Understanding the Modern Periodic Table in IB Chemistry

The modern periodic table represents a fundamental organizing principle in chemistry, arranged by increasing atomic number (Z). Nuclear charge, equivalent to the atomic number, represents the total proton count in an atom's nucleus. The table's structure reflects electron configurations, with groups indicating valence electron count and periods showing principal energy levels.

Definition: Nuclear charge is the total positive charge in an atomic nucleus, determined by the number of protons present.

The periodic table divides elements into distinct blocks (s, p, d, f) based on their electron configurations. The s-block contains alkali metals and alkaline earth metals, while the p-block houses main group elements including halogens and noble gases. Transition metals occupy the d-block, with lanthanoids and actinoids forming the f-block series.

Elements also classify into metals, metalloids, and non-metals based on their properties. Metals typically occupy the left and center portions, while non-metals cluster in the upper right. Metalloids form a diagonal boundary between these regions, exhibiting intermediate characteristics.

Periodic Trends and Atomic Properties in IB Chemistry

Periodic trends demonstrate systematic variations in atomic properties across the periodic table. Atomic radius, a fundamental property, measures the effective size of atoms through different methods depending on bonding type.

Vocabulary: Periodicity refers to the recurring patterns of physical and chemical properties observed across the periodic table.

Atomic radius generally increases down groups due to additional electron shells but decreases across periods due to increasing nuclear charge. This trend results from the interplay between electron shielding and nuclear attraction. Inner electrons shield outer electrons from nuclear attraction, affecting atomic size.

The concept of electron shielding proves crucial in understanding these trends. As atomic number increases across a period, the growing nuclear charge exerts stronger attraction on electrons, while electron shielding remains relatively constant, resulting in smaller atomic radii.

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.