Why Do We Need To Classify Elements ?

By the early 19th century, the number of known elements was increasing rapidly (from 31 in 1800 to 63 by 1865, and currently over 114). With such a large number, studying the individual properties of each element and their countless compounds became extremely challenging.

Scientists recognized the need for a systematic way to organize this vast amount of information. Classifying elements based on similarities in their properties would not only simplify their study but also help in understanding the relationships between different elements and potentially predict the properties of new or undiscovered ones.

The goal was to find a way to organize the elements that would rationalize the known chemical facts and provide a framework for further research and discovery.

Genesis Of Periodic Classification

The idea of classifying elements and observing trends in their properties developed gradually through the efforts of several scientists.

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Early Attempts at Classification

Johann Dobereiner (early 1800s):

• Observed similarities within groups of three elements, which he called Triads.
• Noted that in a Triad, the atomic weight of the middle element was approximately the average of the atomic weights of the other two elements.
• Also found that the properties of the middle element were intermediate between those of the other two.

Element	Atomic weight	Element	Atomic weight	Element	Atomic weight
Li	7	Ca	40	Cl	35.5
Na	23	Sr	88	Br	80
K	39	Ba	137	I	127

($\text{Average wt of Li and K} = (7+39)/2 = 23$, which is the wt of Na. Similarly for other triads). Dobereiner's Law of Triads was limited, as it only applied to a few elements and was initially seen as a coincidence.

A.E.B. de Chancourtois (1862):

• Arranged elements by increasing atomic weight in a cylindrical fashion (telluric helix).
• Observed a periodic repetition of properties along the cylinder's vertical lines.
• His work did not gain widespread recognition at the time.

John Alexander Newlands (1865):

• Proposed the Law of Octaves.
• Arranged elements in increasing order of atomic weights and noticed that properties of every eighth element were similar to those of the first element, analogous to musical 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	...	...	...	...	...

Newlands' law worked well only for lighter elements up to Calcium and was not initially accepted by the scientific community. However, he was later recognized for his contribution.

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Contribution Of Mendeleev And Lothar Meyer

The foundation of the modern Periodic Table is largely credited to Dmitri Mendeleev (Russian chemist) and Lothar Meyer (German chemist), working independently around 1869.

• Both observed that when elements are arranged in increasing order of their atomic weights, similarities in physical and chemical properties recur at regular intervals (periodicity).
• Lothar Meyer plotted physical properties (like atomic volume, melting point, boiling point) against atomic weight and found repeating patterns. He also noticed that the length of the repeating pattern changed for different properties. His table, developed by 1868, closely resembled the modern one but was published later than Mendeleev's.

Dmitri Mendeleev is generally considered the father of the Periodic Table because he published his findings first and made bold predictions based on his classification.

• Mendeleev stated the first Periodic Law: "The properties of the elements are a periodic function of their atomic weights."
• He arranged elements in a table with horizontal rows (called series) and vertical columns (called groups) based on increasing atomic weights, placing elements with similar properties in the same group.
• Mendeleev considered a wider range of physical and chemical properties for classification, particularly focusing on the empirical formulas of compounds formed with hydrogen and oxygen (hydrides and oxides).
• Recognizing the importance of grouping elements with similar properties, he sometimes ignored the strict order of atomic weights (e.g., placing Iodine after Tellurium despite its lower atomic weight) assuming the measured atomic weights might be incorrect.
• Crucially, he left gaps in his table for elements he believed were yet to be discovered. He even predicted the properties of these unknown elements (e.g., Eka-Aluminium and Eka-Silicon, later discovered as Gallium and Germanium, respectively).

Property	Eka-aluminium (predicted)	Gallium (found)	Eka-silicon (predicted)	Germanium (found)
Atomic weight	68	70	72	72.6
Density/(g/cm³)	5.9	5.94	5.5	5.36
Melting point/K	Low	302.93	High	1231
Formula of oxide	E₂O₃	Ga₂O₃	EO₂	GeO₂
Formula of chloride	ECl₃	GaCl₃	ECl₄	GeCl₄

The subsequent discovery of these elements with properties remarkably close to his predictions solidified the acceptance of Mendeleev's Periodic Table.

