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Classification of Elements and Periodicity in Properties

NCERT Class 11 · Chemistry Based on NCERT Class 11 Chemistry textbook · Free CBSE study kit

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Chapter Notes

Overview
This chapter traces the historical development of the Periodic Table from Dobereiner's triads to Mendeleev's Periodic Law and the Modern Periodic Law based on atomic number. It explains how the electronic configuration of elements determines their position in the s, p, d, and f blocks, and explores periodic trends in atomic radius, ionic radius, ionization enthalpy, electron gain enthalpy, and electronegativity. The one big idea a student must walk away with is that the properties of elements are a periodic function of their atomic numbers, arising from the periodic variation in their electronic configurations.
Key Concepts (31)
Periodic Law
The physical and chemical properties of the elements are periodic functions of their atomic numbers.
e.g. Just as seasons repeat every year, the properties of elements repeat after regular intervals when arranged by increasing atomic number.
Modern Periodic Law
The physical and chemical properties of the elements are periodic functions of their atomic numbers.
e.g. Moseley's X-ray studies showed that atomic number, not atomic mass, is the fundamental property governing periodicity.
Atomic Number (Z)
The number of protons in the nucleus of an atom, which is equal to the number of electrons in a neutral atom.
e.g. Carbon has atomic number 6, meaning it has 6 protons and 6 electrons.
Electronic Configuration
The distribution of electrons into orbitals (s, p, d, f) of an atom, which determines its position in the Periodic Table.
e.g. Sodium (Na) has electronic configuration 1s²2s²2p⁶3s¹, placing it in Group 1 and Period 3.
s-Block Elements
Elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) with outermost electronic configuration ns¹ and ns² respectively.
e.g. Lithium (Li) with configuration [He]2s¹ is an s-block element.
p-Block Elements
Elements belonging to Groups 13 to 18 with outermost electronic configuration varying from ns²np¹ to ns²np⁶.
e.g. Chlorine (Cl) with configuration [Ne]3s²3p⁵ is a p-block element.
d-Block Elements (Transition Elements)
Elements of Groups 3 to 12 characterized by the filling of inner d orbitals, with general outer electronic configuration (n-1)d¹⁻¹⁰ns⁰⁻².
e.g. Iron (Fe) with configuration [Ar]3d⁶4s² is a d-block element.
f-Block Elements (Inner-Transition Elements)
Elements in which the last electron enters the f-orbital, comprising the lanthanoids (Ce to Lu) and actinoids (Th to Lr), with outer configuration (n-2)f¹⁻¹⁴(n-1)d⁰⁻¹ns².
e.g. Cerium (Ce) with configuration [Xe]4f¹5d¹6s² is an f-block element.
Ionization Enthalpy (ΔiH)
The energy required to remove an electron from an isolated gaseous atom in its ground state.
e.g. The first ionization enthalpy of sodium is 496 kJ mol⁻¹.
Electron Gain Enthalpy (ΔegH)
The enthalpy change accompanying the addition of an electron to a neutral gaseous atom to form a negative ion.
e.g. Chlorine has a highly negative electron gain enthalpy of -349 kJ mol⁻¹.
Electronegativity
A qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself.
e.g. Fluorine has the highest electronegativity of 4.0 on the Pauling scale.
Atomic Radius
Half the distance between the nuclei of two atoms of the same element bonded together; for non-metals it is covalent radius, for metals it is metallic radius.
e.g. The atomic radius of chlorine is 99 pm (half of Cl₂ bond distance 198 pm).
Ionic Radius
The radius of an ion (cation or anion) measured in an ionic crystal.
e.g. Na⁺ has ionic radius 95 pm, while its parent atom Na has atomic radius 186 pm.
Isoelectronic Species
Two or more species with the same number of electrons, same number of atoms, and same structure.
e.g. O²⁻, F⁻, Na⁺, and Mg²⁺ all have 10 electrons and are isoelectronic.
Law of Triads
Dobereiner's observation that in a triad of three elements, the atomic weight of the middle element is approximately the average of the other two, and its properties are intermediate.
e.g. In the triad Li (7), Na (23), K (39), the atomic weight of Na (23) is roughly the average of 7 and 39.
Law of Octaves
Newlands' observation that when elements are arranged in increasing order of atomic weight, every eighth element has properties similar to the first.
e.g. Lithium (Li) and sodium (Na) are eight elements apart and show similar properties.
Mendeleev's Periodic Law
The properties of the elements are a periodic function of their atomic weights.
e.g. Mendeleev left gaps for undiscovered elements like eka-aluminium (gallium) and eka-silicon (germanium) and predicted their properties.
Eka-aluminium
Mendeleev's name for the undiscovered element under aluminium, later identified as gallium.
