PERIODIC CLASSIFICATION : POINT WISE (CBSE) CLASS XI

 

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.

DEVELOPMENT  OF  THE  PERIODIC  TABLE

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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.
• DOBEREIGNER’S TRIAD

ELEMENTS	ATOMIC WEIGHT	ELEMENTS	ATOMIC WEIGHT	ELEMENTS	ATOMIC WEIGHTS
Li	7	Ca	40	Cl	35.5
Na	23	Sr	88	Br	80
K	39	Ba	137	I	127

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1. 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.
2. 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.
• NEWLAND’S   OCTAVES

ELEMENT	Li	Be	B	C	N	O	F
At. Wt.	7	9	11	12	14	16	19
ELEMENT	Na	Mg	Al	Si	P	S	Cl
At. Wt.	23	24	27	29	31	32	35.5
ELEMENT	K	Ca
At. Wt.	39	40

 

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:

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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.

 

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

Notation for IUPAC Nomenclature of Elements

DIGIT	NAME	Abbreviation	DIGIT	NAME	Abbreviation
0	ni	n	5	pent	p
1	un	u	6	hex	h
2	bi	b	7	sept	s
3	tri	t	8	oct	o
4	quad	q	9	enn	e

 

• 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 (1s¹) and helium (1s²).
• 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.

ELECTRONIC  CONFIGURATIONS OF   ELEMENTS

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.

 

1. Theoretical Foundation - Aufbau Principle:
• The aufbau principle, based on the electronic configuration of atoms, provides the theoretical foundation for the periodic classification of elements.
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2. Chemical Behavior and Groups/Families:
• Elements within the same vertical column or group in the Periodic Table exhibit similar chemical behavior.
• This similarity arises because elements in a group have the same number and distribution of electrons in their outermost orbitals.
3. Classification into Four Blocks:
• Elements are classified into four blocks based on the types of atomic orbitals that are being filled with electrons.
• The four blocks are the s-block, p-block, d-block, and f-block.
4. Exceptions to Categorization:
• There are two exceptions to this categorization:
• Helium, strictly belonging to the s-block, is positioned in the p-block along with other group 18 elements because it has a completely filled valence shell (1s^2) and exhibits properties characteristic of noble gases.
• Hydrogen, with only one s-electron, can be placed in group 1 (alkali metals). It can also gain an electron to achieve a noble gas arrangement and behave similarly to group 17 (halogen family) elements.
5. Positioning of Hydrogen and Helium:
• Hydrogen is placed separately at the top of the Periodic Table due to its unique properties and behavior.
• Helium is also included in the p-block with other noble gases due to its filled valence shell.
6. Brief Overview of Element Types:
• s-block Elements: These elements are found in the s-block of the Periodic Table and typically have their outermost electrons filling the s-orbitals. They include alkali metals and alkaline earth metals.
• p-block Elements: Elements in this block have their outermost electrons filling p-orbitals. They encompass a wide range of elements, including nonmetals, metalloids, and some metals.
• d-block Elements: This block comprises transition metals, which have their outermost electrons filling d-orbitals. Transition metals are known for their characteristic properties, including variable oxidation states and the ability to form colored compounds.
• f-block Elements: Elements in this block include the lanthanides and actinides, which have their outermost electrons filling f-orbitals. They are often referred to as inner transition metals.

 

Characteristics of s-block elements:

1. Classification of s-Block Elements:
• The s-block elements include Group 1 (alkali metals) and Group 2 (alkaline earth metals) elements.
• They share a common outermost electronic configuration, which is either ns¹ for alkali metals or ns² for alkaline earth metals, where "n" represents the principal quantum number of the valence shell.
2. Reactive Metals:
• All s-block elements are highly reactive metals with low ionization enthalpies.
• They readily lose their outermost electron(s) to form ions, resulting in a 1+ ion for alkali metals and a 2+ ion for alkaline earth metals.
• Their reactivity and metallic character increase as you move down the group.
3. Absence in Pure Form in Nature:
• Due to their high reactivity, s-block elements are rarely found in their pure, uncombined form in nature.
• Instead, they are typically found as compounds or minerals.
4. Predominantly Ionic Compounds:
• Compounds formed by s-block elements, with the exception of those containing lithium and beryllium, are predominantly ionic in nature.
• Ionic compounds are characterized by the transfer of electrons from the metal (s-block element) to the non-metal, resulting in the formation of ions held together by electrostatic forces.

s-block elements, encompassing alkali metals and alkaline earth metals, share common characteristics such as high reactivity, metallic character, and the tendency to form ions. They readily lose their outermost electrons to achieve stable electron configurations, resulting in the formation of predominantly ionic compounds.

