Periodic Classification of Elements Class 10 Notes pdf

 

Periodic Classification of Elements 

• Early Attempts at the Classification of Elements 

• Mendeleev's Periodic Table

•  Modern Periodic Table

•  Trends in the Modern Periodic Table

•  Mendeleev's Predictions 

• Metallic and Non-Metallic Behavior

•  Periodic Trends in Physical and Chemical Properties of Elements 

• Anomalous Properties of the Second Period Elements

•  Modern Electronic Configurations and the Periodic Table 

• Uses of the Periodic Table.

• Early attempts at the classification of elements refer to the efforts made by scientists to organize the known elements into a systematic pattern based on their properties.

• Some early attempts at classification include the Dobereiner's Triads, Newlands' Law of Octaves, and Mendeleev's Periodic Table.

• Dobereiner's Triads were groups of three elements with similar properties, where the atomic mass of the middle element was approximately the average of the other two elements.

• Newlands' Law of Octaves stated that every eighth element had similar properties to the first, much like the notes in an octave of music.

• Mendeleev's Periodic Table was the most successful of these early attempts, which arranged the elements in increasing order of atomic mass, and grouped them based on similar chemical and physical properties. He also left gaps for undiscovered elements and made predictions about their properties.

• Despite their limitations, these early attempts laid the foundation for the development of the modern periodic table and helped scientists understand the periodic nature of elements.

Mendeleev's periodic table

• Mendeleev's periodic table is a table that arranges the elements in order of increasing atomic mass and groups them based on their similar chemical and physical properties.

• Mendeleev's periodic table had several rows and columns. The rows were called periods, and the columns were called groups. The elements in the same group had similar chemical and physical properties.

• For example, all the elements in Group 1, also known as the alkali metals, are highly reactive and have only one valence electron in their outermost shell. Sodium (Na), potassium (K), and lithium (Li) are some of the elements in this group.

• Similarly, all the elements in Group 18, also known as the noble gases, are chemically inert and have a complete outermost shell of electrons. Helium (He), neon (Ne), and argon (Ar) are some of the elements in this group.

• Mendeleev's periodic table also left gaps for undiscovered elements and made predictions about their properties based on their position in the table. One such example is germanium (Ge), which was discovered later and matched the properties that Mendeleev had predicted for the element.

• Overall, Mendeleev's periodic table was a significant step forward in the classification of elements and laid the foundation for the modern periodic table.

Modern Periodic Table

• Modern periodic table is a table that arranges the elements in order of increasing atomic number and groups them based on their similar chemical and physical properties.

• The modern periodic table has several rows and columns. The rows are called periods, and the columns are called groups. The elements in the same group have similar chemical and physical properties.

• For example, all the elements in Group 1, also known as the alkali metals, are highly reactive and have only one valence electron in their outermost shell. Sodium (Na), potassium (K), and lithium (Li) are some of the elements in this group.

• Similarly, all the elements in Group 18, also known as the noble gases, are chemically inert and have a complete outermost shell of electrons. Helium (He), neon (Ne), and argon (Ar) are some of the elements in this group.

• In the modern periodic table, the elements are arranged based on their electronic configurations, and their properties repeat in a periodic manner. The horizontal rows, or periods, represent the number of electron shells in the atoms of the elements. The vertical columns, or groups, represent the number of valence electrons in the atoms of the elements.

• The modern periodic table also includes a blockwise classification of elements based on the types of subshells that are being filled in their electronic configuration. The blocks are the s-block, p-block, d-block, and f-block, each containing a different set of elements. The modern periodic table has enabled scientists to predict the properties of elements based on their position in the table and has been instrumental in the discovery and understanding of new elements.

• Overall, it is a powerful tool for organizing and understanding the properties of the elements.

Trends in the Modern Periodic

Table Trends in the modern periodic table refer to the patterns that are observed in the properties of elements as one moves across a period or down a group in the table. These trends can provide valuable information about the behavior and characteristics of elements. The following are some of the significant trends observed in the modern periodic table:

Atomic radius:

The atomic radius of elements generally decreases from left to right across a period and increases down a group. This is because the number of protons in the nucleus increases, which attracts the electrons more strongly and makes the atom smaller. Conversely, going down a group, the number of electron shells increases, making the atom larger.

