Properties of Group 1 and Group 7 Elements

The alkali metals of Group 1 demonstrate distinctive chemical properties due to their single valence electron. These highly reactive metals require storage under liquid paraffin to prevent reaction with oxygen and water vapor. Their reactivity increases down the group as atomic size increases and nuclear attraction to the valence electron decreases, making electron loss progressively easier.

The halogens in Group 7 present contrasting behavior with their seven valence electrons. As strong oxidizing agents, they readily accept electrons to achieve noble gas configuration. Unlike Group 1, their reactivity decreases down the group because larger atoms have reduced ability to attract additional electrons. This pattern becomes evident in displacement reactions where more reactive halogens replace less reactive ones from their compounds.

Example: When chlorine gas is bubbled through potassium bromide solution, it displaces bromine: Cl₂ + 2KBr → 2KCl + Br₂

Specific properties of halogens vary systematically. Chlorine exists as a pale-green gas with a characteristic odor, while bromine appears as a deep-red liquid producing red-brown vapors. Iodine, a grey solid, sublimes to produce distinctive purple vapors. These elements occur naturally in various forms, with chlorine and bromine found predominantly as halides in seawater, while iodine exists in rocks and seaweed.

Understanding Complex Ions and Color in IB Chemistry

The formation of colored complex ions represents a fascinating intersection of electronic structure and visible light in IB Chemistry notes. When ligands approach a transition metal ion, they cause a critical splitting of the metal's d-orbitals into two distinct energy levels - a higher energy set and a lower energy set. This splitting phenomenon forms the foundation for understanding why complex ions exhibit specific colors.

Definition: Complex ions are formed when transition metal ions bond with ligands (molecules or ions that can donate electron pairs), resulting in a characteristic splitting of d-orbitals that determines their color properties.

The color we observe in transition metal complexes results from a sophisticated electronic process. When white light interacts with these complexes, electrons in the lower energy d-orbitals absorb specific wavelengths of light to jump to higher energy levels. The wavelength absorbed corresponds directly to the energy gap between these orbital levels, while the complementary color is transmitted and visible to our eyes. For instance, if a complex absorbs orange light (around 585nm), it appears blue to our eyes.

Example: Copper(II) complexes often appear blue because they absorb orange light (585nm). The energy difference between their split d-orbitals corresponds exactly to the energy of orange light, causing this specific absorption pattern.

Several key factors influence the final color of complex ions, making this topic crucial for Periodic trends IB Chemistry HL notes. The oxidation state of the transition metal significantly affects the energy gap between orbital levels. Additionally, the geometric arrangement of ligands around the metal center influences how strongly the d-orbitals split, while the nature of the transition metal itself - including its nuclear charge and electron configuration - plays a fundamental role in determining color properties.