The Periodic Table

Periodic Table Overview

• Used to classify elements and facilitate their study
• The modern periodic table is also known as the long form

Organization

• Based on Moseley's law: √f α Z (where f is frequency and Z is the atomic number)
• Periodic properties are a function of the atomic number: √f = a(Z-b), where a and b are constants
• Includes 18 groups and 7 periods
• It is divided into 4 blocks: s, p, d, and f

Electronic Configuration by Block

• s-block: G1, G2, ns¹⁻²
• p-block: G13 to G18, ns²np¹⁻⁶
• d-block: G3 to G12, (n-1)d¹⁻¹⁰ns⁰⁻² (Penultimate shell)
• f-block: (n-2)f¹⁻¹⁴(n-1)d⁰⁻¹ns² (Ante-penultimate shell)

F-Block Series

• 4f series (Lanthanoids): Z = 58 to 71, Last e⁻ enters 4f, Group 3, Period 6 and has 14 elements
• 5f series (Actinoids): Z = 90 to 103, Group 3, Period 7 and has 14 elements
• Lanthanum has Z = 57 and is located in Group 3 and is situated in the d-block.
• Actinium has Z = 89 and is located in Group 3and is situated in the d-block.

Representative Elements

• s and p blocks combined are known as representative elements
• Groups 1, 2, and 13-18

Transition Elements

• Elements with incompletely filled d-orbitals in their ground state or most common oxidation state
• Fe: [Ar] 3d⁶4s² is a transition element.
• Cu: [Ar] 4s¹ 3d¹⁰ is a ground state.
• Cu⁺²: [Ar] 4s⁰3d⁹ is a transition element in the oxidation state

Pseudo-Transition Elements

• Group 12 elements (Zn, Cd, Hg) are not transition elements
• Zn: [Ar] 4s²3d¹⁰ is completely filled with d-orbitals
• Zn⁺² [Ar] 4s⁰3d¹⁰

Diagonal Relationships

• Observed between certain elements in adjacent groups and periods.
• Reason: Similar ionic size and charge/size ratio.
• Example: Li and Mg, Be and Al
• Lithium ion (Li⁺) is approximately equal to the size of Magnesium ion (Mg⁺²)

Nomenclature of Elements with Z > 100

• 0 = nil
• 1 = un
• 2 = bi
• 3 = tri
• 4 = quad
• 5 = pent
• 6 = hex
• 7 = sept
• 8 = oct
• 9 = enn
• Suffix: -ium
• Example: Element with Z = 108 is named unnil octium (Uno)

Effective Nuclear Charge (Zeff)

• Zeff = Z - σ, where Z is the atomic number and σ is the Slater's constant (shielding constant)

Screening or Shielding Effect

• Inner electrons shield or screen outer electrons, which decreases the force of attraction
• Zeff also decreases as a result
• Order of screening effect: s > p > d > f meaning that f has a poor sheilding effect

Trends in Zeff

• Along a period, Zeff increases
• Along a group, Zeff decreases or remains the same

Trends Based on Charge

• More positive charge leads to higher Zeff: Fe⁺² < Fe⁺³, Mn < Mn⁺² < Mn⁺⁴ < Mn⁺⁷
• More negative charge leads to lower Zeff: O⁻¹ > O⁻², N > N⁻¹ > N⁻² > N⁻³, H⁺ > H > H⁻

Atomic Radius

• Distance between the nucleus and outermost electron
• It is impossible to find the exact value of atomic radius because Atoms are too small and their exact boundary is not known due to the Heisenberg uncertainity priciple (HUP)

Covalent Radius (Rc)

• Half of the internuclear distance between two atoms bonded by a covalent bond
• For homonuclear molecules (e.g., H₂, F₂, Cl₂): r(c) = d/2 where ‘d’ is the distance between the nucleai
• For heteronuclear molecules: d = r(A) + r(B) - 0.09|χ(A) - χ(B)|

Metallic Radius (Rm)

• Half of the internuclear distance between two atoms in a metallic bond
• r(m) = d/2

Van der Waals Radius (RVR)

• Half of the internuclear distance between two nearest molecules in a non-bonded state
• r(VR) = d/2
• Typically measured in Nobel gases

Trends in Atomic Radius

• Order: RVR > RM > Rc
• Size is inversely proportional to Zeff

Trends in Periodic Table

• Across a period: Zeff increases, size decreases
• Down a group: Zeff decreases, size increases

Ionic Radius

• For cations: Mg > Mg⁺¹ > Mg⁺²
• For anions: O > O⁻¹ > O⁻² Size increases with negative charge

Isoelectronic Species

• Species with the same number of electrons
• Example: N⁻³ > O⁻² > F⁻ > Na⁺ > Mg⁺² > Al⁺³
• Example: S⁻² > Cl⁻ > K⁺ > Ca⁺²

Atomic Size Question

• Arrange elements of the 2nd period in decreasing order of size: Ne > Li > Be > B > C > N > O > F
• Neon is largest due to its Van der Waals radius

Group 13 Exception Question

• Arrange the elemenets of group 13 in increasing order of size: B < Ga < Al < In < Tl
• Poor sheild of d-orbitals i.e., FOA increases and thus leads to size decreases.

Ionization Energy

• Energy required to remove an electron from an isolated gaseous atom
• Always endothermic (ΔH = +ve)
• M(g) → M⁺(g) + e⁻ (IE₁)
• M⁺(g) → M⁺²(g) + e⁻ (IE₂)
• M⁺²(g) → M⁺³(g) + e⁻ (IE₃)
• IE₁ < IE₂ < IE₃

Trends in Ionization Energy

• Ionization energy is proportional to Zeff
• Along a period: Zeff increases, IE increases
• Along a group: Zeff decreases, IE decreases

Exceptions to Ionization Energy Trends in Group 13

• IE sequence for Group 13: B > Tl > Ga > Al > In

Arranging Elements According to Ionization Order

• Li < B <Be < C < O < N < F < Ne

Spartan Trick for Ionization Number

• The valence electron is equal to the biggest difference after number have been subtracted (ie - IE4 - IE3 = 30 →Valence e⁻)

Electron Gain Enthalpy (ΔegH)

• Energy required or released when an electron is added to the outermost shell of an isolated gaseous atom
• Can be endothermic or exothermic
• X(g) + e⁻ → X⁻(g) ΔegH₁
• X⁻(g) + e⁻ → X⁻²(g) ΔegH₂

Trends in Electron Gain Enthalpy

• ΔegH is proportional to Zeff
• Along a period: Zeff increases, ΔegH increases
• Along a group: Zeff decreases, ΔegH decreases

Exceptions

• In the majority of cases, ΔegH₁ is negative (exothermic)
• Exceptions include Be, Mg, N, and noble gases, where ΔegH₁ is positive (endothermic) due to stable electronic configurations

Successive Electron Gain Enthalpies

• Successive electron gain enthalpies (ΔegH₂, ΔegH₃, ...) are always positive (endothermic)
• O(g) + e⁻ → O⁻(g) ΔegH₁ = -ve
• O⁻(g) + e⁻ → O⁻²(g) ΔegH₂ = +ve due to anion repulsion

Electron Gain Enthalpy Exceptions

• Group 16: S > Se > Te > Po > O
• Group 17: Cl > F > Br > I because e-e repulsion factor is greater

Electronegativity

• Tendency of an atom to attract a shared pair of electrons towards itself
• δ⁺ A---B δ⁻ (If electronegativity of A is less than electronegativity of B (ENA