Periodic Table Trends (original) (raw)

Last Updated : 16 May, 2026

The periodic table is arranged according to increasing atomic number, and as we move across a period or down a group, the properties of elements change in a regular pattern. These regular and gradual changes in properties are called periodic trends. Periodic trends help us understand how and why elements differ in size, reactivity, metallic nature, and other chemical properties.

• There were various failed attempts to arrange the known elements so that all the elements with similar properties are arranged together.
• But in 1869, Russian chemist Dmitri Mendeleev gave Mendeleev's Periodic Table which is the most successful attempt of that time. In modern-day, scientists using Mendeleev's Periodic Table as a base created a long-form periodic table, which is called the Modern Periodic Table.
• In this table, Periods are made up of elements that are organised horizontally (from left to right), and the group is made up of elements that are arranged vertically (from top to bottom).

**1. Atomic Radius

Atomic radius is the distance between the centre of the nucleus and the outermost shell of an atom. It tells us the size of an atom.

**a) Trend Across a Period (Left to Right →)

Atomic radius decreases as we move from left to right in a period because:

• The number of protons increases.
• Nuclear charge increases.
• Electrons are added to the same shell.
• A greater attraction pulls electrons closer to the nucleus.

**Example:

Li > Be > B > C > N > O > F

Lithium is largest, Fluorine is smallest in Period 2.

**b) Trend Down a Group (Top to Bottom ↓)

Atomic radius increases as we move down a group because:

• New shells are added.
• Distance between nucleus and outermost electrons increases.
• Shielding effect increases.

**Example:

Li < Na < K < Rb < Cs

Cesium is the largest.

**Exception in Atomic Radius

• In transition elements (d-block), size does not decrease regularly because of poor shielding effect of d-electrons.
• Noble gases have larger atomic radii than expected because their radius is measured as van der Waals radius.

2. Ionisation Enthalpy

Ionisation enthalpy is the energy required to remove the outermost electron from an isolated gaseous atom.

M → M⁺ + e⁻

**a) Trend Across a Period (Left to Right →)

Ionisation enthalpy increases across a period because:

• Atomic size decreases.
• Nuclear charge increases.
• Electrons are held more tightly.
• More energy is required to remove electron.

**Example:

Li < Be < B < C < N < O < F

Fluorine has high ionisation enthalpy.

**b) Trend Down a Group (Top to Bottom ↓)

Ionisation enthalpy decreases down the group because:

• Atomic size increases.
• Outer electrons are farther from nucleus.
• Attraction decreases.
• Less energy needed to remove electron.

**Example:

Li > Na > K > Rb > Cs

Cesium has lowest ionisation enthalpy.

**Exceptions in Ionization Enthalpy

**1. Beryllium and Boron (Be & B)

Be has higher ionisation enthalpy than B.

Reason:
Be = 1s² 2s² (stable filled s-subshell)
B = 1s² 2s² 2p¹ (electron in p-orbital is easier to remove)

So, IE(Be) > IE(B)

**2. Nitrogen and Oxygen (N & O)

N has higher ionisation enthalpy than O.

Reason:
N = 2p³ (half-filled stable configuration)
O = 2p⁴ (one paired electron, more repulsion)

So, IE(N) > IE(O)

**3. Valence Electrons

The electrons in an atom's outermost shell are referred to as the atom's valence electrons. Valence electrons are the electrons present in the outermost shell of an atom.

**a) Trend Across a Period (Left → Right)

Number of valence electrons increases from 1 to 8 as we move across a period because:

• Atomic number increases.
• Electrons are added to the same outer shell.
• So valence electrons increase one by one.

**Example:

Li → 1
Be → 2
B → 3
C → 4
N → 5
O → 6
F → 7
Ne → 8

So, valence electrons increase from 1 to 8.

**b) Trend Down a Group (Top ↓ Bottom)

Number of valence electrons remains the same down a group because:

• Elements in the same group have similar outer electronic configuration.

**Example:

Li = 1
Na = 1
K = 1
Rb = 1

All have 1 valence electron.

**4. Valency

Valency is the combining capacity of an atom.

**a) Trend Across a Period (Left → Right)

Valency first increases from 1 to 4, then decreases from 4 to 0 because:

Atoms try to complete their octet (8 electrons).

**Example:

Li → 1
Be → 2
B → 3
C → 4
N → 3
O → 2
F → 1
Ne → 0

So pattern is:
1 → 2 → 3 → 4 → 3 → 2 → 1 → 0

**b) Trend Down a Group (Top ↓ Bottom)

Valency remains the same down a group.

