Chapter 5 — Periodic Classification of Elements
Welcome to HSLC Guru! This page provides a complete English-medium guide for Class 10 Science Chapter 5 — Periodic Classification of Elements as per the ASSEB (Assam State School Education Board) syllabus. Here you will find a clear chapter summary, every textbook question and answer, additional MCQs, fill-in-the-blanks, true/false questions, and a glossary. The chapter explains how scientists organised more than 118 elements into a meaningful pattern called the Periodic Table, beginning with Dobereiner’s triads and culminating in the Modern Periodic Table based on atomic number. Master this chapter for your HSLC board exam.
Chapter Summary
The need to classify elements arose when the number of known elements grew too large to be studied individually. The earliest attempt was by Johann Wolfgang Dobereiner in 1817, who arranged elements with similar properties in groups of three called triads. The atomic mass of the middle element was approximately the arithmetic mean of the atomic masses of the other two — for example, in the triad Li (7), Na (23), K (39), the mean of 7 and 39 is 23. The limitation was that Dobereiner could identify only three triads, so the system was not universally applicable.
In 1866, John Newlands proposed the Law of Octaves, stating that when elements are arranged in increasing order of atomic mass, every eighth element resembles the first in properties — like the eighth note in a musical octave. This worked up to calcium but failed for heavier elements; Newlands also forced two elements into the same slot, ignoring the discovery of newer elements. Then in 1869, Dmitri Mendeleev arranged 63 known elements in increasing order of atomic mass into a table of periods (rows) and groups (columns), formulating the Periodic Law: “the properties of elements are a periodic function of their atomic masses.” Mendeleev left gaps for undiscovered elements (eka-aluminium=Ga, eka-silicon=Ge) and even predicted their properties accurately. Limitations included the wrong position of hydrogen, no fixed place for isotopes, and certain pairs (Co–Ni, Ar–K) appearing out of mass order.
In 1913, Henry Moseley showed through X-ray studies that atomic number (Z), not atomic mass, is the fundamental property of an element. This led to the Modern Periodic Law: “The properties of elements are a periodic function of their atomic numbers.” The Modern Periodic Table contains 7 horizontal periods and 18 vertical groups. Elements in the same group have the same number of valence electrons and therefore similar chemical properties. Elements in the same period have the same number of shells. The position of an element is decided by its electronic configuration: period number = number of shells; group number reflects valence electrons.
Periodic trends across periods and down groups are: Atomic size decreases across a period (nuclear charge increases) and increases down a group (new shell added). Valency remains the same down a group; across a period it first increases (1 to 4) and then decreases (4 to 0). Metallic character decreases across a period and increases down a group, while non-metallic character shows the opposite trend. Electronegativity (tendency to attract a shared pair of electrons) increases across a period and decreases down a group. Ionisation energy (energy needed to remove the outermost electron) increases across a period and decreases down a group. Metals are placed on the left, non-metals on the right, and metalloids along the zig-zag line.
Textbook Questions and Answers
1-Mark Questions
Q1. Who proposed the Law of Triads?
Answer: Johann Wolfgang Dobereiner proposed the Law of Triads in 1817.
Q2. State the Law of Octaves.
Answer: When elements are arranged in increasing order of atomic mass, the properties of every eighth element are similar to those of the first, like notes in a musical octave.
Q3. On what basis did Mendeleev arrange the elements?
Answer: Mendeleev arranged the elements in increasing order of their atomic masses and similarity in chemical properties.
Q4. State the Modern Periodic Law.
Answer: The properties of elements are a periodic function of their atomic numbers.
Q5. How many groups and periods are there in the Modern Periodic Table?
Answer: The Modern Periodic Table has 18 groups and 7 periods.
Q6. Name the scientist who established that atomic number is more fundamental than atomic mass.
Answer: Henry Moseley established this through his X-ray spectra experiments in 1913.
Q7. Which group contains the noble gases?
Answer: Group 18 (zero group) contains the noble gases — He, Ne, Ar, Kr, Xe, Rn.
Q8. What is the valency of elements in Group 1?
Answer: Elements in Group 1 (alkali metals) have a valency of 1.
Q9. Define ionisation energy.
Answer: The minimum energy required to remove the outermost electron from an isolated gaseous atom is called ionisation energy.