Modern Periodic Law And The Present Form Of The Periodic Table

While Mendeleev's table was based on atomic weight, the understanding of atomic structure at the beginning of the 20th century led to a crucial refinement.

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Modern Periodic Law

In 1913, Henry Moseley studied the characteristic X-ray spectra of elements. He found that the frequency of emitted X-rays was related to the element's atomic number (Z) rather than its atomic mass. Specifically, a plot of $\sqrt{\nu}$ vs Z yielded a straight line.

This experimental evidence showed that the atomic number (Z) is a more fundamental property of an element than its atomic mass.

Based on Moseley's work, the Periodic Law was modified, leading to the Modern Periodic Law:

"The physical and chemical properties of the elements are periodic functions of their atomic numbers."

Since the atomic number equals the number of protons in the nucleus (and the number of electrons in a neutral atom), the Periodic Law is essentially a consequence of the periodic recurrence of similar electronic configurations in the outermost shell of atoms, which are responsible for their chemical and physical properties.

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The Long Form Of The Periodic Table

The most widely used form of the Periodic Table today is the "long form". In this table:

• Elements are arranged in increasing order of their atomic numbers.
• Horizontal rows are called periods. There are seven periods. The period number corresponds to the highest principal quantum number ($n$) of the valence shell electrons in the elements of that period.
• Vertical columns are called groups or families. There are eighteen groups, numbered 1 to 18 according to IUPAC recommendations (replacing the older IA-VIIA, VIII, IB-VIIB, 0 notation). Elements in the same group have similar outermost (valence shell) electronic configurations and thus exhibit similar chemical properties.

The number of elements in each period corresponds to the number of atomic orbitals available in the shell being filled:

• Period 1 ($n=1$): 1s orbital. Holds 2 electrons. 2 elements (H, He).
• Period 2 ($n=2$): 2s, 2p orbitals. Holds 2+6 = 8 electrons. 8 elements.
• Period 3 ($n=3$): 3s, 3p orbitals. Holds 2+6 = 8 electrons. 8 elements.
• Period 4 ($n=4$): 4s, 3d, 4p orbitals. Holds 2+10+6 = 18 electrons. 18 elements.
• Period 5 ($n=5$): 5s, 4d, 5p orbitals. Holds 2+10+6 = 18 electrons. 18 elements.
• Period 6 ($n=6$): 6s, 4f, 5d, 6p orbitals. Holds 2+14+10+6 = 32 electrons. 32 elements.
• Period 7 ($n=7$): 7s, 5f, 6d, 7p orbitals. Can hold up to 32 elements (incomplete). Includes synthetic elements.

The 14 elements filling the 4f orbitals (Lanthanoids) and the 14 elements filling the 5f orbitals (Actinoids) in periods 6 and 7, respectively, are placed separately at the bottom of the table to maintain its structure and highlight the similar chemistry within these series (called inner-transition elements).

Nomenclature Of Elements With Atomic Numbers > 100

Traditionally, the discoverer(s) of a new element had the privilege of naming it. However, for very heavy, unstable elements synthesized in different labs, this led to disputes over discovery claims and naming rights (e.g., element 104, claimed by both American and Soviet scientists).

To avoid controversy, IUPAC developed a systematic nomenclature for newly discovered elements with atomic numbers above 100. This temporary name is derived directly from the atomic number using numerical roots.

IUPAC Numerical Roots:

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

Naming Rule: Combine the roots for the digits of the atomic number in order, and add the suffix "-ium". The symbol is the first letter of each root.

Example: Element with Z = 101

• Digits are 1, 0, 1.
• Roots are un, nil, un.
• Name: unnilunium
• Symbol: Unu

Once a discovery is verified and officially recognized by IUPAC, a permanent name and symbol are assigned, often honoring a scientist or place of discovery.

