• The elements are basic unit of all matters. In 1800, only 31 elements were discovered. But in 1865, 65 elements were discovered. At present, 118 elements are known. Some recently discovered elements are man-made.

FORMATION(or genesis) OF PERIODIC CLASSIFICATION-

  • The idea of trend about properties of elements is first considered by the German chemist, Johann Dobereiner in 1800s.
  • They noticed in 1829, the similar physical and chemical properties of many groups of 3 elements or Triads.
  • He noticed in each case that properties and atomic weight of middle element of each triads is half of the some of other two out of 3 elements, this is called Law of Triads.
  • In 1865, the English chemist, John Alexander Newlands profounded the Law of Octaves.
  • Law of Octaves- It states that every 8th element in periodic table has similar chemical properties when elements are arranged by increasing atomic mass.
  • But Newland ‘s law seemed right for only elements up to calcium, so this idea is not widely accepted at that time. He and his work was later awarded by Davy Medal in 1887 by Royal Society of London.
  1. Development of the Periodic Law:
    • The Periodic Law was developed by Dmitri Mendeleev (Russia) and Lothar Meyer (Germany) independently in 1869.
    • Both proposed that arranging elements in increasing order of atomic weights revealed periodic similarities in their properties.
  2. Lothar Meyer’s Contributions:
    • Meyer plotted physical properties (atomic volume, melting/boiling points) against atomic weights, identifying a periodic pattern.
    • His work, resembling Mendeleev’s, was published later, and he noticed a change in the repeating pattern’s length.
  3. Mendeleev’s Contributions:
    • Mendeleev is credited with the first published Periodic Law: “The properties of the elements are a periodic function of their atomic weights.”
    • Mendeleev arranged elements in rows and columns, based on atomic weights, ensuring elements with similar properties grouped together in vertical columns.
    • He used a broader range of properties for classification, such as empirical formulas and compounds’ properties.
  4. Mendeleev’s Adjustments:
    • Mendeleev adjusted the order when atomic weights didn’t fit, believing some atomic measurements might be incorrect.
    • For example, iodine was placed in Group VII (not VI), due to its similarity to halogens, despite having a lower atomic weight than tellurium.
  5. Prediction of Undiscovered Elements:
    • Mendeleev left gaps in the table for undiscovered elements, such as gallium and germanium.
    • He predicted their properties, which were later confirmed upon their discovery, proving his theory’s accuracy.
  6. Mendeleev’s Fame and Influence:
    • His predictions led to fame, with new elements fitting his table, like gallium and germanium.
    • Mendeleev’s work sparked further research in areas like the noble gases, radioactive decay series, and transuranium elements.
  7. Mendeleev’s Life:
    • Born in Siberia, Mendeleev became a professor in St. Petersburg and wrote Principles of Chemistry, contributing to the development of the Periodic Table.
    • He adjusted the atomic order based on chemical properties and foresaw elements yet to be discovered.
    • He was also involved in other fields like developing an accurate barometer and working on Russia’s natural resources.
  8. Modern Periodic Law:
    • In 1913, Henry Moseley observed that the atomic number, not atomic mass, is the fundamental property of elements.
    • This revised the Periodic Law to: “The physical and chemical properties of the elements are periodic functions of their atomic numbers.”
    • Moseley’s findings led to a better understanding of atomic structure and periodicity.
  9. Significance of the Atomic Number:
    • The atomic number equals the number of protons (and electrons in a neutral atom).
    • Electronic configurations and quantum numbers now explain periodicity in elements’ properties.
  10. Modern Periodic Table:
    • The most commonly used version is the “long form,” which emphasizes electronic configurations.
    • Periods are horizontal rows; groups are vertical columns. Elements with similar outer electron configurations are grouped together.
    • Groups are numbered 1 to 18 (IUPAC recommendation).
  11. Structure of Periods and Groups:
    • There are seven periods in the Periodic Table, with period number corresponding to the highest principal quantum number (n).
    • The first period contains 2 elements, the next periods contain 8, 8, 18, 18, and 32 elements, respectively.
    • The 7th period is incomplete, with a theoretical maximum of 32 elements, similar to the 6th period.
  12. Lanthanoids and Actinoids:
    • The 14 elements of the 6th and 7th periods (lanthanoids and actinoids) are placed in separate panels at the bottom of the Periodic Table.

