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The classification of elements has played a significant role in the advancement of chemistry. The periodic arrangement of elements is a logical outcome of their electronic configurations. In this unit, we explore the historical development of the Periodic Table and the formulation of the Modern Periodic Law. This blog provides you with summarised notes from NCERT Class 11 Chemistry (Part I), Chapter 3: Classification of Elements and Periodicity in Properties.
Contents
- 1 Why Do We Need to Classify Elements?
- 2 Genesis of Periodic Classification
- 3 Modern Periodic Law and the Present Form of the Periodic Table
- 4 Electronic Configurations and the Periodic Table
- 5 Types of Elements: s-, p-, d-, and f-Blocks
- 6 Metals, Non-Metals, and Metalloids
- 7 Periodic Trends in Properties of Elements
- 8 Valency and Oxidation States
- 9 Anomalous Behaviour of Second Period Elements
- 10 Periodic Trends and Chemical Reactivity
- 11 Important Formulas in NCERT Notes Chemistry (Part-I) Chapter 3: Classification of Elements and Periodicity in Properties
- 12 FAQs
Explore Notes of Class 11 Chemistry
Why Do We Need to Classify Elements?
Chemistry deals with the study of elements and their compounds. As the number of known elements increased over time, it became difficult to study each one separately.
Classification provides a systematic way to:
- Understand the properties of elements.
- Predict the properties of unknown elements.
- Organise a large amount of chemical information.
Historical Perspective
- In 1800, only 31 elements were known.
- By 1865, the number had increased to 63.
- As of today, 114 elements are known (including synthetic ones).
- Studying the chemistry of each element individually became complex and impractical.
Need for a Systematic Arrangement
- To manage the growing list of elements, scientists developed ways to classify elements.
- Classification aimed to:
- Group elements with similar properties together.
- Make it easier to study trends in physical and chemical behavior.
- Provide a rational and predictive framework for chemistry.
Foundation of the Periodic Table
- Classification efforts eventually led to the development of the Periodic Table.
- The periodic table reflects:
- Trends in properties.
- Similarities among groups or families of elements.
- It highlights that elements are not randomly distributed, but follow periodic patterns based on atomic properties.
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Genesis of Periodic Classification
Scientists attempted to classify elements based on similar physical and chemical properties to organize the growing number of elements systematically. This historical development laid the foundation for the periodic law and the periodic table.
Dobereiner’s Triads (1829)
Introduced by Johann Dobereiner, a German chemist. He grouped elements into triads – sets of three elements with similar properties.
Key Observations:
- The atomic weight of the middle element was approximately the average of the other two.
- The properties of the middle element were intermediate.
Limitation:
- Worked for only a few elements. It could not be extended to all known elements.
- Dismissed as a coincidence.
Newlands’ Law of Octaves (1865)
Proposed by John Newlands, an English chemist. Elements are arranged in increasing order of atomic weight. Noted that every eighth element had properties similar to the first, like musical octaves
Limitation:
- Valid only up to calcium.
- Failed for heavier elements.
- Not well accepted initially, but recognized later (Newlands received Davy Medal in 1887).
Mendeleev’s Periodic Law (1869)
Developed by Dmitri Mendeleev, a Russian chemist. Proposed the Periodic Law:
“The properties of elements are a periodic function of their atomic weights.”
Mendeleev arranged elements in a tabular form with:
- Horizontal rows (periods).
- Vertical columns (groups) with similar properties.
Key Contributions:
- Used broad physical and chemical properties (e.g., oxides, chlorides) to group elements.
- Prioritized similar properties over atomic weight when needed.
- Left gaps in the table for undiscovered elements and predicted their properties.
Mendeleev’s Predictions
- Mendeleev predicted new elements and left gaps in the table (e.g., Eka-aluminium and Eka-silicon).
- These elements were discovered later as Gallium (Ga) and Germanium (Ge).
Significance:
- Mendeleev’s success in predicting properties of unknown elements validated his Periodic Table.
- His work laid the foundation for the modern Periodic Table, later refined using atomic numbers.
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Modern Periodic Law and the Present Form of the Periodic Table
Mendeleev classified elements based on increasing atomic weight, but some elements did not fit properly. Examples: Iodine (I) is placed after Tellurium (Te) despite a lower atomic weight.
Discovery of the Atomic Number as a Fundamental Property
In 1913, Henry Moseley, an English physicist, conducted experiments with X-ray spectra. He observed a regular pattern when the frequency of X-rays was plotted against atomic number (Z), not atomic mass.