Atomic Number	IUPAC Temporary Name	Symbol	IUPAC Official Name	Official Symbol
101	Unnilunium	Unu	Mendelevium	Md
102	Unnilbium	Unb	Nobelium	No
103	Unniltrium	Unt	Lawrencium	Lr
104	Unnilquadium	Unq	Rutherfordium	Rf
105	Unnilpentium	Unp	Dubnium	Db
106	Unnilhexium	Unh	Seaborgium	Sg
107	Unnilseptium	Uns	Bohrium	Bh
108	Unniloctium	Uno	Hassium	Hs
109	Unnilennium	Une	Meitnerium	Mt
110	Ununnillium	Uun	Darmstadtium	Ds
111	Unununium	Uuu	Rontgenium	Rg
112	Ununbium	Uub	Copernicium	Cn
113	Ununtrium	Uut	Nihonium	Nh
114	Ununquadium	Uuq	Flerovium	Fl
115	Ununpentium	Uup	Moscovium	Mc
116	Ununhexium	Uuh	Livermorium	Lv
117	Ununseptium	Uus	Tennessine	Ts
118	Ununoctium	Uuo	Oganesson	Og

Answer:

Electronic Configurations Of Elements And The Periodic Table

The electronic configuration of an atom, describing the arrangement of electrons in its orbitals, is the fundamental basis for its position and properties in the modern Periodic Table. The quantum numbers ($n, l, m_l, m_s$) characterize the electrons and their orbitals, and the Periodic Table visually organizes elements based on these characteristics.

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Electronic Configurations In Periods

The period number ($n$) in the Periodic Table corresponds to the principal quantum number of the valence shell (outermost energy level) being filled. As you move across a period, electrons are added to orbitals within this valence shell or sometimes to inner shells that become energetically accessible (like 3d filling after 4s).

The number of elements in a period is determined by the total number of electrons that can be accommodated in all the orbitals of the energy levels being filled in that period.

• Period 1 (n=1): Starts with filling the 1s orbital. Maximum 2 electrons in 1s ($1s^1$ to $1s^2$). 2 elements (H, He).
• Period 2 (n=2): Starts with 2s ($2s^1, 2s^2$), followed by 2p ($2p^1$ to $2p^6$). Maximum 2 in 2s + 6 in 2p = 8 electrons. 8 elements.
• Period 3 (n=3): Starts with 3s ($3s^1, 3s^2$), followed by 3p ($3p^1$ to $3p^6$). Maximum 2 in 3s + 6 in 3p = 8 electrons. 8 elements.
• Period 4 (n=4): Starts with 4s ($4s^1, 4s^2$). Then the 3d orbitals become energetically favorable and are filled (3d¹ to 3d¹⁰ - transition series). Finally, 4p orbitals are filled (4p¹ to 4p⁶). Maximum 2 in 4s + 10 in 3d + 6 in 4p = 18 electrons. 18 elements.
• Period 5 (n=5): Similar to Period 4. Starts with 5s ($5s^1, 5s^2$), then 4d (4d¹ to 4d¹⁰ - transition series), and finally 5p (5p¹ to 5p⁶). Maximum 2 in 5s + 10 in 4d + 6 in 5p = 18 electrons. 18 elements.
• Period 6 (n=6): Starts with 6s ($6s^1, 6s^2$). Then 4f orbitals are filled (4f¹ to 4f¹⁴ - Lanthanoids). Then 5d (5d¹ to 5d¹⁰ - remaining transition elements). Finally, 6p (6p¹ to 6p⁶). Maximum 2 in 6s + 14 in 4f + 10 in 5d + 6 in 6p = 32 electrons. 32 elements.
• Period 7 (n=7): Similar to Period 6. Starts with 7s ($7s^1, 7s^2$), then 5f (5f¹ to 5f¹⁴ - Actinoids), then 6d (6d¹ to 6d¹⁰), and finally 7p (7p¹ to 7p⁶). Incomplete, but can theoretically accommodate 32 elements.

The Lanthanoids (4f series) and Actinoids (5f series) are placed separately at the bottom because their filling involves inner f-orbitals, leading to similar properties within each series and keeping the main body of the table structured according to s, p, and d block filling.