These notes summarize key points related to the history, development, and structure of the Periodic Table and its underlying principles

Nomenclature of Elements with Atomic Numbers > 100

  • Traditionally, naming new elements was the privilege of the discoverer(s), with the name ratified by IUPAC. However, controversy arose due to the instability of new elements with high atomic numbers, often only detectable in minute quantities.
  • Challenges in discovering high atomic number elements: These elements are highly unstable, and synthesizing them requires sophisticated and costly equipment. As such, the discovery process is highly competitive and performed in select laboratories.
  • Disputed discoveries: There have been claims from different countries for the same elements (e.g., Rutherfordium), which led to the development of a standardized naming system.
  • IUPAC temporary naming system: The IUPAC recommends a systematic approach for naming elements before their official discovery is confirmed. The name derives from the atomic number, using numerical roots (0-9) for each digit of the atomic number. The temporary name ends with “ium.”
    • Example: For the element with atomic number 104, the IUPAC name would be “Unnilquadium” (Unq).
  • Naming convention: The roots for digits are:
    • 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).
  • Official IUPAC names: Once an element’s discovery is confirmed, it is officially named and assigned a symbol. Some elements with atomic numbers above 113, 115, 117, and 118 have not yet received official names.
    • For example, element 120’s name would be “Unbinilium” (Ubn).

2. Electronic Configurations of Elements and the Periodic Table

  • Electron configuration: The distribution of electrons in an atom’s orbitals defines its electron configuration. This configuration directly correlates with the element’s position in the periodic table.
  • Periods: A period corresponds to a principal energy level (n), and each period in the Periodic Table reflects the filling of a new electron shell.
    • First period: Contains 2 elements (Hydrogen and Helium).
    • Second and third periods: Each contains 8 elements (e.g., second period includes Li, Be, B, C, N, O, F, Ne).
    • Fourth and fifth periods: Each contains 18 elements, with transition elements filling d orbitals.
    • Sixth period: Contains 32 elements, with 4f orbitals filling (Lanthanide series).
    • Seventh period: Also contains 32 elements, with 5f orbitals filling (Actinide series), including synthetic radioactive elements.
  • Filling of orbitals: The order in which orbitals are filled is dictated by the energy levels (e.g., 4s < 3d < 4p).
    • Transition series: For example, the 3d transition series starts from Scandium (Z = 21) and ends at Zinc (Z = 30).
    • Lanthanides and Actinides: Elements in the 4f and 5f series are placed separately to maintain the table’s structure.

3. Justification of 18 Elements in the 5th Period

  • The 5th period has 18 elements because the available orbitals are 4d, 5s, and 5p, which can accommodate 18 electrons in total.

4. Groupwise Electronic Configurations

  • Elements in the same group (vertical columns) have similar valence electron configurations and thus exhibit similar chemical properties.
  • Example of Group 1 elements (Alkali metals): All elements in this group have the same valence shell configuration (ns¹), which leads to their similar chemical behavior.
  • Classification into blocks:
    • s-block: Includes groups 1 and 2.
    • p-block: Includes groups 13-18.
    • d-block: Transition metals.
    • f-block: Lanthanides and Actinides.
    • Exceptions:
      • Helium: Though it is in the s-block (1s²), it is placed with the noble gases due to its full valence shell, exhibiting properties similar to noble gases.
      • Hydrogen: Although it has a single electron in its s-orbital (like alkali metals), it behaves like a halogen in some reactions and is placed separately at the top of the Periodic Table.