Q10. Name the element with atomic number 17 and give its group.
Answer: Chlorine (Cl); it belongs to Group 17 (halogens).
Q11. What is the basic difference between Mendeleev’s Periodic Law and the Modern Periodic Law?
Answer: Mendeleev’s law is based on atomic mass, whereas the Modern Periodic Law is based on atomic number.
Q12. Which element is the most reactive metal in the periodic table?
Answer: Caesium (Cs) is the most reactive metal among naturally occurring elements; Francium is the most reactive but radioactive.
Q13. Why is hydrogen placed separately from other elements?
Answer: Hydrogen shows properties similar to both Group 1 (alkali metals) and Group 17 (halogens), so it does not fit perfectly in either group.
2-3 Marks Questions
Q11. Mention any two limitations of Dobereiner’s Triads.
Answer: (i) Dobereiner could classify only three triads — (Li, Na, K), (Ca, Sr, Ba), (Cl, Br, I); the system was not applicable to all known elements. (ii) The law failed for elements with very low or very high atomic mass; for example, the triad F, Cl, Br does not satisfy the rule perfectly.
Q12. Why did Newlands’ Law of Octaves fail?
Answer: (i) The law was applicable only up to calcium; beyond it, properties did not repeat after every eighth element. (ii) Newlands assumed only 56 elements existed and forced two elements into the same slot. (iii) Discovery of noble gases later disturbed the eight-element pattern.
Q13. What were the achievements of Mendeleev’s Periodic Table?
Answer: (i) Elements with similar properties were grouped together. (ii) Mendeleev left vacant places for elements yet to be discovered, like eka-aluminium (later Gallium) and eka-silicon (later Germanium), and predicted their properties accurately. (iii) Atomic masses of some elements (Be, In, U) were corrected based on their position in the table.
Q14. How does atomic size vary across a period and down a group?
Answer: Across a period (left to right), atomic size decreases because the nuclear charge increases while electrons are added to the same shell, pulling them closer to the nucleus. Down a group (top to bottom), atomic size increases because new shells are added, which increases the distance of valence electrons from the nucleus.
Q15. Why do elements of the same group show similar chemical properties?
Answer: Elements belonging to the same group have the same number of valence electrons in their outermost shell. Since chemical properties are largely determined by valence electrons, they exhibit similar reactions — for example, all Group 1 elements (Li, Na, K) react with water to form alkalis.
Q16. Distinguish between metals and non-metals on the basis of their position in the periodic table.
Answer: Metals are placed on the left side and at the centre of the table (Groups 1, 2 and the d-block). They lose electrons easily and are electropositive. Non-metals are placed on the right side (Groups 14–17 upper part). They gain electrons and are electronegative. A zig-zag line of metalloids (B, Si, Ge, As, Sb, Te) separates the two.
Q17. Why does the metallic character increase down a group?
Answer: As we move down a group, the atomic size increases due to the addition of new shells. The valence electrons are farther from the nucleus, so the nuclear pull on them weakens. Therefore, atoms lose electrons more easily, and metallic character increases down the group.
Q18. Element X has electronic configuration 2, 8, 7. Predict its (i) period, (ii) group, (iii) valency, (iv) nature.
Answer: (i) Period: 3 (three shells). (ii) Group: 17 (7 valence electrons → 10 + 7 = 17). (iii) Valency: 1 (gains 1 electron to complete octet). (iv) Nature: non-metal — element X is Chlorine.
Q19. Why does ionisation energy decrease down a group?
Answer: Down a group, atomic size increases as new shells are added. Outer electrons get farther from the nucleus and experience greater shielding by inner electrons. The effective nuclear pull on the valence electron decreases, so less energy is needed to remove it — ionisation energy decreases down a group.
5-6 Marks Questions
Q17. Explain the Modern Periodic Table on the basis of electronic configuration. Discuss how period and group are decided.
Answer: The Modern Periodic Table is based on the Modern Periodic Law given by Henry Moseley: “the properties of elements are a periodic function of their atomic numbers.” The table consists of 7 periods (horizontal rows) and 18 groups (vertical columns). The number of the period of an element is equal to the number of shells (energy levels) present in its atom. For example, sodium (Na, Z=11) has electronic configuration 2,8,1 — three shells — so it lies in the 3rd period. The group number of an element is decided by the number of valence electrons. For Groups 1 and 2, group number = number of valence electrons; for Groups 13–18, group number = 10 + number of valence electrons. Elements with the same valence-shell configuration fall in the same group and show similar chemical behaviour, while elements with the same number of shells fall in the same period and show a gradation in properties.