Answer:

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Groupwise Electronic Configurations

Elements in the same vertical column (group) of the Periodic Table have similar chemical properties because they have the same number of valence electrons and the same general outer electronic configuration.

For example, Group 1 elements (Alkali Metals) all have one electron in their outermost s subshell ($ns^1$).

Atomic number	Symbol	Electronic configuration	Outer electronic configuration
3	Li	1s² 2s¹	2s¹
11	Na	1s² 2s² 2p⁶ 3s¹	3s¹
19	K	1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹	4s¹
37	Rb	[Kr] 5s¹	5s¹
55	Cs	[Xe] 6s¹	6s¹
87	Fr	[Rn] 7s¹	7s¹

This identical outermost electronic configuration leads to similar chemical behavior throughout the group.

Electronic Configurations And Types Of Elements : S-, P-, D-, F- Blocks

Based on which atomic orbital receives the last electron during the building up of electronic configuration, elements are broadly classified into four blocks: s-block, p-block, d-block, and f-block.

Exceptions:

• Hydrogen (H): Has only one electron ($1s^1$). Can be placed in Group 1 due to $s^1$ configuration or with Group 17 (halogens) as it can gain an electron to achieve a noble gas configuration. Due to its unique properties, it is often placed separately.
• Helium (He): Has configuration $1s^2$. It is in the s-block by electronic configuration, but its properties are very similar to other noble gases (Group 18) because of its completely filled valence shell. Hence, it is placed in the p-block with Group 18 elements.

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The S-Block Elements

• Located on the left side of the Periodic Table.
• Comprise Group 1 (Alkali Metals) and Group 2 (Alkaline Earth Metals).
• General outer electronic configuration is $ns^1$ (Group 1) or $ns^2$ (Group 2), where $n$ is the period number.
• All are reactive metals.
• Have low ionization enthalpies and readily lose their valence electron(s) to form positive ions ($+1$ for Group 1, $+2$ for Group 2).
• Metallic character and reactivity increase down a group.
• Not found pure in nature due to high reactivity.
• Form predominantly ionic compounds (except for some compounds of Li and Be, which show covalent character).

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The P-Block Elements

• Located on the right side of the Periodic Table.
• Comprise Group 13 to Group 18.
• Together with s-block elements, they are called Representative Elements or Main Group Elements.
• General outer electronic configuration varies from $ns^2np^1$ (Group 13) to $ns^2np^6$ (Group 18) in each period.
• Includes metals, non-metals, and metalloids.
• Non-metallic character increases from left to right across a period.
• Metallic character increases down a group.
• Group 18 elements are Noble Gases with a stable $ns^2np^6$ configuration (or $1s^2$ for He). They have high ionization enthalpies and positive electron gain enthalpies, making them very unreactive.
• Group 17 (Halogens, $ns^2np^5$) and Group 16 (Chalcogens, $ns^2np^4$) are important non-metal groups that readily gain one or two electrons, respectively, to achieve stable noble gas configuration, having highly negative electron gain enthalpies.

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The D-Block Elements (Transition Elements)

• Located in the middle of the Periodic Table.
• Comprise Group 3 to Group 12.
• Characterized by the filling of the $(n-1)d$ orbitals (the d subshell of the inner shell).
• General outer electronic configuration is $(n-1)d^{1-10}ns^{0-2}$ (with some exceptions like Pd ($4d^{10}5s^0$) and Cr ($3d^54s^1$), Cu ($3d^{10}4s^1$)).
• All are metals.
• Typically form colored ions, exhibit variable valency (oxidation states), are often paramagnetic, and are used as catalysts.
• Elements in Group 12 (Zn, Cd, Hg) with $(n-1)d^{10}ns^2$ configuration in their common oxidation states do not show typical transition metal properties as their d orbitals are completely filled. They are sometimes considered transition metals, but often grouped separately.
• Transition metals act as a bridge between the highly reactive s-block metals and the less reactive p-block elements.

