the most important step in the history of classification.
• He took atomic mass as the basic property and arranged the 63 elements known then in increasing order of atomic mass.
• He also used physical and chemical properties: formulae of hydrides and oxides, melting points, boiling points and densities of elements and their compounds.
• He wrote each element's details on a separate card and kept rearranging the cards until a pattern appeared.
• Elements with similar properties repeated after definite intervals, so he stated: Properties of elements are a periodic function of their atomic masses.
• Vertical columns are called groups and horizontal rows are called periods. The top of each group shows the general formula of its oxide and hydride, e.g. R₂O for group I, RO₂ and RH₄ for group IV.

4. Merits of Mendeleev's periodic table

• Mendeleev believed that the data of his time was not final. This flexible, scientific attitude is the reason for his table's merits.
• Correction of atomic masses: some atomic masses were revised so that elements fitted their proper places. Beryllium's mass was corrected from 14.09 to 9.4, and it was placed before boron.
• Gaps for undiscovered elements: he left empty places and named the missing elements eka-boron, eka-aluminium and eka-silicon, with predicted atomic masses 44, 68 and 72.
• He even predicted their properties. These were later discovered as scandium (Sc), gallium (Ga) and germanium (Ge), and their properties matched his predictions closely.
• For example, the predicted density of eka-aluminium was 5.9 g/cm³, and gallium's actual density is 5.94 g/cm³; both have an amphoteric oxide.
• Place for noble gases: when helium, neon and argon were discovered later, a new zero group was added without disturbing the rest of the table.

5. Demerits of Mendeleev's periodic table

• Cobalt and nickel have the same whole-number atomic mass (59), so it was unclear which should come first.
• Isotopes were discovered later. They have the same chemical properties but different atomic masses, e.g. Cl-35 and Cl-37. On the basis of atomic mass they would need different places, which is wrong.
• The increase in atomic mass from one element to the next is not uniform, so one could not predict how many elements might exist between two heavy elements.
• Position of hydrogen: hydrogen forms compounds like alkali metals (HCl like NaCl, H₂O like Na₂O) and also exists as diatomic molecules like halogens (H₂ like Cl₂, forms NaH like NaCl).
• So it could not be decided whether hydrogen belongs to group I (alkali metals) or group VII (halogens).

6. Modern periodic law and structure of the modern periodic table

• When Mendeleev made his table, nobody knew the inside of the atom. Atomic number was just a serial number.
• In 1913 Henry Moseley, using X-ray tube experiments, showed that atomic number (Z) equals the positive charge on the nucleus, i.e. the number of protons.
• This made atomic number a more fundamental property than atomic mass. Modern periodic law: Properties of elements are a periodic function of their atomic numbers.
• The modern periodic table (long form) arranges elements in increasing order of atomic number. Most defects of Mendeleev's table disappear, but hydrogen's position is still not settled.
• It has 7 horizontal rows (periods 1 to 7) and 18 vertical columns (groups 1 to 18). Two rows at the bottom are the lanthanide and actinide series. There are 118 boxes in all, and all 118 elements are now known.
• Blocks: s-block = groups 1 and 2; p-block = groups 13 to 18; d-block = groups 3 to 12 (transition elements); f-block = lanthanides and actinides.
• A zig-zag line in the p-block separates metals (left) from non-metals (right). Metalloids lie along this line.
• Elements after uranium (atomic number 92), i.e. 93 to 118, are man-made, radioactive and short-lived.

7. Groups, periods and electronic configuration

• An element's position in the modern table is decided by its electronic configuration.
• Groups: all elements of a group have the same number of valence electrons. Group 1 (alkali metals) have 1, group 2 (alkaline earth metals) have 2, group 17 (halogens) have 7.
• Going down a group, one new shell is added each time, so the number of shells increases.
• Periods: all elements of a period have the same number of shells. Period 2 elements (Li to Ne) use shells K and L; period 3 elements (Na to Ar) use K, L and M.
• Across a period, atomic number and the number of valence electrons increase by one at each step. A new period begins when a new shell starts filling.
• The maximum number of electrons in a shell is 2n². So period 1 has 2 elements and period 2 has 8.
• The M shell can hold 18, but period 3 still has only 8 elements because of the octet rule (the outermost shell holds at most 8).
• Chemical reactivity depends on the number of valence electrons and on which shell they are in, both of which we can read from the element's position.