Q18. Discuss the periodic trends of atomic size, ionisation energy, electronegativity and metallic character across a period and down a group.
Answer: (i) Atomic size: decreases across a period due to increasing nuclear charge and increases down a group due to addition of new shells. (ii) Ionisation energy: increases across a period because the smaller atomic size and stronger nuclear pull make electron removal harder; decreases down a group because outer electrons are farther from the nucleus and are shielded by inner electrons. (iii) Electronegativity: increases across a period (small atoms attract bonding pair more strongly) and decreases down a group (large atoms attract less strongly). Fluorine is the most electronegative element. (iv) Metallic character: decreases across a period because the tendency to lose electrons decreases; increases down a group because atoms become larger and lose outer electrons easily. Non-metallic character shows the opposite trend.
Q19. What were the limitations of Mendeleev’s Periodic Table? How were they removed in the Modern Periodic Table?
Answer: Limitations of Mendeleev’s table: (i) The position of hydrogen was not fixed — it shows properties of both Group 1 and Group 17. (ii) Isotopes of an element have different atomic masses but were placed in the same position, which contradicted Mendeleev’s basis. (iii) Certain pairs were placed out of order of atomic mass — e.g., Co (58.9) was placed before Ni (58.7), Ar (39.9) before K (39.1) — based on chemical similarity. (iv) No separate place was given for lanthanides and actinides. (v) Cause of periodicity could not be explained. The Modern Periodic Table, based on atomic number, removed most of these defects: anomalies of atomic mass disappeared because atomic number increases sequentially; isotopes occupy the same position because they have the same atomic number; lanthanides and actinides are placed in two separate rows below the main table; and periodicity is explained by recurring electronic configurations.
Q20. An element X has atomic number 12. (a) Write its electronic configuration. (b) State its period and group. (c) Is it a metal or non-metal? (d) What is its valency? (e) Predict the formula of its oxide and chloride.
Answer: (a) Electronic configuration of X (Z=12) is 2, 8, 2. (b) Number of shells = 3, so it lies in the 3rd period. Number of valence electrons = 2, so it belongs to Group 2. (c) X is a metal (alkaline earth metal — Magnesium). (d) Valency = 2 (it loses 2 electrons to attain noble gas configuration). (e) Formula of oxide = XO (MgO); formula of chloride = XCl₂ (MgCl₂).
Q21. Compare the properties of Mendeleev’s Periodic Table and the Modern Periodic Table.
Answer: (i) Basis: Mendeleev arranged elements according to atomic mass; the modern table arranges them according to atomic number. (ii) Number of groups: Mendeleev’s table had 8 groups (each split into A and B sub-groups); the modern table has 18 groups. (iii) Number of periods: both tables have 7 periods, but the modern table accommodates 118 known elements. (iv) Position of isotopes: Mendeleev’s table had no fixed place for them; the modern table places isotopes together because they share the same atomic number. (v) Anomalous pairs: Co–Ni and Ar–K appear correctly in the modern table because atomic numbers follow proper order. (vi) Lanthanides and actinides: placed separately in the modern table to maintain symmetry. (vii) Cause of periodicity: explained by recurrence of similar valence-shell electronic configurations in the modern table, while it was unexplained in Mendeleev’s table.
Additional Multiple Choice Questions (MCQs)
Q1. Who is regarded as the father of the Modern Periodic Table?
(a) Newlands (b) Dobereiner (c) Mendeleev (d) Moseley
Answer: (c) Mendeleev.
Q2. The Modern Periodic Law is based on:
(a) atomic mass (b) atomic number (c) atomic volume (d) density
Answer: (b) atomic number.
Q3. Which of the following is an alkali metal?
(a) Mg (b) Ca (c) Na (d) Al
Answer: (c) Na.
Q4. The most electronegative element is:
(a) Cl (b) F (c) O (d) N
Answer: (b) F (Fluorine).