 

Characteristics of p-block elements:

1. p-Block Elements and Main Group Elements:
• The p-Block Elements encompass elements belonging to Groups 13 to 18 in the Periodic Table.
• Together with the s-Block Elements, they are referred to as the Representative Elements or Main Group Elements.
2. Variation in Outermost Electronic Configuration:
• Within the p-block, the outermost electronic configuration varies from ns² np¹ to ns² np⁶ in each period.
• At the end of each period in the p-block is a noble gas element with a closed valence shell ns² np⁶ configuration.
3. Stability of Noble Gases:
• Noble gases, found at the end of each p-block period, have completely filled valence shell orbitals with ns² np⁶ electrons.
• These noble gases exhibit very low chemical reactivity due to the stability of their electron configuration.
4. Halogens and Chalcogens:
• Preceding the noble gas family in the p-block are two chemically important groups of nonmetals: the halogens (Group 17) and the chalcogens (Group 16).
• These groups have highly negative electron gain enthalpies, making them readily accept one or two electrons, respectively, to achieve a stable noble gas configuration.
5. Trends in Chemical Behavior:
• As you move from left to right across a period in the p-block, the non-metallic character of elements increases.
• Conversely, as you go down a group in the p-block, the metallic character of elements increases.

The p-Block Elements, which include the Main Group Elements, display a wide range of chemical behaviors. Elements in this block have varying outermost electronic configurations, leading to different reactivity patterns. Noble gases at the end of each period are extremely stable due to their fully filled valence shell orbitals. Preceding them, the halogens and chalcogens are notable for their electron acceptance tendencies and increasing non-metallic character across a period.

 

 

Characteristics of d-block elements (transition elements):

1. Group Range and Inner d Orbitals:
• The d-Block Elements, also known as Transition Elements, include elements from Group 3 to 12, located in the central part of the Periodic Table.
• These elements are characterized by the filling of inner d orbitals by electrons, hence the name d-Block Elements.
2. General Outer Electronic Configuration:
• Transition elements share a common general outer electronic configuration of (n-1)d^1-10ns^0-2, where "n" represents the principal quantum number of the valence shell.
• Notably, Pd has the electronic configuration 4d^10 5s^0.
3. Metallic Nature:
• All transition elements are metals, characterized by properties such as malleability, ductility, and electrical conductivity.
4. Colorful Ions and Variable Valence States:
• Transition elements typically form colorful ions in their compounds due to the presence of partially filled d orbitals.
• They exhibit variable valence (oxidation) states, meaning they can exist in different states of charge depending on the chemical reaction.
5. Paramagnetism:
• Many transition elements are paramagnetic, meaning they have unpaired electrons in their electron configurations, making them attracted to a magnetic field.
6. Catalytic Properties:
• Transition elements are often used as catalysts in chemical reactions due to their ability to provide an alternate reaction pathway with lower activation energy.
7. Exceptions: Zn, Cd, and Hg:
• Zn (zinc), Cd (cadmium), and Hg (mercury) have the electronic configuration (n-1)d^10ns^2, which differs from other transition elements.
• These three elements do not exhibit many of the typical properties of transition elements, and they behave more like the s-block elements.
8. Bridging Role:
• Transition metals serve as a bridge between the highly reactive metals of the s-block elements and the less reactive elements of Groups 13 and 14.
• This role earned them the name "Transition Elements."

The d-Block Elements, or Transition Elements, are characterized by the filling of inner d orbitals, leading to their distinctive properties. They are typically metals, known for forming colorful ions, displaying variable valence states, and often acting as catalysts. Exceptions include Zn, Cd, and Hg, which have a different electronic configuration and behave more like s-block elements. Transition elements play a vital role in chemistry due to their diverse properties and applications.

 

Characteristics of f-block elements (inner transition elements):

 

1. Lanthanoids and Actinoids:
• The f-Block Elements consist of two rows of elements located at the bottom of the Periodic Table.
• The first row is known as the Lanthanoids, ranging from Ce (Z = 58) to Lu (Z = 71).
• The second row is called the Actinoids, spanning from Th (Z = 90) to Lr (Z = 103).
2. Outer Electronic Configuration:
• These elements are characterized by an outer electronic configuration of (n-2)f^1-14 (n-1)d^0–1 ns².
• The last electron added to each element is filled into the f-orbital, hence the name f-Block Elements.
3. Metals:
• All f-Block Elements are metals, exhibiting typical metallic properties such as electrical conductivity and malleability.
4. Similar Properties within Series:
• Within each series (lanthanoids or actinoids), the properties of the elements are quite similar due to the gradual filling of the f-orbitals.
5. Complex Chemistry of Actinoids:
• The chemistry of actinoid elements, especially the early actinoids, is more complex than that of the lanthanoids.
• Actinoids can exhibit a wide range of oxidation states, making their chemistry intricate and versatile.
6. Radioactive Nature:
• Many of the actinoid elements are radioactive, posing challenges for handling and study.
• These elements often have short half-lives, and some can only be produced in extremely small quantities through nuclear reactions.
7. Transuranium Elements:
• Elements beyond uranium in the actinoid series are referred to as Transuranium Elements.
• These elements have atomic numbers greater than 92 (the atomic number of uranium) and are generally synthetic, created through nuclear reactions.