This proved:
- Atomic number (Z), not atomic weight, is the correct basis for classification.
- Atomic number equals the number of protons (or electrons in a neutral atom).
Modern Periodic Law
The physical and chemical properties of elements are periodic functions of their atomic numbers.
- This revised Mendeleev’s law.
- Explained the periodicity of properties based on electronic configuration.
Structure of the Modern Periodic Table
(A) Groups (Vertical Columns):
- 18 groups, numbered 1 to 18 (as per IUPAC).
- Elements in the same group have:
- Same valence shell configuration.
- Similar chemical properties.
(B) Periods (Horizontal Rows):
- 7 periods, representing the highest principal quantum number (n).
- Number of elements per period:
- 1st: 2 elements
- 2nd & 3rd: 8 elements
- 4th & 5th: 18 elements
- 6th & 7th: 32 elements (including f-block elements)
- 6th and 7th periods include:
- Lanthanides (Z = 58–71)
- Actinides (Z = 90–103)
- These are placed separately at the bottom to preserve table structure.
Electronic Configurations and the Periodic Table
The periodic table is not arranged randomly—each element’s position is based on its electronic configuration. The distribution of electrons in an atom determines both:
- Its period (row) is based on the outermost shell’s principal quantum number (n).
- It’s group (column) – based on the number of electrons in the valence shell.
Period-Wise Configurations
- Each new period begins when electrons start filling a new principal energy level.
- The number of elements in a period is determined by the number of orbitals available in that energy level.
- As we move from left to right in a period:
- Electrons are added successively to the same principal energy level.
- The chemical properties gradually change due to increasing nuclear charge.
Group-Wise Configurations
- Elements in the same group have similar outer electronic configurations.
- This similarity in valence shell structure explains why elements in a group show similar chemical behavior.
- For example, all alkali metals (Group 1) have a single electron in their outermost s-orbital, which gives them similar reactivity and trends.
Relationship with Quantum Numbers
- The electron configuration reflects the quantum numbers of the last electron added.
- The period number corresponds to the highest value of n (principal quantum number).
- The type of orbital being filled (s, p, d, or f) defines the block to which the element belongs.
Classification into Blocks
- Based on the subshell being filled, elements are divided into:
- s-block: Groups 1 and 2 (outer configuration ends in ns¹ or ns²).
- p-block: Groups 13 to 18 (outer configuration ends in np¹ to np⁶).
- d-block: Transition elements (Groups 3 to 12), where (n–1)d orbitals are being filled.
- f-block: Lanthanoids and actinoids, where (n–2)f orbitals are being filled.
Exceptions and Special Cases
- Hydrogen: Although it has a 1s¹ configuration like alkali metals, it also resembles halogens in gaining an electron. Hence, it is placed separately.
- Helium: Though its configuration is 1s² (like Group 2), it is placed in Group 18 because of its noble gas properties and complete valence shell.
Types of Elements: s-, p-, d-, and f-Blocks
Elements are classified into four blocks: s-block, p-block, d-block, and f-block, depending on the type of orbital in which the last electron is added. This classification is supported by the Aufbau principle and the electronic configuration of elements.
s-Block Elements
Electronic Configuration: ns¹ to ns²
Position in the Periodic Table:
- Found in Groups 1 and 2 on the left side of the periodic table.
- Includes: Alkali metals (Group 1) and Alkaline earth metals (Group 2)
Properties:
- Highly reactive metals.
- Low ionization enthalpies.
- Readily lose electrons to form positive ions (M⁺ or M²⁺).
- Metallic character and reactivity increase down the group.
- Never found free in nature due to high reactivity.
- Their compounds (except those of lithium and beryllium) are mostly ionic in nature.
p-Block Elements
Electronic Configuration: General outer configuration: ns² np¹ to ns² np⁶
Position in the Periodic Table:
- Found in Groups 13 to 18, on the right side of the periodic table.
- Along with s-block elements, they form the main group or representative elements.
Properties:
- Show a wide range of properties: metals, non-metals, and metalloids.
- The non-metallic character increases across a period (left to right).
- The metallic character increases down a group.
- Contains important groups like halogens (Group 17) and noble gases (Group 18).
- Noble gases are chemically inert due to a completely filled valence shell.
- Elements like oxygen and fluorine have high electron gain enthalpy, making them highly reactive.
d-Block Elements (Transition Elements)
Electronic Configuration: General outer configuration: (n–1)d¹⁻¹⁰ ns⁰⁻²
Position in the Periodic Table:
- Found in the center, spanning Groups 3 to 12.
