PERIODIC CLASSIFICATION: PART 3

 

p-Block Elements

• The p-Block Elements encompass elements found in Group 13 to 18 on the periodic table.
• When combined with the s-Block Elements, they are referred to as Representative or Main Group Elements.
• The outermost electron arrangement ranges from ns²np¹ to ns²np⁶ within each period.
• At the conclusion of each period lies a noble gas element with a closed valence shell ns²np⁶ configuration.
• Noble gases possess fully populated valence shell orbitals, making their electron arrangements stable and resistant to change.
• Due to this stability, noble gases exhibit minimal chemical reactivity.
• Just before the noble gas series, two significant groups of nonmetals can be found.
• These groups are the halogens (Group 17) and the chalcogens (Group 16).
• Both the halogens and chalcogens have notably negative electron gain enthalpies.
• They readily accept one or two electrons respectively to achieve the stable noble gas configuration.
• Non-metallic tendencies intensify as we progress from left to right across a period.
• Conversely, metallic properties increase as we move down the group.

d-Block Elements

• These elements occupy Group 3 to 12 at the center of the Periodic Table.
• They are identified by their occupation of inner d orbitals by electrons and are termed d-Block Elements.
• These elements typically have the general outer electronic configuration of (n-1)d^1-10ns^0-2, with an exception being Pd which has the electronic configuration 4d¹⁰5s⁰.
• All of these elements are categorized as metals.
• They commonly yield ions with distinct colors, display variable valence (oxidation states), and exhibit paramagnetism.
• These elements often serve as catalysts.
• Notably, Zn, Cd, and Hg possess the electronic configuration (n-1)d¹⁰ns², and consequently, they lack several properties exhibited by typical transition elements.
• Transition metals act as a connection between the more chemically active metals found in the s-block and the comparatively less reactive elements in Groups 13 and 14.

f-Block Elements

• The f-Block Elements, also known as Inner-Transition Elements, encompass two rows at the bottom of the Periodic Table.
• These rows consist of Lanthanoids (Ce(Z = 58) – Lu(Z = 71)) and Actinoids (Th(Z = 90) – Lr (Z = 103)).
• Their outer electronic configuration is defined as (n-2)f^1-14 (n-1)d^0–1 ns².
• The last added electron for each element occupies an f-orbital.
• The Inner-Transition Elements are collectively named due to this arrangement.
• All elements in this category are metals.
• Elements within each series share similar properties.
• Among the Actinoids, chemistry is more intricate due to the multitude of possible oxidation states.
• Actinoid elements are radioactive in nature.
• Many actinoids are produced only in extremely small quantities through nuclear reactions, limiting their chemical study.
• Elements beyond uranium are termed Transuranium Elements.

Metals, Non-metals, and Metalloids:

• Apart from the categorization into s-, p-, d-, and f-blocks, another classification of elements is based on their properties.
• This classification divides elements into Metals and Non-Metals.
• Metals constitute over 78% of known elements and are positioned on the left side of the Periodic Table.
• Typically, metals are solids at room temperature, except for mercury. Gallium and caesium also have notably low melting points (303K and 302K, respectively).
• Metals generally possess high melting and boiling points.
• They excel in conducting heat and electricity.
• Metals display malleability (capable of being flattened into thin sheets) and ductility (can be drawn into wires).
• On the other hand, non-metals are found on the upper right-hand side of the Periodic Table.
• Within a horizontal row, elements transition from metallic on the left to non-metallic on the right.
• Non-metals are primarily solids or gases at room temperature, with exceptions like boron and carbon.
• They have low melting and boiling points.
• Non-metals are poor conductors of heat and electricity.
• Most non-metallic solids are brittle and lack malleability and ductility.
• The metallic nature increases as we descend a group, while non-metallic characteristics intensify from left to right across the Periodic Table.
• The shift from metallic to non-metallic traits is gradual, as denoted by the thick zig-zag line.
• Elements such as silicon, germanium, arsenic, antimony, and tellurium situated along this line and forming a diagonal across the Periodic Table exhibit features that resemble both metals and non-metals.
• These elements are referred to as Semi-metals or Metalloids.

 

Relationship between Ionization Enthalpy and Atomic Radius:

• Ionization enthalpy and atomic radius are interconnected properties.
• Understanding their trends involves considering electron attraction to the nucleus and electron-electron repulsion.
• Two main factors influencing these trends are: (i) electron attraction toward the nucleus, and (ii) electron-electron repulsion.

