Period 3
Module 3: Chemistry of the Elements
Chapter 13: Elements and Periodicity - Period 3
Learning Objectives
Variations in Physical Properties: Describe how physical properties of the period 3 elements are influenced by their structure and bonding.
Element Reactions: Outline the reactions of period 3 elements with oxygen, water, and chlorine.
Chemical Bonding in Compounds: Predict the type of chemical bonding present in the oxides and chlorides of these elements.
Reactions of Oxides and Chlorides: Describe the reactivity of oxides and chlorides with water.
Trends in Acid/Base Behaviour: Explain the trends in acid/base behaviour among oxides and chlorides across period 3.
Introduction
The periodic table illustrates how both physical and chemical properties of elements vary systematically with atomic number, a phenomenon known as 'periodicity'. This periodic nature of the elements has significantly aided chemistry's advancement and our understanding of elemental properties.
Physical Properties of Elements
Physical properties are categorized into:
Atomic Properties: Depend solely on the structure of individual atoms.
Electronic Configuration: Arrangement of electrons in an atom determined by atomic number.
Properties include atomic radii, ionization energy, and electronegativity.
Bulk Properties: Depend not only on atomic structure but also how atoms are packed together.
Properties include melting point, density, and electrical conductivity.
Atomic Properties
Electronic Configuration
The electronic configuration describes the arrangement of electrons within an atom and is displayed in Table 13.1 for the elements in period 3:
Sodium (Na): [Ne] 3s
Magnesium (Mg): [Ne] 3s²
Aluminium (Al): [Ne] 3s² 3p¹
Silicon (Si): [Ne] 3s² 3p²
Phosphorus (P): [Ne] 3s² 3p³
Sulfur (S): [Ne] 3s² 3p⁴
Chlorine (Cl): [Ne] 3s² 3p⁵
Argon (Ar): [Ne] 3s² 3p⁶
The number of valence electrons increases across period 3, which plays a critical role in determining bonding and properties.
Atomic Radius
The atomic radius is defined as the distance from the nucleus to the outermost electron shell.
There are different methods of measuring atomic radius:
Van der Waals Radius: Half the internuclear distance between non-bonded similar atoms.
Covalent Radius: Half the distance between the nuclei of two bonded atoms.
The covalent radius is typically the smallest measurement, while the van der Waals radius is the largest. Trends show that as you move from left to right across period 3, the atomic radius decreases due to increasing nuclear charge, demonstrated in Figure 13.3.
Ionic Radius
Ionic radius is determined from internuclear distances in ionic compounds. Key points include:
Cations: Metals to the left (Na, Mg, Al) lose electrons to form smaller cations compared to their parent atoms.
Anions: Non-metals to the right (P, S, Cl) gain electrons, resulting in larger anions.
The trend in ionic radii shows:Decrease for cations: Na⁺, Mg²⁺, Al³⁺.
Decrease for anions: P³⁻, S²⁻, Cl⁻.
The notable jump occurs between Al³⁺ and P³⁻ due to the addition of an electron shell.
First Ionization Energy
Defined as the energy required to remove the most loosely held electron from a neutral atom, trends across period 3 show an increase in first ionization energy as the atomic radius decreases and nuclear charge increases (Figure 13.5). Notable irregularities occur:
Aluminium: Lower than magnesium due to less nuclear attraction on the unpaired 3p electron.
Sulphur: Lower than phosphorus because of repulsion between paired electrons in the same orbital.
Electronegativity
Electronegativity measures an atom's ability to attract electron pairs in a bond, with fluorine set as the standard (4.0). Electronegativity increases across period 3 as atomic radius decreases (Figure 13.6).
Bulk Properties
Melting Point
The melting point is the temperature at which a solid turns into a liquid, influenced by bonding types. Trends include:
Rapid increase from sodium to silicon due to strong metallic and covalent bonds.
Significant drop from silicon to phosphorus, with non-metals showing lower melting points due to weak van der Waals forces.
Argon's melting point is extremely low as it has negligible bonding forces.
Electrical Conductivity
A measure of a material's ability to conduct electric current, with trends showing increased conductivity from sodium to aluminium due to increased delocalized electrons. Silicon is semiconductive, while phosphorus, sulfur, chlorine, and argon do not conduct electricity due to lack of free-moving charge carriers.
Chemical Properties
Trends in Chemical Properties
Period 3 elements exhibit distinct trends in chemical properties, including in their reactions with oxygen, water, and chlorine, as displayed in Table 13.6.
Oxidation States of Oxides and Chlorides
The highest oxidation states correspond to the number of valence electrons.
Oxides: Na₂O (+1), MgO (+2), Al₂O₃ (+3), SiO₂ (+4), P₄O₁₀ (+5), SO₃ (+6), Cl₂O₇ (+7).
Chlorides: NaCl (+1), MgCl₂ (+2), AlCl₃ (+3), SiCl₄ (+4), PCl₅ (+5). Notably, sulfur's highest oxidation state in SC₁₄ is unstable.
Physical and Chemical Properties of the Oxides
As observed in Table 13.9, the physical properties of period 3 oxides change from ionic to covalent, leading to variances in acid/base behaviour. As such, oxides transition from strongly basic (Na₂O) to amphoteric (Al₂O₃) and end with strongly acidic (SO₃, Cl₂O₇). This behavior correlates with electronegativity differences and structural characteristics.
Physical and Chemical Properties of the Chlorides
Table 13.10 highlights that period 3 chlorides range from ionic (NaCl, MgCl₂) to covalent structures, influencing solubility and acidity in water. Ionic chlorides yield neutral solutions, while covalent chlorides hydrolyze to produce acidic solutions.
Summary
The position of an element in the periodic table correlates directly with its electronic structure.
Physical and chemical properties exhibit clear trends across periods and groups.
Changes in electron structures lead to variations in these properties.
Bulk physical properties exhibit similar trends to atomic properties in period 3.