QCM : Periodic Properties — 28 questions

Explication

Döbereiner grouped three chemically similar elements so that the middle element’s atomic mass was approximately the arithmetic mean of the first and third. Groups of eight describe Newlands’ Law of Octaves.

Explication

Newlands observed that certain properties repeated at every eighth element when elements were ordered by increasing atomic mass. The three-element pattern belongs to Döbereiner’s Triads.

Explication

Mendeleev arranged elements by increasing atomic mass and proposed that their physical and chemical properties were periodic functions of atomic masses. Atomic number became the basis of the modern law later.

Explication

Moseley showed that atomic number is the more fundamental basis for periodic classification. This corrected the limitations of using atomic mass as Mendeleev had done.

Explication

The modern periodic law states that the physical and chemical properties of elements are periodic functions of their atomic numbers. Atomic mass was the basis of Mendeleev’s earlier formulation.

Explication

The long-form periodic table arranges elements in increasing atomic number and contains 18 groups and 7 periods. Groups are vertical columns, while periods are horizontal rows.

Explication

The s-block contains elements whose last electron enters an s-orbital and occupies groups 1 and 2. The p-block occupies groups 13–18.

Explication

The d-block is associated with variable valency, coloured compounds, and catalytic properties. Inner-transition elements with f-orbital filling belong to the f-block.

Explication

Metals are mainly located on the left side of the periodic table and are generally solid, malleable, ductile, and good conductors. The other choices describe non-metals, metalloids, or incorrect combinations of properties.

Explication

Non-metals may occur in all three physical states: solid, liquid, or gas. This distinguishes them from metals, which are generally solid.

Explication

Shielding is the reduction in nuclear attraction experienced by an outer electron because inner-shell electrons partially screen the nucleus. The attraction remaining after shielding is described by effective nuclear charge.

Explication

Using Z_eff = Z − σ, the effective nuclear charge is 12 − 4.5 = 7.5. The nuclear charge is represented directly by Z, whereas shielding is subtracted from it.

Explication

Effective nuclear charge increases across a period because nuclear charge rises while shielding does not increase proportionally. Down a group, it increases only slowly and remains approximately constant for many outer electrons.

Explication

Across a period, electrons are added to the same principal shell while nuclear charge and effective nuclear charge increase, pulling the electrons closer to the nucleus. Therefore, atomic radius decreases.

Explication

Moving down a group adds electron shells, placing the outer electrons farther from the nucleus. Effective nuclear charge does not increase dramatically, so the atomic radius generally increases.

Explication

Adding electrons to form an anions increases electron–electron repulsion, making the ion larger. Removing electrons to form a cation leaves the same nuclear charge acting on fewer electrons, making the ion smaller.

Explication

In an isoelectronic series, all species have the same number of electrons, so increasing nuclear charge pulls those electrons more strongly and decreases size. Thus the order is O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺ > Al³⁺.

Explication

Ionisation enthalpy is the energy needed to remove an electron from an isolated gaseous atom or ion, expressed in kJ mol⁻¹. Electron gain enthalpy instead concerns electron addition.

Explication

The second ionisation removes an electron from X⁺ to produce X²⁺. The first ionisation begins with neutral X.

Explication

After the first electron is removed, the species is positively charged, so its remaining electrons are held more strongly. Consequently, IE₂ > IE₁ and successive ionisation enthalpies increase.

Explication

Ionisation enthalpy generally increases across a period because nuclear attraction becomes stronger and decreases down a group as atomic size and shielding increase.

Explication

Electron gain enthalpy is the enthalpy change for X(g) + e⁻ → X⁻(g). It involves electron addition, unlike ionisation enthalpy, which involves electron removal.

Explication

Electron gain enthalpy generally becomes more negative across a period and less negative down a group. These trends reflect changes in nuclear attraction, shielding, and atomic size.

Explication

Noble gases possess stable ns²np⁶ configurations, so adding another electron is unfavorable and their electron gain enthalpies are practically zero or positive.

Explication

Electronegativity is the tendency of an element in a molecule to attract the shared pair of electrons toward itself. It is measured on the Pauling scale.

Explication

Representative-element valence is generally related to the number of outermost electrons; for H-type valency, it is commonly eight minus that number. This differs from oxidation state, which is treated as a charge.

Explication

Groups 15, 16, and 17 commonly show valencies of 3/5, 2/6, and 1/7, respectively. These paired values distinguish them from the simpler pattern for groups 1, 2, 13, and 14.

Explication

Oxidation state can be treated as the charge acquired by an atom when bonding is considered using electronegative relationships. It is distinct from valence, which relates primarily to outermost electrons.