Chemistry of Main Group and Transition Elements

Chemistry of Main Group and Transition Elements

Course Overview

• C 12012/CHE 12012 (30 h) – 2025
  • Chemistry of Main Group Elements (13 h)
  • Chemistry of Transition Elements (10 h)
  • Chemistry of Lanthanoids & Actinoids (7 h)

Objectives

• To teach the Periodic Table & Periodicity
• Chemistry of s-block elements
• Chemistry of p-block elements

Intended Learning Outcomes

• Discuss classification of elements in the modern Periodic Table.
• Discuss trends in physical and chemical properties of elements down a Group and across a period.
• Explain the term diagonal relationship.
• Explain 3-centre 2-electron bonding.

Recommended Textbooks

• Cartherine E Housecroft and Alan G Sharpe (2012), Pearson, England, Inorganic Chemistry, 4th Edition.
• J. D. Lee (2006), Blackwell science, Concise Inorganic Chemistry, 5th Edition.
• Shriver and Atkins Inorganic Chemistry, (2010) P.W. Atkins, T.L. Overton, J.P. Rourke, M.T. Weller and F.A. Armstrong, 5th Edition
• The Periodic Table and Periodicity (2015), The Open University of Sri Lanka

The Periodic Table & Periodicity

• Elements arranged in ↑ order of atomic #.
• # of elements = 118.
• Groups are numbered from 1-18.
• #s given for periods = 1 – 7.

IUPAC Recommended Names for Groups

• Gp 1: Alkali metals
• Gp 2: Alkaline earth metals
• Gp 16: Chalcogens
• Gp 17: Halogens
• Gp 18: Noble gases

Periods and Electronic Configuration

• 1st period: 1s (2 elements)
• 2nd period: 2s 2p (8 elements)
• 3rd period: 3s 3p (8 elements)
• 4th period: 4s 3d 4p (18 elements)
• 5th period: 5s 4d 5p (18 elements)
• 6th period: 6s 4f 5d 6p (32 elements)

Group Number and Valence Configuration

• 1: $ns^1$
• 2: $ns^2$
• 13: $ns^2 np^1$
• 14: $ns^2 np^2$
• 15: $ns^2 np^3$
• 16: $ns^2 np^4$
• 17: $ns^2 np^5$
• 18: $ns^2 np^6$

s, p, d, and f Blocks

• Helium is not a p-block element.
• Except He, remaining elements in Gp 18 belong to p-block elements.
• In the four blocks (s, p, d, and f), s, p, d, and f levels are filled.

f-block

• Lanthanoids and Actinoids

Triads

• Vertical groups of d-block elements.

Electronic Configuration

• Aufbau principle, Hund’s rule, Screening (shielding) and penetration.

Periodicity (Periodic Law)

• When elements are arranged in increasing order of atomic #s, a periodic repetition of their physical and chemical properties is seen.

Metallic and Non-metallic Character

• Metallic character: Tendency to lose electrons.
• Non-metallic character: Tendency to gain electrons.

Trends in Metallic Character Down a Group

• Group ↓ radius ↑, IE ↓, become more metallic.
• This trend is most noticeable in Groups 13 to 16.
• e.g., Group 14: C - nonmetal, Si & Ge - metalloids, Sn & Pb - metals.

Atomic Size (Atomic Radius)

• Properties such as B.Pt, M.Pt, IE & EN depend on the atomic size.
• Elements can have covalent, ionic, metallic radii.

Atomic Radii

• Covalent radii & Metallic radii are jointly considered as atomic radii.
• Covalent radii - for nonmetallic elements.
• Metallic radii - for metallic elements.

Covalent Radius $(r_{cov})$

• For a homonuclear X-X single bond.
• $r<em>{cov} = (1/2) (2r</em>{cov})$

Metallic Radius $(r_m)$

• $r_m = ½$ (distance btn two nearest neighboring atoms of the solid metal)

Van der Waals Radius

• Van der Waals radius of an element $(r<em>v) >$ its covalent radius $(r</em>{cov})$
• e.g., $r<em>v$ and $r</em>{cov}$ for I are 2.15 Å and 1.33 Å respectively.