e.g. Mendeleev predicted eka-aluminium would have atomic weight 68 and density 5.9 g/cm³; gallium was found to have atomic weight 70 and density 5.94 g/cm³.
Eka-silicon
Mendeleev's name for the undiscovered element under silicon, later identified as germanium.
e.g. Mendeleev predicted eka-silicon would have atomic weight 72 and density 5.5 g/cm³; germanium was found to have atomic weight 72.6 and density 5.36 g/cm³.
Shielding Effect (Screening Effect)
The reduction in effective nuclear charge on the valence electron due to the repulsion from inner core electrons.
e.g. In lithium, the 2s electron is shielded from the +3 nuclear charge by the 1s² core, experiencing a net charge less than +3.
Effective Nuclear Charge
The net positive charge experienced by an electron in an atom after accounting for shielding by inner electrons.
e.g. Across a period, effective nuclear charge increases, pulling electrons closer and decreasing atomic radius.
Metals
Elements on the left side of the Periodic Table that are solids at room temperature (except mercury), good conductors of heat and electricity, malleable and ductile.
e.g. Sodium (Na) is a metal that conducts electricity and can be cut with a knife.
Non-metals
Elements on the top right side of the Periodic Table that are usually solids or gases at room temperature, poor conductors, and brittle.
e.g. Sulfur (S) is a non-metal that is brittle and does not conduct electricity.
Metalloids (Semi-metals)
Elements that border the zig-zag line in the Periodic Table and show properties intermediate between metals and non-metals.
e.g. Silicon (Si) and germanium (Ge) are metalloids used in semiconductors.
IUPAC Nomenclature for Elements with Z > 100
A systematic naming system using numerical roots (nil, un, bi, tri, quad, pent, hex, sept, oct, enn) combined with 'ium' to name elements until an official name is approved.
e.g. Element with Z = 120 is named unbinilium (Ubn).
Transuranium Elements
Elements with atomic number greater than 92 (uranium), which are man-made and radioactive.
e.g. Plutonium (Pu, Z=94) is a transuranium element discovered by Glenn T. Seaborg.
Lanthanoids
The 4f-inner transition series of elements from cerium (Z=58) to lutetium (Z=71), placed separately at the bottom of the Periodic Table.
e.g. Cerium (Ce) is the first lanthanoid, with electron configuration [Xe]4f¹5d¹6s².
Actinoids
The 5f-inner transition series of elements from thorium (Z=90) to lawrencium (Z=103), placed separately at the bottom of the Periodic Table.
e.g. Thorium (Th) is the first actinoid, with electron configuration [Rn]6d²7s².
Valence Shell
The outermost shell of an atom, which determines the chemical properties and the group of the element.
e.g. For sodium (Na), the valence shell is n=3 with one electron (3s¹).
Period
A horizontal row in the Periodic Table; the period number corresponds to the highest principal quantum number (n) of the elements in that period.
e.g. Period 2 contains elements from Li (n=2) to Ne (n=2).
Group (Family)
A vertical column in the Periodic Table; elements in the same group have similar valence shell electronic configurations and similar properties.
e.g. Group 1 elements (alkali metals) all have ns¹ configuration.
Formulas & Equations (6)
X(g) → X⁺(g) + e⁻
Reads as: A gaseous atom X loses an electron to form a gaseous cation X⁺ and an electron.
Variables: X = element symbol; g = gaseous state; e⁻ = electron
Use when: Defining first ionization enthalpy (ΔiH) for an element.
⚠ Students forget to specify 'gaseous' state and 'ground state' — the definition requires an isolated gaseous atom in its ground state.
X⁺(g) → X²⁺(g) + e⁻
Reads as: A gaseous cation X⁺ loses an electron to form a gaseous dication X²⁺ and an electron.
Variables: X⁺ = singly charged cation; X²⁺ = doubly charged cation; g = gaseous state; e⁻ = electron
Use when: Defining second ionization enthalpy.
⚠ Students confuse second ionization enthalpy with the energy to remove a second electron from a neutral atom — it is from the +1 ion.
X(g) + e⁻ → X⁻(g)
Reads as: A gaseous atom X gains an electron to form a gaseous anion X⁻.
Variables: X = element symbol; g = gaseous state; e⁻ = electron; X⁻ = anion
Use when: Defining electron gain enthalpy (ΔegH).
⚠ Students write the equation backwards or forget the negative sign convention — if energy is released, ΔegH is negative.
ΔegH = –Ae – 5/2 RT
Reads as: Electron gain enthalpy equals negative of electron affinity minus five halves times the gas constant times temperature.
Variables: ΔegH = electron gain enthalpy (kJ mol⁻¹); Ae = electron affinity (kJ mol⁻¹); R = gas constant (8.314 J mol⁻¹ K⁻¹); T = temperature (K)
Use when: Relating electron gain enthalpy at temperature T to electron affinity at absolute zero.
⚠ Students ignore the –5/2 RT term and treat electron gain enthalpy as equal to electron affinity.
Bond distance in Cl₂ = 198 pm; Covalent radius of Cl = 99 pm
Reads as: The bond distance in chlorine molecule is 198 picometers, so the covalent radius of chlorine is half of that, 99 picometers.
