Periodic Classification of Elements for Class 10

Dobereiner's Triads

Dobereiner's triads were an early attempt to classify elements based on their chemical properties and atomic masses. They were proposed by the German chemist Johann Wolfgang Dobereiner in the early 19th century, specifically in 1829. Dobereiner noticed that certain groups of three elements shared similar chemical properties and had a pattern related to their atomic masses.

1. Formation of Triads: Dobereiner identified several groups of three elements, which he called "triads." Each triad consisted of three elements with similar chemical properties.
2. Similar Chemical Properties: The key criterion for elements to be part of a triad was that they should exhibit similar chemical properties. These properties included how the elements reacted with other substances, their valency (the number of chemical bonds they could form), and the types of compounds they could produce.
3. Atomic Mass Relationship: The most significant aspect of Dobereiner's triads was the relationship between the atomic masses of the elements within a triad. Specifically, the atomic mass of the middle element (the second element in the triad) was approximately equal to the arithmetic mean (average) of the atomic masses of the first and third elements in the triad.

Examples of Dobereiner's Triads

1. Alkali Metal Triad

1. One of the most famous examples of a Dobereiner's triad is the alkali metal triad, which includes lithium (Li), sodium (Na), and potassium (K).
2. Properties shared by these elements included:

• They were all metals.
• They had a valency of 1 (monovalent), meaning they could form compounds with other elements by losing one electron.
• They reacted with water to form alkaline solutions (alkalis) and released hydrogen gas.
• The atomic mass of sodium (the middle element) was approximately the average of the atomic masses of lithium and potassium.

2. Alkaline Earth Metal Triad

1. Another example of a Dobereiner's triad was the alkaline earth metal triad, which included calcium (Ca), strontium (Sr), and barium (Ba).
2. Properties shared by these elements included:

• They were all metals.
• They had a valency of 2 (divalent), meaning they could form compounds by losing two electrons.
• Their oxides were alkaline in nature.
• The atomic mass of strontium (the middle element) was approximately the average of the atomic masses of calcium and barium.

3. Halogens Triad

1. Dobereiner also identified a triad for the halogens, which included chlorine (Cl), bromine (Br), and iodine (I).
2. Properties shared by these elements included:

• They were all non-metals.
• They had a valency of 1 (monovalent).
• They could react with metals to form salts.
• The atomic mass of bromine (the middle element) was approximately the average of the atomic masses of chlorine and iodine.

Limitations:

1. The main limitation of Dobereiner's classification was that it could only accommodate a limited number of elements, and not all known elements at the time fit into triads with similar properties and atomic mass relationships.
2. Dobereiner's classification system was not comprehensive enough to encompass all the known elements and their variations.

Newlands' Law of Octaves

John Alexander Newlands, an English chemist, proposed the "Law of Octaves" in 1864 as a way to classify and arrange the known elements at that time. His idea was based on the observation that when elements were arranged in order of increasing atomic masses, every eighth element displayed similar chemical properties to the first element, just like the repetition of musical notes in an octave.

Key points of Newlands' Law of Octaves:

1. Arrangement by Increasing Atomic Mass: Newlands arranged the elements in order of increasing atomic masses, starting with hydrogen (H) as the first element.
2. Repetition of Properties: He observed that after arranging elements in this way, the properties of the eighth element (counting from a given element) were similar to those of the first element. This repetition continued periodically.
3. Analogous to Music Octaves: Newlands likened this pattern to the octaves in music, where the eighth note in a musical octave repeats the properties of the first note.
4. Formation of Rows: Newlands divided the elements into horizontal rows, with each row containing seven elements. The eighth element in each row would have properties similar to the first element, following the pattern.

Simplified representation of Newlands' arrangement of elements:

1. In this arrangement, lithium (Li) was the first element, and sodium (Na) was the eighth element from it.
2. Sodium displayed similar properties to lithium.
3. The same pattern continued for other elements in their respective rows.

Limitations:

However, Newlands' Law of Octaves had several limitations:

1. Limited to Lighter Elements: The law worked well for lighter elements up to calcium (Ca). Beyond calcium, it failed to predict the properties of elements accurately.
2. Assumed Limited Elements: Newlands incorrectly assumed that only 56 elements existed, and he didn't account for the discovery of new elements in the future.
3. Arbitrary Grouping: In some cases, he grouped elements together that had very different properties. For instance, cobalt (Co) and nickel (Ni) were placed in the same slot with elements like fluorine, chlorine, and bromine, despite having dissimilar properties.
4. Lack of Predictive Power: The law did not provide a systematic way to predict the properties of unknown elements.

