What is meant by Newland's Law of Octaves?

Newland's Law of Octaves stated that when elements were arranged by increasing atomic weight, every eighth element had similar properties, like musical octaves.

What is Mendeleev's Periodic Law?

Mendeleev's Periodic Law states that the properties of elements are a periodic function of their atomic weights.

What is the Modern Periodic Law?

The Modern Periodic Law states that the properties of elements are a periodic function of their atomic numbers ($Z$).

Explain Moseley's experiment and the results of it.

Moseley bombarded elements with electrons, measuring the emitted X-ray frequencies ($\nu$). He found $\sqrt{\nu} \propto Z$ (atomic number). This showed atomic number, not atomic weight, determines element properties, leading to the Modern Periodic Law.

Write down nomenclature of the below mentioned elements having atomic number 101 to 118.

Explain the periodic properties of Lanthanoid and Actinoid series.

Lanthanoids:

• Electrons fill 4f orbitals.
• Silvery-white metals, tarnish in air.
• High melting and boiling points.
• Mostly +3 oxidation state.
• Form colored ions and complexes.
• Used as catalysts.

Actinoids:

• Electrons fill 5f orbitals.
• Radioactive.
• Early members have multiple oxidation states, later ones mostly +3.
• High density and melting points.
• Form complexes.
• Used in nuclear reactors and weapons.

How would you justify the presence of 18 elements in the 5th period of the periodic table?

The 5th period starts with filling the 5s orbital (2 electrons). Then, the 4d orbitals are filled (10 electrons), followed by the 5p orbitals (6 electrons). This totals 2 + 10 + 6 = 18 electrons, hence 18 elements in the 5th period.

Write down general group electronic configuration of s block, p block, d block and f block?

• s-block: $ns^{1-2}$ where n = 2-7
• p-block: $ns^2np^{1-6}$ where n = 2-7
• d-block: $(n-1)d^{1-10}ns^{0-2}$ where n = 4-7
• f-block: $(n-2)f^{0-14}(n-1)d^{0-1}ns^2$ where n = 6-7

What is meant by transuranium element?

Transuranium elements are elements with atomic numbers greater than 92 (Uranium). They are all synthetic and radioactive.

Arrange the following elements in the increasing order of metallic chracter: Si, Be, Mg, Na, P

P < Si < Be < Mg < Na.

Metallic character increases down a group and decreases across a period. Na and Mg are in the same period with Na being more metallic (left of Mg). Be is above Mg, so Be is less metallic. Si and P are to the right of Mg with P farthest to the right, thus it is least metallic, with Si less metallic than Be.

What is meant by (i) covalent radius, (ii) Metallic Radius, (iii) Van der Waal's radius and (iv) Ionic Radius.

(i) Covalent radius: Half the distance between nuclei of two covalently bonded atoms of the same element in a molecule.

(ii) Metallic Radius: Half the distance between nuclei of two adjacent atoms in a metallic crystal.

(iii) Van der Waals radius: Half the distance between nuclei of two non-bonded atoms of adjacent molecules in the solid state.

(iv) Ionic Radius: The effective distance from the nucleus of an ion to the point up to which it has an influence in the ionic bond.

How atomic radius varies along the period and down the group?

Along the period: Atomic radius decreases. Increased nuclear charge pulls electrons closer.

Down the group: Atomic radius increases. Electrons fill new shells, increasing size.

Why the atomic radii of Noble gases are very high?

Noble gases have their atomic radii measured as van der Waals radii, which are larger than covalent radii. Since they don't form covalent bonds, only weak van der Waals forces exist, resulting in larger interatomic distances and thus, larger atomic radii.

What do you understand by isoelectronic species? Give examples

Isoelectronic species are atoms or ions that have the same number of electrons.

Examples:

• $N^{3-}, O^{2-}, F^-, Ne, Na^+, Mg^{2+}, Al^{3+}$ (all have 10 electrons)
• $S^{2-}, Cl^-, Ar, K^+, Ca^{2+}$ (all have 18 electrons)

Comparing the Ionic Radii of O^2-, F^-, Na⁺, and Mg²⁺

O^2- > F^- > Na⁺ > Mg²⁺.

All are isoelectronic (10 electrons). Ionic radius decreases with increasing nuclear charge (greater positive charge or smaller negative charge).

Which of the following species will have the largest and smallest size? [Mg, Mg2+, Al, Al2+]

Largest: Mg

Smallest: Al³⁺

Neutral atoms are larger than their cations. Mg > Al, and greater positive charge means smaller size: Mg²⁺ > Al³⁺, Mg > Mg²⁺ and Al > Al³⁺.

What is meant by Ionisation Enthalpy?

Ionisation enthalpy is the energy required to remove an electron from an isolated gaseous atom in its ground state, forming a cation.

$A(g) \rightarrow A^+(g) + e^-$

Why second I.E. is higher than that of First I.E.?

Second I.E. is higher because it's harder to remove an electron from a positively charged ion than from a neutral atom. The increased nuclear attraction makes it more difficult.

Why Noble Gas elements have the highest IE along the same periods?

Noble gases have the highest I.E. because of their stable, filled electron shells ($ns^2np^6$). This configuration requires a large amount of energy to remove an electron.