The f-Block Elements, comprising the Lanthanoids and Actinoids, are characterized by their unique electron configurations and metallic nature. Elements within each series exhibit similar properties. The chemistry of actinoid elements is notably complex due to their ability to adopt various oxidation states. Many of these elements are radioactive, and those beyond uranium are considered Transuranium Elements, created artificially in small quantities.

 

 

METALS,     NON-METALS      AND     METALLOIDS

 

Metals:

1. Abundance: Metals make up more than 78% of all known elements and are primarily found on the left side of the Periodic Table.
2. State at Room Temperature: Typically, metals are solids at room temperature, with the exception of mercury (a liquid), and gallium and caesium (with very low melting points).
3. High Melting and Boiling Points: Metals generally have high melting and boiling points, contributing to their solid state at room temperature.
4. Conductivity: They are excellent conductors of heat and electricity, making them essential materials in electrical and thermal applications.
5. Malleability and Ductility: Metals are malleable, meaning they can be flattened into thin sheets by hammering, and ductile, allowing them to be drawn into wires.

Non-Metals:

1. Location: Non-metals are predominantly found on the top right-hand side of the Periodic Table.
2. State at Room Temperature: They are typically either solids or gases at room temperature, with a few exceptions like boron and carbon.
3. Low Melting and Boiling Points: Non-metals generally have low melting and boiling points, making them exist in various physical states.
4. Poor Conductors: They are poor conductors of both heat and electricity.
5. Brittle Nature: Most non-metallic solids are brittle and lack the malleability and ductility observed in metals.
6. Trend: The metallic character of elements increases as one goes down a group, while non-metallic character intensifies as one moves from left to right across the Periodic Table.

Metalloids (Semi-metals):

1. Location: Metalloids are situated along the thick zig-zag line in the Periodic Table, bordering the transition from metals to non-metals.
2. Properties: These elements, such as silicon, germanium, arsenic, antimony, and tellurium, display properties that exhibit characteristics of both metals and non-metals.
3. Transitional Nature: Metalloids have properties that make them intermediate between metals and non-metals, contributing to their classification as semi-metals.

Elements are broadly classified into Metals and Non-Metals based on their properties. Metals are typically solid, have high melting points, and are good conductors of heat and electricity. Non-Metals, on the other hand, are often either solids or gases, possess lower melting points, and are poor conductors. Metalloids, also known as Semi-metals, exhibit a mix of properties from both categories and are situated along the boundary line between metals and non-metals in the Periodic Table.

 

1. Patterns in Chemical Reactivity:

• Within a Period (Horizontal Row):
• Reactivity tends to be high in Group 1 metals (alkali metals) on the far left.
• It decreases as you move towards the middle of the Periodic Table.
• It then increases again to a maximum in Group 17 non-metals (halogens) on the right.
• Within a Group (Vertical Column):
• For representative metals (e.g., alkali metals), reactivity increases as you move down the group.
• For non-metals (e.g., halogens), reactivity decreases as you move down the group.

2. Explanation in Terms of Atomic Structure:

• Electron Configuration and Energy Levels:
• Periodic trends are explained by the distribution of electrons in an atom's energy levels and orbitals.
• Elements within a group have the same number of valence electrons (outermost electrons), leading to similar chemical behavior.
• As you move across a period, the number of valence electrons increases by one from left to right.
• Effective Nuclear Charge:
• The effective nuclear charge, experienced by valence electrons, increases across a period due to an increase in the number of protons in the nucleus.
• This stronger attraction between the nucleus and valence electrons makes it harder for elements to lose or gain electrons, affecting their reactivity.
• Atomic Size (Atomic Radius):
• Atomic size generally decreases from left to right across a period.
• This is because the increasing positive charge in the nucleus pulls electrons closer to the nucleus, reducing the atomic radius.
• Ionization Energy:
• Ionization energy, the energy required to remove an electron from an atom, generally increases across a period.
• As atomic size decreases, electrons are held more tightly, making it harder to remove them.
• Electronegativity:
• Electronegativity, the ability of an atom to attract electrons in a chemical bond, increases across a period.
• Smaller atoms with higher effective nuclear charges have greater electronegativity.
• Metallic Character:
• Metallic character decreases across a period.
• Metals are on the left side of the Periodic Table, and as you move to the right, elements become less metallic and more non-metallic in character.
• Group Trends:
• In a group, elements have the same valence electron configuration, resulting in similar chemical properties.
• Reactivity trends in a group are largely due to the increase in energy levels as you move down the group, making valence electrons farther from the nucleus.