Despite its limitations, Newlands' Law of Octaves was an early step in understanding the periodicity of elements and contributed to the development of the periodic table.

Mendeleev’s Classification of Elements

Mendeleev's Classification of Elements, also known as the Mendeleev Periodic Table, was a pioneering system for organising chemical elements based on their properties and atomic masses. Dmitri Mendeleev, a Russian chemist, introduced this classification system in 1869.

Overview of Mendeleev's classification:

1. Periodic Law: Mendeleev's classification was built upon the Periodic Law, which states that the properties of elements are periodic functions of their atomic masses. This means that when elements are arranged in order of increasing atomic masses, their properties repeat at regular intervals.
2. Periods and Groups: Mendeleev's periodic table consisted of seven horizontal rows called "periods" and eight vertical columns called "groups." Each group had specific properties and similarities.
3. Elements with Similar Properties: Mendeleev organised elements into groups based on their similar chemical properties. He considered various properties, such as the chemical formulas of oxides and hydrides, to determine these groupings.
4. Use of Gaps: Mendeleev realised that not all elements were known at the time. To account for undiscovered elements, he left gaps in his periodic table and predicted the properties of these yet-to-be-discovered elements. He gave them names like eka-boron, eka-aluminium, and eka-silicon.

Merits of Mendeleev’s Classification:

Mendeleev's classification of elements had several notable merits:

1. Predictability: One of the most remarkable aspects of Mendeleev's classification was his ability to predict the properties of undiscovered elements. When these elements were eventually found, their properties closely matched Mendeleev's predictions. This predictive power demonstrated the validity of the periodic law.
2. Organisation of Elements: Mendeleev's periodic table provided a systematic and organised way to study the properties and relationships among elements. It helped chemists understand the similarities and differences between elements more effectively.
3. Basis for Future Discoveries: Mendeleev's classification provided a framework for future discoveries of elements. As new elements were found, they could easily be incorporated into the existing periodic table.
4. Identification of Gaps: By leaving gaps in his periodic table, Mendeleev identified areas where new elements were expected. This encouraged further research and discovery in the field of chemistry.

Limitations:

Despite its success, Mendeleev's classification had some limitations:

1. Isotopes: Mendeleev's periodic table did not account for isotopes, which are atoms of the same element with different atomic masses. Isotopes were placed in the same position in the table, even though they had different masses and, in some cases, different properties.
2. Wrong Order of Atomic Masses: To maintain chemical similarities within groups, Mendeleev occasionally placed elements in the wrong order of atomic masses. For example, he put cobalt (with a higher atomic mass) before nickel (with a lower atomic mass) because of their similar properties.
3. Hydrogen's Placement: The placement of hydrogen in the periodic table remained a challenge. It exhibits properties similar to both alkali metals and halogens, making its exact position a subject of debate.
4. Lack of Atomic Number: Mendeleev's classification was based on atomic masses, not atomic numbers. The modern periodic table, based on atomic number, resolved many of the anomalies present in Mendeleev's classification.

Modern Periodic Law

The Modern Periodic Law is a fundamental principle in chemistry that states:

"The properties of elements are a periodic function of their atomic numbers."

Key Points:

1. Atomic Numbers: The atomic number is a unique identifier for each chemical element and represents the number of protons in the nucleus of an atom. It determines the identity of an element.
2. Periodic Function: The term "periodic" means that certain properties of elements exhibit regular patterns or repetitions as you move across or down the periodic table.
3. Arrangement by Atomic Number: Elements are arranged in the periodic table in increasing order of their atomic numbers. This arrangement ensures that elements with similar chemical properties are grouped together.
4. Periods and Groups: The periodic table is organised into rows called "periods" and columns called "groups" or "families." Elements in the same group share similar chemical properties due to having the same number of valence electrons, which determines their reactivity.
5. Electronic Configurations: The arrangement of electrons in an atom's electron shells (electron configuration) plays a crucial role in determining an element's chemical behaviour. Elements in the same group often have similar electron configurations, leading to similarities in their chemical properties.
6. Exceptions and Anomalies: While the periodic table generally follows periodic trends, there can be exceptions and anomalies due to factors such as electron shielding, nuclear charge, and orbital shapes. These exceptions are typically found in the transition metals and other regions of the periodic table.

Modern Periodic Table

The Modern Periodic Table is a visual representation of the organisation of chemical elements based on the Modern Periodic Law. It is a tabular arrangement of elements that provides information about each element.

