Water is one of the most complex fluids on Earth. Even after intense study, there are many aspects regarding the structure, properties, and chemistry of water that are not well understood. In this chapter, we highlight the attributes of water that dictate many of the reactions that take place between water and oxides. We start with a single water molecule and progress to water clusters, then finally to extended liquid and solid phases. This chapter provides a baseline for evaluating what happens when water encounters simple ions, soluble oxide complexes called hydrolysis products, and extended oxide phases. The primary phenomenon highlighted in this chapter is hydrogen bonding. Hydrogen bonding dominates the structure and properties of water and influences many water–oxide interactions. A single water molecule has eight valence electrons around a central oxygen anion. These electrons are contained in four sp3-hybridized molecular orbitals arranged as lobes that extend from the oxygen in a tetrahedral geometry. Each orbital is occupied by two electrons. Two of the lobes are bonded to protons; the other two lobes are referred to as lone pairs of electrons. The H–O–H bond angle of 104.5° is close to the tetrahedral angle of 109.5°. The O–H bond length in a single water molecule is 0.96 Ǻ. It is important to recognize that this bond length is really a measure of the electron density associated with the oxygen lone pair bonded to the proton. This is because a proton is so incredibly small (with an ionic radius of only 1.3·10−5 Ǻ) that it makes no contribution to the net bond length. The entire water molecule has a hard sphere diameter of 2.9 Ǻ, which is fairly typical for an oxygen anion. This means the unoccupied lone pairs are distended relative to the protonated lone pairs, extending out to roughly 1.9 Ǻ. The unequal distribution of charges introduces a dipole within the water molecule that facilitates electrostatic interactions with other molecules.