These periodic trends in properties can be explained by considering the atomic structure, including electron configuration, effective nuclear charge, and atomic size. These factors play a crucial role in determining how elements react and behave within the Periodic Table.

 

 

Periodic trends in physical properties of elements:

1. Atomic Radii:

• Across a Period (Left to Right):
• Atomic radii generally decrease as you move from left to right across a period.
• This is due to the increase in effective nuclear charge, which pulls electrons closer to the nucleus, resulting in smaller atomic sizes.
• Down a Group (Top to Bottom):
• Atomic radii generally increase as you move down a group.
• This is because of the addition of new energy levels (shells) as you descend the group, leading to larger atomic sizes.

2. Ionic Radii:

• Cations (Positive Ions):
• Cations are smaller than their parent atoms.
• When an atom loses electrons to become a cation, the electron-electron repulsion decreases, causing the remaining electrons to be pulled closer to the nucleus.
• Anions (Negative Ions):
• Anions are larger than their parent atoms.
• When an atom gains electrons to become an anion, the increased electron-electron repulsion causes electrons to spread out, leading to larger ionic sizes.

3. Ionization Enthalpy (Ionization Energy):

• Across a Period (Left to Right):
• Ionization enthalpy generally increases as you move from left to right across a period.
• It becomes harder to remove electrons because of the stronger effective nuclear charge.
• Down a Group (Top to Bottom):
• Ionization enthalpy generally decreases as you move down a group.
• Electrons are farther from the nucleus, experiencing less attraction, making them easier to remove.

4. Electron Gain Enthalpy:

• Across a Period (Left to Right):
• Electron gain enthalpy tends to become more negative (more exothermic) as you move from left to right across a period.
• Atoms have a higher affinity for gaining electrons, except for noble gases.
• Down a Group (Top to Bottom):
• Electron gain enthalpy generally becomes less negative (less exothermic) as you move down a group.
• Atoms have a lower affinity for gaining electrons due to increased atomic size.

5. Electronegativity:

• Across a Period (Left to Right):
• Electronegativity generally increases as you move from left to right across a period.
• Atoms become more electronegative due to the increasing effective nuclear charge.
• Down a Group (Top to Bottom):
• Electronegativity generally decreases as you move down a group.
• Atoms become less electronegative as the atomic size increases and valence electrons are farther from the nucleus.

These trends in physical properties provide insights into how the size of atoms, the ease of ionization, electron affinity, and electronegativity change across the Periodic Table, influencing the chemical behavior of elements.

 

PERIODIC  TRENDS  IN  ATOMIC RADIUS

 

1. Atomic Radius Estimation:

• Measuring the size of an individual atom is challenging due to its tiny size (~1.2 Å).
• Atomic size is estimated by measuring the distance between atoms in a covalent molecule (covalent radius) or in a metallic crystal (metallic radius).

2. Covalent Radius for Non-Metals:

• Covalent radius is half the distance between two atoms bound by a single covalent bond in a molecule.
• Example: Chlorine (Cl2) molecule has a bond distance of 198 pm, so the covalent radius of chlorine is 99 pm.

3. Metallic Radius for Metals:

• Metallic radius is half the internuclear distance between metal cores in a metallic crystal.
• Example: In solid copper, the distance between adjacent copper atoms is 256 pm, so the metallic radius of copper is 128 pm.

4. Atomic Radii Measurement:

• Atomic radii can be determined using various methods, including X-ray and spectroscopic techniques.

5. Periodic Trends in Atomic Radius:

• Across a Period (Left to Right):
• Atomic radius generally decreases as you move from left to right across a period.
• This is because the effective nuclear charge increases due to a greater number of protons, leading to stronger attraction of electrons to the nucleus.
• Down a Group (Top to Bottom):
• Atomic radius generally increases as you move down a group.
• This is because electrons occupy higher energy levels (shells) as you descend the group, causing them to be farther from the nucleus.
• Inner energy levels shield outer electrons from the nucleus's pull, contributing to larger atomic sizes.

6. Exception for Noble Gases:

• Noble gases have extremely large atomic radii compared to other elements.
• Their radii are not considered covalent radii but should be compared with van der Waals radii of other elements.

These trends in atomic radius reflect how the size of atoms changes across the Periodic Table due to variations in nuclear charge and energy levels.

 

CONCEPTS  OF  IONIC  RADIUS  AND  ITS  TRENDS

Concept of ionic radius and its trends:

1. Ionic Radius Estimation:

• Ionic radii are determined by measuring the distances between cations and anions in ionic crystals.
• They can be thought of as the sizes of ions formed when electrons are either removed (cation) or added (anion) to an atom.

