Grade 9
  1. Matter and its nature  1.1. Introduction to matter  1.2. Physical and chemical properties of matter  1.3. States of matter  1.3.1. Solid state  1.3.2. Fluid state  1.3.3. Gaseous state  1.3.4. Plasma and Bose–Einstein condensate  1.4. Changes in the states of matter  1.4.1. Melting and freezing  1.4.2. Evaporation and Condensation  1.4.3. Sublimation and deposition  1.5. Classification of matter  1.6. Elements, Compounds, and Mixtures  1.7. Pure substances and impurities  1.8. Separation Techniques  1.8.1. Filtration  1.8.2. Distillation  1.8.3. Crystallization  1.8.4. Chromatography  1.8.5. Centrifugation  1.8.6. Sublimation  1.8.7. Decantation and sedimentation
  2. Atomic Structure  2.1. Discovery of atoms  2.2. Dalton's atomic theory  2.3. Structure of the atom  2.4. Subatomic particles  2.5. Atomic number and mass number  2.6. Isotopes and Isobars  2.7. Electronic configuration  2.8. Bohr's atomic model  2.9. Modern Atomic Model (Quantum Mechanical Model - Introduction)  2.10. Valency and chemical bonding
  3. C hemical reactions and equations  3.1. Introduction to Chemical Reactions  3.2. Chemical Equation Formulation  3.3. Balancing Chemical Equations  3.4. Types of Chemical Reactions  3.4.1. Combination Reactions  3.4.2. Decomposition reactions  3.4.3. Displacement reactions  3.4.4. Double displacement reactions  3.4.5. Redox reactions  3.4.6. Endothermic and Exothermic Reactions  3.5. Effects of chemical reactions  3.6. Energy Changes in Chemical Reactions  3.7. Catalysts and their role in reactions
  4. Periodic table and periodicity  4.1. History of Periodic Classification  4.2. Mendeleev's periodic table  4.3. Modern Periodic Table  4.4. Periods and groups in the periodic table and periodicity  4.5. Trends in the Periodic Table  4.5.1. Atomic size  4.5.2. Ionization Energy  4.5.3. Electron affinity  4.5.4. Electronegativity  4.5.5. Metallic and Non-Metallic Properties  4.6. Noble gases and their properties
  5. Chemical bond  5.1. Introduction to Chemical Bonding  5.2. Types of chemical bonds  5.2.1. Ionic bond  5.2.2. Covalent bond  5.2.3. Metallic bond  5.2.4. Hydrogen bonding  5.2.5. Van der Waals force  5.3. Properties of Ionic and Covalent Compounds  5.4. Molecular Structure and Geometry (VSEPR Theory - Basics)
  6. Acids, Bases and Salts  6.1. Introduction to Acids and Bases  6.2. Properties of Acids  6.3. Properties of bases  6.4. pH scale and its importance  6.5. Indicators and their uses  6.6. Neutralization reactions  6.7. Strength of Acids and Bases (Strong vs. Weak)  6.8. Preparation of salts  6.9. Uses of acids, bases and salts in daily life
  7. Metals and Nonmetals  7.1. Physical Properties of Metals  7.2. Physical Properties of Nonmetals  7.3. Chemical properties of metals  7.4. Chemical Properties of Nonmetals  7.5. Reactivity Series of Metals  7.6. Extraction of metals  7.7. Corrosion and its prevention  7.8. uses of metals and nonmetals
  8. Carbon and its compounds  8.1. Introduction to Carbon  8.2. Allotropes of carbon (diamond, graphite, fullerene)  8.3. Hydrocarbons  8.3.1. Hydrocarbons  8.3.2. Alkene  8.3.3. Alkynes  8.4. Functional groups in organic chemistry  8.5. Isomerism in organic compounds  8.6. Alcohols and Carboxylic Acids  8.7. Soaps and detergents  8.8. Polymers (Introduction to Natural and Synthetic Polymers)
  9. Solutions and Mixtures  9.1. Types of mixtures  9.2. Homogeneous and Heterogeneous Mixtures  9.3. Concentration of solutions  9.4. Solubility and factors affecting solubility  9.5. Suspensions and Colloids  9.6. Saturated, unsaturated and supersaturated solutions
  10. Air and atmosphere  10.1. Composition of air  10.2. Role of gases in the atmosphere  10.4. Acid rain and its effects  10.5. Greenhouse effect and global warming  10.6. Ozone layer and its depletion  10.7. Air pollution control measures
  11. Water and its importance  11.1. Composition and properties of water  11.2. Hardness of water (temporary and permanent hardness)  11.3. Causes of water pollution  11.4. Effects of Water Pollution  11.5. Water purification methods (filtration, boiling, chlorination)
  12. Industrial Chemistry  12.1. Introduction to Industrial Chemistry  12.2. Important industrial chemicals (sulfuric acid, ammonia, sodium hydroxide)  12.3. Chemical processes in industry  12.4. Uses of Industrial Chemistry in Daily Life  12.5. Environmental impact of industrial chemical processes

Grade 9 → Periodic table and periodicity ↓


Modern Periodic Table


The Modern Periodic Table is a remarkable tool in chemistry that arranges all known chemical elements in a way that displays patterns and relationships between them. Understanding the Modern Periodic Table is fundamental for students as it helps them understand the behaviour and characteristics of elements, which is important in chemistry. The aim of this lesson is to explain the concept of the Modern Periodic Table in a simple manner, thereby providing a comprehensive guide for students.

Development of the periodic table

The development of the periodic table has a very long history, starting with early attempts to classify elements by scientists such as Johann Wolfgang Dobereiner, who grouped elements into triads based on their similarities. However, the most recognized and influential contribution was that of Dmitry Mendeleev in 1869, who arranged the elements in order of increasing atomic weight and grouped them according to their properties. He left gaps for elements yet to be discovered, accurately predicting their properties.