Variables: pm = picometer (10⁻¹² m)
Use when: Explaining how atomic radius is estimated from bond distances in covalent molecules.
⚠ Students think atomic radius is the same as the distance between two nuclei — it is half that distance.
Metallic radius of Cu = 128 pm (half of 256 pm internuclear distance in solid copper)
Reads as: The metallic radius of copper is 128 picometers, which is half the distance between two adjacent copper atoms in the metallic crystal.
Variables: pm = picometer
Use when: Explaining how metallic radius is estimated for metals.
⚠ Students confuse metallic radius with covalent radius — they are measured differently.
How It Works
Determining the Number of Elements in a Period
  1. Identify the principal quantum number (n) for the period.
  2. List the orbitals being filled in that period in order of increasing energy (e.g., for n=4: 4s, 3d, 4p).
  3. Count the total number of atomic orbitals available (e.g., 1 s-orbital + 5 d-orbitals + 3 p-orbitals = 9 orbitals for n=4).
  4. Multiply the number of orbitals by 2 (each orbital holds 2 electrons) to get the maximum number of electrons, which equals the number of elements in that period.
💡 Students often forget that d and f orbitals from lower principal quantum numbers are filled in higher periods — e.g., 3d orbitals are filled in Period 4, not Period 3.
Classifying Elements into s, p, d, f Blocks
  1. Write the electronic configuration of the element.
  2. Identify the last orbital that is being filled with electrons (the differentiating electron).
  3. If the last electron enters an s-orbital, the element belongs to the s-block.
  4. If the last electron enters a p-orbital, the element belongs to the p-block.
  5. If the last electron enters a d-orbital, the element belongs to the d-block.
  6. If the last electron enters an f-orbital, the element belongs to the f-block.
💡 Helium (1s²) is an exception — it is an s-block element but placed in the p-block (Group 18) because it has a completely filled valence shell like noble gases.
Naming Elements with Z > 100 Using IUPAC Nomenclature
  1. Write the atomic number (Z) in digits.
  2. Replace each digit with its corresponding root from the IUPAC table: 0=nil, 1=un, 2=bi, 3=tri, 4=quad, 5=pent, 6=hex, 7=sept, 8=oct, 9=enn.
  3. Combine the roots in the order of the digits.
  4. Add the suffix 'ium' at the end.
  5. The symbol is formed from the first letter of each root (e.g., Ubn for unbinilium).
💡 Students often forget that the roots are combined without spaces and that 'ium' is added only at the end, not after each root.
Trend in Atomic Radius Across a Period
  1. As we move from left to right across a period, the atomic number increases.
  2. The added electrons go into the same valence shell (same n).
  3. The nuclear charge (number of protons) increases.
  4. The shielding effect from inner core electrons does not increase significantly.
  5. The effective nuclear charge on the valence electrons increases.
  6. This pulls the valence electrons closer to the nucleus, decreasing the atomic radius.
💡 Students often think atomic radius increases across a period because more electrons are added — but the increased nuclear charge outweighs electron-electron repulsion.
Trend in Atomic Radius Down a Group
  1. As we move down a group, the principal quantum number (n) increases.
  2. New energy levels (shells) are added.
  3. The valence electrons are farther from the nucleus.
  4. The inner filled shells shield the valence electrons from the nuclear charge.
  5. The increase in shielding outweighs the increase in nuclear charge.
  6. The atomic radius increases down the group.
💡 Students sometimes think atomic radius decreases down a group because nuclear charge increases — but the addition of new shells dominates.
Trend in Ionization Enthalpy Across a Period
  1. As we move from left to right across a period, atomic number increases.
  2. Effective nuclear charge increases because shielding does not increase much.
  3. The valence electrons are held more tightly by the nucleus.
  4. More energy is required to remove an electron.
  5. Ionization enthalpy generally increases across a period.
💡 There are exceptions: Be has higher ionization enthalpy than B (because Be removes an s-electron, B removes a p-electron which is easier), and N has higher than O (because O has electron-electron repulsion in a doubly occupied p-orbital).
Trend in Ionization Enthalpy Down a Group
  1. As we move down a group, atomic size increases.
  2. The valence electron is farther from the nucleus.
  3. Shielding by inner electrons increases.
  4. The effective nuclear charge on the valence electron decreases.
  5. Less energy is required to remove the electron.
  6. Ionization enthalpy decreases down a group.