Effective Nuclear Charge and Shielding:

• Valence electrons experience an effective nuclear charge due to shielding by inner core electrons.
• Shielding or screening reduces the net positive charge experienced by valence electrons.
• Shielding is more effective when inner shell orbitals are fully filled.
• Alkali metals exhibit effective shielding with a single outermost ns-electron following a noble gas electronic configuration.

Periodic Trend - Across a Period:

• Moving from lithium to fluorine across the second period, successive electrons enter the same principal quantum level.
• Shielding doesn't increase significantly to offset the stronger attraction between electrons and nucleus.
• Increasing nuclear charge dominates over shielding, leading to tighter hold on outermost electrons.
• Ionization enthalpy increases across a period due to stronger electron-nucleus attraction.

Group Trend - Down a Group:

• Moving down a group, outermost electrons are farther from the nucleus.
• Increased shielding by inner electrons outweighs the rising nuclear charge.
• Outermost electron removal requires less energy down a group.

Ionization Enthalpy Anomalies:

• First ionization enthalpy of boron (Z = 5) is slightly less than beryllium (Z = 4) due to electron configuration differences.
• Beryllium's ionization removes an s-electron, while boron's removes a p-electron.
• Penetration of 2s-electron is higher than 2p-electron, leading to greater shielding for boron's p-electron.
• Oxygen's smaller first ionization enthalpy compared to nitrogen's arises from electron configuration differences.
• In oxygen, electron-electron repulsion increases due to pairing in 2p-orbitals, making removal of the fourth 2p-electron easier.

Electron Gain Enthalpy:

• Electron Gain Enthalpy (∆_egH) measures the enthalpy change when a neutral gaseous atom gains an electron to form a negative ion (anion).
• This process is represented by the equation: X(g) + e– → X–(g).

Exothermic and Endothermic Processes:

• The addition of an electron to an atom can be either exothermic (energy released) or endothermic (energy absorbed), depending on the element.
• Elements like halogens (group 17) release energy when gaining an electron due to reaching stable noble gas configurations.
• Noble gases, however, have positive electron gain enthalpies as the added electron enters a higher principal quantum level, resulting in an unstable electronic configuration.

Trends Across the Periodic Table:

• Electron gain enthalpy is more negative towards the upper right of the periodic table before the noble gases.
• Generally, electron gain enthalpy becomes more negative as atomic number increases across a period.
• Increasing effective nuclear charge across a period makes it easier to add an electron to smaller atoms due to stronger attraction to the nucleus.

Trends Down a Group:

• Electron gain enthalpy becomes less negative as you move down a group.
• The larger atomic size results in the added electron being farther from the nucleus.

Anomalies for Oxygen and Fluorine:

• Electron gain enthalpy of oxygen (O) and fluorine (F) is less negative than that of the following element in the period.
• Adding an electron to O or F places it in the smaller n = 2 quantum level, causing significant repulsion from other electrons present.
• For Sulfur (S) or Chlorine (Cl) in the n = 3 quantum level, the added electron occupies more space, leading to less electron-electron repulsion.

 

 

Electronegativity:

• Electronegativity is a qualitative measure of an atom's ability in a chemical compound to attract shared electrons towards itself.
• It's not a directly measurable quantity but can be estimated using various numerical scales.
• Notable scales include Pauling scale, Mulliken-Jaffe scale, and Allred-Rochow scale, with Pauling scale being the most widely used.
• Linus Pauling assigned an arbitrary value of 4.0 to fluorine as a reference point for electronegativity.

Variation in Electronegativity:

• Electronegativity varies based on the element it's bound to.
• It offers insights into the nature of the bonding force between atoms.

Trends in Electronegativity:

• Across a period (left to right), electronegativity generally increases (e.g., lithium to fluorine).
• Down a group (top to bottom), electronegativity generally decreases (e.g., fluorine to astatine).
• This trend is related to atomic radii: electronegativity increases across periods as atomic radii decrease, and it decreases down groups as atomic radii increase.

Relationship with Non-Metallic Properties:

• Non-metallic elements tend to gain electrons, making their electronegativity high.
• Electronegativity correlates with non-metallic properties of elements.
• Electronegativity is inversely related to metallic properties: higher electronegativity corresponds to lower metallic properties.
• Across a period, as electronegativity increases, non-metallic properties increase (metallic properties decrease).
• Down a group, as electronegativity decreases, non-metallic properties decrease (metallic properties increase).

