Chapter 3: Classification of Elements and Periodicity in Properties Application Questions | chemistry Class 11 CBSE

Chapter 3: Classification of Elements and Periodicity in Properties Application Questions | chemistry Class 11 CBSE | ByteNotes.in

Application Question

Medium difficulty • Concept in a practical situation

Medium

Question 1

Applied Concept

A student is given an unknown element X that reacts vigorously with water, is always found in combined form in nature, and has a very low first ionization enthalpy. Based on periodic trends, identify the probable block and group of element X and justify your answer.

Element X belongs to the s-block, specifically Group 1 (alkali metals), because alkali metals have the lowest ionization enthalpies in any period and react vigorously with water (e.g., Na + H2O → NaOH + H2).
The very low ionization enthalpy indicates that the outermost ns1 electron is very loosely held, characteristic of alkali metals where a single electron is shielded by a complete inner noble gas configuration.
Being found only in combined form (never as free element) is consistent with high chemical reactivity — a direct consequence of low ionization enthalpy; highly reactive elements lose electrons readily to form stable ionic compounds.
As we move down Group 1 (Li to Cs), ionization enthalpy decreases further and reactivity increases, so element X is likely a heavier alkali metal such as K, Rb, or Cs based on the extreme vigour of its reaction.

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Application Question

Medium difficulty • Concept in a practical situation

Medium

Question 2

Applied Concept

With reference to the IUPAC system of nomenclature, write the temporary IUPAC name and symbol for element 115. If this element were to be officially recognised, predict the group and period it belongs to.

For element 115, the digits are 1, 1, 5 giving numerical roots un + un + pent + ium = Ununpentium, with symbol Uup (first letter of each root: U, U, p).
The official IUPAC name given to element 115 is Moscovium (Mc, symbol), named after the Moscow region of Russia where it was discovered.
Element 115 belongs to period 7 (since elements from Z=87 to 118 are in the 7th period) and Group 15 (similar to nitrogen, phosphorus, arsenic family) based on its position in the periodic table.
Its expected outer electronic configuration would be [Rn]5f146d107s27p3, placing it in the p-block as it fills 7p orbitals; it would be expected to show properties similar to other Group 15 elements though with significant differences due to relativistic effects.

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Application Question

Medium difficulty • Concept in a practical situation

Medium

Question 3

Applied Concept

With reference to the periodic table, explain why silicon is used as a semiconductor while sodium is a good electrical conductor and sulphur is a poor conductor.

Sodium is an s-block metal (Group 1) with one loosely held valence electron (ns1) and very low ionization enthalpy; in metallic sodium, this electron is delocalised and free to move, making sodium an excellent electrical conductor.
Sulphur is a p-block non-metal (Group 16) with configuration ns2np4; it has high ionization enthalpy and high electronegativity, so its electrons are tightly bound and not free to carry current, making it a poor conductor.
Silicon is a metalloid (semi-metal) bordering the zig-zag line between metals and non-metals in the periodic table; it has intermediate ionization enthalpy and electronegativity, showing properties of both metals and non-metals.
Silicon's intermediate electrical conductivity makes it a semiconductor; its conductivity can be controlled by adding impurities (doping), which is the basis of semiconductor devices; metalloids like silicon and germanium are technologically vital for electronics.

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Application Question

Medium difficulty • Concept in a practical situation

Medium

Question 4

Applied Concept

In a science experiment, a student observes that element A reacts more violently with water than element B, both in the same group. Element A also has a larger atomic radius. What can the student conclude about the relative positions of A and B in the group and their ionization enthalpies?

Since element A has a larger atomic radius than B and both belong to the same group, element A must be below element B in the group; atomic radius increases as we descend a group due to addition of principal energy levels.
Element A reacts more violently with water, indicating higher chemical reactivity; for metals in the same group, reactivity increases on moving down the group as ionization enthalpy decreases.
Element A must therefore have a lower ionization enthalpy than element B; as we go down the group, valence electrons are farther from the nucleus and better shielded by inner electrons, making them easier to remove.
Conclusion: A is positioned below B in the same group; A has lower ionization enthalpy, greater metallic character, and higher reactivity consistent with the downward trend for Group 1 (alkali metal) behaviour.

