Chapter 8.1 – Periodic Table | Chapter Solution Class 10

At the time of the publication of Mendeleev’s periodic table in 1869, a total of 63 elements were known.

Dobereiner’s Law of Triads states that the atomic weight of the middle element of a triad is approximately equal to the arithmetic mean of the atomic weights of the other two elements in the triad.

Mendeleev’s periodic law states that the physical and chemical properties of elements are the periodic functions of their atomic weights.

Inert gas elements were discovered after the publication of Mendeleev’s periodic table. Initially, Mendeleev’s table had vacant places, and noble gases were later accommodated in a separate group 18 (zero group).

• Subgroup A and B elements of the same group exhibit some similarities in properties.
• The only major similarity is that some of their valencies are identical.
• Example: Alkali metals (1A) exhibit a valency of 1, and coinage metals (IB) also exhibit a valency of 1 in many of their compounds.

However, subgroup B elements often differ in other chemical properties compared to subgroup A.

Reasons are:

1. Mendeleev placed potassium (K, atomic mass 39) after argon (Ar, atomic mass 40) despite its lower atomic mass.
2. This was done to keep elements with similar chemical properties in the same group.
3. Potassium is an alkali metal and has properties similar to lithium and sodium, so it was placed in Group 1.
4. Argon, being a noble gas, was placed in Group 18.
5. This showed that atomic weight alone was not the sole criterion for arranging elements; properties were given more importance.

Moseley’s periodic law states that the physical and chemical properties of elements are the periodic functions of their atomic number.

• Moseley’s X-ray experiments proved that atomic number (not atomic weight) is the fundamental property of an element.
• The modern periodic table is based on atomic number instead of atomic weight.

Alkaline earth metals are placed in Group 2 (IIA) of the periodic table.

The term halogen comes from the Greek word “halos”, meaning “salt-forming”.

They are placed in Group 17 (VIIA) of the periodic table.

• Examples: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At).

Noble gases are Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), and Radon (Rn).

They are called inert gases because:

• They have completely filled valence shells, making them chemically stable.
• They do not easily react with other elements under normal conditions.

Periodicity in properties refers to the recurrence of similar properties at regular intervals when elements are arranged in the increasing order of their atomic numbers.

This happens because of the repetition of similar electronic configurations at regular intervals.

Radioactivity is not a periodic property.

It does not follow a regular trend in the periodic table.

Some periodic properties of elements include:

1. Atomic radius – Decreases across a period, increases down a group.
2. Ionization energy – Increases across a period, decreases down a group.
3. Electronegativity – Increases across a period, decreases down a group.
4. Oxidizing and Reducing Properties – Oxidizing power increases across a period, reducing power increases down a group.

(d) 118

Explanation:

As of 2021, there are 118 confirmed elements in the periodic table. The most recent elements to be added are element 113, Nihonium (Nh), element 115, Moscovium (Mc), element 117, Tennessine (Ts), and element 118, Oganesson (Og).

(d) 0

Explanation : 

At the time of publication of Mendeleev’s periodic table, no inert gases had been discovered yet. The discovery of the first noble gas, helium, did not occur until 1895, two years after the publication of Mendeleev’s periodic table.

(a) 7

Explanation : 

There are 7 periods in Mendeleev’s periodic table. The table is arranged in rows and each row is called a period.

(b) Chemical Properties

Explanation : 

The similarity between the subgroup A and B elements is with respect to their chemical properties. The elements in the same group, also known as the vertical columns of the periodic table, have similar chemical properties because they have the same number of valence electrons, which determines their reactivity and chemical behaviour.

(a) Atomic Number

Explanation:

The long form of the periodic table is based on the atomic number of the elements. The elements are arranged in order of increasing atomic number, which is the number of protons in the nucleus of an atom.

Dobereiner’s Law of Triads states that the atomic weight of the middle element in a triad is approximately equal to the arithmetic mean of the atomic weights of the other two elements.

Mendeleev’s periodic law states that the physical and chemical properties of elements are the periodic functions of their atomic weights.

Inert gas elements (noble gases) are placed in Group 18 (Zero Group) because they have completely filled valence shells (outermost electron shells).

Moseley’s periodic law states that the physical and chemical properties of elements are the periodic functions of their atomic number.

Alkali metals (Group 1A) and alkaline earth metals (Group 2A) are strong reducing agents.

1. The second period of the periodic table contains 8 elements.
2. Halogens are strong oxidizing agents.
3. Name of a group 13 element is boron.
4. Along a period from left to right atomic radii decrease.
5. Name of a transuranic element is neptunium.

1. The property which is not periodic in nature is radioactivity.
2. The covalent radius of atoms is less than Vanderwaal’s radius.
3. There are 18 groups in the long form of the periodic table.
4. An example of a light metal is K (potassium).
5. Group and period numbers of ₁₉K are 1 and 4.

Answer

The periodicity of properties of elements refers to the recurring trends and patterns in the chemical and physical properties of the elements as they are arranged in the periodic table.

Moseley’s experiment using X-ray spectra showed that the atomic number (not atomic weight) is the fundamental property of an element. This led to the formulation of Moseley’s periodic law, which states:

Alkali metals (Li, Na, K, Rb, Cs, Fr) are placed in Group 1A because they have one electron in their outermost shell and exhibit similar chemical properties.