💡 Students often confuse this with electron gain enthalpy trend — ionization enthalpy always decreases down a group, while electron gain enthalpy becomes less negative (but not always monotonically).
Trend in Electron Gain Enthalpy Across a Period
  1. As we move from left to right across a period, atomic size decreases.
  2. Effective nuclear charge increases.
  3. The added electron is closer to the nucleus and experiences stronger attraction.
  4. More energy is released when the electron is added.
  5. Electron gain enthalpy becomes more negative across a period.
💡 Oxygen has less negative electron gain enthalpy than sulfur — this is an exception because the added electron in oxygen goes into the smaller n=2 level and experiences greater repulsion.
Trend in Electron Gain Enthalpy Down a Group
  1. As we move down a group, atomic size increases.
  2. The added electron is farther from the nucleus.
  3. Electron-electron repulsion decreases in larger orbitals.
  4. Less energy is released (or more energy is required) when the electron is added.
  5. Electron gain enthalpy becomes less negative down a group.
💡 Fluorine has less negative electron gain enthalpy than chlorine — this is an exception because the added electron in fluorine goes into the smaller n=2 level and experiences greater repulsion.
Trend in Electronegativity Across a Period and Down a Group
  1. Across a period: atomic size decreases, effective nuclear charge increases, so the atom attracts shared electrons more strongly — electronegativity increases.
  2. Down a group: atomic size increases, shielding increases, so the atom attracts shared electrons less strongly — electronegativity decreases.
💡 Electronegativity is not a measurable quantity — it is a qualitative scale (Pauling scale) with fluorine assigned the highest value of 4.0.
Comparing Sizes of Isoelectronic Species
  1. Identify the isoelectronic species (same number of electrons).
  2. Determine the nuclear charge (Z) of each species.
  3. The species with the higher nuclear charge will attract the electrons more strongly.
  4. The species with the higher nuclear charge will have the smaller radius.
  5. Cations are smaller than anions among isoelectronic species.
💡 Students often forget that isoelectronic species have different nuclear charges — e.g., Mg²⁺ (Z=12) is smaller than Na⁺ (Z=11) even though both have 10 electrons.
Exam Traps — Don't Make These Mistakes
⚠ Confusing Atomic Radius with Ionic Radius
✗ The atomic radius of sodium is 95 pm.
→ 95 pm is the ionic radius of Na⁺, not the atomic radius of Na. The atomic radius of Na is 186 pm. A cation is always smaller than its parent atom because it has fewer electrons but the same nuclear charge.
✓ The atomic radius of sodium is 186 pm, while the ionic radius of Na⁺ is 95 pm.
⚠ Mistaking the Trend in Electron Gain Enthalpy for Halogens
✗ Fluorine has the most negative electron gain enthalpy because it is the most electronegative element.
→ Fluorine actually has a less negative electron gain enthalpy (-328 kJ mol⁻¹) than chlorine (-349 kJ mol⁻¹). This is because the added electron in fluorine goes into the smaller n=2 level and experiences significant electron-electron repulsion.
✓ Chlorine has the most negative electron gain enthalpy among all elements (-349 kJ mol⁻¹), not fluorine.
⚠ Ignoring Exceptions in Ionization Enthalpy Trends
✗ Ionization enthalpy increases uniformly from Li to Ne across Period 2.
→ There are two exceptions: Be (1s²2s²) has higher ionization enthalpy than B (1s²2s²2p¹) because Be removes an s-electron while B removes a p-electron which is easier to remove. Also, N (1s²2s²2p³) has higher ionization enthalpy than O (1s²2s²2p⁴) because O has electron-electron repulsion in a doubly occupied p-orbital.
✓ Ionization enthalpy generally increases across Period 2, but with dips at B (lower than Be) and O (lower than N) due to electronic configuration effects.
⚠ Confusing Mendeleev's Periodic Law with Modern Periodic Law
✗ Mendeleev's Periodic Law states that properties of elements are a periodic function of their atomic numbers.
→ Mendeleev's Periodic Law was based on atomic weights, not atomic numbers. The Modern Periodic Law (based on atomic numbers) was proposed after Moseley's work in 1913.
✓ Mendeleev's Periodic Law: The properties of the elements are a periodic function of their atomic weights. Modern Periodic Law: The physical and chemical properties of the elements are periodic functions of their atomic numbers.
⚠ Misidentifying the Block of Helium
✗ Helium is a p-block element because it is placed in Group 18.
→ Helium has electronic configuration 1s², so its last electron enters the s-orbital, making it an s-block element. However, it is placed in the p-block (Group 18) because it has a completely filled valence shell like other noble gases.
✓ Helium is an s-block element (1s²) but is placed in Group 18 (p-block) due to its inert nature and completely filled valence shell.
⚠ Forgetting the 'Gaseous' and 'Ground State' Conditions in Definitions