- Called transition elements because they form a bridge between s-block and p-block elements.
Properties:
- All are metals.
- Often show variable oxidation states.
- Tend to form colored compounds.
- Exhibit paramagnetism due to unpaired electrons.
- Frequently act as catalysts.
- The last three elements (Zn, Cd, Hg) are exceptions and don’t show typical transition properties due to their completely filled d-subshells.
f-Block Elements (Inner Transition Elements)
Electronic Configuration: General outer configuration: (n–2)f¹⁻¹⁴ (n–1)d⁰⁻¹ ns²
Position in the Periodic Table:
- Placed separately at the bottom to keep the table compact.
- Includes two series:
- Lanthanoids (Z = 58 to 71)
- Actinoids (Z = 90 to 103)
Properties:
- All are metals.
- Within each series, elements show similar chemical properties.
- Actinoids exhibit variable oxidation states and complex chemistry.
- Many actinoids are radioactive, and those beyond uranium are called transuranium elements.
- Only minute quantities of some actinoids are known due to their instability and synthetic nature.
Metals, Non-Metals, and Metalloids
In this section, we have discussed about the metals, non-metals, and metalloids.
Metals
Metals are described below.
- Make up over 78% of known elements.
- Usually exist as solids at room temperature (exceptions: mercury is liquid; gallium and cesium have low melting points).
- Show properties such as:
- High melting and boiling points
- Good conductivity of heat and electricity
- Malleability (can be hammered into sheets)
- Ductility (can be drawn into wires)
Non-Metals
Non-metals are described below.
- Mostly found on the top right of the periodic table.
- Exist as solids or gases at room temperature (e.g., oxygen, nitrogen).
- Have low melting and boiling points (exceptions: boron and carbon).
- Poor conductors of heat and electricity.
- Brittle in solid form; neither malleable nor ductile.
Metalloids (Semi-Metals)
Elements that border the zig-zag line (or “stair-step”) between metals and non-metals in the periodic table. Show intermediate properties—some characteristics of both metals and non-metals. Examples include: silicon, germanium, arsenic, antimony, and tellurium. Useful in industries, especially in semiconductors.
Periodic Trends in Properties of Elements
Physical and chemical properties of elements change periodically across periods and down groups. These trends are explained by:
- Electronic configuration
- Effective nuclear charge
- Atomic size
- Shielding effect
Valency and Oxidation States
Valency is the combining capacity of an atom. It is usually equal to the number of electrons in the outermost shell or eight minus that number (based on the octet rule).
Periodicity in Valency:
- Across a period:
- Valency first increases from 1 to 4, then decreases back to 0.
- Example: In the 2nd period → Li (1), Be (2), B (3), C (4), N (3), O (2), F (1), Ne (0).
- Down a group:
- Valency remains constant, as elements have the same number of valence electrons.
Oxidation State
The oxidation state (oxidation number) is the apparent charge assigned to an atom in a compound, assuming electrons in each bond are assigned to the more electronegative atom.
- Determined using the concept of electronegativity.
- Helps in writing chemical formulae and understanding redox reactions.
Periodic Trends in Oxidation States
The periodic trends in oxidation states, across a period and down a group, are discussed below.
Across a Period:
- For p-block elements, the highest oxidation state equals the group number.
- The lowest oxidation state is generally equal to group number minus 8.
Example:- In Group 15 (e.g., nitrogen): Oxidation states range from –3 to +5.
- In Group 16 (e.g., oxygen, sulphur): Range from –2 to +6.
Down a Group:
- Heavier p-block elements exhibit variable oxidation states.
- Due to the inert pair effect, the lower oxidation states become more stable down the group.
Example:- In Group 13: Boron shows +3 only; Thallium prefers +1 over +3.
- In Group 14: Carbon and silicon show +4; lead shows +2 more commonly than +4.
Examples of Oxidation States in Compounds
Hydrides:
- Group 1: All form MH-type hydrides, where oxidation state is +1.
- Group 15: Form hydrides like NH₃, PH₃, showing –3 oxidation state.
Oxides:
- Group 1: Form M₂O, with metal in +1 oxidation state.
- Group 16: Form oxides like H₂O, SO₂, SO₃, showing oxidation states of –2, +4, +6 depending on the element and compound.
Anomalous Behaviour of Second Period Elements
Elements of the second period (Li, Be, B, C, N, O, F) show anomalous properties when compared with other elements in their respective groups. These differences are systematic and significant, especially for lithium and beryllium.