Briefly explain the group trend and periodic trend of I.E.

Group Trend: I.E. decreases down a group. Increased atomic size and shielding reduce the attraction between the nucleus and outer electrons.

Periodic Trend: I.E. generally increases across a period. Increased nuclear charge and smaller atomic size result in a stronger attraction, requiring more energy to remove an electron.

I.E. depends on which factors?

I.E. depends on:

1. Atomic size: Larger atoms, lower I.E.
2. Nuclear charge: Higher charge, higher I.E.
3. Shielding effect: More shielding, lower I.E.
4. Electronic configuration: Stable configurations (half-filled, fully-filled), higher I.E.

What is meant by Shielding or Screening effect? Explain with an example.

The shielding effect is the reduction in the effective nuclear charge on the electron cloud, due to the difference in the attraction forces on the electrons in the atom. Inner electrons repel outer electrons, shielding them from the full nuclear charge.

Example: In potassium (K), the 19th electron is shielded by the 18 inner electrons, reducing the effective nuclear charge felt by the outermost electron.

Why I.E.1 of B<Be?

Boron's electron is removed from a 2p orbital, while Beryllium's is removed from a 2s orbital. 2s electrons penetrate closer to the nucleus than 2p electrons, are more tightly held, and thus Be has a higher I.E.1. Also, Be has a filled 2s orbital (stable configuration).

Why I.E.1 of O< N?

Oxygen's electron is removed from a doubly occupied 2p orbital, while nitrogen's is removed from a half-filled 2p orbital. The electron-electron repulsion in the doubly occupied orbital makes it easier to remove an electron from oxygen, thus O has a lower I.E.1. Also, N has a stable half-filled configuration.

What is meant by Electron Gain Enthalpy? Explain with an element X. What is its Unit?

Electron gain enthalpy is the energy change when an electron is added to a neutral gaseous atom, forming a negative ion.

For element X: $X(g) + e^- \rightarrow X^-(g)$

Unit: kJ mol^-1 or eV/atom.

Why noble gages have very large positive electron gain enthalpy, whereas Group 17 elements have very high negative?

Noble gases have a stable, filled-shell electron configuration. Adding an electron requires a large amount of energy (positive electron gain enthalpy) as it disrupts this stability.

Group 17 elements have one electron less to attain a stable, noble gas configuration. Adding an electron releases a large amount of energy (high negative electron gain enthalpy) because of increased stability.

Briefly explain group trend and periodic trend of electron gain enthalpy?

Group Trend: Electron gain enthalpy generally becomes less negative down a group. Increased atomic size reduces the attraction for the added electron.

Periodic Trend: Electron gain enthalpy generally becomes more negative across a period. Increased nuclear charge and smaller atomic size increase the attraction for the added electron.

Why is the Electron Gain Enthaly of O less than S and F less than Cl?

Due to their small size, O and F experience greater electron-electron repulsion when adding an electron compared to the larger S and Cl atoms. This increased repulsion makes the electron gain enthalpy of O less negative than S, and F less negative than Cl.

What is electronegativity?

Electronegativity is the tendency of an atom to attract shared electrons towards itself in a chemical bond.

How many numerical scales are there to measure to electronegativity of different elements?

There are several scales, but the most common are:

1. Pauling scale
2. Mulliken scale
3. Allred-Rochow scale

Briefly explain group trend and periodic trend of electronegativity?

Group Trend: Electronegativity decreases down a group. Increased atomic size and shielding reduce the attraction for shared electrons.

Periodic Trend: Electronegativity increases across a period. Increased nuclear charge and smaller atomic size increase the attraction for shared electrons.

What is meant by diagonal relationship? Briefly explain with examples.

Diagonal relationship is the similarity in properties between certain diagonally adjacent elements in the second and third periods.

Examples:

• Li and Mg: Similar size, polarizing power, form normal oxides, form nitrides.
• Be and Al: Amphoteric hydroxides, form covalent compounds, form complex ions.
• B and Si: Similar electronegativity, form acidic oxides, form covalent hydrides.

Briefly explain some of the anomalous properties of second period elements. (Include diagonal relationship)

Second-period elements show anomalous properties due to:

1. Small size: Higher ionization enthalpy, more covalent character.
2. High electronegativity: Formation of stronger bonds.
3. Absence of d-orbitals: Limited covalency, inability to form certain complexes.
4. Ability to form pπ-pπ multiple bonds: C=C, C≡C, N=N unlike heavier congeners.

These factors also contribute to the diagonal relationship, where they resemble the third-period elements diagonally across from them. For example, Li resembles Mg due to comparable size and polarizing power, and Be resembles Al due to similar charge/radius ratio.

On the basis of QM calculations, justify that the sixth period should have 32 elements.

The sixth period begins with filling the 6s orbital (2 electrons), followed by the 4f orbitals (14 electrons), then the 5d orbitals (10 electrons), and finally the 6p orbitals (6 electrons). This totals 2 + 14 + 10 + 6 = 32 electrons, thus the sixth period should have 32 elements based on the filling of these orbitals as determined by quantum mechanical calculations.

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Written by Kaavje Sahé