✗ Ionization enthalpy is the energy required to remove an electron from an atom.
→ The definition requires the atom to be in the gaseous state and in its ground state. Without these conditions, the definition is incomplete and will lose marks in board exams.
✓ Ionization enthalpy is the energy required to remove an electron from an isolated gaseous atom in its ground state.
Compare & Contrast
Basis of classification
Mendeleev's Periodic Law
Modern Periodic Law
Mendeleev's Periodic Law arranges elements by increasing atomic weights, while the Modern Periodic Law arranges them by increasing atomic numbers, which is more fundamental as shown by Moseley's X-ray studies.
Energy change during electron removal vs. addition
Ionization Enthalpy
Electron Gain Enthalpy
Ionization enthalpy is always positive (energy required to remove an electron), while electron gain enthalpy can be negative (energy released when adding an electron) or positive (energy required to add an electron), depending on the element.
Orbital being filled
s-Block Elements
p-Block Elements
s-block elements have their last electron in an s-orbital (ns¹ or ns²) and are highly reactive metals, while p-block elements have their last electron in a p-orbital (ns²np¹ to ns²np⁶) and include metals, non-metals, and metalloids.
Orbital being filled and radioactivity
Lanthanoids
Actinoids
Lanthanoids involve filling of 4f orbitals and are non-radioactive (except promethium), while actinoids involve filling of 5f orbitals and are all radioactive, with many being man-made transuranium elements.
Exam Focus
Marks pattern: Typically 2-3 MCQs (1 mark each), 1-2 short answer questions (2-3 marks), 1 long answer question (5 marks) on periodic trends or electronic configuration and block classification. Total approximately 8-12 marks.
Most asked: Modern Periodic Law and its significance · Electronic configuration and block classification of elements · Periodic trends in atomic radius, ionization enthalpy, electron gain enthalpy, and electronegativity · Exceptions in periodic trends (e.g., Be vs B, N vs O, F vs Cl) · Isoelectronic species and their size comparison · IUPAC nomenclature of elements with Z > 100 · Mendeleev's predictions and their success
Don't skip: Explain the trends in ionization enthalpy across a period and down a group with reasons and exceptions.; Classify elements into s, p, d, f blocks based on electronic configuration and justify.; Compare and contrast Mendeleev's Periodic Table with the Modern Periodic Table.; Arrange isoelectronic species in order of increasing ionic radius and justify.; Explain why chlorine has a more negative electron gain enthalpy than fluorine.
One-Line Revision (read aloud before the exam)
  1. The Modern Periodic Law states that the physical and chemical properties of elements are periodic functions of their atomic numbers.
  2. Mendeleev's Periodic Law was based on atomic weights, while the Modern Periodic Law is based on atomic numbers as shown by Moseley's X-ray studies.
  3. Elements are classified into s, p, d, and f blocks based on the orbital into which the last electron enters.
  4. s-block elements have ns¹ or ns² configuration and are highly reactive metals.
  5. p-block elements have ns²np¹ to ns²np⁶ configuration and include metals, non-metals, and metalloids.
  6. d-block elements (transition elements) have general configuration (n-1)d¹⁻¹⁰ns⁰⁻² and are all metals.
  7. f-block elements (inner-transition elements) include lanthanoids (4f) and actinoids (5f) and are placed separately at the bottom.
  8. Atomic radius decreases across a period due to increasing effective nuclear charge.
  9. Atomic radius increases down a group due to addition of new energy levels and increased shielding.
  10. Ionization enthalpy generally increases across a period and decreases down a group.
  11. Beryllium has higher ionization enthalpy than boron because Be removes an s-electron while B removes a p-electron.
  12. Nitrogen has higher ionization enthalpy than oxygen because oxygen has electron-electron repulsion in a doubly occupied p-orbital.
  13. Electron gain enthalpy becomes more negative across a period and less negative down a group.
  14. Chlorine has the most negative electron gain enthalpy (-349 kJ mol⁻¹), not fluorine.
  15. Electronegativity increases across a period and decreases down a group; fluorine has the highest value of 4.0 on the Pauling scale.
  16. Cations are smaller than their parent atoms, while anions are larger than their parent atoms.
  17. Among isoelectronic species, the one with higher nuclear charge has a smaller radius.
  18. Helium is an s-block element (1s²) but placed in Group 18 due to its inert nature.
  19. The IUPAC name for element with Z=120 is unbinilium (Ubn).
  20. Metals are on the left side of the Periodic Table, non-metals on the top right, and metalloids border the zig-zag line.