2. Trends in Ionic Radii:

• Cations (Positive Ions):
• Cations are smaller than their parent atoms.
• This is because they have fewer electrons while maintaining the same nuclear charge (protons).
• The reduced electron-electron repulsion leads to a more compact ionic structure.
• For example, Na+ is smaller than neutral sodium (Na).
• Anions (Negative Ions):
• Anions are larger than their parent atoms.
• Adding one or more electrons results in increased electron-electron repulsion within the electron cloud.
• The increased repulsion outweighs the effective nuclear charge, causing the ion to expand.
• For example, F- is larger than neutral fluorine (F).

3. Isoelectronic Species:

• Isoelectronic species are atoms and ions that have the same number of electrons.
• Despite having the same number of electrons, their ionic radii differ due to varying nuclear charges.
• Greater positive charge (cation) results in a smaller radius, while greater negative charge (anion) leads to a larger radius.
• Example: O^2-, F^-, Na⁺, and Mg²⁺ all have 10 electrons, but their ionic radii vary due to differences in nuclear charge.

These trends in ionic radii help explain the size variations among cations and anions and are important in understanding the properties and behavior of ionic compounds.

 

 

PERIODIC    TRENDS    IN    IONIZATION   ENTHALPY

 

1. Ionization Enthalpy:

• Ionization enthalpy is a quantitative measure of an element's tendency to lose electrons.
• It represents the energy required to remove an electron from an isolated gaseous atom in its ground state.
• The first ionization enthalpy refers to the energy needed to remove the first electron from the atom.
• Ionization enthalpies are always positive, indicating that energy is required to remove electrons.

2. Trends in Ionization Enthalpy:

• Periodic Trend Across Periods :
• Ionization enthalpy generally increases as you move across a period (from left to right).
• This is due to the increasing effective nuclear charge as more protons are added to the nucleus.
• Shielding or screening by inner core electrons remains relatively constant across a period, leading to stronger electron-nucleus attraction.
• As a result, outermost electrons are held more tightly.
• Periodic Trend Down Groups :
• Ionization enthalpy generally decreases as you move down a group (from top to bottom).
• This is because the outermost electron is farther from the nucleus in lower energy levels (higher principal quantum number, n).
• The increased distance and increased shielding by inner electrons reduce the effective nuclear charge on the outermost electron.
• Removal of the outermost electron requires less energy.

3. Factors Influencing Ionization Enthalpy:

• Penetration of Orbitals:
• When comparing elements within the same principal quantum level, s-electrons are more strongly attracted to the nucleus than p-electrons.
• For example, boron (Z = 5) has a slightly lower first ionization enthalpy than beryllium (Z = 4) because the electron removed from boron is a p-electron, which is more shielded from the nucleus than the s-electron in beryllium.
• Electron-Electron Repulsion:
• Anomalies in ionization enthalpy trends can occur due to electron-electron repulsion within atomic orbitals.
• For example, oxygen (Z = 8) has a lower first ionization enthalpy than nitrogen (Z = 7) because oxygen's fourth 2p-electron experiences increased repulsion since it must occupy the same 2p-orbital as the other three 2p-electrons.

These trends in ionization enthalpy are influenced by both the nuclear charge and the electron configuration of the elements, shedding light on their reactivity and behavior in chemical reactions.

 

 

PERIODIC  TRENDS   IN  ELECTRON   GAIN   ENTHALPY

 

1. Electron Gain Enthalpy (Δ_egH):

• Electron gain enthalpy measures the energy change when an electron is added to a neutral gaseous atom (X) to form a negative ion (X⁻).
• It can be represented by the equation: X(g) + e⁻ → X⁻(g).

2. Nature of Electron Gain Enthalpy:

• Electron gain enthalpy can be either endothermic or exothermic, depending on the element.
• For many elements, energy is released (exothermic) when an electron is added to the atom, resulting in negative electron gain enthalpy.
• Group 17 elements (halogens) have highly negative electron gain enthalpies because they can attain stable noble gas electronic configurations by gaining an electron.
• Noble gases have large positive electron gain enthalpies because adding an electron leads to an unstable electronic configuration.

3. Variation Across Periods:

• Electron gain enthalpies generally become more negative (exothermic) as you move from left to right across a period.
• This trend is due to the increasing effective nuclear charge (proton count) across a period.
• Smaller atoms have more negative electron gain enthalpies because the added electron is closer to the positively charged nucleus.

4. Variation Down Groups:

• Electron gain enthalpies tend to become less negative (endothermic) as you move down a group.
• This is because the atomic size increases down a group, and the added electron is farther from the nucleus, requiring more energy.
• Electron gain enthalpies of elements generally follow this trend.