Vertical Columns (Groups):

1. The periodic table is divided into 18 vertical columns known as groups or families.
2. Elements in a group are assigned a group number.
3. Elements in the same group share similar outer shell configurations, which leads to similarities in their chemical properties.
4. Elements in the same group have the same number of valence electrons, which are the electrons in the outermost shell of an atom.
5. Elements with up to two valence electrons have a group number equal to their valence electrons.
6. Elements with more than two valence electrons have a group number equal to their valence electrons plus 10.

Number of Shells and Valence Electrons:

1. Elements in the same group have similar numbers of electron shells, but their valence electrons may vary.
2. Valence electrons are the electrons in the outermost shell of an atom and play a crucial role in determining chemical behaviour.

Horizontal Rows (Periods):

1. The periodic table has seven horizontal rows called periods.
2. Elements in a period do not have the same number of valence electrons but do contain the same number of electron shells.
3. The period number of an element corresponds to the number of electron shells in its atom.

Changing Chemical Properties Across a Period:

1. As you move from left to right along a period, the chemical properties of elements change.
2. This change occurs because the number of valence shell electrons increases by one unit from one element to the next.

Maximum Electron Capacity in Shells:

1. The maximum number of electrons that can occupy a shell is determined by the formula 2n², where 'n' represents the shell number, counting from the nucleus.
2. For example, the first shell (K-shell) can hold a maximum of 2 electrons, while the second shell (L-shell) can hold a maximum of 8 electrons.

Block Classification:

1. Elements are classified into different blocks, including s-block, p-block, d-block, and f-block, based on the subshells in which their valence electrons are found.

Periodic Trends:

1. The periodic table provides a visual representation of periodic trends such as atomic size, electronegativity, and metallic character.
2. These trends are evident when moving across periods or down groups.

Explanation of Anomalies:

1. The Modern Periodic Table accounts for several anomalies or exceptions, including:

• Hydrogen: Hydrogen holds a unique position in the periodic table because it does not fit neatly into any one group due to its unique properties. It has one valence electron.
• Isotopes: Isotopes of the same element are placed in the same group of the periodic table. This is because isotopes have the same number of protons (same atomic number) but different numbers of neutrons.

Trends in the Modern Periodic Table

1. Valence electrons

Valence electrons are the electrons found in the outermost energy level (shell) of an atom. They are the electrons involved in chemical reactions and bonding.

Variation Down a Group:

1. As you move down a group from the top to the bottom, the number of valence electrons remains the same for all the elements within that group.
2. For example, in Group 1 (alkali metals) of the periodic table, elements like lithium (Li), sodium (Na), and potassium (K) all have 1 valence electron each in their outermost shell.

Variation Along a Period:

1. The number of valence electrons in elements within a period increases as you move from left to right across that period.
2. So, as you move from left to right across the third period, the number of valence electrons increases from 1 to 8.

2. Valency

Valency is a chemical property that describes the combining capacity of an element. It determines how many other atoms an element can bond with to form compounds. Valency is usually expressed as a positive or negative integer, representing the number of electrons an atom can gain, lose, or share when it forms chemical bonds. Understanding valency is crucial in predicting how elements will react with each other to create various compounds.

Valency based on Valence Electrons:

1. If an element has 1, 2, 3, or 4 valence electrons, its valency is equal to the number of valence electrons. For example, elements in Group 1 (like hydrogen and sodium) have one valence electron, so their valency is 1.
2. If an element has 5, 6, 7, or 8 valence electrons, its valency is calculated as 8 minus the number of valence electrons. For example, elements in Group 16 (like oxygen and sulphur) have 6 valence electrons, so their valency is 8 - 6 = 2.

Variation Down a Group:

1. Within a group (vertical column), all the elements have the same number of valence electrons, which means they all have the same valency. For instance, all the elements in Group 1 have one valence electron and a valency of 1.

Variation Along a Period:

As you move along a period (horizontal row) from left to right, the number of valence electrons in the elements increases by one unit with each element.

The valency follows a trend along the period:

1. It starts at 1 for elements in Group 1 (Period 1) because they have one valence electron.
2. It increases to 2 for elements in Group 2 (Period 2) because they have two valence electrons.
3. It further increases to 3 for elements in Group 13 (Period 3) and 4 for elements in Group 14 (Period 4).
4. Then, it starts decreasing back to 3 for elements in Group 15 (Period 5), 2 for elements in Group 16 (Period 6), 1 for elements in Group 17 (Period 7), and finally, 0 for elements in Group 18 (Period 8).