MCQs — 10 Questions with Answers

Q1. Dobereiner's Law of Triads applies to which set of elements?

  • A. Li, Na, K (with Na as middle element) ✓
  • B. Be, Mg, Ca (with Mg as middle element)
  • C. C, Si, Ge (with Si as middle element)
  • D. N, P, As (with P as middle element)

Answer: A — Li (7), Na (23), K (39): Na's atomic weight (23) is approximately the average of Li (7) and K (39), exemplifying Dobereiner's Triad principle exactly.

Q2. Newlands' Law of Octaves failed because it worked only up to which element?

  • A. Oxygen
  • B. Nitrogen
  • C. Calcium ✓
  • D. Potassium

Answer: C — Newlands observed the octave pattern (every 8th element similar to the 1st) held true only until calcium; beyond calcium, the pattern broke down.

Q3. What key decision did Mendeleev make that Newlands did not?

  • A. He arranged elements by atomic weight in horizontal rows
  • B. He left gaps for undiscovered elements and violated strict atomic weight order for chemical similarity ✓
  • C. He used a cylindrical table arrangement
  • D. He recognized only s and p block elements

Answer: B — Mendeleev prioritized chemical properties (similar empirical formulas) over strict atomic weight order, placed iodine before tellurium, and predicted gallium and germanium to fill gaps—revolutionary bold predictions.

Q4. Which statement correctly explains why the Modern Periodic Law uses atomic number instead of atomic weight?

  • A. Atomic number is smaller and easier to calculate
  • B. Atomic number uniquely determines electronic configuration, which governs all periodic properties and trends ✓
  • C. Isotopes have the same atomic weight but different atomic numbers, causing confusion
  • D. Atomic weight was discovered after atomic number

Answer: B — Electronic configuration (determined by atomic number alone) directly causes periodic repetition of chemical properties; atomic weight does not determine electron configuration uniquely (isotopes).

Q5. The prediction of Eka-Aluminium by Mendeleev was later confirmed when which element was discovered?

  • A. Germanium
  • B. Gallium ✓
  • C. Scandium
  • D. Indium

Answer: B — Mendeleev left a gap under aluminium (Group 13) and called it Eka-Aluminium; this was confirmed as gallium when discovered, validating Mendeleev's predictive power.

Q6. Assertion: Mendeleev placed iodine (At. wt. 127) before tellurium (At. wt. 128) in his periodic table. Reason: Iodine's chemical properties are similar to fluorine and chlorine (halogens).

  • A. Both assertion and reason are correct; reason explains assertion ✓
  • B. Both assertion and reason are correct; reason does not explain assertion
  • C. Assertion is correct; reason is incorrect
  • D. Assertion is incorrect; reason is correct

Answer: A — Mendeleev did place I before Te despite I having higher atomic weight because I belongs to Group VII (halogens like F and Cl), demonstrating that chemical property similarity trumped weight order.

Q7. Which of the following is NOT a key characteristic that Mendeleev used to classify elements?

  • A. Empirical formulas of compounds formed by elements
  • B. Physical and chemical properties of elements
  • C. Atomic weight (with flexibility when properties clashed)
  • D. Mass number of the element ✓

Answer: D — Mendeleev relied on atomic weight, empirical formulas, and physical/chemical properties; mass number (total protons + neutrons) was not used and is largely irrelevant for periodic classification.

