Periodic Table & Trends - Free MCQ Practice Test with solutions, NEET Chemistry

Lanthanides, Ce(Z = 58) ñ Lu(Z = 71) and Actinoids comprises of f block elements.

Elements characterized by the filling of 4f orbitals are known as lanthanides or lanthanoids. These elements belong to the f-block of the periodic table and include:

• Cerium (Ce)
• Praseodymium (Pr)
• Neodymium (Nd)
• Promethium (Pm)
• Samarium (Sm)
• Europium (Eu)
• Gadolinium (Gd)
• Terbium (Tb)
• Dysprosium (Dy)
• Holmium (Ho)
• Erbium (Er)
• Thulium (Tm)
• Ytterbium (Yb)
• Lutetium (Lu)

The term 'lanthanides' specifically refers to this series, ensuring clarity in understanding their position in the periodic table.

group is 15 so total valence electrons will be 5 while period is 3rd so n value of valence shell will be 3.

Modern periodic law is essentially the result of periodic variation and electronic configuration.

Metalloids exhibit properties that are intermediate between metals and nonmetals. They share some characteristics with metals, such as:

• Being solid at room temperature
• Having a metallic appearance

However, they also exhibit nonmetallic properties, including:

• Poor electrical conductivity compared to metals

Examples of metalloids include:

• Boron
• Silicon
• Germanium

These elements display this dual nature in their physical and chemical behaviours.

Among the alkali metals cesium is the most reactive because the outermost electron is more loosely bound than the outermost electron of the other alkali metals.

In XeF₂O₃, xenon forms a total of five sigma (σ) bonds – two bonds with fluorine atoms and three bonds with oxygen atoms. Moreover, three of the Xe–O bonds also form π bonds due to the effective overlap between xenon’s d orbitals and oxygen’s p orbitals. 

In the modern periodic table, the horizontal rows are known as periods, and the vertical columns are referred to as groups.

This structure organises elements based on their:

• Atomic number
• Chemical properties

In this context:

• Periods represent the principal quantum number.
• Groups indicate similar valence electron configurations.

Alkali metals typically form ionic bonds with non-metals like oxygen, resulting in ionic compounds such as oxides (e.g., Na₂O). They do not generally form covalent bonds with oxygen.

Therefore, statement A is false because alkali metals primarily form ionic bonds when reacting with oxygen, not covalent ones.

The correct order of decreasing radius is:

• Van der Waals' radius > Metallic radius > Covalent radius.

Explanation:

• Van der Waals' Radius: This is the largest as it includes the distance between two atoms when they are in contact, encompassing both covalent or metallic radii and additional space due to electron clouds.
• Metallic Radius: Smaller than Van der Waals', it measures the radius of an atom in a metallic crystal structure, where electrons are delocalised but not as tightly packed as in covalent bonds.
• Covalent Radius: The smallest, as it measures the distance from the nucleus to the bonding electron pair shared with another atom in a covalent bond.

Thus, the correct order is Van der Waals' radius > Metallic radius > Covalent radius.

Density: Li < Na < K < Rb.
The density of K is lower than that of Na. Thus, option D is incorrect. The correct trend is
Li < K < Na < Rb

Dmitri Mendeleev and Lothar Meyer proposed arranging elements in 1. increasing order of their atomic weights. This arrangement eventually led to the development of the periodic table of elements.

The groups in the modern periodic table are numbered from 1 to 18 according to IUPAC recommendations. This includes:

• Group 1: Alkali metals
• Group 18: Noble gases

as we move from left to right in a period atomic radii decreases and due to effective nuclear charge outermost electrons are held more closely to the nucleus.

1. N₂​: Triple bond → Shortest bond length.
2. O₂: Double bond → Shorter bond length than single bonds.
3. F₂: Single bond, but shorter than Cl2 because fluorine atoms are smaller.
4. Cl₂​: Single bond → Longest bond length due to larger atomic size.

Correct Order (Increasing Bond Length): 

The stability of an oxidation state depends on electronic configuration after ionization.

For transition metals, when forming +2 oxidation state, they lose two 4s electrons first.

• Cr (Z = 24): Ground state = [Ar] 3d⁵ 4s¹
  Cr²⁺ = [Ar] 3d⁴ → unstable (not half-filled)

• Mn (Z = 25): Ground state = [Ar] 3d⁵ 4s²
  Mn²⁺ = [Ar] 3d⁵ → half-filled d⁵ (most stable)

• Fe (Z = 26): Ground state = [Ar] 3d⁶ 4s²
  Fe²⁺ = [Ar] 3d⁶ → fairly stable

• Co (Z = 27): Ground state = [Ar] 3d⁷ 4s²
  Co²⁺ = [Ar] 3d⁷ → less stable compared to Fe²⁺

When a transition metal ion is in highest oxidation state it tends to lower its oxidation state by gaining electrons. thus it itself reduces and oxidizes the substance from which it accepts electron. So, it acts as oxidising agent. 

Smaller the size of cation, higher will be the hydration and its effective size will increase and hence mobility in aqueous solution will decrease. Larger size ions have more ionic mobility due to less hydration. Thus the degree of hydration of M⁺ ions decreases from Li⁺ to Cs⁺. Consequently the radii of the hydrated ion decreases from Li⁺ to Cs⁺. Hence the ionic conductance of these hydrated ions increases from Li⁺ to Cs⁺ 

electronegativity increases on moving from left to right while decreases as we move down the group.

Sodium (Na) has a lower ionisation energy than Magnesium (Mg). The reasons for this are as follows:

• The removal of an electron from Na requires less energy.
• This is due to Na's lower effective nuclear charge.
• Generally, ionisation energy increases across a period.

Therefore, this trend indicates that Na is not an exception to the rule.

Cl just need one more electrons to get fully filled electronic configuration so it has most negative electron gain enthalpy while P has d3 configuration which is already stable so it cannot accept one electron extra easily so have large positive electron gain enthalpy.

sept is for 7.