3. Atomic Size

Atomic size, also known as atomic radius, refers to the physical size of an atom. It is a measure of the distance from the nucleus of an atom to its outermost electron shell, or in simpler terms, it tells us how large an atom is. The atomic size is typically described in picometres (pm), where 1 picometre is equal to one trillionth of a metre (10^-12 metres).

Variation Down a Group:

1. When you move down a group (vertical column) in the periodic table, the atomic size or atomic radius increases.
2. This trend occurs because, as you move down a group, a new energy level (shell) is added for each element.
3. Each additional energy level further away from the nucleus increases the atomic size.
4. The outermost electrons are located in higher energy levels, and they are farther from the nucleus, which weakens the attraction between the nucleus and electrons.
5. As a result, the electrons are held less tightly, and the atomic size increases.
6. For example, in Group 1 (alkali metals) of the periodic table, as you move down from lithium (Li) to potassium (K) and then to caesium (Cs), the atomic radius increases. Cesium (Cs) has the largest atomic radius in this group.

Variation Along a Period:

1. When you move from left to right across a period (horizontal row) in the periodic table, the atomic size or atomic radius decreases.
2. This trend occurs because, as you move across a period, the number of protons in the nucleus increases (due to increasing atomic number).
3. The increased positive charge in the nucleus exerts a stronger attraction on the electrons in the outermost energy level (valence electrons), pulling them closer to the nucleus.
4. As a result, the electrons are held more tightly, and the atomic size decreases.
5. For example, consider the third period (row) of the periodic table, which includes sodium (Na), magnesium (Mg), aluminium (Al), silicon (Si), phosphorus (P), sulphur (S), chlorine (Cl), and argon (Ar). In this period, as you move from left (Na) to right (Cl), the atomic radius decreases. Sodium (Na) has the largest atomic radius in this period, while chlorine (Cl) has the smallest atomic radius.

4. Metallic Character

The tendency of elements to lose electrons is called metallic character. Metals are found on the left side and centre of the periodic table. They tend to lose electrons to form positive ions. Metals typically have 1, 2, or 3 electrons in their valence shells.

Variation Down a Group:

1. When you move down a group (column) in the periodic table, the metallic character of elements increases.
2. The top of a group typically contains non-metals, while the bottom contains the most metallic elements.
3. For example, in Group 1 (alkali metals), lithium (Li) is the least metallic, while francium (Fr) is the most metallic.

Variation Along a Period:

1. When you move from left to right across a period (row) in the periodic table, the metallic character of elements decreases, while the non-metallic character increases.
2. The left side of a period consists of metals, the middle contains metalloids, and the right side contains non-metals.
3. For example, in the third period (Na, Mg, Al, Si, P, S, Cl), sodium (Na), magnesium (Mg), and aluminium (Al) are metals, silicon (Si) is a metalloid, and phosphorus (P), sulphur (S), and chlorine (Cl) are non-metals.
4. The most metallic element in a period is found on the far left, while the most non-metallic element is found on the far right. In this example, sodium (Na) is the most metallic, and chlorine (Cl) is the most non-metallic.

5. Non-Metallic Character

The tendency of an element to gain electrons is referred to as the non-metallic character. Non-metals are typically found on the right side of the periodic table. They tend to gain electrons to form negative ions. Non-metals generally have 4 to 8 electrons in their valence shells.

Variation Down a Group:

1. Non-metallic character decreases down a group (column) in the periodic table.
2. This decrease occurs because, as you move down a group, the distance between the nucleus and the valence electron increases.
3. With a greater distance, the effective nuclear pull experienced by the valence electron decreases, reducing the atom's tendency to gain an additional electron in its valence shell.

Variation Along a Period:

1. Non-metallic character increases along a period (row) in the periodic table.
2. This increase happens because, as you move across a period, the effective nuclear charge acting on the valence electron increases.
3. The increasing nuclear charge enhances the atom' s tendency to gain electrons in the valence shell, thus increasing its non-metallic character.

6. Electronegativity

Electronegativity is the measure of an element's tendency to attract the shared pair of electrons towards itself in a covalently bonded molecule.

Variation Down a Group:

1. Electronegativity decreases down a group (column) in the periodic table.
2. This decrease occurs because, as you move down a group, the non-metallic character of elements decreases.
3. Elements in lower rows of the periodic table are generally less electronegative.

