Test: Classification of Elements and Periodicity in Properties - Chemistry MCQ

The electronic arrangement of atoms, which is used to study the physical and chemical properties of the elements, has a logical effect on the periodic classification of elements.

The sum of all the electrons in the given electronic configuration will determine the atomic number. The element that is present just below the given element will have an outermost electronic configuration as 4s²4p³, hence, its full electronic configuration will be 1s²,2s²2p⁶,3s²3p⁶,4s², 3d¹⁰,4p². Thus, its atomic number is 33.

The amount of energy needed to remove one electron from an isolated gaseous atom's outermost orbit is known as the ionization enthalpy. The ionization enthalpy is lowest for the electronic structure 1s²2s²2p⁶3s¹.

Atomic number 15 (Z) belongs to P, this can be written as [Ne] 3s²3p³
Phosphorus is in group 15 of elements. In the ground state, the valence electron count for phosphorus is 5 and the valency is 3, from 3s²3p³.

The cation's ionic radius is always less than its atomic radius. The loss of electrons produces cations. Ionic radius decreases as a result of the rise in effective nuclear charge.

A group's electron affinity declines as its atomic number rises. Se, O, and S are all group 16 elements. Thus, the expected order of their increasing electron affinities would be O > S > Se. However, this is not the observed case as the oxygen atom is extremely small. So, oxygen cannot hold a negative charge in such a small space. Therefore, the electron affinity of oxygen is the lowest in the group. 

SiO₂ is slightly acidic, CaO is basic, and CO₂ is acidic. SnO₂ is amorphous.

The ability of an atom in a molecule to draw electrons toward it is known as electronegativity, according to Pauling. When we examine the periodic table's trend, we can observe that while electronegativity increases between periods, it drops as we move down the groups. Pauling estimated the differences in electronegativity between the atoms in a covalent bond based on the bond energies and assigned the most electronegative element, fluorine, a value of 4, and other elements were calculated concerning that value. So, on Pauling's scale, oxygen is the element following fluorine.

The atomic size increases down the periodic table. Due to a rise in atomic size, a set of elements in the periodic table have ionization enthalpies that drop from top to bottom.

Their charge density, or the ratio of charge to size, is very similar. Even though they are grouped differently due to a diagonal relationship, lithium and magnesium have remarkably comparable chemical properties.

Atomic radii increase as you move down in a group as new shells are added one at a time, which is why O < S < Se. Additionally, As is in group 15 and has one fewer electron in its p orbital than group 16 elements, giving it a bigger atomic radius. Thus, O < S < Se < As.

The chemical elements are arranged in a periodic table according to their atomic number (the number of protons in the nucleus), electron configurations, and characteristic chemical properties. A period in the Modern Periodic Table is indicated by the value of the primary quantum number (n) for the outermost shell or the valence shell.

A trait of elements in the f-block is the decrease in atomic size with the increase in atomic number. Both lanthanoid and actinoid contractions refer to it. This is brought on by the f subshell's electrons' inadequate shielding.

The typical oxides of elements become more acidic as you move from left to right in a period as their electronegativity rises.

The term "representative elements" refers to substances with complete inner shells but incomplete outer shells, i.e., substances with fewer than 8 electrons in the outermost shell. Except for inert gas, the s and p block elements are referred to as representative elements.