Q5. The total number of periods in the Modern Periodic Table is:
(a) 6 (b) 7 (c) 8 (d) 18
Answer: (b) 7.
Q6. An element with atomic number 20 belongs to which group?
(a) 1 (b) 2 (c) 17 (d) 18
Answer: (b) 2 (Calcium).
Q7. Across a period, atomic size:
(a) increases (b) decreases (c) remains same (d) first increases then decreases
Answer: (b) decreases.
Q8. Which of the following is a noble gas?
(a) H (b) N (c) Ne (d) Cl
Answer: (c) Ne.
Q9. The element predicted by Mendeleev as eka-silicon is:
(a) Gallium (b) Germanium (c) Scandium (d) Tellurium
Answer: (b) Germanium.
Q10. Down a group, metallic character:
(a) decreases (b) increases (c) remains constant (d) becomes zero
Answer: (b) increases.
Q11. The number of valence electrons in an element of Group 16 is:
(a) 4 (b) 5 (c) 6 (d) 7
Answer: (c) 6.
Q12. Which of the following has the largest atomic size?
(a) Li (b) Na (c) K (d) Rb
Answer: (d) Rb.
Fill in the Blanks
1. Dobereiner classified elements into groups of three called __________. Answer: triads.
2. Newlands proposed the Law of __________. Answer: Octaves.
3. The Modern Periodic Table is based on increasing order of __________. Answer: atomic number.
4. Group 17 elements are called __________. Answer: halogens.
5. Ionisation energy generally __________ down a group. Answer: decreases.
True or False
1. Mendeleev’s periodic law states that properties of elements are periodic functions of their atomic numbers. Answer: False (it was atomic mass).
2. Atomic size decreases across a period from left to right. Answer: True.
3. Noble gases belong to Group 1. Answer: False (Group 18).
4. The valency of elements in the same group is generally the same. Answer: True.
5. Fluorine is the most electronegative element in the periodic table. Answer: True.
Quick Recap: Periodic Trends at a Glance
| Property | Across a Period (left to right) | Down a Group (top to bottom) |
|---|---|---|
| Atomic size | Decreases | Increases |
| Valency (1 to 4 then 4 to 0) | Variable | Same |
| Metallic character | Decreases | Increases |
| Non-metallic character | Increases | Decreases |
| Electronegativity | Increases | Decreases |
| Ionisation energy | Increases | Decreases |
| Electron affinity | Generally increases | Generally decreases |
| Nuclear charge | Increases | Increases |
| Number of shells | Same | Increases by one |
| Number of valence electrons | Increases (1 to 8) | Same |
Important Triads and Predicted Elements
| Mendeleev’s Predicted Name | Actual Element Discovered | Discoverer / Year |
|---|---|---|
| Eka-Aluminium | Gallium (Ga) | Lecoq de Boisbaudran, 1875 |
| Eka-Boron | Scandium (Sc) | Lars Fredrik Nilson, 1879 |
| Eka-Silicon | Germanium (Ge) | Clemens Winkler, 1886 |
Common Group Names
| Group Number | Common Name | Examples |
|---|---|---|
| Group 1 | Alkali Metals | Li, Na, K, Rb, Cs, Fr |
| Group 2 | Alkaline Earth Metals | Be, Mg, Ca, Sr, Ba, Ra |
| Group 11 | Coinage Metals | Cu, Ag, Au |
| Group 16 | Chalcogens | O, S, Se, Te, Po |
| Group 17 | Halogens | F, Cl, Br, I, At |
| Group 18 | Noble Gases | He, Ne, Ar, Kr, Xe, Rn |
Glossary
| Term | Meaning |
|---|---|
| Triad | A group of three elements with similar properties in which the atomic mass of the middle element is the mean of the other two. |
| Law of Octaves | Newlands’ law stating that every eighth element has properties similar to the first when arranged by atomic mass. |
| Periodic Law (Mendeleev) | Properties of elements are a periodic function of their atomic masses. |
| Modern Periodic Law | Properties of elements are a periodic function of their atomic numbers. |
| Period | Horizontal row in the periodic table; equals the number of electron shells in the atom. |
| Group | Vertical column in the periodic table; elements share the same number of valence electrons. |
| Valency | Combining capacity of an atom, decided by valence electrons. |
| Atomic Size | Distance between the nucleus and the outermost shell of an atom. |
| Ionisation Energy | Minimum energy required to remove the outermost electron from a gaseous atom. |
| Electronegativity | Tendency of an atom to attract a shared pair of electrons in a covalent bond. |
| Metallic Character | Tendency of an atom to lose electrons and form positive ions. |
| Non-metallic Character | Tendency of an atom to gain electrons and form negative ions. |
| Metalloid | Element showing properties of both metals and non-metals (e.g., Si, Ge, As). |
| Noble Gases | Group 18 elements with completely filled outermost shells; chemically inert. |
| Alkali Metals | Group 1 elements (Li, Na, K, Rb, Cs, Fr) — highly reactive metals. |
| Halogens | Group 17 elements (F, Cl, Br, I, At) — highly reactive non-metals. |
| Eka-element | Term used by Mendeleev for an undiscovered element predicted by his table (e.g., eka-silicon = germanium). |
Important Numerical / Configuration Practice
P1. Find the period and group of the element with atomic number 19.