Q8. If an element X has atomic number 19 and element Y has atomic number 20, how many valence electrons does Y have if X is potassium (Group 1)?

  • A. 1 valence electron
  • B. 2 valence electrons ✓
  • C. 19 electrons total
  • D. 20 electrons total

Answer: B — Potassium (K, Z=19) is in Group 1 with 1 valence electron; the next element (Z=20) is calcium in Group 2 with 2 valence electrons, following the periodic trend of increasing valence electrons across a period.

Q9. Between Newlands and Mendeleev, which scientist's work was eventually awarded the Davy Medal by the Royal Society, London?

  • A. Only Newlands (in 1887) ✓
  • B. Only Mendeleev (in 1906)
  • C. Both scientists in the same year
  • D. Neither scientist received this award

Answer: A — Newlands' Law of Octaves, initially dismissed, was recognized 22 years later when he was awarded the Davy Medal in 1887; Mendeleev received other honors but not specified here as the primary one.

Q10. HOTS: A new element with atomic number 115 is discovered. Based on the Modern Periodic Law and periodic trends, predict which group it most likely belongs to and justify.

  • A. Group 13 (p-block), because 115 = 5d + 6s + 6p partially filled configuration fits p-block
  • B. Group 15 (p-block), because 115 = 7s² + 7p³ electronic configuration places it in p-block with 5 valence electrons ✓
  • C. Group 17 (halogens), because elements with atomic numbers ending in 5 are always halogens
  • D. Cannot be predicted without experimental data on its chemical properties

Answer: B — Element 115 (Moscovium) has configuration [Rn]5f¹⁴6d¹⁰7s²7p³, with 5 valence electrons in 7p, placing it in Group 15; this demonstrates that atomic number determines electronic configuration, which determines group placement via periodicity.

Flashcards

What was Dobereiner's Law of Triads and when was it proposed?

In 1829, Dobereiner noted that groups of three elements showed similar properties, with the middle element's atomic weight approximately halfway between the other two.

State Newlands' Law of Octaves.

Every eighth element in increasing order of atomic weights possessed properties similar to the first element, resembling octaves in music (true only up to calcium).

What is the Modern Periodic Law?

The properties of elements are a periodic function of their atomic number (not atomic weight, as Mendeleev originally stated).

Why did Mendeleev ignore atomic weight order in his periodic table?

He placed elements with similar properties together and predicted undiscovered elements would fill gaps, prioritizing chemical similarities over strict weight order.

Name the two elements Mendeleev predicted and left gaps for in his table.

Mendeleev predicted Eka-Aluminium (gallium) and Eka-Silicon (germanium) and left gaps under aluminium and silicon respectively.

What is the significance of atomic number in modern periodic classification?

Atomic number determines electronic configuration, which repeats periodically and causes the periodic repetition of chemical and physical properties.

How does electronic configuration justify periodic trends?

Elements in the same group have identical valence electron configurations, explaining similar chemical properties; periodicity arises from repeating patterns in filling electron shells.

Name the four blocks of the periodic table and what they represent.

s, p, d, and f blocks represent the subshells being filled: s (2 electrons max), p (6 electrons), d (10 electrons), and f (14 electrons).

Why is the periodic table called 'the everyday support for students'?

It organizes all 114 known elements into predictable groups and families, allowing rapid prediction of properties without studying each element individually.

How did Lothar Meyer contribute differently from Mendeleev?

Meyer plotted physical properties (atomic volume, melting point, boiling point) against atomic weight and discovered periodically repeated patterns, but Mendeleev published first and is credited.

Important Board Questions

State Mendeleev's Periodic Law and explain why he violated the strict order of atomic weights in his periodic table with one example. [2 marks]

Define the periodic law clearly; then explain that Mendeleev prioritized chemical properties (similar compounds) over weight order and cite iodine (127) placed before tellurium (128) because iodine forms halides like fluorine and chlorine.

Explain how the Modern Periodic Law (based on atomic number) overcomes the limitations of Mendeleev's Periodic Law (based on atomic weight). Illustrate with the case of iodine and tellurium. [5 marks]

Show that atomic number uniquely determines electronic configuration and valence electrons; explain that isotopes have different mass numbers but same atomic number (same chemistry); prove that I (Z=53, Group 17) and Te (Z=52, Group 16) now fit correctly by atomic number, not weight—atomic number reflects electron configuration which truly governs periodicity. Use the example to show why modern approach is superior.