5. Periodic Classification

Electronic Configurations and Types of Elements:

S-, P-, D-, F- Block Elements:

  • Aufbau Principle: This principle helps in understanding the electronic configuration of atoms and their placement in the Periodic Table.
  • The elements in the Periodic Table are categorized into s-, p-, d-, and f-blocks based on the orbitals being filled by electrons.

3.6.1 S-Block Elements:

  • Group 1 (Alkali metals) and Group 2 (Alkaline earth metals).
  • Their outer electronic configurations are ns¹ (Group 1) and ns² (Group 2).
  • Properties:
    • Highly reactive metals.
    • Low ionization enthalpies, making them lose electrons easily to form 1⁺ (alkali metals) or 2⁺ (alkaline earth metals) ions.
    • Reactivity increases as we move down the group.
    • These elements are rarely found pure in nature due to their high reactivity.
    • The compounds of s-block elements (except lithium and beryllium) are predominantly ionic.

3.6.2 P-Block Elements:

  • Groups 13 to 18; also known as Representative Elements or Main Group Elements.
  • Outer electronic configurations range from ns²np¹ to ns²np⁶.
  • Noble gases (Group 18) have a stable, fully filled ns²np⁶ configuration and exhibit very low chemical reactivity.
  • Halogens (Group 17) and Chalcogens (Group 16) are highly reactive non-metals.
  • Trends:
    • Non-metallic character increases from left to right across a period.
    • Metallic character increases as we move down a group.

3.6.3 D-Block Elements (Transition Elements):

  • Groups 3 to 12: These elements fill their d-orbitals and are thus called d-Block Elements.
  • General Electronic Configuration: (n-1)d¹⁰ns² or similar variations.
  • Properties:
    • All transition elements are metals.
    • They exhibit variable oxidation states, paramagnetism, and often act as catalysts.
    • Zn, Cd, and Hg do not show typical transition metal properties due to filled d-orbitals.
    • Transition elements form a bridge between the reactive metals of the s-block and less reactive metals of Groups 13-14.

3.6.4 F-Block Elements (Inner-Transition Elements):

  • The Lanthanides (Ce to Lu) and Actinides (Th to Lr).
  • Outer electronic configuration: (n-2)f¹⁴(n-1)d⁰-¹ns².
  • Properties:
    • Lanthanides have similar properties.
    • Actinides are more complicated due to multiple oxidation states and radioactivity.
    • Transuranium elements are those beyond Uranium and are synthesized artificially.

3.6.5 Metals, Non-Metals, and Metalloids:

  • Metals: Majority of the elements, usually solid at room temperature (except mercury). Good conductors of heat and electricity, malleable, and ductile.
  • Non-Metals: Found on the top-right side of the periodic table. They are typically solids or gases with poor conductivity, low melting/boiling points, and are brittle.
  • Metalloids: Elements bordering the zig-zag line (e.g., silicon, germanium) show properties of both metals and non-metals.

Classification based on periodic trends:

  • Metallic Character: Increases down a group and decreases from left to right across a period.

3.7 Periodic Trends in Properties:

  • Periodic Trends: Observable patterns in the properties of elements as you move across periods or down groups.

3.7.1 Trends in Physical Properties:

  • Atomic Radius:
    • Across a period: Atomic radius decreases as nuclear charge increases, pulling electrons closer.
    • Down a group: Atomic radius increases because the number of electron shells increases, distancing the outer electrons from the nucleus.
  • Ionic Radius:
    • Cations are smaller than their parent atoms due to loss of electrons.
    • Anions are larger than their parent atoms because added electrons increase electron-electron repulsion.
    • Isoelectronic species (atoms and ions with the same electron number) have different radii based on nuclear charge.