Ionic Radius

• $R$ = distance between the centers of neighboring cations and anions in an ionic compound.
• $r^+ + r^- = R$
• $r^+ = R - r^-$

Effective Nuclear Charges (Zeff)

• $Z_{eff} = Z – σ$
  • Z = nuclear charge
  • σ = screening (or shielding) constant

Trends in Atomic Radii - Down a Group

• Group ↓ # of inner shells with electrons ↑
• valence electrons are shielded by the electrons in those inner shells.
• So attraction of valance electrons to the nucleus ↓ & atomic radius ↑

Trends in Atomic Radii – Across a Period

• Left to right across a period: Atomic radii ↓

Radius of Monoatomic Ions

• Monoatomic anions > their parent atoms (electron–electron repulsion ↑).
• Monoatomic cations < their parent atoms (attraction ↑).

Ionization Energy (IE)

• Energy required to remove the least tightly bound electron from the gas phase atom.
  • 1st IE: $M(g) ➝ M^+ (g) + e^–$
  • 2nd IE: $M^+ (g) → M^{2+} (g) + e^-$

Factors Affecting IE

• Atomic Radius ↓, IE ↑
• Electronic configuration (filled or half-filled shells → IE high).
• $Z_{eff}$ ↑, IE ↑
• Shielding by inner electrons ↑, IE ↓ (Shielding effect s > p > d > f)

Trends in IE Down a Group

• Group ↓ normally IE ↓

Trends in IE: Left to Right of a Period

• Left to right, normally IE ↑

Electron Affinity (EA)

• Amount of energy released when an electron is added to a neutral gaseous atom.
• $A(g) + e^- ➝ A^- (g)$

Factors Affecting EA

• $Z_{eff}$ of the valance shell receiving the electron↑, EA ↑
• Electronic configuration (e.g., half-filled orbital).

Trends in EA: Left to Right of a Period

• Effective nuclear charge ↑, EA ↑

Electronegativity (EN)

• Power of an atom to attract electrons when it is covalently bonded to other atom(s).

Factors Affecting EN

• EN ↑ when:
  • Size of the atom ↓
  • Close to having a filled shell of electrons
  • High EA (favorable - gaining electrons)

Pauling’s Method

• Fluorine (F) is given an arbitrary value of 4.0.

Periodic Trends in Electronegativity

• Across a period, EN ↑
• Group ↓ EN ↓

Importance of EN

• Useful for predicting bond type, dipole moments, and bond energies.

Variation of Ionic Character

• If $ꭓ<em>A - ꭓ</em>B = 1.7$:
  • 50 % ionic.
• If ꭓA - ꭓB < 1.7:
  • < 50 % ionic.
• If ꭓA - ꭓB > 1.7:
  • > 50 % ionic.

Electropositive Nature

• Describes element’s ability to lose electrons.
• Low IE, low electronegativity.

Chemical Properties of Metals & Nonmetals

• Metals with nonmetals → hard, nonvolatile solids (e.g., NaCl).
• Nonmetals → volatile molecular compounds (e.g., $PCl_3$).
• Metals combine → alloys (e.g., brass from copper and zinc).

Enthalpy of Atomization of Elements

• Energy required to form gaseous atoms.
• For solids - enthalpy change associated with the atomization of the solid.
• For molecular species - enthalpy of dissociation of the molecules.

Melting Points

• M. Pt depends on the strength of these forces.
• e.g., For a metal, MPt depends on the strength of metallic bonding

Factors Affecting Strength of Metallic Bonding

• Strength of metallic bonding ↑ with:
  • ↑ # of electrons in the delocalized sea of electrons
  • ↑ Charge of the cation
  • ↓ atomic radius

Trends in M.Pts Down a Group

• For alkali metals: Group-1 ↓ M. Pts ↓
• For halogens: Group ↓ M. Pts ↑ (van der Waals forces).

Oxidation States of Elements

• Maximum positive oxidation number = # of s- and p- electrons in the valance shell
  • E.g. Gp 13 elements $s^2p^1$: +3.

Inert Pair Effect

• Heavier elements of the p-block elements form more stable compounds when the: oxidation # = (Group oxidation #) - (2)

Diagonal Relationship

• Li, Be & B of the 2nd period show a relationship with the diagonally opposite member in the 3rd period (Mg, Al & Si).

Diagonal Relationship - Reason

• close similarity of the:
  • atomic radii
  • charge densities
  • electronegativities