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 configuration1s 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 configuration1s 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.
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.
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.
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.