Ionization Enthalpy (Ionization Energy)

  • Definition: Ionization enthalpy is the energy required to remove an electron from an isolated gaseous atom in its ground state.
  • First Ionization Enthalpy: The enthalpy change (∆iH) for removing one electron from an atom (x) = X(g) → X+(g) + e–
  • Units: Ionization enthalpy is measured in kJ/mol.
  • Second Ionization Enthalpy: The energy required to remove the second most loosely bound electron from a positive ion: X+(g) →X+2 (g) + e–
  • Trend in Ionization Enthalpy:
    1. Increases across a period: As we move across a period, the nuclear charge increases, attracting the outer electrons more strongly, making them harder to remove.
    2. Decreases down a group: As we move down a group, the outer electron is farther from the nucleus, and more electron shielding occurs, making it easier to remove the electron.

Factors Affecting Ionization Enthalpy:

  1. Attraction to the nucleus: Electrons closer to the nucleus are harder to remove because the nucleus exerts a stronger pull.
  2. Electron repulsion: Electrons in the same shell experience repulsion, which can affect how easily they are removed.
  3. Shielding: Inner shell electrons shield the outer electrons from the full charge of the nucleus. This is significant in alkali metals, where the outermost electron is shielded by a filled inner shell (e.g., noble gas configuration).

Trends and Exceptions:

  • Across a period: As atomic number increases, ionization enthalpy generally increases due to a higher nuclear charge without much increase in shielding.
    • Example: Ionization enthalpy of boron is less than that of beryllium because boron’s outer electron is in a 2p orbital, which is more shielded than the 2s electron of beryllium.
    • Exception: Oxygen has a lower ionization enthalpy than nitrogen because of electron-electron repulsion in the doubly occupied 2p orbitals of oxygen.
  • Down a group: Ionization enthalpy decreases because the outer electron is farther from the nucleus and experiences more shielding.

Electron Gain Enthalpy

  • Definition: Electron gain enthalpy is the energy change when an electron is added to a neutral gaseous atom to form a negative ion: (x) = X(g) + e– →X−(g) .
  • Exothermic vs Endothermic:
    • Exothermic (negative value): For elements like halogens (Group 17), energy is released when an electron is added, making electron gain enthalpy negative.
    • Endothermic (positive value): For noble gases, adding an electron would require energy because the next available orbital is much higher in energy.
  • Trends in Electron Gain Enthalpy:
    1. Across a period: It becomes more negative from left to right across a period due to increased nuclear charge, making it easier to attract an additional electron.
    2. Down a group: Electron gain enthalpy becomes less negative because the atom becomes larger, and the added electron is farther from the nucleus.
  • Exceptions:
    • Oxygen and Fluorine: These elements have less negative electron gain enthalpies compared to their subsequent elements because the added electron in O or F experiences more repulsion due to being added to a small 2p orbital.

Summary:

  • Ionization Enthalpy increases across a period and decreases down a group due to changes in nuclear charge and electron shielding.
  • Electron Gain Enthalpy is negative for elements that gain an electron easily (like halogens) and positive for noble gases, reflecting their stability. It generally becomes more negative across a period and less negative down a group, with specific exceptions due to electron repulsion in small orbitals.

Here are the detailed notes on electronegativity and its related concepts, covering basic principles and important knowledge to build concepts and solve competitive questions:

1. Definition of Electronegativity:

  • Electronegativity is a qualitative measure of an atom’s ability to attract shared electrons in a chemical compound.
  • Unlike ionization enthalpy and electron gain enthalpy, electronegativity is not a directly measurable quantity.

2. Electronegativity Scales:

  • Several scales have been developed to quantify electronegativity, including:
    • Pauling Scale (most widely used): Developed by Linus Pauling in 1922, fluorine is arbitrarily assigned a value of 4.0 as the most electronegative element.
    • Mulliken-Jaffe Scale.
    • Allred-Rochow Scale.
  • Pauling scale values for some elements (given in tables in the passage) are useful for comparison.

3. Variation of Electronegativity:

  • Across a Period: Electronegativity generally increases from left to right due to decreasing atomic radius and increasing nuclear charge (e.g., from lithium to fluorine).
  • Down a Group: Electronegativity generally decreases because of increasing atomic radius and shielding effect (e.g., from fluorine to astatine).