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Application Question

Hard difficulty • Concept in a practical situation

Hard

Question 5

Applied Concept

Using the concept of isoelectronic species, arrange the following ions in order of increasing ionic radius and justify: N3-, O2-, F-, Na+, Mg2+, Al3+.

All six species are isoelectronic with 10 electrons each (same as neon); their sizes differ based solely on their nuclear charges: N has Z=7, O has Z=8, F has Z=9, Na has Z=11, Mg has Z=12, Al has Z=13.
For isoelectronic species with the same number of electrons, a higher nuclear charge means the nucleus attracts all 10 electrons more strongly, resulting in a smaller ionic radius.
Therefore, the order of increasing ionic radius (smallest to largest) is: Al3+ < Mg2+ < Na+ < F- < O2- < N3-; Al3+ has the highest nuclear charge and smallest radius, while N3- has the lowest nuclear charge and largest radius.
This example from the chapter illustrates that the size of isoelectronic species is determined entirely by nuclear charge (Z); more protons pulling the same number of electrons = smaller size.

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Application Question

Medium difficulty • Concept in a practical situation

Medium

Question 6

Applied Concept

With reference to periodic trends, explain why fluorine is stored in special containers, while argon is used in light bulbs as an inert filling gas.

Fluorine (Group 17, Period 2) has the highest electronegativity (4.0 on Pauling scale) and one of the most negative electron gain enthalpies; it is the strongest oxidising agent and reacts with almost all materials including glass and most metals, requiring special containers such as those made of fluoropolymers.
Fluorine's extreme reactivity is a direct consequence of its high effective nuclear charge, small atomic radius, and the large energy released when it gains an electron to achieve the stable noble gas configuration.
Argon (Group 18, Period 3) has a completely filled valence shell (3s23p6); all orbitals are filled, it has very high ionization enthalpy, and its electron gain enthalpy is positive (energy must be supplied to add an electron), making it chemically inert.
Argon does not react with the hot tungsten filament or other components of a light bulb; this inertness makes it ideal as an inert atmosphere to prevent filament oxidation and extend bulb life.

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Application Question

Medium difficulty • Concept in a practical situation

Medium

Question 7

Applied Concept

A chemist testing two unknown oxides finds that one turns red litmus blue (oxide P) and the other turns blue litmus red (oxide Q). Using the concept of periodic trends in chemical reactivity, predict the probable position of the parent elements of P and Q in the periodic table.

Oxide P turns red litmus blue, indicating it is a basic oxide; according to periodic trends, basic oxides are formed by metallic elements located on the extreme left of a period (e.g., Group 1 or 2 metals like Na forming Na2O).
Oxide Q turns blue litmus red, indicating it is an acidic oxide; acidic oxides are formed by non-metallic elements located on the extreme right of a period (e.g., Group 16 or 17 non-metals like Cl forming Cl2O7).
The parent element of P is likely an alkali or alkaline earth metal (s-block, Groups 1 or 2) with low ionization enthalpy and highly metallic character; Na2O + H2O → 2NaOH is the classic example.
The parent element of Q is likely a non-metal from the p-block (right side of a period); elements with high electronegativity and non-metallic character form acidic oxides: Cl2O7 + H2O → 2HClO4.

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Application Question

Easy difficulty • Concept in a practical situation

Easy

Question 8

Applied Concept

Predict the formula of the compound formed between silicon (Group 14) and bromine (Group 17) using the concept of valence, and explain how you used the periodic table to arrive at the formula.

Silicon belongs to Group 14; its number of valence electrons is 4; the valence of silicon is 4 (equal to the number of outermost electrons for groups 1-4) since it can share all four outer electrons in bond formation.
Bromine belongs to Group 17 (halogens); its number of valence electrons is 7; the valence of bromine is 1 (= 8 minus 7, the number of outermost electrons), as it needs only one more electron to complete its octet.
To balance valences: silicon (valence 4) combines with 4 bromine atoms (valence 1 each); the formula of the compound is SiBr4.
This is the same approach used in the chapter (Problem 3.8a): silicon is a Group 14 element with valence 4, bromine is a halogen with valence 1, so the compound formed is SiBr4, consistent with the systematic use of valence from the periodic table.

