Fiche de révision : Periodic Properties

📋 Course Outline

1. Historical Periodic Classification
2. Modern Periodic Law
3. Periodic Table Blocks
4. Metals Non-metals Metalloids
5. Shielding and Effective Nuclear Charge
6. Atomic and Ionic Radii
7. Ionisation Enthalpy
8. Electron Gain and Electronegativity
9. Valence and Oxidation States

📖 1. Historical Periodic Classification

🔑 Key Concepts & Definitions

• Döbereiner's Triads : Groups of three elements with similar chemical properties in which the atomic mass of the middle element was approximately the arithmetic mean of the first and third.
• Newlands' Law of Octaves : Stated in 1866 that, after arranging elements by increasing atomic mass, certain properties repeated at every eighth element.

★ Must-know

• Mendeleev arranged elements by increasing atomic mass and stated that their physical and chemical properties were periodic functions of atomic masses.

Further detail

• Mendeleev predicted new elements including scandium, gallium and germanium, enabled systematic study of elements, and helped correct some atomic masses.

• Mendeleev's major defects were anomalous pairs such as Ar/K, Co/Ni and Te/I, the position of hydrogen, and difficulty placing isotopes.

💡 Memory Hook

Triads → Octaves → Mendeleev

📖 2. Modern Periodic Law

🔑 Key Concepts & Definitions

• Modern Periodic Law : States that the physical and chemical properties of elements are periodic functions of their atomic numbers.

★ Must-know

• Henry Moseley showed that atomic number is more fundamental than atomic mass for periodic classification.

• The long-form periodic table arranges elements in increasing atomic number across 18 groups and 7 periods.

Further detail

• Elements in the same group have similar, though not identical, electronic configurations and properties.

💡 Memory Hook

Atomic number, not atomic mass

📖 3. Periodic Table Blocks

🔑 Key Concepts & Definitions

• s-block : Contains elements whose last electron enters an s-orbital, has the general configuration ns¹ or ns², and occupies groups 1–2.
• f-block : Contains elements whose last electron enters an f-orbital, has the general configuration (n−2)f¹–¹⁴(n−1)d⁰–¹ns², and includes lanthanides and actinides.

📝 Essential Points

• The d-block is characterized by variable valency, coloured compounds and catalytic properties.

📖 4. Metals Non-metals Metalloids

★ Must-know

• Metals are mainly on the left of the periodic table and are generally solid, malleable, ductile and conducting.

• Non-metals are mainly on the right of the periodic table and may be solid, liquid or gas.

Further detail

• Metalloids occur along the zig-zag boundary and include B, Si, Ge, As, Sb and Te.

📖 5. Shielding and Effective Nuclear Charge

🔑 Key Concepts & Definitions

• Shielding Effect : The reduction of nuclear attraction experienced by an outer electron because of inner-shell electrons.

★ Must-know

📐 Formula — Effective nuclear charge is given by $Z_{\mathrm{eff}} = Z - \sigma$, where $Z$ is nuclear charge or atomic number and $\sigma$ is the shielding constant.

• Effective nuclear charge increases across a period and increases only slowly down a group, remaining approximately constant for many outer electrons.

Further detail

• For a 2p electron of fluorine, the shielding constant is σ = 3.80 and the effective nuclear charge is $Z_{\mathrm{eff}} = 9 - 3.80 = 5.20$.

📖 6. Atomic and Ionic Radii

🔑 Key Concepts & Definitions

• Atomic Radius : Decreases across a period because the shell number remains constant while nuclear charge and effective nuclear charge increase, pulling electrons closer to the nucleus.
• Isoelectronic Species : Have the same number of electrons, as in F⁻, Na⁺ and Mg²⁺, which each have 10 electrons.

★ Must-know

• Atomic radius increases down a group because the number of electron shells increases while effective nuclear charge does not increase dramatically.

📌 Anions are larger than their parent atoms because added electrons increase electron–electron repulsion, whereas cations are smaller because they have fewer electrons while nuclear charge remains the same.

• In an isoelectronic series, ionic size decreases as nuclear charge increases; the order is O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺ > Al³⁺.

Further detail

• Metallic radii in the third-row d-block are similar to those in the second-row d-block because of lanthanide contraction associated with poor shielding by f-orbitals.

💡 Memory Hook

Cations contract; anions expand

📖 7. Ionisation Enthalpy

🔑 Key Concepts & Definitions

• Ionisation Enthalpy : The energy required to remove an electron from an isolated gaseous atom or ion, expressed in kJ mol⁻¹.

★ Must-know

📐 Formula — The first and second ionisation processes are $X(g) \rightarrow X^+(g) + e^-$ and $X^+(g) \rightarrow X^{2+}(g) + e^-$, respectively.

• Successive ionisation enthalpies increase because removing an electron from an increasingly positive ion is more difficult:

  • IE₂ > IE₁
  • IE₃ > IE₂
  • so on

• Ionisation enthalpy generally increases across a period and decreases down a group.

Further detail

• The important ionisation-enthalpy exceptions are Be > B and N > O because B loses a 2p electron more easily than Be loses a 2s electron, while paired 2p electrons in O create extra repulsion.

📖 8. Electron Gain and Electronegativity

🔑 Key Concepts & Definitions

• Electron Gain Enthalpy : The enthalpy change when a gaseous atom gains an electron to form an anion, represented as $X(g) + e^- \rightarrow X^-(g) + energy$.
• Electronegativity : The tendency of an element in a molecule to attract the shared pair of electrons toward itself, measured on the Pauling scale.

★ Must-know

• Electron gain enthalpy generally becomes more negative across a period and less negative down a group.

• Noble gases have practically zero or positive electron gain enthalpies because they already possess stable ns²np⁶ configurations, while halogens have a strong tendency to gain an electron.

• Electronegativity increases across a period and decreases down a group because effective nuclear charge and attraction increase across a period, while shells and atomic radius increase down a group.

Further detail

• The electron gain enthalpy of oxygen is less negative than that of sulfur, and fluorine's is less negative than chlorine's, because the added electron enters the compact n = 2 level and experiences greater repulsion.

• Fluorine is assigned an electronegativity value of 4 on the Pauling scale in the notes and is the most electronegative element.

📖 9. Valence and Oxidation States

🔑 Key Concepts & Definitions

• Oxidation State : Can be treated as the charge acquired by an atom based on electronegative considerations.

★ Must-know

• For representative elements, valence is generally related to the number of outermost electrons or, for H-type valency, to eight minus that number.

• The common valencies for groups 1, 2, 13 and 14 are 1, 2, 3 and 4, while groups 15, 16 and 17 commonly show valencies 3/5, 2/6 and 1/7, respectively.

Further detail

• Groups commonly show characteristic positive or negative oxidation states, with gradual changes across a period.

📊 Synthesis Tables

Periodic Trends

Block Characteristics