Answer: Z = 19; configuration 2, 8, 8, 1. Number of shells = 4, so period = 4. Valence electrons = 1, so group = 1. The element is Potassium (K).
P2. Identify the element with electronic configuration 2, 8, 8.
Answer: Configuration 2, 8, 8 corresponds to atomic number 18 — Argon (Ar). It belongs to period 3 and group 18 (noble gas).
P3. Arrange the following in increasing order of atomic size: Mg, Al, Na, Si.
Answer: All four lie in period 3. Atomic size decreases left to right; so Si < Al < Mg < Na. Increasing order: Si < Al < Mg < Na.
P4. Out of Cl and Br, which has higher electronegativity and why?
Answer: Chlorine (Cl) has higher electronegativity than Bromine (Br). Both are in Group 17, but Cl has fewer shells, so its valence electrons are closer to the nucleus and the atom attracts the bonding pair more strongly.
P5. An element belongs to period 3 and group 2. Identify the element and write its valency.
Answer: Period 3 (3 shells) and group 2 (2 valence electrons) → configuration 2, 8, 2 → Z = 12 → Magnesium (Mg). Valency = 2.
Exam Tips for Chapter 5
- Memorise the names and symbols of the first 20 elements with their electronic configurations — most numerical questions are based on them.
- Practise locating an element on the periodic table from its atomic number and vice versa.
- Remember the four key trends — atomic size, ionisation energy, electronegativity, and metallic character — and the reason behind each.
- Learn at least two limitations of Mendeleev’s table and how the modern table corrected each.
- Be ready with one example each: a triad, an octave failure, and an eka-element prediction.
- For 5-mark questions, draw a small portion of the periodic table to support your answer wherever asked.
Frequently Asked Short Questions (Board Style)
Q1. Why do all alkali metals show similar chemical properties?
Answer: Because all alkali metals have one valence electron in their outermost shell, which they readily lose to form +1 ions, leading to similar reactions.
Q2. Why are noble gases chemically inert?
Answer: Noble gases have completely filled outermost shells (octet, except He which has duplet). They are stable and do not need to gain, lose or share electrons.
Q3. Why is the second period called a short period?
Answer: The second period contains only 8 elements (Li to Ne) because the second shell can hold a maximum of 8 electrons.
Q4. Why does atomic size decrease from Li to F across period 2?
Answer: The number of shells remains the same (2) but the nuclear charge increases as protons are added. The increased pull on the electrons brings them closer, decreasing atomic size.
Q5. Why is fluorine the most electronegative element?
Answer: Fluorine has a small atomic size and high effective nuclear charge, allowing it to attract a bonding pair of electrons more strongly than any other element.
Q6. Which group elements are called transition elements?
Answer: Elements of Groups 3 to 12 in the d-block of the periodic table are called transition elements. Examples include Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn.
Q7. Where are lanthanides and actinides placed in the modern periodic table?
Answer: Lanthanides (atomic numbers 58–71) and actinides (atomic numbers 90–103) are placed in two separate rows at the bottom of the periodic table to maintain its compact shape.
Q8. What is meant by “periodicity of properties”?
Answer: The repetition of similar properties of elements at regular intervals when arranged in increasing order of atomic number is called periodicity, and is caused by the periodic recurrence of similar valence-shell electronic configurations.
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