8. Periodic trends: valency and atomic size

• Periodic trends are the regular changes in properties seen as we move along a period or down a group.
• Valency depends on valence electrons. For 1 to 4 valence electrons, valency equals that number; for 5 to 8, valency = 8 − valence electrons.
• Across period 3: Na 1, Mg 2, Al 3, Si 4, P 3, S 2, Cl 1, Ar 0. Valency first rises from 1 to 4 and then falls to 0.
• Down a group, valency stays the same because the number of valence electrons is the same (group 1 all 1, group 2 all 2, group 18 all 0).
• Atomic radius is the distance between the nucleus and the outermost shell. It is measured in picometre (1 pm = 10^-12 m).
• Across a period, atomic radius decreases. Nuclear charge rises by one each step, but electrons enter the same shell, so they are pulled closer.
• Down a group, atomic radius increases. A new shell is added each time, so the outermost electrons are farther away even though nuclear charge also increases.

9. Periodic trends: metallic and non-metallic character

• In each period, metals (like Na, Mg) are on the left, non-metals (like S, Cl) on the right, and metalloids (like Si) between them.
• Metals have 1 to 3 valence electrons that are held weakly, so they lose them easily and form cations. This tendency is called electropositivity, the metallic character.
• Inner shells shield the valence electrons, so the effective nuclear charge felt by them is less than the actual nuclear charge.
• Down a group, the valence shell is farther from the nucleus, the effective pull is weaker, and the cation left behind has a stable complete octet. So metallic character increases down a group.
• Across a period, nuclear charge increases and size decreases, so the effective nuclear charge rises. Electrons are held more tightly and metallic character decreases.
• The tendency of an atom to attract and accept electrons to form an anion is electronegativity, the non-metallic character. It increases across a period and decreases down a group.
• The more electropositive or electronegative an element is, the more reactive it is.
• Gradation in group 17: all halogens exist as X₂. F₂ and Cl₂ are gases, Br₂ is a liquid and I₂ is a solid.
• Gradation in group 2: Be does not react with water, Mg reacts with steam, and Ca, Sr, Ba react with water at room temperature, faster and faster down the group.

Formulas at a glance

Key terms

Periodic Classification of Elements – Flashcards

Read the question, answer it in your head, then tap the card to check. 20 cards.

Periodic Classification of Elements – MCQ Quiz

15 questions. Tap an option — you will see at once if it is right and why.

How to study Periodic Classification of Elements

⏱ About 3 hours over 3 days

1. Day 1 — History of classification (60 min)
   Make a timeline: Dobereiner 1817, Newlands 1866, Mendeleev 1869–72, Moseley 1913. Under each, write the basis, the law and two or three limitations. Check two triads yourself with (a + c)/2. List Mendeleev's merits and demerits in two columns.
2. Day 2 — Modern periodic table (60 min)
   Draw a blank outline of the modern periodic table, mark groups 1–18, periods 1–7, the s, p, d, f blocks and the zig-zag line. Write configurations of the first 20 elements and, from each, read off the period, group and valency. Learn 2n² and why period 3 has only 8 elements.
3. Day 3 — Trends and self-test (60 min)
   Make a trend chart with arrows: across a period and down a group for atomic radius, valency, metallic and non-metallic character, with one-line reasons. Practise 'give scientific reasons' answers aloud. Then do our 20 flashcards and the 15-question quiz, and revise every wrong answer.

✍️ Diagrams to practise

• Outline of the modern periodic table showing groups, periods, s, p, d, f blocks and the zig-zag line separating metals, metalloids and non-metals
• Arrow diagram of periodic trends (atomic radius, electropositivity, electronegativity) across a period and down a group
• Table of electronic configurations of elements of periods 2 and 3 (new period, new shell)
• Shell capacity table: K, L, M, N with n and 2n²