They are called “alkali metals” because they react with water to form strong alkaline (basic) solutions (hydroxides) such as NaOH and KOH.

The term halogen comes from the Greek word “halos”, meaning “salt-forming”.

Halogens are highly reactive non-metals found in Group 17 (VIIA) of the periodic table.

Properties of Halogens:

1. They exist as diatomic molecules (F₂, Cl₂, Br₂, I₂).
2. They have seven valence electrons and need one electron to achieve stability.
3. They form salts with metals (e.g., NaCl, KBr).
4. Their reactivity decreases down the group (F > Cl > Br > I).
5. They are strong oxidizing agents.

The covalent radius of an atom is half the distance between the nuclei of two identical atoms bonded by a single covalent bond.

• It is a measure of atomic size in covalently bonded molecules.
• Covalent radius decreases across a period due to increased nuclear attraction.
• It increases down a group as additional electron shells are added.

Fluorine (9F):

• Atomic number = 9
• Electronic configuration = 2, 7
• Group = 17 (VIIA) (Halogens)
• Period = 2 (since it has two electron shells)

Phosphorus (15P):

• Atomic number = 15
• Electronic configuration = 2, 8, 5
• Group = 15 (VA)
• Period = 3 (since it has three electron shells)

Transuranic elements are elements with atomic numbers greater than uranium (92U). These elements are man-made (synthetic) and are produced through nuclear reactions.

They belong to the actinide series in the periodic table.

Examples of Transuranic Elements:

1. Neptunium (Np, atomic number 93)
2. Plutonium (Pu, atomic number 94)
3. Americium (Am, atomic number 95)
4. Curium (Cm, atomic number 96)

Mendeleev’s periodic table was the first systematic arrangement of elements based on their atomic masses.

Key Features of Mendeleev’s Periodic Table:

Periodic Law: Mendeleev stated that the physical and chemical properties of elements are periodic functions of their atomic weights.

Groups and Periods:

• The table consisted of 7 periods (horizontal rows) and 8 groups (vertical columns).
• Groups I to VII were further divided into A and B subgroups.
• Group VIII contained transition elements.

Prediction of New Elements:

• Mendeleev left gaps in his table for undiscovered elements and predicted their properties.
• Examples: Gallium (Eka-Aluminium) and Germanium (Eka-Silicon) were later discovered.

The periodic table consists of different types of elements categorized based on their chemical properties and group numbers.

The ionic radius refers to the size of an ion, which depends on:

• The nuclear charge (proton number).
• The number of electrons removed or gained.

Variation of Ionic Radii Along a Period:

Cations (Positive Ions) Shrink:

• Moving from left to right, metals form cations (positive ions) by losing electrons.
• As electrons are removed, the remaining electrons experience a stronger nuclear attraction, decreasing the ionic radius.
• Example: Na⁺ > Mg²⁺ > Al³⁺ (Ionic size decreases across the period).

Anions (Negative Ions) Expand Initially, Then Decrease:

• Non-metals (right side of the period) gain electrons to form anions (negative ions).
• The addition of electrons increases repulsion, causing anions to be larger than neutral atoms.
• However, moving from left to right, nuclear charge increases, pulling the outer electrons closer, reducing ionic size.
• Example: O²⁻ > F⁻ > N³⁻ (Ionic size decreases across the period).

Thus, cations become smaller, while anions first increase and then decrease in size across a period.

Oxidizing Property (Electron Gain):

• Increases from left to right in a period.
• Non-metals (right side) have high electronegativity and readily gain electrons, making them strong oxidizing agents.

Example:

• Halogens (F, Cl, Br) are the strongest oxidizing agents.
• Fluorine (F) is the strongest oxidizer due to its high electron affinity.

Reducing Property (Electron Loss):

• Decreases from left to right in a period.
• Metals (left side) lose electrons easily, making them strong reducing agents.

Example:

• Alkali metals (Na, K, Rb) are strong reducing agents because they readily donate electrons.
• Sodium (Na) is a stronger reducer than Aluminum (Al).

Atomic radius is the distance from the center of the nucleus to the outermost shell of an atom. It determines the size of an atom and is measured using different methods.

Types of Atomic Radius:

1. Covalent Radius: Half the distance between the nuclei of two covalently bonded atoms.
2. Metallic Radius: Half the distance between the nuclei of two adjacent metal atoms in a metallic structure.
3. Ionic Radius: Radius of cations or anions in an ionic compound.
4. Van der Waals Radius: The radius of non-bonded atoms in a molecule.

In Mendeleev’s periodic table, elements were arranged by atomic weight, so placing potassium (39K) after argon (40Ar) violated the periodic law.

However, this anomaly was resolved in the modern periodic table, where elements are arranged by atomic number, not atomic mass.

Justification in the Modern Periodic Table:

Atomic Number Determines Position:

• Argon (Ar) has atomic number 18, so it is placed in Group 18 (Noble gases).
• Potassium (K) has atomic number 19, so it is placed in Group 1 (Alkali metals).

Chemical Properties Take Precedence:

• Potassium shares properties with alkali metals (Li, Na), so it is placed in Group 1.
• Argon is a noble gas, so it is placed in Group 18.