🧪 Chemistry · Class 11 · NEET & JEE
Classification of Elements and Periodicity in Properties - Practice Questions with Answers
75 free MCQs on Classification of Elements and Periodicity in Properties, each with its own worked answer and explanation. The periodic table organises all 118 elements by atomic number and similar properties. Study periodic trends like atomic radius, ionization energy, and electronegativity, and learn how an element's position predicts its behaviour.
Take the timed Classification of Elements and Periodicity in Properties chapterwise test →75 practice questions on Classification of Elements and Periodicity in Properties, sorted Easy → Hard. Try each one first, then open its answer page for the worked explanation. Want the full theory first? Read the Classification of Elements and Periodicity in Properties notes.
The complete modern periodic table (118 elements). Image: Cepheus et al., Public Domain, via Wikimedia Commons.
Easy - 25 questions
Q1.
In Dobereiner's triads, the atomic mass of the middle element is approximately:
- A the average of the other two elements
- B the sum of the other two elements
- C twice that of the lightest element
- D half that of the heaviest element
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Q2.
Mendeleev arranged the elements in his periodic table in the increasing order of their:
- A atomic number
- B atomic mass
- C density
- D valency
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Q3.
The modern periodic law states that the physical and chemical properties of elements are a periodic function of their:
- A atomic mass
- B number of neutrons
- C atomic number
- D density
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Q8.
The modern periodic table is based on the atomic number. This was proposed by:
- A Moseley
- B Mendeleev
- C Dobereiner
- D Newlands
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Q9.
Elements in the same group have the same:
- A Number of valence electrons
- B Atomic mass in most textbook accounts
- C Number of neutrons during normal conditions
- D Atomic radius as generally observed
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Q11.
Atomic radius generally decreases across a period because:
- A Nuclear charge increases while electrons are added to the same shell
- B The number of occupied electron shells increases steadily in typical laboratory settings
- C Electron shielding from inner shells increases sharply under usual circumstances
- D Atomic mass decreases steadily across the period according to most researchers
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Q12.
Atomic radius generally increases down a group because:
- A New electron shells are added
- B Nuclear charge decreases
- C Electronegativity increases
- D Fewer protons
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Q17.
Which property increases across a period (left to right)?
- A Ionisation energy
- B Atomic radius
- C Metallic character
- D Electropositive character
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Q18.
Mendeleev predicted the existence of undiscovered elements by leaving gaps. He called them by names like:
- A Eka-aluminium, Eka-silicon
- B Proto-metal in the majority of cases studied
- C Pseudo-metals as widely reported
- D Quasi-elements in standard practice
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Q23.
Period 3 elements include:
- A Na, Mg, Al, Si, P, S, Cl, Ar
- B Li, Be, B, C, N, O, F, Ne
- C K, Ca, Sc...
- D H, He
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Q24.
Electron affinity is the energy released when:
- A A neutral atom gains an electron
- B A cation loses an electron
- C An atom forms a cation
- D Two atoms bond
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Q25.
Which element has the lowest first ionisation energy among noble gases?
- A Radon (Rn)
- B Helium (He)
- C Neon (Ne)
- D Argon (Ar)
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Medium - 25 questions
Q27.
Beryllium shows a diagonal relationship in the periodic table with which element?
- A Boron
- B Magnesium
- C Aluminium
- D Silicon
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Q28.
The increasing order of ionic radius for the isoelectronic species Na<sup>+</sup>, Mg<sup>2+</sup>, F<sup>-</sup> and O<sup>2-</sup> is:
- A O<sup>2-</sup> < F<sup>-</sup> < Na<sup>+</sup> < Mg<sup>2+</sup>
- B Na<sup>+</sup> < Mg<sup>2+</sup> < F<sup>-</sup> < O<sup>2-</sup>
- C F<sup>-</sup> < O<sup>2-</sup> < Na<sup>+</sup> < Mg<sup>2+</sup>
- D Mg<sup>2+</sup> < Na<sup>+</sup> < F<sup>-</sup> < O<sup>2-</sup>
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Q29.
The correct order of first ionisation enthalpy for the elements B, C, N and O is:
- A B < C < O < N
- B B < C < N < O
- C O < N < C < B
- D C < B < O < N
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Q30.