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The F-Block Elements (Inner-Transition Elements)

• Located in two separate rows at the bottom of the Periodic Table.
• Comprise the Lanthanoids (Ce to Lu, filling 4f orbitals) and Actinoids (Th to Lr, filling 5f orbitals).
• Characterized by the filling of the $(n-2)f$ orbitals (the f subshell two shells inside the outermost shell).
• General outer electronic configuration is $(n-2)f^{1-14}(n-1)d^{0-1}ns^2$.
• All are metals.
• Properties within each series are quite similar due to the filling of inner f-orbitals (which shield the outer electrons effectively).
• Actinoids are radioactive, and many are synthetic. Their chemistry is more complex than lanthanoids due to the larger number of oxidation states.
• Elements after Uranium (Z=92) are called Transuranium elements.

Answer:

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Metals, Non-Metals And Metalloids

Elements can also be broadly classified based on their general physical and chemical properties.

• Metals:
  • Comprise over 78% of elements, located on the left side of the Periodic Table.
  • Usually solid at room temperature (except Mercury). Ga and Cs have low melting points.
  • Generally have high melting and boiling points.
  • Good conductors of heat and electricity.
  • Are malleable (can be hammered into sheets) and ductile (can be drawn into wires).
  • Tend to lose electrons easily to form positive ions.
• Non-metals:
  • Located on the top right side of the Periodic Table.
  • Usually solids or gases at room temperature (except Bromine, which is liquid).
  • Generally have low melting and boiling points (except Boron and Carbon).
  • Poor conductors of heat and electricity.
  • Solid non-metals are often brittle.
  • Tend to gain electrons easily to form negative ions.
• Metalloids (Semi-metals):
  • Elements located along the zigzag line separating metals and non-metals (e.g., Si, Ge, As, Sb, Te, Po, At).
  • Exhibit properties intermediate between those of metals and non-metals.
  • Often act as semiconductors.

The metallic character of elements increases as you move down a group and decreases as you move from left to right across a period. Conversely, non-metallic character increases across a period and decreases down a group.

Answer:

Periodic Trends In Properties Of Elements

Many physical and chemical properties of elements show gradual changes and predictable patterns as you move across periods or down groups in the Periodic Table. These systematic variations are called periodic trends and are directly linked to the electronic configuration of the atoms.

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Trends In Physical Properties

We will discuss periodic trends in atomic and ionic radii, ionization enthalpy, electron gain enthalpy, and electronegativity.

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Atomic Radius

Defining the exact size of an atom is difficult because the electron cloud doesn't have a sharp boundary. Atomic size is often estimated from the distance between the centers of bonded atoms.

• Covalent Radius: Half the distance between the nuclei of two identical atoms joined by a single covalent bond (e.g., in a Cl₂ molecule, Cl-Cl distance is 198 pm, so covalent radius of Cl is 99 pm).
• Metallic Radius: Half the internuclear distance between two adjacent metal atoms in a metallic crystal lattice (e.g., Cu-Cu distance in solid copper is 256 pm, so metallic radius of Cu is 128 pm).
• Van der Waals Radius: Half the internuclear distance between two identical non-bonded atoms in adjacent molecules in the solid state. This is typically larger than covalent or metallic radii.

Generally, the term "Atomic Radius" is used to refer to covalent or metallic radius.

Periodic Trends in Atomic Radius:

• Across a Period (Left to Right): Atomic radius generally decreases.

  Explanation: As you move across a period, electrons are added to the same valence shell. The nuclear charge (number of protons) increases, leading to a stronger attraction between the nucleus and the valence electrons. While inner electrons do provide some shielding, the effect of increased nuclear charge is more significant, pulling the electron cloud closer and reducing the atomic size.

• Down a Group (Top to Bottom): Atomic radius generally increases.

  Explanation: As you move down a group, the principal quantum number ($n$) of the valence shell increases, meaning the valence electrons are in shells further away from the nucleus. Additionally, the number of inner core electrons increases, providing increased shielding of the valence electrons from the nuclear charge. This shielding effect outweighs the increase in nuclear charge, reducing the effective nuclear attraction on the valence electrons and causing the atomic size to increase.

