PERIODIC CLASSIFICATION

 

PERIODIC   CLASSIFICATION

1. Basic Building Blocks: Elements serve as the fundamental building blocks of matter. Understanding their properties and behavior is essential for comprehending the properties and reactions of all substances in the universe.
2. Increasing Number of Elements: Over time, the number of known elements has grown significantly. In 1800, only 31 elements were known, but by 1865, this number had more than doubled to 63. Today, we have identified 114 elements, with ongoing efforts to create new ones.
3. Complexity of Individual Study: Investigating each element and its countless compounds individually is a daunting task due to the sheer number of elements and the numerous possible combinations. This complexity makes it impractical to study each element in isolation.
4. Systematic Organization: Scientists recognized the need for a systematic way to organize their knowledge about elements. This led to the development of the periodic table, which arranges elements in a structured manner, simplifying the study of their properties and relationships.
5. Grouping by Similarities: The periodic table groups elements with similar properties together. Elements within the same group often share common characteristics, which makes it easier to understand their behavior and predict their reactions.
6. Predictive Power: The periodic table's arrangement enables scientists to predict the properties of unknown elements based on the trends and patterns observed among neighboring elements. This predictive capability aids in the discovery and study of new elements.
7. Rationalizing Chemical Facts: The periodic table not only organizes existing knowledge about elements but also rationalizes known chemical facts. It provides a framework for explaining why certain elements exhibit specific behaviors and reactions.
8. Encouraging Further Study: By highlighting gaps and patterns in the periodic table, it encourages scientists to explore and study elements that may be missing or poorly understood. This drives further research and exploration in the field of chemistry.
9. Education and Communication: The periodic table is an invaluable tool for teaching and learning chemistry. It simplifies the subject by offering a visual representation of the relationships between elements, making it accessible to students and educators.
10. Scientific Advancements: The classification of elements through the periodic table has played a crucial role in advancing the field of chemistry. It has guided research, inspired new theories, and led to the development of innovative technologies and materials.

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Development of the Periodic Table and the contributions of various scientists, including Johann Dobereiner, A.E.B. de Chancourtois, John Alexander Newlands, Dmitri Mendeleev, and Lothar Meyer:

1. Johann Dobereiner (Early 1800s):
• Dobereiner was the first to suggest the idea of trends among the properties of elements.
• He noticed similarities among the physical and chemical properties of groups of three elements called "Triads."
• In each Triad, the middle element had an atomic weight approximately halfway between the other two.
• His idea, known as the "Law of Triads," was dismissed as a coincidence since it only applied to a few elements.
2. A.E.B. de Chancourtois (1862):
• A French geologist, he arranged known elements in order of increasing atomic weights and created a cylindrical table to show the periodic recurrence of properties.
• His work did not gain much attention at the time.
3. John Alexander Newlands (1865):
• Newlands proposed the "Law of Octaves," arranging elements in order of increasing atomic weights.
• He noticed that every eighth element had properties similar to the first element, akin to musical octaves.
• Newlands' law worked for elements only up to calcium and was not widely accepted initially.
• However, he was later awarded the Davy Medal in 1887 by the Royal Society, London, for his work.
4. Dmitri Mendeleev and Lothar Meyer (1869):
• Working independently, both Mendeleev and Meyer proposed that when elements are arranged by increasing atomic weights, similarities in properties occur at regular intervals.
• Meyer plotted physical properties like atomic volume, melting point, and boiling point against atomic weight and observed a periodically repeating pattern.
• Unlike Newlands, Meyer noted changes in the length of this repeating pattern.
• Mendeleev, a Russian chemist, is generally credited with developing the Modern Periodic Table.
• Mendeleev arranged elements in horizontal rows (periods) and vertical columns (groups) in order of increasing atomic weights.
• Elements with similar properties occupied the same vertical column.
• Mendeleev's system was more elaborate and relied on a broader range of physical and chemical properties.
• He occasionally ignored strict atomic weight order to maintain similarities in properties within groups.
• Mendeleev left gaps in the table for undiscovered elements and successfully predicted the existence and properties of elements like gallium and germanium.
• His bold quantitative predictions and their eventual success made him and his Periodic Table famous.
5. Mendeleev's Periodic Table (1905):
• Mendeleev's Periodic Table, published in 1905, is considered a milestone in the development of the Periodic Table.