The modern periodic table has evolved since Mendeleev's time and has been refined by subsequent discoveries, including the arrangement of the elements by Henry Moseley, who developed the concept of atomic number in 1913. This replaced atomic weight as the organizing principle, leading to improved accuracy and understanding of the arrangement of the elements.

Structure of the modern periodic table

The modern periodic table is arranged in rows and columns called periods and groups respectively. It is arranged as follows:

Period

The rows in the periodic table are called periods. The modern periodic table has seven periods, numbered from 1 to 7. Elements are arranged in a period in order of increasing atomic number from left to right. Each period begins with a new principal energy level.

For example, period 1 contains hydrogen (H) and helium (He), while period 2 begins with lithium (Li) and ends with neon (Ne).

Period 1: H, He
Period 2: Li, Be, B, Si, N, O, F, Ne
Period 3: Na, Mg, Al, Si, P, S, Cl, Ar

Group

The columns of the periodic table are called groups. The modern periodic table has 18 groups, numbered 1 to 18. Elements in the same group have similar chemical properties because they have the same number of electrons in their outer shell, known as valence electrons.

Some of the major groups are:

  • Group 1 (Alkali Metals): It includes Lithium (Li), Sodium (Na), Potassium (K), etc. These elements are highly reactive metals.
  • Group 2 (Alkaline Earth Metals): It includes Beryllium (Be), Magnesium (Mg), Calcium (Ca), etc.
  • Group 17 (Halogens): It includes fluorine (F), chlorine (Cl), bromine (Br), etc. These are very reactive non-metals.
  • Group 18 (Noble gases): It includes helium (He), neon (Ne), argon (Ar), etc. These elements are known to be inert.

Blocks of the periodic table

The periodic table is divided into blocks based on the electron configuration of the elements. These blocks are named after the orbitals being filled.

S Block

The s-block includes groups 1 and 2 and the element helium. Characteristics of the s-block elements include:

  • Highly reactive metals like alkali and alkaline earth metals
  • their outermost electrons are in the s orbital
  • Example: Lithium (Li) with electron configuration 1s 2 2s 1

P-block

The p-block includes groups 13 to 18. Characteristics of the p-block elements include:

  • It includes metals, nonmetals, and metalloids
  • their outermost electrons are in the p orbital
  • Example: Carbon (C) electron configuration 1s 2 2s 2 2p 2

D-block

The d-block is located in the center of the periodic table, containing the transition metals. Characteristics include:

  • Metallic Properties
  • Their outermost electrons are in the d orbital
  • Example: Iron (Fe) electron configuration [Ar] 3d 6 4s 2

F Block

The f-block contains the lanthanides and actinides. These are located below the main body of the periodic table. Characteristics include:

  • Rare earth elements and actinides, many of which are radioactive
  • their outermost electrons are in the f orbital
  • Example: Uranium (U) electron configuration [Rn] 5f 3 6d 1 7s 2

Periodicity of elements

Periodicity refers to recurring trends that are observed in the properties of elements. These trends are particularly visible when elements are arranged according to their atomic number and are important for predicting element behavior.

Atomic radius

The atomic radius is the distance from the nucleus of an atom to its outermost shell of electrons. As you move from left to right across a period, the atomic radius decreases. This is due to the increasing nuclear charge which pulls the electron cloud closer to the nucleus. Conversely, as you move down a group, the atomic radius increases due to the addition of more energy levels.

Period (from left to right): Atomic radius decreases Group (top to bottom): Atomic radius increases

Ionization energy

Ionization energy is the energy needed to remove an electron from a gaseous atom. Ionization energy generally increases as you move across a period from left to right. This is because atoms are smaller and hold on to their electrons more tightly due to the stronger nuclear charge. Moving down the group, ionization energy generally decreases because the outer electrons are farther from the nucleus and shielded by inner shells.

Electron affinity

Electron affinity is the amount of energy released when an electron is added to a gaseous atom. Across a period, electron affinity becomes more negative, indicating a stronger attraction for the additional electrons. Moving down the group, electron affinity becomes less negative due to larger atomic size.

Less High

Salient features of modern periodic table

Electronic configuration

The periodic table helps determine the electronic configuration of elements. The distribution of electrons among the orbitals of an atom determines its electronic configuration and ultimately its chemical behaviour.

For example, the electronic configuration of the element sodium (Na) is 1s 2 2s 2 2p 6 3s 1.

Valency

Valency refers to the combining capacity of an element. It is determined by the number of valence electrons present in the outer shell of an atom. Elements in the same group usually have the same valency.

  • Group 1 elements such as lithium (Li) and sodium (Na) have one valence electron, giving them a valency of 1.
  • Group 17 elements, such as chlorine (Cl), have seven valence electrons, which usually gives them a valency of 1 when combining with metals.

Metallic and non-metallic properties

In a period, metallic character decreases and non-metallic character increases due to increase in ionization energy and electron affinity.

For example, in period 2, as you move from lithium (Li) to fluorine (F), lithium is a metal, while fluorine is a nonmetal.

Metalloids

Metalloids have properties intermediate between metals and nonmetals. They are found along the ladder line on the periodic table. Elements such as boron (B), silicon (Si), and arsenic (As) are metalloids.

Metals Metalloids Non metallic

Conclusion

The modern periodic table is not just a static chart, but a dynamic tool that helps chemists understand the principles that govern the behavior of the elements. Its organization by atomic number, electron configuration, and recurring chemical properties reveals a fascinating regularity among the elements, allowing predictions of chemical behavior and interactions. Understanding the modern periodic table equips students with the fundamental knowledge needed to explore the complex world of chemistry.


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Modern Periodic Table

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