Noble gases' radii are typically given as van der Waals radii, which are larger than covalent radii, making direct comparison within a period tricky.

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Ionic Radius

The ionic radius is the effective distance from the nucleus of an ion to its outermost electrons. It is determined from distances between ions in ionic crystals.

• Cations: Positively charged ions formed by losing electrons. A cation is always smaller than its parent neutral atom.

  Explanation: The nucleus has the same positive charge, but there are fewer electrons. The effective nuclear charge per electron increases, pulling the remaining electrons closer to the nucleus.

• Anions: Negatively charged ions formed by gaining electrons. An anion is always larger than its parent neutral atom.

  Explanation: The nucleus has the same positive charge, but there are more electrons. The increased electron-electron repulsion within the same shell (or addition to an outer shell) outweighs the nuclear attraction, causing the electron cloud to expand.

Isoelectronic Species: These are atoms or ions that have the same number of electrons. For isoelectronic species, ionic size is determined by the nuclear charge (number of protons).

• Among isoelectronic species, the one with the highest positive charge (largest Z) will have the smallest size (electrons are pulled more strongly by the nucleus).
• Among isoelectronic species, the one with the highest negative charge (lowest Z) will have the largest size (greatest electron-electron repulsion).

Example of isoelectronic species with 10 electrons: O²⁻, F⁻, Ne, Na⁺, Mg²⁺, Al³⁺.

Their increasing order of nuclear charge (Z): O (8) < F (9) < Ne (10) < Na (11) < Mg (12) < Al (13).

Their decreasing order of size: O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺ > Al³⁺.

Answer:

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Ionization Enthalpy

Ionization Enthalpy ($\Delta_i H$) is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state. It is a measure of how difficult it is to remove an electron.

$$ \text{X(g)} \rightarrow \text{X⁺(g)} + \text{e⁻} \quad (\text{First Ionization Enthalpy, } \Delta_i H_1) $$

Subsequent ionization enthalpies ($\Delta_i H_2$, $\Delta_i H_3$, etc.) are the energies required to remove the second, third, etc., electrons. $\Delta_i H_2 > \Delta_i H_1$ because removing an electron from a positively charged ion is harder.

Ionization enthalpies are always positive (endothermic process).

Periodic Trends in Ionization Enthalpy:

• Across a Period (Left to Right): Ionization enthalpy generally increases.

  Explanation: As atomic number increases across a period, the effective nuclear charge increases while the valence electrons are in the same shell. This stronger attraction to the nucleus makes it harder to remove an electron, requiring more energy.

  Exceptions exist:

  • Group 13 elements (e.g., B) have lower $\Delta_i H_1$ than Group 2 (e.g., Be). This is because Group 2 has a filled $ns^2$ configuration, making it stable. Group 13 has an electron in the $np^1$ orbital, which is slightly higher in energy and less penetrating than the $ns^2$ electrons, making it easier to remove.
  • Group 16 elements (e.g., O) have lower $\Delta_i H_1$ than Group 15 (e.g., N). Group 15 has a stable half-filled $np^3$ configuration (according to Hund's rule, each p orbital has one electron). Group 16 has $np^4$, meaning one p orbital has paired electrons, creating electron-electron repulsion that makes it slightly easier to remove one of the paired electrons.
• Down a Group (Top to Bottom): Ionization enthalpy generally decreases.

  Explanation: As you move down a group, the valence electrons are in higher energy shells, further from the nucleus. The increased number of inner core electrons provides greater shielding of the valence electrons from the nuclear charge. This increased distance and shielding reduce the effective nuclear attraction, making it easier to remove the outermost electron, requiring less energy.

Alkali metals (Group 1) have the lowest ionization enthalpies in their periods, correlating with their high reactivity as they easily lose an electron. Noble gases (Group 18) have the highest ionization enthalpies due to their very stable closed electron shell configurations.