Periodic Trends in Chemical Properties

• Valence and Oxidation States:
• Valence is a key characteristic property of elements and is linked to their electronic configurations.
• For representative elements, valence is often (though not always) equal to the number of electrons in their outermost orbitals, or eight minus the number of outermost electrons.
• The term "oxidation state" is commonly used interchangeably with valence.
• Oxidation States in Compounds:
• Consider compounds OF₂ and Na₂O with the elements F, O, and Na.
• Electronegativity order: F > O > Na.
• In OF₂, each fluorine (F) atom shares one electron with oxygen (O), resulting in F having an oxidation state of -1 due to its high electronegativity.
• Oxygen in OF₂ shares two electrons with fluorine atoms, leading to an oxidation state of +2.
• In Na2O, oxygen accepts two electrons (oxidation state -2) from two sodium (Na) atoms.
• Sodium loses an electron to oxygen, resulting in an oxidation state of +1.
• Oxidation state of an element in a compound is defined as the charge it acquires based on electronegativity considerations from other atoms in the molecule.
• Periodic Trends in Valence:
• Various periodic trends are observed in the valence of elements, specifically in hydrides and oxides.
• Table 3.9 displays some of these periodic trends in the valence of elements.
• The book discusses other periodic trends in the chemical behavior of elements separately.
• Variable Valence:
• Several elements exhibit variable valence.
• This variability is particularly notable in transition elements and actinoids.
• Detailed study of variable valence in these elements will be covered later in the material.

 

Anomalous Properties of Second Period Elements:

1. Covalent Character of First Elements in Groups 1 and 2:

• Lithium (Group 1) and beryllium (Group 2) exhibit covalent character in their compounds, unlike other alkali and alkaline earth metals.
• Contrast with the other group members that tend to form predominantly ionic compounds.

2. Diagonal Relationship:

• Similar behavior observed between lithium and magnesium, and between beryllium and aluminum.
• This relationship is referred to as diagonal relationship in the periodic properties.

3. Reasons for Different Chemical Behavior:

• Attributed to small size, large charge/radius ratio, and high electronegativity of the first group member.
• The small size and higher electronegativity influence bonding and reactivity patterns.

4. Valence Orbitals:

• First member of a group has only four valence orbitals (2s and 2p) available for bonding.
• Second member of the group possesses nine valence orbitals (3s, 3p, 3d).

5. Maximum Covalency:

• Due to limited valence orbitals, the first group member's maximum covalency is 4 (e.g., boron can only form BF₄).
• Other group members can expand their valence shells, accommodating more than four electron pairs (e.g., aluminum forms AlF₆³⁻).

6. Formation of Multiple Bonds:

• First member of p-block elements displays enhanced ability to form pπ - pπ multiple bonds.
• Examples include C=C, C≡C, N=N, N≡N.
• Also forms multiple bonds with other second period elements, like C=O, C=N, C≡N, N=O.

 

1. Atomic Radii Change in Transition Metals (3d Series):
• The change in atomic radii among transition metals (3d series) is notably smaller compared to representative elements in the same period.
• This trend also holds true for inner-transition metals (4f series), where the change in atomic radii is even smaller.
2. Ionization Enthalpies and Electropositivity:
• Transition metals' ionization enthalpies fall between those of s-block and p-block elements.
• Consequently, these metals exhibit lower electropositivity when compared to group 1 and 2 metals.
3. Group Trends: Increase in Atomic and Ionic Radii:
• Within a group, there's a consistent increase in atomic and ionic radii as the atomic number rises.
• This increase leads to a gradual decrease in ionization enthalpies.
• Electron gain enthalpies also generally decrease (with exceptions in some third period elements), particularly in the case of main group elements.
4. Group Trends: Metallic and Non-Metallic Character:
• Down a group, the trend shows an increase in metallic character and a decrease in non-metallic character.
• This shift in properties can be linked to the elements' reducing and oxidizing behaviors, which will be explained later.
5. Transition Elements' Exceptional Trend:
• In contrast to main group elements, transition elements exhibit a reverse trend.
• This anomalous trend can be rationalized by considering factors like atomic size and ionization enthalpy.

 

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