Ch8: Periodic Properties of the Elements – Structured Study Notes

Chapter 8: Periodic Properties of the Elements

8.2 The Development of the Periodic Table

The periodic table is a foundational tool in chemistry, organizing elements according to their properties and atomic structure.

• Dmitri Mendeleev is credited with the modern arrangement of the periodic table, grouping elements with similar characteristics.

• Originally, elements were arranged by increasing atomic mass, but this led to inconsistencies (e.g., Tellurium and Iodine).

• The current table is organized by increasing atomic number (number of protons), which aligns with electron configuration and chemical properties.

• This arrangement allows prediction of the number of outer (valence) electrons, which largely determine chemical behavior.

8.3 Electron Configurations, Valence Electrons, and the Periodic Table

Electron configuration describes the distribution of electrons among atomic orbitals, which is key to understanding element properties.

• Valence electrons are the outermost electrons, crucial for chemical bonding and reactivity.

• For main group elements (groups 1–2, 13–18), valence electrons occupy the highest principal quantum number orbital; all others are core electrons.

• Orbital blocks:

  • s-block: Groups 1–2 (2 electrons in s orbital)

  • p-block: Groups 13–18 (6 electrons in p orbital)

  • d-block: Groups 3–12 (10 electrons in d orbital)

  • f-block: Lanthanoids and Actinoids (14 electrons in f orbital)

• Electron configurations can be written relative to the previous noble gas (e.g., Cl: $[Ne]3s^2 3p^5$).

• For d- and f-block elements, valence electrons include those in the highest principal quantum number and the outermost d or f electrons.

Example: Silicon (Si, atomic number 14) has the configuration $1s^2 2s^2 2p^6 3s^2 3p^2$; the valence electrons are in the 3s and 3p orbitals.

8.4 Explanatory Power of the Quantum Mechanical Model

The quantum mechanical model explains the arrangement and behavior of electrons in atoms, which underlies periodic trends.

• Electrons fill orbitals in a restricted manner, following the Aufbau principle.

• Noble gases (group 18) have full quantum levels and are generally stable and unreactive.

• Elements tend to gain or lose electrons to achieve a noble gas configuration, resulting in low-energy, stable states.

• Group 17 elements (halogens) have 7 valence electrons and readily gain one to achieve stability; group 16 elements gain two.

• Group 1 and 2 elements lose electrons to become cations with noble gas configurations.

8.5 Periodic Trends: Size & Effective Nuclear Charge

Atomic size and effective nuclear charge are key periodic trends that influence element properties.

• Atomic radius is determined by the average bonding radii in compounds; it is smaller than the van der Waals radius.

• Nonbonding atomic radius (van der Waals radius): Size of atoms not bound to others (e.g., noble gases).

• Bonding atomic radius (covalent radius): Half the distance between nuclei in a covalent bond.

• Across a period, atomic radius decreases due to increasing nuclear charge attracting electrons more strongly.

• Down a group, atomic radius increases as additional electron shells are added.

• Effective nuclear charge ($Z_{eff}$): The net positive charge experienced by valence electrons, calculated as $Z_{eff} = Z - S$ (where $Z$ is the number of protons and $S$ is the shielding constant).

• Slater's rules provide a method to estimate $Z_{eff}$ by accounting for electron shielding.

Example Table:

Element	Na	Mg	Al	Si	P	S	Cl	Ar
$Z_{eff}$	2.20	2.85	3.50	4.15	4.80	5.45	6.10	6.75
Radius (pm)	166	141	121	111	107	102	106	106

As $Z_{eff}$ increases (left to right), atomic radius decreases.

8.6 Ionic Radii

Ionic radius refers to the size of an ion, which changes when atoms gain or lose electrons.

• Cations (positive ions) are smaller than their neutral atoms because loss of electrons reduces electron-electron repulsion and increases effective nuclear charge.

• Anions (negative ions) are larger than their neutral atoms due to increased electron-electron repulsion and unchanged nuclear charge.

• Trends:

  • Ionic radius increases down a group.

  • Ionic radius decreases across a period (for cations and anions separately).

Example: Cl$^-$ has the configuration $[Ne]3s^2 3p^6$ and is larger than neutral Cl.

8.7 Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom.

• First ionization energy ($IE_1$): Energy to remove the first electron.

• Second ionization energy ($IE_2$): Energy to remove a second electron from the resulting cation.

• Trends:

  • $IE_1$ increases across a period (atoms hold electrons more tightly).

  • $IE_1$ decreases down a group (outer electrons are farther from the nucleus).

• Exceptions:

  • Group 2 elements have higher $IE_1$ than group 13 due to filled s orbitals.

  • Group 15 elements have higher $IE_1$ than group 16 due to electron pairing and repulsion in p orbitals.

Example Equation: $Na(g) \rightarrow Na^+(g) + e^- \quad IE_1 = 496\ \text{kJ mol}^{-1}$

8.8 Electron Affinities and Metallic Character

Electron affinity is the energy change when an atom gains an electron; metallic character describes how readily an element loses electrons.

• Electron affinity (EA): Generally increases across a period, but trends are less clear than for radius or ionization energy.

• Noble gases have low or negative EA values (do not readily gain electrons).

• Metallic character: Increases down a group and decreases across a period; metals lose electrons easily and form cations.

8.9 Examples of Periodic Chemical Behaviour

Representative groups show characteristic trends and reactivity.

• Alkali metals (Group 1):

  • Low ionization energies, large atomic radii, high reactivity.

  • Density increases down the group; melting point decreases.

  • React vigorously with water and halogens.

• Alkaline earth metals (Group 2):

  • Higher ionization energies and melting points than Group 1.

  • Reactivity increases down the group.

  • React with water and halogens, but less vigorously than Group 1.

• Halogens (Group 17):

  • High electron affinities, strong oxidizing agents.

  • Reactivity and boiling point increase down the group.

  • React with metals to form halides and with hydrogen to form hydrogen halides.

• Noble gases (Group 18):

  • Very high ionization energies, low reactivity.

  • Boiling point increases down the group.

  • Some heavier noble gases (Kr, Xe) can form compounds under extreme conditions.

Additional info: These notes are based on slides and textbook references from "Chemistry: A Molecular Approach, 4th Canadian Edition" by Tro, Fridgen, and Shaw (2023).