Elements and Periodicity in Properties: Comprehensive NEET Chemistry Notes

1. Introduction to Periodicity

The periodic table is a fundamental tool in chemistry, organizing elements based on their properties and periodic trends. Understanding periodicity helps predict element behavior and chemical reactions.

1.1 Importance of Classification

Elements are classified to systematize the study of their properties and behaviors. Historically, classification led to the development of the periodic table, allowing chemists to predict properties of elements.

Did You Know?

Dmitri Mendeleev, the father of the periodic table, predicted the properties of undiscovered elements based on periodic trends.

Real-life Application:

Periodic trends are used to design alloys and other materials with specific properties.

2. Genesis of Periodic Classification

The classification of elements evolved over time through the work of many scientists.

2.1 Early Classifications

Johann Dobereiner's triads grouped elements with similar properties. John Newlands proposed the Law of Octaves, noting that every eighth element had similar properties.

2.2 Mendeleev’s Periodic Table

Mendeleev arranged elements by increasing atomic weight, predicting the existence and properties of undiscovered elements. His table organized elements into groups with similar chemical properties.

Mnemonic:

"Mendeleev's Magic: Properties repeat at intervals!"

3. Modern Periodic Law and Table

Henry Moseley revised Mendeleev's table by arranging elements in order of increasing atomic number, leading to the modern periodic law.

3.1 Modern Periodic Law

The modern periodic law states that the properties of elements are a periodic function of their atomic numbers.

3.2 Structure of the Modern Periodic Table

The table is divided into periods and groups. Elements in the same group have similar outer electron configurations, leading to similar chemical properties.

NEET Tip:

Focus on understanding the trends within groups and periods, as these are frequently tested in NEET exams.

4. Nomenclature of Elements with Atomic Numbers >100

Elements with atomic numbers greater than 100 are named systematically using a combination of numerical roots.

Example:

The element with atomic number 120 would be named Unbinilium (Ubn).

5. Electronic Configurations and Periodic Table

The periodic table reflects the electronic configurations of elements.

5.1 Periods and Electronic Configurations

Each period corresponds to the filling of a principal energy level (n). The number of elements in each period is determined by the number of available orbitals.

5.2 Groups and Valence Shell Configurations

Elements in the same group have similar valence shell configurations, leading to similar chemical properties.

NEET Problem-Solving Strategy:

Use the electronic configuration to predict an element's chemical behavior and reactivity.

6. Classification of Elements: s, p, d, f Blocks

Elements are classified into blocks based on the type of atomic orbitals being filled.

6.1 s-Block Elements

Groups 1 and 2 elements have outermost s orbitals. They are highly reactive metals with low ionization enthalpies.

6.2 p-Block Elements

Groups 13 to 18 elements have outermost p orbitals. This block includes metals, metalloids, and non-metals with varied properties.

6.3 d-Block Elements

Transition metals (Groups 3 to 12) have d orbitals. They form colored compounds, exhibit variable oxidation states, and are often catalysts.

6.4 f-Block Elements

Lanthanides and actinides have f orbitals. These elements are metals, with actinides being mostly radioactive.

Concept Connection:

Transition metals are essential in biological processes and industrial catalysis.

7. Periodic Trends in Properties

Understanding periodic trends helps predict element behavior.

7.1 Atomic and Ionic Radii

Atomic size decreases across a period and increases down a group. Cations are smaller than their parent atoms, while anions are larger.

7.2 Ionization Enthalpy

Ionization enthalpy increases across a period and decreases down a group. It is the energy required to remove an electron from an atom.

7.3 Electron Gain Enthalpy

Electron gain enthalpy becomes more negative across a period and less negative down a group. It is the energy change when an electron is added to an atom.

7.4 Electronegativity

Electronegativity increases across a period and decreases down a group. It measures an atom's ability to attract shared electrons.

Common Misconception:

Ionization enthalpy is not the same as electron affinity; the former is the energy to remove an electron, while the latter is the energy change when an electron is added.

Quick Recap

• Periodic classification organizes elements based on properties.
• Mendeleev's periodic table laid the foundation for the modern periodic law.
• Elements are classified into s, p, d, and f blocks.
• Periodic trends include atomic and ionic radii, ionization enthalpy, electron gain enthalpy, and electronegativity.

Practice Questions

1. Question: Define the modern periodic law. Solution: The modern periodic law states that the properties of elements are a periodic function of their atomic numbers.
2. Question: What is the trend in atomic radius across a period? Solution: Atomic radius decreases across a period due to increased nuclear charge attracting electrons closer.
3. Question: Predict the electronic configuration of the element with atomic number 120. Solution: The electronic configuration would be [Uuo]8s2.
4. Question: Explain why ionization enthalpy generally increases across a period. Solution: Ionization enthalpy increases due to increased nuclear charge, making it harder to remove an electron.
5. Question: What is the significance of electron gain enthalpy? Solution: Electron gain enthalpy indicates how easily an atom can gain an electron to form an anion.
6. Question: Differentiate between s-block and p-block elements. Solution: s-block elements have outermost s orbitals, while p-block elements have outermost p orbitals.
7. Question: Describe the trend in electronegativity down a group. Solution: Electronegativity decreases down a group due to increased atomic size and shielding effect.
8. Question: Calculate the number of elements in the fourth period. Solution: The fourth period has 18 elements, including the 3d transition series.
9. Question: Why do cations have smaller radii than their parent atoms? Solution: Cations have fewer electrons, leading to reduced electron-electron repulsion and a smaller radius.
10. Question: Explain the diagonal relationship with an example. Solution: Lithium shows properties similar to magnesium, despite being in different groups, due to similar charge/radius ratio.

Quick Reference Guide and Glossary

• Periodic Law: Properties of elements are periodic functions of their atomic numbers.
• s-Block Elements: Elements with outermost s orbitals (Groups 1 and 2).
• p-Block Elements: Elements with outermost p orbitals (Groups 13 to 18).
• d-Block Elements: Transition metals with d orbitals.
• f-Block Elements: Lanthanides and actinides with f orbitals.
• Ionization Enthalpy: Energy required to remove an electron from an atom.
• Electron Gain Enthalpy: Energy change when an electron is added to an atom.
• Electronegativity: Ability of an atom to attract shared electrons.