 

 

Development of the Modern Periodic Law and the present form of the Periodic Table:

1. Background of Early Periodic Table Development:
• Mendeleev developed his Periodic Table without knowledge of the internal structure of the atom.
• At the start of the 20th century, significant advancements in subatomic particle theories began to emerge.
2. Henry Moseley's Contribution (1913):
• English physicist Henry Moseley observed regularities in the X-ray spectra of elements.
• Plotting the square root of the frequency of X-rays (ν) against the atomic number (Z) resulted in a straight line, rather than a plot against atomic mass.
• Moseley's findings established that the atomic number (number of protons or electrons in a neutral atom) is a more fundamental property than atomic mass.
• This discovery led to the modification of Mendeleev's Periodic Law, giving rise to the Modern Periodic Law:
• "The physical and chemical properties of the elements are periodic functions of their atomic numbers."
3. Significance of Atomic Number and Quantum Numbers:
• Atomic number, being equivalent to the nuclear charge or the number of electrons in a neutral atom, became crucial in understanding the periodicity of elements.
• Quantum numbers and electronic configurations play a pivotal role in explaining the periodic variation in the properties of elements.
4. Development of Various Forms of Periodic Tables:
• Over time, numerous forms of the Periodic Table have been devised.
• Some emphasize chemical reactions and valence, while others stress the electronic configuration of elements.
• The "long form" of the Periodic Table is the most widely used today.
5. Elements in the Modern Periodic Table:
• In the Modern Periodic Table, horizontal rows are referred to as "periods," and vertical columns as "groups" or "families."
• Elements with similar outer electronic configurations are grouped in vertical columns.
• The International Union of Pure and Applied Chemistry (IUPAC) recommends numbering the groups from 1 to 18, replacing the older notations.
6. Number of Periods and Elements:
• There are a total of seven periods in the Periodic Table.
• The period number corresponds to the highest principal quantum number (n) of the elements within that period.
• The first period contains 2 elements, while the subsequent periods contain 8, 8, 18, 18, and 32 elements, respectively.
• The seventh period is incomplete and, like the sixth period, would theoretically contain a maximum of 32 elements based on quantum numbers.
7. Lanthanoids and Actinoids:
• In the Modern Periodic Table, elements from both the sixth and seventh periods, known as lanthanoids and actinoids, respectively, are placed in separate panels at the bottom of the table.

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Nomenclature of elements with atomic numbers greater than 100:

1. Traditional Naming Privileges:
• Historically, the privilege of naming newly discovered elements rested with the discoverer or discoverers.
• The suggested names for new elements were ratified by the International Union of Pure and Applied Chemistry (IUPAC).
2. Controversy Surrounding Naming:
• New elements with very high atomic numbers are highly unstable and often exist in extremely small quantities, sometimes just a few atoms.
• The synthesis and characterization of these elements require sophisticated and costly equipment and laboratories.
• Due to the competitive nature of scientific research, disputes can arise, with different scientists and laboratories claiming credit for the discovery of the same element.
• For instance, both American and Soviet scientists claimed to have discovered element 104, leading to naming conflicts (Rutherfordium vs. Kurchatovium).
3. IUPAC's Recommendation for Systematic Nomenclature:
• To address naming disputes and controversies, IUPAC made a recommendation that, until a new element's discovery is fully proven and its name officially recognized, a systematic nomenclature should be used.
• This systematic nomenclature is derived directly from the element's atomic number, utilizing numerical roots for zero and numbers one to nine.
4. Numerical Roots and Nomenclature for Elements with Z > 100:
• The numerical roots for elements with atomic numbers above 100 are shown in Table

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• These roots are combined in the order of the digits that make up the atomic number.
• The suffix "ium" is added to these roots to form a temporary name for the new element.
4. Temporary Naming and Symbol:
• Initially, the new element is given a temporary name, often represented by a symbol consisting of three letters.
5. Permanent Naming Process:
• Later, a permanent name and symbol are assigned through a formal vote involving IUPAC representatives from different countries.
• The permanent name might reflect the country or state where the element was discovered or pay tribute to a notable scientist.
6. Current Status:
• As of now, elements with atomic numbers up to 118 have been discovered.
• IUPAC has officially announced the names of all these elements, following the systematic nomenclature and permanent naming process.