5. Anomalies in Electron Gain Enthalpy:

• There are some exceptions to the general trend, such as the electron gain enthalpy of oxygen (O) or fluorine (F) being less negative than that of the succeeding element in the same period.
• This occurs because when an electron is added to O or F, it enters the smaller n = 2 quantum level, causing significant repulsion from other electrons in the same level.
• In contrast, for the n = 3 quantum level (S or Cl), the added electron occupies a larger region of space, resulting in less electron-electron repulsion and more negative electron gain enthalpy.

These trends in electron gain enthalpy provide insights into an element's ability to gain electrons and form anions, impacting its chemical reactivity.

 

PERIODIC  TRENDS   IN  ELECTRONEGATIVITY

 

1. Electronegativity:

• Electronegativity is a qualitative measure of an atom's ability in a chemical compound to attract shared electrons towards itself.
• It is not a directly measurable quantity but can be quantified using various scales, with the Pauling scale being the most widely used.

2. Pauling Scale:

• Linus Pauling assigned a value of 4.0 to fluorine, considering it to have the greatest electron-attracting ability.
• Electronegativity values for other elements are determined relative to this scale.

3. Variation Across Periods:

• Electronegativity generally increases from left to right across a period (e.g., from lithium to fluorine) in the periodic table.
• This trend is explained by the increasing attraction between the outer electrons and the nucleus as the atomic radius decreases across a period.

4. Variation Down Groups:

• Electronegativity generally decreases as you move down a group (e.g., from fluorine to astatine).
• This trend is associated with the increase in atomic radii down a group, resulting in weaker electron-nucleus attraction.

5. Relationship Between Electronegativity and Atomic Radius:

• Electronegativity and atomic radius exhibit an inverse relationship: as atomic radius increases, electronegativity decreases, and vice versa.

6. Electronegativity and Non-Metallic Properties:

• Non-metallic elements have a strong tendency to gain electrons, and this tendency is directly related to electronegativity.
• The increase in electronegativity across a period corresponds to an increase in non-metallic properties (or a decrease in metallic properties) of elements.
• Conversely, the decrease in electronegativity down a group corresponds to a decrease in non-metallic properties (or an increase in metallic properties) of elements.

Understanding electronegativity helps predict chemical behavior, especially in the formation of ionic and covalent compounds, and provides insights into the relative tendencies of elements to attract electrons.

 

 

PERIODIC  TRENDS   IN  CHEMICAL  PROPERTIES  OF  ELEMENTS

 

The periodic trends in chemical properties of elements, focusing on valence states and anomalies in the second period elements:

1. Valence States:

• Elements in the periodic table exhibit periodicity in the valence states they can achieve.
• Valence states are associated with the number of electrons an element can gain or lose to attain a stable electron configuration.
• The periodic table provides a framework for understanding the possible valence states of elements within each group.

2. Diagonal Relationships:

• Diagonal relationships refer to similarities in properties and behaviors between elements in diagonally related positions in the periodic table.
• For example, lithium (Li) and magnesium (Mg) show some similarities due to their diagonal relationship.

3. Inert Pair Effect:

• The inert pair effect is observed in heavier elements, particularly in the p-block elements.
• It refers to the tendency of some elements to preferentially lose or share the s-electrons in their outermost electron shell, leaving the p-electrons unaltered.
• This results in the formation of ions with a lower oxidation state than expected based on the group number.

4. Effects of Lanthanoid Contraction:

• Lanthanoid contraction is a phenomenon observed in the transition metals, specifically in the lanthanide series.
• It involves a decrease in atomic and ionic radii across the lanthanide series due to the poor shielding of inner electrons.
• This contraction can affect the chemical properties and reactivity of elements in this series.

5. Anomalies in Second Period Elements:

• Second period elements from lithium (Li) to fluorine (F) exhibit several anomalies in their properties compared to elements in other periods.
• These anomalies are often attributed to the small size of second period atoms and their high effective nuclear charge.
• Examples include the relatively low ionization enthalpy of boron compared to beryllium and the electron-electron repulsion in oxygen compared to nitrogen.

These trends and phenomena in chemical properties provide valuable insights into the behavior of elements and their compounds, helping chemists understand and predict the outcomes of chemical reactions.

PERIODICITY  OF   VALENCE  OR   OXIDATION   STATES

1. Valence and Electronic Configurations:

• Valence refers to the most characteristic property of elements, and it can be understood in terms of their electronic configurations.
• For representative elements, valence is often equal to the number of electrons in the outermost orbitals (valence electrons).
• Alternatively, valence can be calculated as eight minus the number of outermost electrons.

2. Oxidation States:

• The term "oxidation state" is frequently used interchangeably with valence, especially in the context of chemical reactions and compounds.
• Oxidation states are numerical values that represent the charge an atom would have in a compound or ion.
• Oxidation states can be positive, negative, or zero, depending on the electron transfer in a chemical reaction.