4. Relationship with Atomic Radius:

  • As atomic radius decreases across a period, electronegativity increases.
  • As atomic radius increases down a group, electronegativity decreases.

5. Relation to Non-Metallic and Metallic Properties:

  • Non-metallic elements tend to gain electrons and thus have higher electronegativity.
  • Electronegativity is inversely related to metallic properties.
  • As electronegativity increases across a period, non-metallic properties increase, and metallic properties decrease.
  • Conversely, as electronegativity decreases down a group, non-metallic properties decrease, and metallic properties increase.

6. Predicting Bond Nature:

  • Electronegativity differences between two elements help predict the nature of the bond:
    • Covalent bond: If the electronegativity difference is small.
    • Ionic bond: If the electronegativity difference is large.

7. Periodic Trends in Valence States and Oxidation States:

  • The valence state of an element is related to its electronic configuration. This can often be linked to the number of electrons in the outermost shell.
  • The oxidation state is defined as the charge an atom takes on due to electronegativity differences in a molecule (e.g., in OF2, oxygen exhibits oxidation state +2 due to fluorine’s high electronegativity).

8. Anomalous Properties of Second Period Elements:

  • Lithium (Li) and Beryllium (Be) exhibit anomalous behavior compared to other elements in their respective groups.
    • Lithium forms covalent compounds, unlike other alkali metals, which form ionic compounds.
    • Beryllium has similar properties to magnesium and aluminum due to its small size and high electronegativity.
  • Diagonal Relationship: The similarities in properties between the first element in one group and the second element in the adjacent group (e.g., Li and Mg, Be and Al) are referred to as diagonal relationships.

9. Trends in Chemical Reactivity:

  • Metallic and Non-metallic Character:
    • Metallic character increases down a group and decreases across a period (e.g., from Na to Cl).
    • Non-metallic character increases across a period and decreases down a group.
  • Elements at the extremes of a period (alkali metals and halogens) exhibit the highest chemical reactivity due to their tendency to lose or gain electrons respectively.
    • Alkali metals form basic oxides (e.g., Na₂O), while halogens form acidic oxides (e.g., Cl₂O₇).
  • Amphoteric Oxides: Oxides formed by elements in the middle of a period can behave as either acidic or basic (e.g., Al₂O₃, As₂O₃).

10. Chemical Reactivity with Oxygen:

  • Basic oxides are formed by elements with low electronegativity (e.g., Na₂O).
  • Acidic oxides are formed by elements with high electronegativity (e.g., Cl₂O₇).
  • Amphoteric oxides: Elements with intermediate electronegativity form oxides that can act both as acids and bases (e.g., Al₂O₃).

11. Reactivity of Transition Metals:

  • Transition metals show less variation in atomic radii compared to main group elements, and their ionization enthalpies are intermediate between s- and p-block elements.
  • They are less electropositive than alkali and alkaline earth metals but more electropositive than non-metals.

12. Periodic Trends in Chemical Reactivity and Ionization:

  • As atomic and ionic radii decrease across a period, ionization energy generally increases.
  • Electron gain enthalpies also become more negative across a period, except for noble gases, which have positive electron gain enthalpies due to their stable electron configuration.
  • The chemical reactivity is highest at the extremes of the periodic table (e.g., alkali metals and halogens) and lowest at the center.
  • Reactivity trends are linked to oxidizing (non-metals) and reducing (metals) properties.

Conclusion:

  • Electronegativity plays a key role in understanding chemical bonding, reactivity, and trends in the periodic table.
  • The periodicity of elements is closely tied to their electronic configurations, influencing their oxidation states, valence states, and chemical reactivity.
  • Analyzing electronegativity alongside other periodic properties helps predict the behavior of elements in chemical reactions and bonding.

These insights will aid in solving competitive chemistry questions, as they highlight the underlying concepts related to periodicity, electronegativity, and reactivity.

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