The shielding (screening) effect exerted by electrons in different subshells follows the order:
- A f > d > p > s
- B s > p > d > f
- C p > s > d > f
- D d > p > s > f
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Q31.
The effective nuclear charge (Zeff) increases across a period because:
- A Shielding remains roughly constant while nuclear charge increases
- B Shielding from inner electrons increases faster than nuclear charge
- C The total electron count decreases steadily across the period
- D Atomic mass decreases steadily across the period
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Q32.
Ionisation energy of sulfur (S) is lower than phosphorus (P) even though S has higher atomic number. This is because:
- A Phosphorus has a half-filled 3p subshell (extra stability)
- B Sulfur simply has more total electrons than phosphorus
- C Phosphorus has a lower nuclear charge than sulfur
- D Sulfur has a noticeably larger atomic radius than phosphorus
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Q33.
Which element has the highest first ionisation energy in period 2?
- A Neon (Ne)
- B Fluorine (F)
- C Nitrogen (N)
- D Oxygen (O)
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Q34.
The electron affinity of fluorine is less than that of chlorine because:
- A F's small size causes electron-electron repulsion in the compact 2p subshell
- B Fluorine generally has fewer protons in its nucleus than chlorine as widely reported
- C Fluorine is less electronegative than chlorine on the Pauling scale in standard practice
- D Fluorine has a noticeably larger atomic radius than chlorine under most conditions encountered
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Q35.
Diagonal relationship in the periodic table exists between:
- A Li & Mg, Be & Al, B & Si
- B Li & Na, Be & Ca as frequently observed in practice
- C H & Li in many documented cases
- D C & Si according to conventional understanding
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Q36.
The shielding effect of orbitals follows the order:
- A s > p > d > f
- B f > d > p > s
- C p > s > d > f
- D d > f > s > p
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Q37.
Screening constant in Slater's rules: electrons in the same shell contribute a shielding constant of:
- A 0.35 per electron
- B 0.85 per electron
- C 1.00 per electron
- D 0 per electron
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Q38.
Second ionisation energy is always greater than first ionisation energy because:
- A Electron is removed from a positive ion with higher nuclear charge attraction
- B The electron configuration rearranges into a higher-energy state in routine practice
- C The surrounding temperature rises during ionisation overall in most cases
- D The resulting ion becomes larger than the neutral atom under typical conditions
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Q39.
The period number corresponds to the:
- A Principal quantum number of the outermost electron
- B The total number of valence electrons in the outer shell
- C The group number assigned to that element's column
- D The atomic number of the first element in that row
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Q40.
Why do noble gases have very high ionisation energies?
- A They have completely filled valence shells with high Zeff
- B They have unusually large atomic radii compared to their period
- C They possess very few electrons in their outer shell
- D They are inherently radioactive elements
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Q41.
Metallic character decreases across a period because:
- A Ionisation energy increases and the tendency to lose electrons decreases
- B Atomic mass steadily decreases from left to right according to standard textbooks
- C The total number of electrons decreases across the period in general practice
- D Valence electrons are progressively removed across the period as frequently described
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Q42.
The oxidising power of halogens follows the order:
- A F<sub>2</sub> > Cl<sub>2</sub> > Br<sub>2</sub> > I<sub>2</sub>
- B I<sub>2</sub> > Br<sub>2</sub> > Cl<sub>2</sub> > F<sub>2</sub>
- C Cl<sub>2</sub> > F<sub>2</sub> > Br<sub>2</sub> > I<sub>2</sub>
- D All equal
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Q43.
Which element in period 3 has the highest ionisation energy?
- A Argon (Ar)
- B Chlorine (Cl)
- C Phosphorus (P)
- D Silicon (Si)
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Q44.
Electronegativity increases across a period and decreases down a group. This is due to:
- A Changes in atomic radius and effective nuclear charge
- B Atomic mass alone, independent of nuclear charge or radius
- C The number of neutrons present in the nucleus
- D Electron configuration alone, independent of radius or charge
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Q45.
The oxide of which period 3 element is amphoteric?
- A Al<sub>2</sub>O<sub>3</sub>
- B Na<sub>2</sub>O
- C P<sub>2</sub>O<sub>5</sub>
- D SiO<sub>2</sub>
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Q46.
Anomalous properties of Li compared to other alkali metals are due to:
- A Its very small size giving high charge density and polarising power
- B Its position as the most reactive metal in Group 1
- C Its unusually low melting point relative to the other alkali metals
- D Its comparatively low first ionisation energy among Group 1 metals
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Q47.
