Free Study Material for ICSE Chemistry- Grade 10- “Periodic Properties and Variations of Properties

Understanding the Modern Periodic Table: Trends and Concepts Simplified

The periodic table is one of the most important tools in chemistry, helping us predict and understand the properties of elements based on their atomic structure. Over time, our understanding has evolved — from Mendeleev’s arrangement to the modern version used today.

🧪 From Mendeleev to Modern: The Evolution of Periodic Law

Mendeleev’s Periodic Law:“The physical and chemical properties of elements are periodic functions of their atomic weights.”
Modern Periodic Law:“The physical and chemical properties of elements are periodic functions of their atomic numbers.”

🔁 Why Periodicity Occurs:
It’s due to the repetition of similar electronic configurations, especially the valence electrons (electrons in the outermost shell), which influence how elements react chemically.

📊 Structure of the Long Form Periodic Table

🔹 Vertical Columns – Groups (18 Total)

Elements in the same group share similar valence electron configurations.
Group 1 and 2 (left side): Metals
Groups 13 to 17 (right side): Non-metals
Group 18: Noble gases (Inert gases – unreactive)
Groups 3 to 12: Transition elements – show intermediate properties between metals and non-metals.

🔸 Horizontal Rows – Periods (7 Total)

Each period corresponds to a principal energy level (or shell).
As you move from left to right in a period, electrons are added to the same shell.

📈 Periodic Trends Explained

1️⃣ Atomic Radius

➕ Increases down a group (due to addition of new shells)
➖ Decreases across a period (due to stronger nuclear attraction)

Isoelectronic Series Tip:
In atoms or ions with the same number of electrons, more protons mean smaller size.

Cations are smaller than their parent atoms
Anions are larger than their parent atoms

2️⃣ Metallic and Non-Metallic Character

In a period:

Metallic character decreases
Non-metallic character increases

In a group:

Metallic character increases
Non-metallic character decreases

3️⃣ Ionization Energy / Enthalpy / Potential

Definition:
The minimum energy required to remove the most loosely held electron from an isolated gaseous atom to form a positive ion.

🔼 Increases across a period (due to smaller atomic size and higher nuclear charge)
🔽 Decreases down a group (due to larger atomic size and shielding effect)

📌 Successive Ionization Energies:
Always increase: IE₁ < IE₂ < IE₃

Factors affecting ionization energy:

Atomic number
Atomic size
Electron configuration

4️⃣ Electron Gain Enthalpy (Electron Affinity)

Definition:
The energy released when a gaseous atom gains an electron to form a negative ion.

🔼 Becomes more negative across a period
🔽 Becomes less negative down a group

5️⃣ Electronegativity

Definition:
The ability of an atom to attract shared electrons in a chemical bond.

🔼 Increases left to right across a period
🔽 Decreases top to bottom in a group

Group 1: The Alkali Metals – Trends, Reactions, and Properties

The alkali metals—Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), and Cesium (Cs)—are located in Group 1 of the periodic table. These metals are known for their high reactivity, softness, and unique flame coloration. Let’s explore the fascinating trends and chemical behavior of these elements.

🔸 1. Atomic and Ionic Radii

Trend: Both atomic and ionic radii increase as we move down the group from lithium to cesium.
Why? Due to the addition of new electron shells, the size of the atom and its corresponding ion grows larger.

🔸 2. Softness

Alkali metals are very soft and can often be cut with a knife.
Trend: Softness increases down the group. Lithium is relatively harder, while cesium is extremely soft.

🔸 3. Melting and Boiling Points

These metals have low melting and boiling points, which decrease down the group.
This is due to weak cohesive forces between the atoms in their metallic lattice.

🔸 4. Flame Colouration

Alkali metals impart distinct colours to flames due to their low ionization energies, which allow outer electrons to be easily excited.

Element	Flame Colour
Lithium	Crimson Red
Sodium	Golden Yellow
Potassium	Pale Violet
Cesium	Blue

🔸 5. Reactivity

Alkali metals are highly reactive, especially with water and air.
Reactivity increases from Li to Cs, due to decreasing ionization energy.

🔸 6. Hydrides Formation

Alkali metals react with dry hydrogen gas to form ionic metallic hydrides:

LiH + H₂O → LiOH + H₂  
NaH + H₂O → NaOH + H₂

These hydrides are strong reducing agents, and their reducing power increases down the group.

🔸 7. Oxides and Hydroxides

Alkali metals react with oxygen to form basic oxides.
Their hydroxides (MOH) are:

Strong bases
Highly soluble in water
Thermally stable
Exception: Lithium hydroxide is less soluble and less basic than others.

🔸 8. Reaction with Water

Alkali metals react vigorously with water, liberating hydrogen gas and forming alkaline solutions:

2Na + 2H₂O → 2NaOH + H₂

They also react with hydrogen halides (HX) in a similar fashion:

2Na + 2HX → 2NaX + H₂

(X = halogen)

🔸 9. Reaction with Halogens

Alkali metals react directly with halogens to form ionic metal halides:

M + X → MX

These halides are crystalline solids and follow the general formula M⁺X⁻.

🔸 10. Reaction with Non-Metals

When heated, alkali metals react with non-metals such as sulfur and phosphorus to form sulfides and phosphides:

2M + S → M₂S  
3M + P → M₃P

Group 17 Elements – The Halogen Family

The halogens—Fluorine, Chlorine, Bromine, and Iodine—are some of the most reactive non-metals on the periodic table. Belonging to Group 17, these elements display unique trends in physical and chemical properties that become more pronounced down the group.

1. Characteristic Colours

Each halogen has a distinct appearance:

Fluorine (F₂): Pale yellow gas
Chlorine (Cl₂): Greenish-yellow gas
Bromine (Br₂): Dark red liquid
Iodine (I₂): Violet solid

These vivid colours reflect changes in molecular structure and electronic transitions as we move down the group.

2. Nature of Bonding

With Metals: Halogens form ionic halides by gaining one electron (e.g., NaCl, CaF₂).
With Non-Metals: They form covalent compounds by sharing electrons (e.g., HCl, CCl₄).

3. Ionisation Energy (IE)

Halogens exhibit very high first ionisation energies (IE₁) due to their small atomic size and high nuclear charge.
Trend: IE₁ values decrease down the group as atomic size increases and shielding effect reduces nuclear attraction.

4. Electronegativity

These elements are highly electronegative, with fluorine being the most electronegative element in the periodic table.
Trend: Electronegativity decreases down the group:

Fluorine: 4.0
Chlorine: 3.0
Bromine: 2.9
Iodine: 2.7

5. Electron Affinity (EA)

Halogens show high electron affinities, reflecting their strong tendency to gain electrons.
Interestingly, chlorine has a slightly higher EA than fluorine, due to smaller electron repulsion in its outer shell.
Order of EA₁: Cl > F > Br > I

6. Hydrides of Halogens

All halogens form volatile hydrides with the general formula HX:

Examples: HF, HCl, HBr, HI
Acid Strength: The acidic character increases from fluorine to iodine:
HI > HBr > HCl > HF

7. Oxidizing Power

Halogens are strong oxidizing agents, capable of extracting electrons from other species.
Trend: Oxidizing strength decreases down the group:

Fluorine is the most powerful oxidizer.
Iodine is the weakest among the halogens.

Understanding these periodic trends helps explain the reactivity, industrial uses, and biological roles of halogens. Whether you're studying for an exam or exploring chemical behavior, Group 17 offers a fascinating glimpse into the power of periodicity.

Free Study Material for ICSE Chemistry- Grade 10- “Periodic Properties and Variations of Properties – Physical & Chemical”