Electronegativity:

Electronegativity is a measure of an atom's ability to attract electrons to itself in a chemical bond. Electronegativity generally increases from left to right across a period and decreases down a group. This is because the atomic radius decreases across a period, making it harder for electrons to escape, and the nuclear charge increases. Conversely, going down a group, the distance between the nucleus and the valence electrons increases, reducing the attraction between the nucleus and the electrons.

Ionization energy:

Ionization energy is the energy required to remove an electron from an atom. Ionization energy generally increases from left to right across a period and decreases down a group. This is because the atomic radius decreases across a period, making it harder for electrons to escape, and the nuclear charge increases. Conversely, going down a group, the distance between the nucleus and the valence electrons increases, reducing the attraction between the nucleus and the electrons.

Metallic character:

Metallic character refers to the tendency of an element to lose electrons and form positive ions. Metallic character generally increases down a group and decreases across a period. This is because the atomic radius increases down a group, making it easier for electrons to escape, and the number of valence electrons increases. Conversely, going across a period, the atomic radius decreases, making it harder for electrons to escape, and the number of valence electrons remains constant.

Mendeleev's Predictions

• Mendeleev's periodic table was a significant step forward in the classification of elements and laid the foundation for the modern periodic table. One of the most remarkable aspects of Mendeleev's work was his ability to use the periodic table to make predictions about the properties of undiscovered elements.

• Mendeleev left gaps in his periodic table for elements that had not yet been discovered. He predicted the properties of these elements based on the properties of the elements around the gaps. For example, he predicted the existence and properties of germanium, which was discovered later and matched the properties he had predicted.

• Mendeleev also predicted the existence of elements that were not discovered until much later, such as gallium and scandium. He even predicted the properties of an element he called ekasilicon, which he believed would have similar properties to silicon but with a higher atomic weight. This element was later discovered and named germanium.

• Mendeleev's predictions were based on the observation of periodicity in the properties of elements. He believed that the properties of elements were related to their atomic weight and that there was a fundamental order to the elements that could be used to predict their behavior. His predictions were a testament to the power of the periodic table as a tool for understanding the properties of elements and predicting the behavior of undiscovered elements.

Metallic and Non-Metallic Behavior

• Metallic and non-metallic behavior refers to the characteristic properties of elements that determine their ability to conduct electricity, heat, and their tendency to lose or gain electrons in chemical reactions.

• Metals are generally good conductors of heat and electricity, and they tend to lose electrons to form positively charged ions in chemical reactions. They are also usually lustrous, malleable, and ductile. Examples of metals include iron, copper, gold, and aluminum.

• Non-metals, on the other hand, are generally poor conductors of heat and electricity, and they tend to gain electrons to form negatively charged ions in chemical reactions. They are also often dull, brittle, and non-malleable. Examples of non-metals include carbon, oxygen, nitrogen, and sulfur.

• Metallic behavior tends to increase as one moves down a group in the periodic table, while non-metallic behavior tends to increase as one moves across a period. This is because as one moves down a group, the atomic radius increases, making it easier for electrons to move around and for metals to lose electrons. Conversely, as one moves across a period, the number of valence electrons increases, making it more difficult for non-metals to gain electrons.

However, there are also elements that exhibit properties of both metals and non-metals, known as metalloids. These elements have intermediate properties and can behave either as metals or non-metals depending on the conditions. Examples of metalloids include boron, silicon, and arsenic.

Periodic Trends in Physical and Chemical Properties of Elements

Atomic radius: The atomic radius of elements generally increases as one moves down a group in the periodic table, and decreases as one moves across a period from left to right. This is because the number of electrons in the outermost shell increases down a group, while the effective nuclear charge increases across a period.

Ionization energy: The ionization energy of an element is the energy required to remove an electron from the outermost shell of an atom. The ionization energy generally decreases as one moves down a group in the periodic table, and increases as one moves across a period from left to right. This is because the outermost electrons are farther from the nucleus down a group and are more strongly attracted to the nucleus across a period.

Electronegativity: Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond. Electronegativity generally decreases as one moves down a group in the periodic table and increases as one moves across a period from left to right.

Metallic character: Metallic character refers to the degree to which an element exhibits metallic properties such as luster, ductility, and conductivity. Metallic character generally increases as one moves down a group in the periodic table and decreases as one moves across a period from left to right.

Some of the important periodic trends in chemical properties of elements include:

• Reactivity: The reactivity of elements generally increases as one moves down a group in the periodic table. This is because the outermost electrons are farther from the nucleus and are more easily lost or gained in chemical reactions.