3. Example: OF₂ and Na₂O:

• Consider the compounds OF₂ and Na₂O.
• The electronegativity order of the elements involved is F > O > Na.
• In OF2, each fluorine atom tends to gain one electron to achieve a stable noble gas configuration (F: 1s²2s²2p⁵ → F⁻: 1s²2s²2p⁶).
• Therefore, each fluorine atom has an oxidation state of -1 in OF₂.
• In Na2O, sodium tends to lose one electron to achieve a noble gas configuration (Na: 1s²2s²2p⁶3s¹ → Na⁺: 1s²2s²2p⁶).
• Oxygen, on the other hand, tends to gain two electrons to attain a noble gas configuration (O: 1s²2s²2p⁴ → O²⁻: 1s²2s²2p⁶).
• Therefore, sodium has an oxidation state of +1, and oxygen has an oxidation state of -2 in Na₂O.

Understanding oxidation states and valence is crucial in predicting and explaining chemical reactions and the formation of compounds. It helps chemists balance chemical equations and determine the electron transfer in reactions.

 

 

ANOMALOUS  PROPERTIES  OF  SECOND  PERIOD  ELEMENTS

1. Anomalous Properties of First Group Elements:
• The first element in each group, such as lithium in Group 1 and beryllium in Group 2, and groups 13 to 17 (boron to fluorine), exhibits distinctive differences compared to the other members within their respective groups.
2. Covalent Character in Lithium and Beryllium:
• Unlike the other alkali metals, lithium and the other alkaline earth metal beryllium tend to form compounds with significant covalent characteristics. This is in contrast to the predominantly ionic compounds formed by the other members of these groups.
3. Similarity with Second Elements in Adjacent Group:
• Interestingly, the chemical behavior of lithium resembles that of the second element in the following group, magnesium, and beryllium's behavior is akin to that of aluminum, the second element in its group. This similarity is referred to as a "diagonal relationship" in periodic properties.
4. Reasons for Anomalous Behavior:
• The anomalous behavior of the first group element compared to its group members is primarily attributed to several factors:
• Small size: The first member has a smaller atomic size.
• Large charge-to-radius ratio: It possesses a higher charge-to-radius ratio.
• High electronegativity: These elements tend to be highly electronegative.
5. Limited Valence Orbitals:
• The first member of each group has only four valence orbitals (2s and 2p) available for bonding.
6. Covalency Limitation:
• Consequently, the maximum covalency of the first group member is limited to 4. For example, boron can form compounds like BF₄⁻, but it cannot expand its valence shell beyond four pairs of electrons.
7. Greater Ability for π Bonds:
• Additionally, the first member of the p-block elements shows a greater propensity to form pπ – pπ multiple bonds with itself (e.g., C = C, C ≡ C, N = N, N ≡ N) and with other second-period elements (e.g., C = O, C = N, C ≡ N, N = O). This ability diminishes in subsequent group members.
8. Expanding Valence Shell:
• In contrast, the second member of each group, with nine valence orbitals (3s, 3p, 3d), can expand its valence shell to accommodate more than four pairs of electrons, resulting in a wider range of bonding possibilities (e.g., aluminum can form AlF₆³⁻).

 

 

PERIODIC  TRENDS   AND  CHEMICAL  REACTIVITY

1. Fundamental Properties and Periodicity:
• All chemical and physical properties of elements are linked to their electronic configurations. Understanding periodic trends in properties such as atomic and ionic radii, ionization enthalpy, electron gain enthalpy, and valence is essential.
2. Periodic Trends in Radii and Enthalpies:
• Across a period, atomic and ionic radii generally decrease from left to right. Consequently, ionization enthalpies tend to increase. Exceptions exist, as mentioned earlier.
3. Chemical Reactivity and Enthalpies:
• Elements on the far left and right ends of a period exhibit the highest chemical reactivity. The leftmost element loses electrons to form cations, while the rightmost element gains electrons to form anions. This behavior relates to the metallic and non-metallic character of elements.
4. Metallic and Non-Metallic Character:
• Moving from left to right across a period, metallic character decreases, and non-metallic character increases. Alkali metals on the far left are highly metallic, while halogens on the far right are non-metals.
5. Reactivity with Oxygen:
• Chemical reactivity is evident in reactions with oxygen. Elements at the extremes of a period readily form oxides. Leftmost elements produce basic oxides (e.g., Na₂O), while rightmost elements produce acidic oxides (e.g., Cl₂O₇). Central elements produce amphoteric (e.g., Al₂O₃, As₂O₃) or neutral (e.g., CO, NO, N₂O) oxides.
6. Transition Metals:
• Transition metals (3d series) exhibit smaller changes in atomic radii across a period compared to representative elements. Their ionization enthalpies are intermediate, making them less electropositive than group 1 and 2 metals.
7. Group Trends:
• Within a group, as atomic and ionic radii increase with atomic number, ionization enthalpies generally decrease. This leads to an increase in metallic character and a decrease in non-metallic character down the group.
8. Transition Elements in Groups:
• In contrast, transition elements show a reverse trend within groups, primarily due to atomic size and ionization enthalpy considerations.