The first element in each period (alkali metals, except Period 1) is in which block?
- A s-block
- B p-block
- C d-block
- D f-block
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Q48.
Ionisation energy shows a dip between Group 2 and Group 13 because:
- A The p orbital electron is easier to remove than the s orbital electron
- B Group 13 elements simply have more protons in the nucleus
- C Group 2 elements experience unusually higher electron shielding
- D The extra electron in Group 13 is added to an inner shell
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Q49.
Which period 2 element forms the most stable +3 ion?
- A Aluminium (period 3 actually)
- B Nitrogen
- C Carbon
- D Oxygen
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Q50.
Isoelectronic species have the same:
- A Number of electrons
- B Number of protons
- C Atomic mass
- D Atomic number
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Hard - 25 questions
Q51.
The successive ionisation enthalpies (kJ mol<sup>-1</sup>) of an element are 738, 1451, 7733 and 10540. The element most likely belongs to:
- A Group 1
- B Group 13
- C Group 2
- D Group 14
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Q53.
As a direct consequence of the lanthanide contraction, which pair of elements has almost identical atomic radii?
- A Zr and Hf
- B Fe and Co
- C Na and K
- D Cu and Ag
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Q54.
The general formula of the highest (most common) oxide formed by an element of Group 15 is:
- A EO<sub>2</sub>
- B E<sub>2</sub>O<sub>3</sub>
- C E<sub>2</sub>O<sub>5</sub>
- D EO<sub>3</sub>
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Q55.
According to IUPAC nomenclature for elements with atomic number greater than 100, the element with atomic number 104 is named:
- A Unniltrium
- B Unnilquadium
- C Unnilpentium
- D Ununbium
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Q56.
The successive ionisation energies of magnesium show a large jump between the 2nd and 3rd ionisation energies. This indicates:
- A Mg has 2 valence electrons (after removing 2, the 3rd comes from a noble gas core)
- B Mg actually has 3 valence electrons available for removal
- C The third electron occupies a higher principal energy level than the first two
- D Mg behaves as a transition metal with variable oxidation states
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Q57.
Using Slater's rules, the effective nuclear charge for a 3p electron in chlorine (Z=17) is approximately:
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Q58.
Fluorine's electron affinity (-328 kJ/mol) is less negative than chlorine's (-349 kJ/mol). This anomaly is best explained by:
- A High electron-electron repulsion in the compact 2p subshell of fluorine
- B Fluorine's notably higher electronegativity value on the Pauling scale
- C Fluorine's higher first ionisation energy compared to chlorine
- D Chlorine's greater nuclear charge compared to fluorine
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Q59.
The second electron affinity of oxygen is endothermic (+744 kJ/mol) because:
- A Adding an electron to O- (a negative ion) requires energy to overcome repulsion
- B Oxygen's atomic radius is generally too small to accept another electron in typical laboratory settings
- C Oxygen has an unusually low electronegativity for its period under usual circumstances
- D The process instead forms an unstable O<sup>3-</sup> trianion intermediate according to most researchers
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Q60.
Elements of the 6th and 7th period in Groups 4-10 have nearly identical atomic radii to their 5th period counterparts. This is called:
- A Lanthanide contraction
- B Actinide expansion
- C Diagonal relationship
- D Inert pair effect
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Q61.
The 'inert pair effect' in Group 14 and 15 heavy elements (Pb, Bi) refers to:
- A Reluctance of the 6s<sup>2</sup> electrons to participate in bonding
- B The complete loss of all valence electrons during bonding
- C The unusually high reactivity of the inner d electrons
- D The preferential formation of purely metallic bonds in these elements
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Q62.
Why does phosphorus form PCl<sub>5</sub> but nitrogen does not form NCl<sub>5</sub>?
- A N has no available d orbitals to expand beyond octet; P uses 3d orbitals
- B Nitrogen is less electronegative than phosphorus in the majority of cases studied
- C Phosphorus has a higher first ionisation energy than nitrogen as widely reported
- D Nitrogen behaves as a metal rather than a non-metal in standard practice
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Q63.
Across period 3 (Na to Cl), the oxides change from basic to acidic. The amphoteric oxide is:
- A Al<sub>2</sub>O<sub>3</sub>
- B MgO
- C P<sub>2</sub>O<sub>5</sub>
- D SiO<sub>2</sub>
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Q64.