• Oxidation states: The oxidation state of an element is the number of electrons that it gains or loses in a chemical reaction. The range of oxidation states that an element can exhibit generally increases as one moves down a group in the periodic table.

• Chemical behavior: The chemical behavior of elements can be predicted based on their position in the periodic table. For example, alkali metals are highly reactive and tend to form ionic compounds with non-metals, while halogens are highly reactive non-metals that tend to form covalent compounds with metals.

Anomalous Properties of the Second Period Elements

The second period of the periodic table consists of elements from lithium (Li) to neon (Ne). These elements exhibit certain anomalous properties that make them different from other elements in the periodic table. Some of the anomalous properties of the second period elements are:

• Small atomic size: The second period elements have small atomic radii compared to elements in other periods. This is because the second period elements have only two electron shells, and the outermost shell is very close to the nucleus.

• High electronegativity: The second period elements have high electronegativity values compared to elements in other periods. This is because the second period elements have a small atomic size, and the outermost electrons are held tightly by the nucleus, making it difficult to remove them.

• Low melting and boiling points: The second period elements have low melting and boiling points compared to elements in other periods. This is because the second period elements have weak metallic bonding due to their small size and high electronegativity.

• Large ionization energy: The second period elements have large ionization energies compared to elements in other periods. This is because the second period elements have a small atomic size and high electronegativity, making it difficult to remove electrons from them.

• Anomalous electron configurations: The electron configurations of the second period elements are anomalous due to the stability provided by the half-filled and fully-filled subshells. For example, nitrogen (N) has a half-filled p subshell, which makes it more stable than oxygen (O), which has one electron in the p subshell.

• Diagonal relationship: The second period elements exhibit a diagonal relationship with elements in the third period. For example, lithium (Li) has similar properties to magnesium (Mg), and beryllium (Be) has similar properties to aluminum (Al). This is because the second period elements and third period elements have similar atomic sizes and electronegativities.

Modern Electronic Configurations and the Periodic Table

• The electronic configuration of an atom describes the distribution of electrons among the energy levels or orbitals within the atom. The arrangement of electrons in an atom plays a crucial role in determining the chemical and physical properties of the element.

• In the modern periodic table, the elements are arranged in order of increasing atomic number, and the electronic configuration of an element is related to its position in the periodic table. Each element has a unique electronic configuration that determines its chemical and physical properties. The electronic configuration of an element can be written using the noble gas notation, which represents the electron configuration of the noble gas that precedes the element in the periodic table.

• The periodic table is divided into rows and columns, known as periods and groups, respectively. The number of electrons in the outermost shell of an atom determines its position in the periodic table. The elements in the same group or column have similar electronic configurations and exhibit similar chemical properties.

• For example, the elements in Group 1 (the alkali metals) have one electron in their outermost shell, and their electronic configuration is [noble gas] ns1. The elements in Group 17 (the halogens) have seven electrons in their outermost shell, and their electronic configuration is [noble gas] ns2 np5. The elements in Group 18 (the noble gases) have a completely filled outermost shell, and their electronic configuration is [noble gas] ns2 np6.

• The modern periodic table also explains the concept of periodicity, which refers to the repetition of properties of elements in a periodic manner. The periodic table exhibits periodic trends in physical and chemical properties, such as atomic radius, ionization energy, electronegativity, and metallic character. These trends are related to the electronic configuration of elements and can be used to predict the properties and behavior of elements.

Uses of the Periodic Table

Major uses of the periodic table are:

• Predicting the chemical and physical properties of elements: The periodic table provides a systematic way of organizing elements based on their atomic structure and electronic configuration. This arrangement helps to predict the chemical and physical properties of elements, such as atomic radius, electronegativity, ionization energy, and metallic character.
• Identifying and classifying elements: The periodic table allows us to identify and classify elements based on their properties, such as metal, non-metal, or metalloid. This classification is based on the position of the elements in the periodic table and is used to predict their behavior in chemical reactions.
• Designing new materials: The periodic table is used to design new materials with specific properties. By knowing the properties of elements and how they react with other elements, scientists can create new materials with unique characteristics.
• Understanding the behavior of compounds: The periodic table provides information on the properties of individual elements, which can be used to understand the behavior of compounds. For example, the electronegativity values of elements can be used to predict the polarity of chemical bonds in a compound.
• Developing new technologies: The periodic table is used in the development of new technologies, such as semiconductors, batteries, and solar cells. By understanding the properties of elements and how they interact, scientists can create new materials and technologies that are more efficient and effective.