The oxidation states of Period 3 elements can provide insights into the solubility of their oxides and chlorides in water, although solubility is also influenced by other factors like lattice energy and hydration energy. Generally, compounds with lower oxidation states, like sodium oxide (Na₂O) and magnesium oxide (MgO), are highly soluble in water. This is because these basic oxides, formed by elements in lower oxidation states (+1 and +2 respectively), readily react with water to form hydroxides. As we move across the period to elements with higher oxidation states, the solubility patterns change. Aluminium oxide (Al₂O₃) is amphoteric and less soluble, while silicon dioxide (SiO₂) is virtually insoluble due to its macromolecular structure. The acidic oxides of phosphorus, sulfur, and chlorine (e.g., P₄O₁₀, SO₂, Cl₂O) are generally soluble, as they react with water to form acidic solutions. For chlorides, those of elements with lower oxidation states (like NaCl and MgCl₂) are highly soluble due to their ionic nature, which allows them to dissociate easily in water. However, the solubility decreases for chlorides with covalent character, such as SiCl₄, which is hydrolysed by water. Therefore, while oxidation states offer a guide, solubility in water is a complex trait that depends on multiple factors.

Oxidation states play a significant role in determining the bonding and structure of Period 3 chlorides. Starting with sodium chloride (NaCl), where sodium has a +1 oxidation state, the compound forms an ionic bond due to the complete transfer of an electron from sodium to chlorine. As we move to magnesium chloride (MgCl₂) with magnesium in a +2 oxidation state, the ionic character strengthens due to the increased charge difference between magnesium and chlorine ions. In aluminium chloride (AlCl₃), with aluminium in a +3 oxidation state, the compound exhibits ionic bonding at lower temperatures but shows covalent character at higher temperatures due to the smaller size and higher charge of the aluminium ion. Silicon tetrachloride (SiCl₄), where silicon has a +4 oxidation state, forms covalent bonds, as silicon shares electrons with chlorine atoms. Phosphorus and sulfur chlorides, like PCl₅ and S₂Cl₂, where these elements have higher oxidation states, also exhibit covalent bonding. The higher the oxidation state, the more electronegative the central atom becomes, influencing the type of bonding and consequently the molecular structure. Thus, the oxidation state directly impacts the nature of bonding (ionic or covalent) and the resultant structural attributes of these chlorides.

The concept of oxidation states in Period 3 elements is crucial in understanding their reactivity with water. Elements with lower oxidation states, like sodium and magnesium, react vigorously with water. Sodium, with a +1 oxidation state, reacts to form sodium hydroxide (NaOH) and hydrogen gas, indicating its high reactivity due to the ease of losing its single valence electron. Magnesium, in a +2 oxidation state, also reacts with water, albeit more slowly, to form magnesium hydroxide (Mg(OH)₂) and hydrogen. However, as we move across the period and the oxidation states increase, the reactivity with water decreases. Aluminium, with a +3 oxidation state, forms a protective oxide layer, making it less reactive with water under normal conditions. Silicon, phosphorus, sulfur, and chlorine, with even higher oxidation states, do not react with water in the same way. Silicon dioxide (SiO₂) is essentially unreactive with water, while phosphorus, sulfur, and chlorine oxides react differently, often forming acidic solutions. Understanding the oxidation states provides insights into the energy required to remove electrons from an element, which is a key factor in determining its reactivity with water.

Elements in Period 3 form different types of oxides due to variations in their electronic configurations and the resulting differences in their abilities to lose or share electrons. Starting with sodium, which forms a basic oxide (Na₂O), the nature of the oxides changes across the period. Magnesium oxide (MgO) is also basic, while aluminium oxide (Al₂O₃) is amphoteric, displaying both acidic and basic properties. Silicon dioxide (SiO₂) is a giant covalent structure and is acidic. Moving further, the oxides of phosphorus, sulfur, and chlorine (such as P₄O₁₀, SO₂, and ClO₂) are increasingly acidic. This transition from basic to acidic oxides is a direct result of the increasing oxidation states of the elements. Basic oxides are formed by elements with lower oxidation states (like +1 in Na₂O and +2 in MgO), where the element donates electrons to oxygen. In contrast, the acidic oxides are formed by elements with higher oxidation states (like +5 in P₄O₁₀), where the element has a stronger pull on the shared electrons in its bonds with oxygen. This trend illustrates how oxidation states influence the acid-base nature of oxides, reflecting the underlying electronic structure of the elements.

The trend in electronegativity across Period 3 significantly affects the oxidation states of its elements in their compounds. Electronegativity increases from left to right across the period, meaning that elements towards the right have a stronger tendency to attract electrons. Sodium and magnesium, being less electronegative, easily lose their valence electrons, resulting in lower oxidation states (+1 for Na and +2 for Mg). As we move towards elements like aluminium and silicon, the increasing electronegativity reflects in their higher oxidation states (+3 and +4 respectively), as these elements still lose electrons but with less ease compared to their predecessors. Phosphorus, sulfur, and chlorine, with even higher electronegativities, exhibit even higher oxidation states in their compounds. These elements are more capable of attracting and sharing electrons, leading to the formation of compounds with oxidation states of +5, +4, +6 (for sulfur), and varying states for chlorine, including +1, +3, +5, and +7. The increasing electronegativity thus correlates with the ability of an element to maintain a higher oxidation state by either sharing more electrons (in covalent bonding) or by having a stronger pull on the electrons in ionic compounds.