Introduction In this Chapter, we will discuss how the nature of acid-base chemistry, which governs reactions involving proton (H + ) transfer, affecting the pH of systems ranging from biological fluids to the Earth’s oceans. The pH of a solution is a measure of its acidity or alkalinity and is determined by the balance between acids and bases existing in a state of dissociation equilibrium. In aqueous solutions, acids and bases establish an equilibrium between the dissociated (ionized) and undissociated (molecular) forms. This reversible reaction reaches a state of balance, described by an equilibrium constant K a , which measures the strength of the acid. Additionally, certain proportions of acidic and basic ions in solution can lead to a buffer solution, which play a crucial role in stabilizing pH by resisting changes when acids or bases are added. One of the most pressing environmental issues involving acid-base chemistry is ocean acidification . The increasing concentration of atmospheric CO 2 dissolving in seawater forms carbonic acid H 2 CO 3 : CO 2 (g) + H 2 O (l) ⇌ H 2 CO 3 (aq) ⇌ H + (aq) + H 2 O (l) + HCO 3 − (aq) ⇌ H 3 O + (aq) + CO 3 2 − (aq) This lowers ocean pH and reduces the availability of carbonate ions CO 3 2 − , essential for marine organisms like corals and shellfish to build calcium carbonate CaCO 3 skeletons. Coral reefs, which support diverse ecosystems, are particularly vulnerable as reduced pH weakens their structural integrity and hinders growth. Seawater’s natural buffering capacity, primarily from bicarbonate and carbonate ions, moderates pH changes. However, the scale of modern CO 2 emissions overwhelms these systems, threatening marine life and biodiversity. Figure 4.1: Healthy coral reefs rely on a symbiotic relationship with photosynthetic algae called zooxanthellae , which provide corals with energy and their vibrant colors. Under stress from elevated sea temperatures, acidification, or pollution, corals expel these algae, resulting in a stark white appearance known as bleaching. Prolonged bleaching weakens corals, leaving them vulnerable to disease and reducing their ability to support marine ecosystems. (CC BY 3.0; 2015, James Gilmore via Australian Institute of Marine Sciences) Ocean acidification illustrates the delicate balance of natural systems and the far-reaching impact of chemical changes on global ecosystems. Understanding acid-base chemistry provides the knowledge needed to mitigate environmental concerns like ocean acidification. For example, current work is being done to create artificial alkalinity enhancements that aim to restore buffering capacity of seawater in at-risk ecosystems. 4.1: Acids and Bases In chemistry, acids and bases have been defined differently by three sets of theories: One is the Arrhenius definition defined above, which revolves around the idea that acids are substances that ionize (break off) in an aqueous solution to produce hydrogen (H+) ions while bases produce hydroxide (OH-) ions in solution. The other two definitions are discussed in detail alter in the chapter and include the Brønsted-Lowry definition and the Lewis theory. 4.1.1: Definitions 4.1.1.1: Lewis Acids and Bases 4.2: Dissociation Equilibrium and Acid Strength Acid–base reactions always contain two conjugate acid–base pairs. Each acid and each base has an associated ionization constant that corresponds to its acid or base strength. Two species that differ by only a proton constitute a conjugate acid–base pair. The magnitude of the equilibrium constant for an ionization reaction can be used to determine the relative strengths of acids and bases. 4.2.1: Polyprotic Acids 4.2.2: Molecular Shape and Dissociation 4.3: The Autoionization of Water and pH Water is amphiprotic: it can act as an acid by donating a proton to a base to form the hydroxide ion, or as a base by accepting a proton from an acid to form the hydronium ion ( H3O+ ). The autoionization of liquid water produces OH− and H3O+ ions. The equilibrium constant for this reaction is called the ion-product constant of liquid water (Kw) and is defined as Kw=[H3O+][OH−] . At 25°C, Kw is 1.01×10−14 ; hence pH+pOH=pKw=14.00 . 4.3.1: Finding pH from Ka or Kb 4.4: Buffer Solutions Buffers are solutions that resist a change in pH after adding an acid or a base. Buffers contain a weak acid ( HA ) and its conjugate weak base (A−). Adding a strong electrolyte that contains one ion in common with a reaction system that is at equilibrium shifts the equilibrium in such a way as to reduce the concentration of the common ion. Buffers are characterized by their pH range and buffer capacity. 4.4.1: The Henderson-Hasselbalch Approximation 4.5: Titration Stoichiometry The shape of a titration curve, a plot of pH versus the amount of acid or base added, provides important information about what is occurring in solution during a titration. The shapes of titration curves for weak acids and bases depend dramatically on the identity of the compound. The equivalence point of an acid–base titration is the point at which exactly enough acid or base has been added to react completely with the other component. 4.5.1: Acid-Base Indicators