 

Electronic configurations of elements and the long form of the Periodic Table:

1. Electronic Configurations Defined by Quantum Numbers:
• An electron in an atom is defined by a set of four quantum numbers.
• The principal quantum number (n) determines the main energy level or shell in which the electron is located.
• Electrons are distributed into different subshells, often referred to as orbitals (s, p, d, f) within an atom.
• The arrangement of electrons into these orbitals is known as the electronic configuration of the element.
2. Periods Reflect Valence Shell Energy Levels:
• In the Periodic Table, each period corresponds to the value of n for the outermost or valence shell.
• Successive periods are associated with the filling of the next higher principal energy level (n = 1, n = 2, and so on).
3. Number of Elements in Each Period:
• The number of elements in each period is twice the number of atomic orbitals available in the energy level being filled.
• The first period (n = 1) begins with the filling of the lowest level (1s) and contains two elements: hydrogen (1s^1) and helium (1s^2).
• The second period (n = 2) starts with lithium, and the third electron enters the 2s orbital, followed by the filling of 2p orbitals, resulting in 8 elements.
• The third period (n = 3) starts at sodium, and the added electron enters a 3s orbital. The filling of 3s and 3p orbitals gives 8 elements.
• The fourth period (n = 4) starts at potassium, and before the 4p orbital is filled, the filling of 3d orbitals becomes energetically favorable. This leads to the 3d transition series starting at scandium (Z = 21) and ending at zinc (Z = 30). The fourth period contains 18 elements.
• The fifth period (n = 5) is similar to the fourth, with the addition of the 4d transition series, beginning at yttrium (Z = 39). The period ends at xenon, with 18 elements.
• The sixth period (n = 6) consists of 32 elements and includes the filling of 6s, 4f, 5d, and 6p orbitals. The 4f-inner transition series starts at cerium (Z = 58) and ends at lutetium (Z = 71), known as the lanthanoid series.
• The seventh period (n = 7) resembles the sixth, with the filling of 7s, 5f, 6d, and 7p orbitals. This period includes most of the man-made radioactive elements and will end with the element having atomic number 118, belonging to the noble gas family.
• The filling of 5f orbitals starts after actinium (Z = 89) and gives rise to the 5f-inner transition series known as the actinoid series.
4. Separate Placement of Inner Transition Series:
• The 4f and 5f-inner transition series of elements (lanthanoids and actinoids) are placed separately in the Periodic Table to maintain its structure and classification principles, preserving elements with similar properties within a single column.

 

1. Groupwise Electronic Configurations:
• Elements within the same vertical column or group in the Periodic Table share similar valence shell electronic configurations.
• A group is characterized by having the same number of electrons in the outermost orbitals, leading to similar chemical properties among group members.
2. Example: Group 1 Elements (Alkali Metals):
• Group 1 elements, known as alkali metals, all possess ns^1 valence shell electronic configurations.
• The ns¹ configuration signifies that these elements have one electron in their outermost s orbital.
• The electronic configuration of alkali metals can be expressed as follows:
• Li (Lithium): 1s²2s¹
• Na (Sodium): 1s²2s²2p⁶3s¹
• K (Potassium): 1s²2s²2p⁶3s²3p⁶4s¹
• Cs (Cesium): 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s¹
• These elements exhibit similar chemical behavior due to their common valence shell configuration, such as readily donating their outermost electron to form cations with a +1 charge.
3. Periodic Dependence on Atomic Number:
• The properties of elements show a periodic dependence on their atomic number, which reflects the number of protons (and electrons) in the nucleus.
• This periodicity in properties is a fundamental principle of the Periodic Table and is not based on the relative atomic mass of elements.

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