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Application Question

Medium difficulty • Concept in a practical situation

Medium

Question 9

Applied Concept

A geologist finds a mineral containing only one type of metal whose oxide reacts with both dilute HCl and NaOH solution. Identify the nature of the oxide and predict where in the periodic table the metal is likely to be located.

An oxide that reacts with both acids (HCl) and bases (NaOH) is called an amphoteric oxide; the chapter gives Al2O3 and As2O3 as examples of amphoteric oxides.
Amphoteric oxides are formed by elements located in the centre of a period, where neither purely metallic nor purely non-metallic character dominates; this intermediate position gives the oxide dual reactivity.
The metal is likely located in the middle of a period in the p-block or early d-block, transitioning between the strongly metallic left and strongly non-metallic right; aluminium (Group 13) is a classic example given in the chapter.
The presence of an amphoteric oxide tells the geologist that the metal element has moderate electronegativity and ionization enthalpy — neither as low as alkali metals (which form basic oxides) nor as high as halogens (which form acidic oxides).

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Application Question

Hard difficulty • Concept in a practical situation

Hard

Question 10

Applied Concept

With reference to the periodic table, explain why beryllium is used in alloys requiring high strength and rigidity while barium compounds are used in rat poison and pyrotechnics, linking to their positions in the periodic table.

Beryllium is a Group 2 element in Period 2 with small atomic radius (111 pm) and high charge-to-radius ratio; its anomalous properties include forming compounds with significant covalent character, high melting point, and exceptional hardness compared to other Group 2 metals.
Beryllium's small size and high electronegativity (for a metal) give its compounds and alloys unusual strength and stiffness; it is placed in Period 2 where first-member anomalies give it unique character compared to heavier alkaline earth metals.
Barium is a Group 2 element in Period 6 with much larger atomic radius; it is far down the group where ionization enthalpy is very low, making it highly reactive; barium sulfate and other barium compounds are used medicinally and in pyrotechnics due to characteristic flame colours.
The contrast between Be and Ba illustrates the periodic trend: as we descend Group 2, atoms become larger, ionization enthalpy decreases, metallic and ionic character increases, and first-member anomalous behaviour (Be's covalent character) disappears.

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Chapter 3: Classification of Elements and Periodicity in Properties | Application Questions

A student is given an unknown element X that reacts vigorously with water, is always found in combined form in nature, and has a very low first ionization enthalpy. Based on periodic trends, identify the probable block and group of element X and justify your answer.

With reference to the IUPAC system of nomenclature, write the temporary IUPAC name and symbol for element 115. If this element were to be officially recognised, predict the group and period it belongs to.

With reference to the periodic table, explain why silicon is used as a semiconductor while sodium is a good electrical conductor and sulphur is a poor conductor.

In a science experiment, a student observes that element A reacts more violently with water than element B, both in the same group. Element A also has a larger atomic radius. What can the student conclude about the relative positions of A and B in the group and their ionization enthalpies?

Using the concept of isoelectronic species, arrange the following ions in order of increasing ionic radius and justify: N3-, O2-, F-, Na+, Mg2+, Al3+.

With reference to periodic trends, explain why fluorine is stored in special containers, while argon is used in light bulbs as an inert filling gas.

A chemist testing two unknown oxides finds that one turns red litmus blue (oxide P) and the other turns blue litmus red (oxide Q). Using the concept of periodic trends in chemical reactivity, predict the probable position of the parent elements of P and Q in the periodic table.

Predict the formula of the compound formed between silicon (Group 14) and bromine (Group 17) using the concept of valence, and explain how you used the periodic table to arrive at the formula.

A geologist finds a mineral containing only one type of metal whose oxide reacts with both dilute HCl and NaOH solution. Identify the nature of the oxide and predict where in the periodic table the metal is likely to be located.

With reference to the periodic table, explain why beryllium is used in alloys requiring high strength and rigidity while barium compounds are used in rat poison and pyrotechnics, linking to their positions in the periodic table.

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