The Bohr-Bury rule states that the maximum number of electrons in any shell is:
- A 2n<sup>2</sup>
- B n<sup>2</sup>
- C 2n
- D 4n<sup>2</sup>
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Q65.
Which period 4 element has an anomalous electronic configuration [Ar]3d<sup>5</sup> 4s<sup>1</sup> instead of [Ar]3d<sup>4</sup> 4s<sup>2</sup>?
- A Chromium (Cr)
- B Calcium (Ca)
- C Titanium (Ti)
- D Vanadium (V)
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Q66.
The Newlands' law of octaves failed because:
- A It did not work for elements beyond calcium and did not account for undiscovered elements
- B It used the atomic mass values of the elements incorrectly under most conditions encountered
- C It largely ignored the existence of isotopes of each element as frequently observed in practice
- D It deliberately placed metals and non-metals in the same octave group in many documented cases
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Q67.
In which group does the element with configuration [Xe] 4f<sup>14</sup> 5d<sup>6</sup> 6s<sup>2</sup> belong?
- A Group 8 (d-block, 5d series = Osmium Os)
- B Group 6, matching the chromium-molybdenum-tungsten family
- C Group 18, the noble gas family with filled valence shells
- D Group 2, the alkaline earth metal family
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Q68.
Pauling's electronegativity scale is based on:
- A Bond energy data (extra ionic resonance energy)
- B Ionisation energy values alone, without bond energy data
- C Electron affinity values alone, without bond energy data
- D Atomic radius measurements alone, without bond energy data
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Q69.
The isoelectronic series O<sup>2-</sup>, F-, Ne, Na+, Mg<sup>2+</sup> all have 10 electrons. Their ionic radii decrease because:
- A Increasing nuclear charge attracts the same 10 electrons more strongly
- B Each successive species actually has fewer than 10 electrons according to conventional understanding
- C Electron shielding increases steadily across the series in routine practice
- D Each species occupies a different number of electron shells overall
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Q70.
Which factor primarily determines the position of hydrogen in the periodic table as ambiguous?
- A It can lose one electron like alkali metals (1s<sup>1</sup>) but also gain one like halogens (1s<sup>2</sup>)
- B It exists as a diatomic gas under ordinary atmospheric conditions like the halogens
- C Its most abundant isotope, protium, contains no neutron in its nucleus in general practice
- D It is generally the lightest and smallest element found on the entire periodic table
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Q71.
The period 4 transition metals show irregular filling because 3d and 4s energies are close. Copper's configuration is [Ar]3d<sup>10</sup> 4s<sup>1</sup> (not 3d<sup>9</sup> 4s<sup>2</sup>) because:
- A A completely filled 3d<sup>10</sup> is extra stable, as is d<sup>5</sup>
- B 3d orbitals are always lower in energy than the 4s orbital
- C The 4s orbital is filled before the 3d orbital in every transition metal
- D Copper's odd atomic number forces an unpaired 4s electron
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Q72.
Moseley's experiment showed a linear relationship between sqrt(frequency of X-ray) and:
- A Atomic number (Z)
- B Atomic mass
- C Number of neutrons
- D Group number
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Q73.
The reason the 3d subshell fills before 4p but after 4s in order of energy is explained by:
- A n+l rule: subshell with lower n+l fills first; if equal, lower n fills first
- B A simple rule that the lowest n value always fills first regardless of l
- C Shielding provided specifically by the 4s electrons pushes 3d lower
- D Unpredictable quantum fluctuations that vary by element
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Q74.
Van der Waals radius is always larger than covalent radius for the same element because:
- A It represents the distance in non-bonded contact (weak attraction), while covalent radius is within a bond
- B The two radii are conventionally reported using largely different units of measurement in most cases
- C The comparison is actually being made between two chemically different elements under typical conditions
- D The two radii are generally measured at very different experimental temperatures according to standard textbooks
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Q75.
Predicted by Mendeleev, which element confirmed his predictions most spectacularly?
- A Germanium (eka-silicon, discovered 1886)
- B Silicon, whose properties were already well known before Mendeleev
- C Gallium, discovered independently without matching a gap prediction
- D Scandium, discovered before Mendeleev published his periodic law
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