Answer:

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Electron Gain Enthalpy

Electron Gain Enthalpy ($\Delta_{eg} H$) is the enthalpy change that occurs when an electron is added to an isolated gaseous atom in its ground state to form a gaseous anion.

$$ \text{X(g)} + \text{e⁻} \rightarrow \text{X⁻(g)} \quad (\Delta_{eg} H) $$

Electron gain enthalpy can be exothermic (energy released, $\Delta_{eg} H$ is negative) or endothermic (energy absorbed, $\Delta_{eg} H$ is positive). A more negative $\Delta_{eg} H$ indicates that the atom has a greater tendency to accept an electron.

Examples:

• Halogens (Group 17) have very negative $\Delta_{eg} H$ because gaining one electron gives them a stable noble gas configuration.
• Noble gases (Group 18) have large positive $\Delta_{eg} H$ because adding an electron requires it to enter a new, higher energy shell, resulting in an unstable configuration and significant electron-electron repulsion.

Periodic Trends in Electron Gain Enthalpy:

• Across a Period (Left to Right): Electron gain enthalpy generally becomes more negative.

  Explanation: As effective nuclear charge increases across a period, the nucleus has a stronger attraction for an incoming electron. This makes it easier for the atom to gain an electron, and more energy is released (more negative $\Delta_{eg} H$).

• Down a Group (Top to Bottom): Electron gain enthalpy generally becomes less negative.

  Explanation: As atomic size increases down a group, the incoming electron is further from the nucleus and experiences weaker attraction. The increased shielding also reduces the effective nuclear charge felt by the incoming electron. This makes it harder to add an electron, and less energy is released (less negative $\Delta_{eg} H$).

  Exceptions exist, particularly for the second period elements (O, F). Their $\Delta_{eg} H$ values are less negative than those of the corresponding third period elements (S, Cl). This is because the second period elements are smaller. When an electron is added to a small 2p orbital, it experiences significant repulsion from the existing electrons in that relatively compact shell. For the larger 3p orbitals in the third period, the added electron occupies a larger space, and electron-electron repulsion is reduced, making electron addition more favorable (more negative $\Delta_{eg} H$). Thus, Cl has a more negative $\Delta_{eg} H$ than F, and S has a more negative $\Delta_{eg} H$ than O.

Answer:

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Electronegativity

Electronegativity is a qualitative measure of the ability of an atom in a chemical compound to attract a shared pair of electrons towards itself. It is not a directly measurable quantity but is calculated based on various scales (e.g., Pauling scale, Mulliken scale).

On the widely used Pauling scale, Fluorine (F) is assigned the highest electronegativity value (4.0).

Electronegativity values are not constant for an element but vary depending on the atoms it is bonded to and its oxidation state.

Periodic Trends in Electronegativity:

• Across a Period (Left to Right): Electronegativity generally increases.

  Explanation: As atomic size decreases and effective nuclear charge increases across a period, the nucleus has a stronger pull on the shared electrons in a bond.

• Down a Group (Top to Bottom): Electronegativity generally decreases.

  Explanation: As atomic size increases and the distance between the nucleus and valence shell increases down a group, the nucleus's pull on the shared electrons weakens due to increased distance and shielding.

Relationship with Metallic and Non-metallic Character:

• Non-metals tend to gain electrons and are generally highly electronegative. Electronegativity is directly related to non-metallic character.
• Metals tend to lose electrons and have low electronegativity. Electronegativity is inversely related to metallic character.

As electronegativity increases across a period, non-metallic character increases, and metallic character decreases. As electronegativity decreases down a group, metallic character increases, and non-metallic character decreases.

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Periodic Trends In Chemical Properties

The chemical properties of elements are determined by their electronic configurations, particularly the valence electrons. Periodic trends in fundamental properties like ionization enthalpy and electron gain enthalpy directly influence chemical reactivity and the nature of compounds formed.

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Periodicity Of Valence Or Oxidation States

The valence of a representative element is often related to the number of electrons in its outermost shell or eight minus that number. The term oxidation state is frequently used and refers to the charge an atom appears to have in a compound based on electronegativity differences.