Variation Along a Period:

1. Electronegativity increases along a period (row) in the periodic table.
2. This increase happens because, as you move across a period, the non-metallic character of elements increases.
3. Elements on the right side of the periodic table (non-metals) are typically more electronegative than those on the left side (metals).

7. Chemical reactivity

Chemical reactivity refers to the tendency of an element or substance to undergo chemical reactions and form new compounds when it comes into contact with other substances. It is a fundamental property that helps us understand how elements interact and combine with one another.

Variation Down a Group:

1. In a group (column) of the periodic table, elements have similar electronic configurations (the same number of valence electrons), leading to similar chemical properties within the group.
2. The chemical reactivity of metals tends to increase as you move down a group. This is because the size of metal atoms increases down the group, making it easier for valence electrons to be removed and participate in chemical reactions.
3. Conversely, the chemical reactivity of non-metals decreases as you move down a group. The larger size of non-metal atoms reduces their ability to attract and gain electrons in chemical reactions.

Variation Along a Period:

1. In a period (row) of the periodic table, the chemical reactivity of elements often exhibits a pattern where it first decreases and then increases.
2. As you move from left to right in a period, the chemical reactivity of elements generally decreases initially and then increases.
3. For example, in the third period, sodium (Na) is very reactive because it has only one valence electron, which it can easily lose. As you move to magnesium (Mg), which has two valence electrons, its reactivity decreases. Aluminium (Al) and silicon (Si) have even more valence electrons, making them less reactive.
4. On the non-metal side, elements like phosphorus (P), sulphur (S), and chlorine (Cl) in the third period show increasing reactivity from left to right because they need to gain electrons to complete their outermost electron shells, and it becomes progressively easier to do so.

8. Nature of Oxides

The nature of oxides refers to whether oxides of an element are acidic, basic, or amphoteric (exhibiting both acidic and basic properties) when they react with water. This property can be used to understand how the oxides of elements behave chemically. Here are the trends in the nature of oxides based on their position in the periodic table:

Variation Along a Period (Top to Bottom):

1. When you move down in a group of the periodic table, there is generally no significant change in the nature of oxides of elements within the same group.
2. Elements within a group have similar chemical properties due to having the same number of valence electrons. For example, all the elements in Group 1 (e.g., Li, Na, K) form basic oxides, and all the elements in Group 17 (e.g., F, Cl, Br) form acidic oxides.

Variation Along a Period (Left to Right):

On moving from left to right in a period of the periodic table, the basic nature of oxides decreases, and the acidic nature of oxides increases.

1. Basic Oxides: Elements on the left side of the periodic table (metals) typically form basic oxides. These oxides react with water to produce alkaline (basic) solutions. For example, sodium oxide (Na₂O) is highly basic.
2. Acidic Oxides: Elements on the right side of the periodic table (non-metals) generally form acidic oxides. These oxides react with water to produce acidic solutions. For example, sulphur dioxide (SO₂) and chlorine dioxide (Cl₂O₇) are acidic oxides.
3. Amphoteric Oxides: Some elements, especially those close to the dividing line between metals and nonmetals, form amphoteric oxides. These oxides can exhibit both acidic and basic properties depending on the conditions. For example, aluminium oxide (Al₂O₃) is amphoteric.

Frequently Asked Questions

1. Can Dobereiner's triads be applied to modern elements?

While Dobereiner's idea was important historically, it does not fit well with the modern understanding of the periodic table. However, it highlighted the concept of recurring patterns, which paved the way for more sophisticated classification systems.

2. How does the Modern Periodic Law differ from Mendeleev's Periodic Law?

Mendeleev's Periodic Law was based on atomic mass, while the Modern Periodic Law is based on atomic number. The atomic number is more fundamental because it defines the element and its properties, whereas atomic mass does not always follow a strict periodic pattern.

3. Why are noble gases placed in Group 18?

Noble gases are placed in Group 18 because they have completely filled outer electron shells, making them highly stable and chemically inert. Their lack of reactivity distinguishes them from other groups.

4. How does the Periodic Table help in predicting chemical reactions?

The tendency of elements to receive, lose, or share electrons is illustrated by the Periodic Table, which helps anticipate chemical reactions. Because they have the same valence electron configuration, elements in the same group frequently show comparable reactions.

5. What is the trend for electronegativity in the periodic table?

Electronegativity increases as you move from left to right across a period and decreases as you move down a group. Elements in the upper right corner of the table (like fluorine) have the highest electronegativity because they attract electrons strongly.