Lernzettel: 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

Property	Across a period	Down a group
Atomic radius	Decreases	Increases
Effective nuclear charge	Increases	Increases slowly
Ionisation enthalpy	Generally increases	Generally decreases
Electron gain enthalpy	Generally becomes more negative	Generally becomes less negative
Electronegativity	Increases	Decreases

Block Characteristics

Block	Last electron	Groups or elements
s-block	s-orbital	Groups 1–2
p-block	p-orbital	Groups 13–18
d-block	d-orbital	Groups 3–12; transition metals
f-block	f-orbital	Lanthanides and actinides