General Trend in Valence (for Representative Elements):

Group	1	2	13	14	15	16	17	18
Number of valence electrons	1	2	3	4	5	6	7	8
Valence (Common)	1	2	3	4	3,5	2,6	1,7	0,8

Elements in Group 1 typically show a valence of 1 (losing one $ns^1$ electron). Group 2 shows 2 (losing two $ns^2$ electrons). Group 13 shows 3 (losing $ns^2np^1$). Group 14 shows 4 (losing $ns^2np^2$). Groups 15, 16, and 17 often show valence equal to 8 minus the number of valence electrons (gaining electrons to achieve octet), or variable valency by involving d orbitals.

Oxidation state considers the charge based on electronegativity differences. In OF₂, Oxygen is less electronegative than Fluorine, so Oxygen has a positive oxidation state (+2), while F is -1. In Na₂O, Oxygen is more electronegative than Sodium, so Oxygen is -2, and Na is +1.

Transition elements (d-block) and inner-transition elements (f-block) often exhibit variable valency/oxidation states due to the involvement of inner d or f electrons in bonding.

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Anomalous Properties Of Second Period Elements

The first element of each group in the s-block and p-block (Li, Be, B, C, N, O, F) exhibits some properties that are different from the other members of their respective groups. This is known as the anomalous behavior of second period elements.

Examples:

• Lithium (Li) and Beryllium (Be) form compounds with more covalent character compared to other alkali and alkaline earth metals, which form predominantly ionic compounds.
• The properties of the first element of a group are often more similar to the second element of the next group. This is called the diagonal relationship (e.g., Li resembles Mg, Be resembles Al).

Reasons for anomalous behavior:

• Small Size: Second period elements are significantly smaller than other elements in their group.
• Large Charge/Radius Ratio: Due to small size, they have a high concentration of charge relative to their size, leading to greater polarizing power.
• High Electronegativity: They are generally more electronegative than heavier elements in their group.
• Absence of d-orbitals: The valence shell of second period elements (n=2) has only s and p orbitals (2s, 2p), totaling 4 valence orbitals. This limits their maximum covalency to 4. Heavier elements in the same groups have access to vacant d-orbitals in their valence shell (e.g., 3s, 3p, 3d for Period 3), allowing them to expand their valence shell and show covalencies greater than 4 (e.g., P can form PCl₅).
• Ability to form $p\pi - p\pi$ multiple bonds: Second period elements (especially C, N, O) can form stable double and triple bonds with themselves (e.g., C=C, N≡N) and with other second period elements (e.g., C=O, C≡N). Heavier elements form weaker $p\pi - p\pi$ bonds and prefer single bonds or $d\pi - p\pi$ or $d\pi - d\pi$ bonding.

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Periodic Trends And Chemical Reactivity

The chemical reactivity of an element is closely related to its tendency to lose or gain electrons, which in turn depends on properties like ionization enthalpy and electron gain enthalpy.

Trends in Reactivity:

• Across a Period: Chemical reactivity is generally high at the extremes (left and right) and lowest in the center.
  • Left side (Alkali Metals): Highly reactive, easily lose electrons (low $\Delta_i H$). Reactivity increases down the group.
  • Right side (Halogens): Highly reactive (except noble gases), easily gain electrons (highly negative $\Delta_{eg} H$). Reactivity decreases down the group.
  • Center (e.g., Group 14): Lower reactivity compared to extremes.
• Down a Group:
  • For Metals (e.g., Alkali Metals, Alkaline Earth Metals): Reactivity increases down the group due to decreasing ionization enthalpy (easier to lose electrons).
  • For Non-metals (e.g., Halogens): Reactivity decreases down the group due to less negative electron gain enthalpy (harder to gain electrons).

Metals on the left side tend to form basic oxides (e.g., Na₂O + H₂O $\rightarrow$ NaOH). Non-metals on the right side tend to form acidic oxides (e.g., Cl₂O₇ + H₂O $\rightarrow$ HClO₄). Elements in the center can form amphoteric oxides (react with both acids and bases, e.g., Al₂O₃) or neutral oxides (e.g., CO, NO).

Summary of Major Periodic Trends:

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