Reasons for Anomalous Behaviour
Reasons for anomalous behaviour are discussed below.
Small Atomic and Ionic Size: Second-period elements have much smaller atomic and ionic radii than other group members.
High Electronegativity: Their nuclei attract bonding electrons more strongly.
High Ionization Enthalpy: It requires more energy to remove electrons from these atoms.
Absence of d-Orbitals: These elements cannot expand their octet because they lack d-orbitals in their valence shell. This limits their ability to form complex or multiple bonds, unlike heavier elements.
Diagonal Relationship
Certain second-period elements resemble the properties of the elements diagonally placed to them in the third period.
Examples:
- Lithium (Li) shows similarities with Magnesium (Mg).
- Beryllium (Be) resembles Aluminium (Al).
Reasons for Diagonal Relationship:
- Similar charge-to-radius ratio.
- Similar polarizing power, leading to comparable chemical behavior.
Key Characteristics
Lithium vs. Magnesium:
- Both form nitrides on reaction with nitrogen.
- Both form normal oxides (Li₂O, MgO) instead of peroxides or superoxides.
- Both form carbonates that decompose on heating.
Beryllium vs. Aluminium:
- Both form amphoteric oxides (can act as acids or bases).
- Both form covalent chlorides that hydrolyze in water.
- Both exhibit a tendency to form complex compounds.
Periodic Trends and Chemical Reactivity
Chemical reactivity of elements is influenced by their:
- Atomic and ionic size
- Ionization enthalpy
- Electron gain enthalpy
- Electronegativity
- Elements react by losing, gaining, or sharing electrons, and this ability is determined by these periodic properties.
Reactivity Across a Period
From Left to Right:
- Metals (e.g., alkali and alkaline earth metals) on the left tend to lose electrons easily due to low ionization enthalpy → high reactivity.
- Non-metals (e.g., halogens) on the right tend to gain electrons easily due to high electron gain enthalpy → also highly reactive.
- Noble gases (far right) are least reactive due to complete octet and very high ionization enthalpy.
Reactivity Down a Group
In Metals (e.g., Group 1 – Alkali Metals):
- Reactivity increases down the group.
- Ionization enthalpy decreases → electrons are more easily lost → increased metallic reactivity.
In Non-metals (e.g., Group 17 – Halogens):
- Reactivity decreases down the group.
- Electron gain enthalpy becomes less negative → lower tendency to gain electrons → decreased non-metallic reactivity.
Examples of Reactivity
With Oxygen:
- Alkali metals form basic oxides (e.g., Na₂O, K₂O).
- Non-metals form acidic oxides (e.g., SO₂, Cl₂O₇).
- Amphoteric oxides are formed by elements like Aluminium and Zinc, showing dual nature.
With Halogens:
- Reactivity of metals with halogens leads to formation of halides (e.g., NaCl, KBr).
- Alkali metals react vigorously with halogens to form ionic halides.
- Transition metals form covalent halides with variable oxidation states.
Important Formulas in NCERT Notes Chemistry (Part-I) Chapter 3: Classification of Elements and Periodicity in Properties
Here are the important formulas from Chapter 3: Classification of Elements and Periodicity in Properties.
- Effective Nuclear Charge (Z_eff):
Z_eff = Z – S
Where:- Z = atomic number
- S = shielding or screening constant
- Ionization Enthalpy (ΔiH):
X(g) → X⁺(g) + e⁻
(Energy required to remove an electron from a gaseous atom) - Electron Gain Enthalpy (ΔegH):
X(g) + e⁻ → X⁻(g)
(Enthalpy change when a gaseous atom gains an electron) - Oxidation Number (based on electronegativity):
In a binary compound:- The more electronegative element is assigned a negative oxidation state.
- The less electronegative element is assigned a positive oxidation state.
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FAQs
Answer: Atomic number represents the number of protons in an atom and determines the element’s identity and electronic configuration. Modern Periodic Law is based on atomic number because it correctly explains the periodic repetition of properties, resolves anomalies in Mendeleev’s table, and aligns with observed chemical behavior.
Answer: Lithium and beryllium show properties different from other group members due to their small size, high ionization enthalpy, high electronegativity, and absence of d-orbitals. These factors lead to stronger covalent character and the formation of compounds unlike those of other alkali or alkaline earth metals.
Answer: In alkali metals, ionization enthalpy decreases down the group, making it easier to lose electrons and thus increasing reactivity. In contrast, halogens have less negative electron gain enthalpy down the group, reducing their tendency to gain electrons and thus decreasing reactivity.
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