Mendeleev left gaps in his periodic table and predicted the existence of undiscovered elements with specific properties. Critically analyze his predictions for Eka-Aluminium and Eka-Silicon: (a) Why did Mendeleev leave these gaps? (b) Which elements confirmed his predictions? (c) What does this success reveal about the predictive power of the periodic table as a scientific tool? [6 marks]

Part (a): He observed gaps between known elements with similar properties, so he hypothesized undiscovered elements would fill them. Part (b): Eka-Al = Ga, Eka-Si = Ge. Part (c): Explain that periodic classification is not arbitrary but reflects a deep underlying order (electronic configuration); the periodic table becomes predictive because properties recur systematically—this validates that the periodic law is a fundamental principle of chemistry, not just a convenient catalog. Connect to modern synthesis of new elements beyond uranium.

Assertion-Reason Questions for Classification of Elements and Periodicity in Properties

Q1. Assertion (A):

Dobereiner's Law of Triads was eventually dismissed as coincidence despite identifying a genuine periodic relationship.

Reason (R):

The Law of Triads successfully classified all elements known in the early 1800s into groups with predictable properties.

Answer: C — A is true but R is false — The assertion is true—Dobereiner's work was dismissed due to its limited scope, but it did reveal a genuine pattern. The reason is false because the Law only worked for a few elements, not all known elements. A true reason would explain why limited success led to dismissal.

Q2. Assertion (A):

Mendeleev prioritized chemical property similarities over strict adherence to atomic weight order when arranging elements.

Reason (R):

This priority allowed him to place chemically similar elements in the same group and leave gaps for undiscovered elements with predictable properties.

Answer: A — Both A and R are true, and R is the correct explanation of A — Both the assertion and reason are true. The reason directly explains why this prioritization was important—it preserved chemical grouping and enabled predictions. The example of iodine's placement exemplifies this principle perfectly.

Q3. Assertion (A):

The Modern Periodic Law states that properties of elements depend on their atomic weights.

Reason (R):

Atomic number, not atomic weight, is the true basis of periodicity because it directly reflects the electronic configuration of atoms.

Answer: D — A is false but R is true — The assertion is false—the Modern Periodic Law uses atomic number, not atomic weight (that was Mendeleev's original statement). The reason is true but irrelevant since the assertion itself is incorrect. This reflects the historical correction made to the periodic law.

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Case-Based Questions (CBQ) for Classification of Elements and Periodicity in Properties

Case Study:

In 1869, a research scientist was studying the properties of newly discovered elements and noticed some inconsistencies with the then-proposed atomic weight-based ordering. She observed that an element with lower atomic weight than another showed stronger chemical similarity to a third element that had a much higher atomic weight. She consulted Mendeleev's newly published Periodic Table and found that his arrangement explained this anomaly perfectly. The element with lower atomic weight had been placed in a group with chemically similar elements, suggesting that chemical properties were more reliable than strict atomic weight order for organizing the periodic table.

Q1. Which principle from Mendeleev's classification system best explains why the element with lower atomic weight was placed separately from strict weight ordering?

  • A. Mendeleev strictly followed atomic weight order without exceptions
  • B. Mendeleev prioritized grouping chemically similar elements together, even if it meant violating strict atomic weight sequence ✓
  • C. The element in question was a transition metal with unique properties
  • D. Atomic weight measurements were more accurate in Mendeleev's table

Answer: B

Q2. The scientist's observation that chemical similarity transcended atomic weight order would later support the development of:

  • A. The Law of Triads
  • B. Newlands' Law of Octaves
  • C. The Modern Periodic Law based on atomic number ✓
  • D. Chancourtois' helical table

Answer: C

Q3. If this same scientist had been working in 1829 with Dobereiner's classification system instead, what limitation might she have encountered?

  • A. Dobereiner's system could organize all elements into predictable triads
  • B. Dobereiner's system would have successfully explained the anomaly she observed
  • C. Dobereiner's system was too limited to address the element grouping she was investigating ✓
  • D. Dobereiner had already solved the problem of atomic weight inconsistencies

Answer: C

Q4. The scientist's confidence in Mendeleev's table was based on its ability to account for anomalies. Which additional feature of Mendeleev's system would have further strengthened her confidence?

  • A. The presence of gaps with predicted properties for undiscovered elements ✓
  • B. The use of atomic numbers instead of atomic weights
  • C. The grouping of elements based solely on atomic volume
  • D. The division of the table into exactly eight groups

Answer: A

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