**Example:

Group 1 → All have valency 1
Group 2 → All have valency 2
Group 17 → All have valency 1

**Exception in Valency

In transition elements (d-block), valency can vary.

**Example:

• Iron (Fe) → Valency 2 and 3
• Copper (Cu) → Valency 1 and 2

This is called variable valency.

**5. Melting and Boiling point

• Melting point is the temperature at which a solid changes into liquid.
• Boiling point is the temperature at which a liquid changes into gas.

**a) Trend Across a Period (Left → Right)

Melting and boiling points first increase, then decrease across a period because:

**From Group 1 to Group 14:
Melting point generally increases because bonding becomes stronger (metallic → covalent network).

**After Group 14:
Melting point decreases because elements form simple molecules with weak intermolecular forces.

**Example:

Li < Be < B < C → Increasing
(C has very high melting point due to strong covalent network)

After Carbon:
N < O < F < Ne → Decreasing

Neon has very low melting and boiling point because it is a noble gas with weak forces.

**b) Trend Down a Group (Top ↓ Bottom)

The trend depends on the type of element.

**In Metals (Group 1 & 2)

• Melting and boiling points generally decrease down the group.
• Atomic size increases.
• Metallic bonding becomes weaker.

**Example :

Li > Na > K > Rb > Cs

Lithium has highest melting point in Group 1.

In Non-metals (Like Halogens)

• Melting and boiling points generally increase down the group.
• Molecular size increases.
• Intermolecular forces become stronger.

**Example:

F₂ < Cl₂ < Br₂ < I₂

Iodine has higher melting and boiling point than fluorine.

**Exceptions in Melting and Boiling point

• Carbon has extremely high melting point because of strong covalent network bonding.
• In transition elements, melting points are very high due to strong metallic bonding.
• Noble gases have very low melting and boiling points because they have very weak intermolecular forces.

**Example: Iron has high melting point.

6. Metallic and Non Metallic Character

• Metallic character is the tendency of an atom to lose electrons and form positive ions (cations).
• Non-metallic character is the tendency of an atom to gain electrons and form negative ions (anions).

**a) Trend Across a Period (Left → Right)

Metallic character decreases across a period and Non-metallic character increases across a period because:

• Atomic size decreases.
• Nuclear charge increases.
• Ionisation energy increases.
• Atoms find it harder to lose electrons.
• Instead, they tend to gain electrons.

**Example:

Na → Mg → Al → Si → P → S → Cl

• Sodium (Na) is highly metallic.
• Chlorine (Cl) is highly non-metallic.

So, metallic nature decreases and non-metallic nature increases from Na to Cl.

**b) Trend Down a Group (Top ↓ Bottom)

Metallic character increases down a group and Non-metallic character decreases down a group because:

• Atomic size increases.
• Ionisation energy decreases.
• It becomes easier to lose electrons.

**Example:

Li < Na < K < Rb < Cs

Cesium (Cs) is more metallic than Lithium (Li).

Example (Group 17):

F > Cl > Br > I

Fluorine is more non-metallic than iodine.

**Exception in Metallic and Non Metallic Character

These exceptions occur due to special reasons such as small atomic size, unusual electronic configuration, metalloids nature, and the presence of transition elements. Because of these factors, certain elements show properties that are slightly different from the expected periodic trend. These are:

• Hydrogen is placed in Group 1 but it is a non-metal.
• Some elements like silicon and germanium show both metallic and non-metallic properties. These are called metalloids.
• Transition elements show moderate metallic character and do not follow simple trends.

7. Electronegativity

Electronegativity is the tendency of an atom to attract shared electrons towards itself in a chemical bond. The higher the electronegativity, the stronger the atom attracts electrons.

**a) Trend Across a Period (Left → Right)

Electronegativity increases across a period because:

• Atomic size decreases.
• Nuclear charge increases.
• Attraction between nucleus and bonding electrons increases.
• So atoms pull electrons more strongly.

**Example :

Li < Be < B < C < N < O < F

Fluorine has the highest electronegativity in the periodic table.

**b) Trend Down a Group (Top ↓ Bottom)

Electronegativity decreases down a group because:

• Atomic size increases.
• Distance between nucleus and valence electrons increases.
• Shielding effect increases.
• Attraction for bonding electrons decreases.

**Example:

F > Cl > Br > I

Fluorine is more electronegative than iodine.

List of All Periodic Trends

The following table is the summary of all the periodic trends in the properties of different elements.

**Related Articles:

• Periodic Classification of Elements
• Merits of the Modern Periodic Table
• Characteristics of the Periods and Groups of the Periodic Table