 

 

QUESTION  AND  ANSWERS

3.1 What is the basic theme of organisation in the periodic table?

3.2 Which important property did Mendeleev use to classify the elements in his periodic table and did he stick to that?

3.3 What is the basic difference in approach between the Mendeleev’s Periodic Law and the Modern Periodic Law?

3.4 On the basis of quantum numbers, justify that the sixth period of the periodic table should have 32 elements.

3.5 In terms of period and group where would you locate the element with Z =114?

3.6 Write the atomic number of the element present in the third period and seventeenth group of the periodic table.

3.7 Which element do you think would have been named by (i) Lawrence Berkeley Laboratory (ii) Seaborg’s group?

3.8 Why do elements in the same group have similar physical and chemical properties?

3.9 What does atomic radius and ionic radius really mean to you?

3.10 How do atomic radius vary in a period and in a group? How do you explain the variation?

3.1 Basic Theme of Organization in the Periodic Table:

• The basic theme of organization in the periodic table is the arrangement of chemical elements in a systematic and ordered manner based on their atomic number, electron configuration, and recurring chemical properties. Elements are grouped into periods (horizontal rows) and groups (vertical columns) to emphasize similarities and trends in their properties.

3.2 Mendeleev's Use of Property in Classification:

• Dmitri Mendeleev classified elements in his periodic table primarily based on their atomic mass. He organized the elements in order of increasing atomic mass and noticed that elements with similar chemical properties occurred at regular intervals. However, Mendeleev did not strictly stick to atomic mass; he adjusted the order of elements when it better fit their chemical properties, sometimes even leaving gaps for undiscovered elements.

3.3 Difference between Mendeleev's and Modern Periodic Law:

• Mendeleev's Periodic Law was based on atomic mass and chemical properties, and he left gaps for undiscovered elements. The key difference is that Mendeleev did not have a theoretical explanation for his periodic table, while the Modern Periodic Law is based on the atomic number (number of protons) and is explained by the electronic configuration of elements. In the modern table, elements are arranged in order of increasing atomic number, and there are no gaps.

3.4 Quantum Numbers and the Sixth Period:

• According to the quantum mechanical model, each period of the periodic table corresponds to the filling of a new energy level or shell. The maximum number of electrons that can occupy a given energy level is determined by the formula 2n², where 'n' is the principal quantum number. For the sixth period, 'n' is equal to 6, so the maximum number of elements that can be accommodated is 2 × 6² = 72. Since there are already 18 elements in the sixth period, 72 - 18 = 54 elements could potentially fill the sixth period, resulting in a total of 18 + 54 = 72 elements in the sixth period.

3.5 Locating Element with Z = 114:

• An element with Z = 114 would be located in the seventh period (row) of the periodic table and in the fourteenth group (column), often referred to as Group 14.

3.6 Atomic Number in the Third Period and Seventeenth Group:

• The element in the third period and seventeenth group of the periodic table has an atomic number of 17. This element is chlorine (Cl).

3.7 Naming Elements by Research Institutions:

• (i) The element with Z = 102, which is No. 102 on the periodic table, was named "Nobelium" by the Lawrence Berkeley Laboratory.
• (ii) Elements beyond No. 102 were discovered by Seaborg's group, and they are known as transuranium elements. For example, element No. 106 was named "Seaborgium" in honor of Glenn T. Seaborg, who played a significant role in the discovery of these elements.

3.8 Similar Properties in the Same Group:

• Elements in the same group have similar physical and chemical properties because they have the same number of valence electrons. Valence electrons are responsible for an element's chemical reactivity, and elements in the same group have the same outer electron configuration, leading to similar chemical behavior.

3.9 Atomic and Ionic Radius:

• Atomic radius refers to the size of an atom, typically measured as the distance from the nucleus to the outermost electron cloud. It reflects the size of the electron cloud.
• Ionic radius refers to the size of an ion, which can be larger or smaller than the atomic radius depending on whether the ion has gained or lost electrons.

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3.10 Variation of Atomic Radius:

• In a period (horizontal row), atomic radius generally decreases from left to right. This is because, within a period, the number of protons in the nucleus increases, leading to a greater nuclear charge pulling the electrons closer to the nucleus.
• In a group (vertical column), atomic radius generally increases as you move down the group. This is because each new element in a group has an additional energy level (shell) of electrons, making the electron cloud larger.
• The variation in atomic radius can be explained by the balance between increasing nuclear charge and increasing electron shells as you move across periods and down groups in the periodic table.