THE GEOCHEMISTRY OF NATURAL WATERS

1. 1 THE GEOCHEMISTRY OF NATURAL WATERS ACID-BASE REACTIONS AND THE CARBONATE SYSTEM CHAPTER 3a - Kehew (2001)

2. 2 LEARNING OBJECTIVES Define acids and bases. Learn to express the strength of an acid in terms of the dissociation constant (pKa). Understand acid-base equilibria in the carbonate system. Define and be able to sketch a Bjerrum plot. Many reactions that occur in natural waters involve the interactions of acids and bases. For example, chemical weathering reactions involve rainwater containing the weak acid, carbonic acid (H2CO3), with minerals in rocks acting as bases. Acid-base reactions are relevant to a number of environmental problems, including acid precipitation and acid-mine drainage. In the second part of this lecture, we will define acids and bases and we will learn how to express the strengths of acids in terms of their dissociation constants. We will make a start at understanding acid-base equilibria in the H2O-CO2 system, and we will learn how to depict acid-base equilibria on the so-called Bjerrum plot. Many reactions that occur in natural waters involve the interactions of acids and bases. For example, chemical weathering reactions involve rainwater containing the weak acid, carbonic acid (H2CO3), with minerals in rocks acting as bases. Acid-base reactions are relevant to a number of environmental problems, including acid precipitation and acid-mine drainage. In the second part of this lecture, we will define acids and bases and we will learn how to express the strengths of acids in terms of their dissociation constants. We will make a start at understanding acid-base equilibria in the H2O-CO2 system, and we will learn how to depict acid-base equilibria on the so-called Bjerrum plot.

3. 3 ACID-BASE REACTIONS Acid-base reactions involve the transfer of a proton (H+) between reactants and products. Many significant reactions occurring in natural waters, including reactions involving CO2, are acid-base reactions. Acid: any compound that donates a proton. Base: any compound that accepts a proton. Generalized acid-base reaction: HA + B ? A + HB A is the conjugate base of HA, and HB is the conjugate acid of B. The definition of acids and bases given in this slide is known as the Bronsted definition. This definition covers many of the acids and bases of interest to study of natural waters. However, a broader definition, the Lewis definition, is sometime useful. In the Lewis definition, acids are electron-pair acceptors, and bases are electron-pair donors. Thus for example, BF3, which contains no protons, is an acid in the Lewis sense because it is capable of accepting electrons from a base. Thus, in the Lewis definition, the following is an acid-base reaction: NH3 + BF3 ? NH3BF3 In this case, an electron pair on the N atom in ammonia is donated to the B atom in BF3. When an acid HA reacts with a base B, an exchange of protons occurs to form A and HB. We say that A is the conjugate base of acid HA, and HB is the conjugate acid of base B. The definition of acids and bases given in this slide is known as the Bronsted definition. This definition covers many of the acids and bases of interest to study of natural waters. However, a broader definition, the Lewis definition, is sometime useful. In the Lewis definition, acids are electron-pair acceptors, and bases are electron-pair donors. Thus for example, BF3, which contains no protons, is an acid in the Lewis sense because it is capable of accepting electrons from a base. Thus, in the Lewis definition, the following is an acid-base reaction: NH3 + BF3 ? NH3BF3 In this case, an electron pair on the N atom in ammonia is donated to the B atom in BF3. When an acid HA reacts with a base B, an exchange of protons occurs to form A and HB. We say that A is the conjugate base of acid HA, and HB is the conjugate acid of base B.

4. 4 THE HYDRONIUM ION The proton does not actually exist in aqueous solution as a bare H+ ion. The proton exists as the hydronium ion (H3O+). Consider the acid-base reaction: HCO3- + H2O ? H3O+ + CO32- Here water acts as a base, producing the hydronium ion as its conjugate acid. For simplicity, we often just write this reaction as: HCO3- ? H+ + CO32- Bare protons do not exist in aqueous solutions. The H+ ion, being quite small, has a strong tendency to attract the negative end of polar water molecules to it; in other words, the proton is strongly hydrated (see lecture 1). We often write the hydrated proton as the hydronium ion, H3O+. In fact, the proton is almost certainly hydrated by more than one water molecule, and it might be more appropriate to write the hydrated proton with four water molecules, e.g., H9O4+, or perhaps even a larger number of waters of hydration. However, as long as we keep the hydration of the proton in mind, it is permissible to write reactions in terms of H+ for simplicity. From a thermodynamic point of view, it does not matter whether H+ is hydrated or not, because thermodynamics deals only with macroscopic properties. On the other hand, the hydration of the proton needs to be taken into account when considering reaction mechanisms or kinetics. Bare protons do not exist in aqueous solutions. The H+ ion, being quite small, has a strong tendency to attract the negative end of polar water molecules to it; in other words, the proton is strongly hydrated (see lecture 1). We often write the hydrated proton as the hydronium ion, H3O+. In fact, the proton is almost certainly hydrated by more than one water molecule, and it might be more appropriate to write the hydrated proton with four water molecules, e.g., H9O4+, or perhaps even a larger number of waters of hydration. However, as long as we keep the hydration of the proton in mind, it is permissible to write reactions in terms of H+ for simplicity. From a thermodynamic point of view, it does not matter whether H+ is hydrated or not, because thermodynamics deals only with macroscopic properties. On the other hand, the hydration of the proton needs to be taken into account when considering reaction mechanisms or kinetics.

5. 5 AMPHOTERIC SUBSTANCE Now consider the acid-base reaction: NH3 + H2O ? NH4+ + OH- In this case, water acts as an acid, with OH- its conjugate base. Substances that can act as either acids or bases are called amphoteric. Bicarbonate (HCO3-) is also an amphoteric substance: Acid: HCO3- + H2O ? H3O+ + CO32- Base: HCO3- + H3O+ ? H2O + H2CO30 Some substances can either donate or accept a proton, depending on the pH of the solution. Such substances are termed amphoteric. If pH is low (i.e., the activity of H+ is high), an amphoteric substance will act as a base and accept a proton. However, if pH is high (i.e., H+ ions are scarce), an amphoteric substance will act as an acid and donate a proton. Examples of amphoteric substances include H2O and HCO3- as shown above, as well as HSO4-, H2PO4-, HPO42-, etc. Acid: HSO4- ? SO42- + H+ Base: HSO4- + H+ ? H2SO40 Acid: H2PO4- ? HPO42- + H+ Base: H2PO4- + H+ ? H3PO40 Acid: HPO42- ? PO43- + H+ Base: HPO42- + H+ ? H2PO4- Some substances can either donate or accept a proton, depending on the pH of the solution. Such substances are termed amphoteric. If pH is low (i.e., the activity of H+ is high), an amphoteric substance will act as a base and accept a proton. However, if pH is high (i.e., H+ ions are scarce), an amphoteric substance will act as an acid and donate a proton. Examples of amphoteric substances include H2O and HCO3- as shown above, as well as HSO4-, H2PO4-, HPO42-, etc. Acid: HSO4- ? SO42- + H+ Base: HSO4- + H+ ? H2SO40 Acid: H2PO4- ? HPO42- + H+ Base: H2PO4- + H+ ? H3PO40 Acid: HPO42- ? PO43- + H+ Base: HPO42- + H+ ? H2PO4-

6. 6 STRENGTH OF ACIDS - I The strength of an acid is expressed by the value of the equilibrium constant for its dissociation reaction. Consider: HCO3- ? H+ + CO32- The dissociation constant for this reaction at 25°C is: This can also be expressed as pK a = 6.35. The larger the pK a, the weaker the acid. Bicarbonate is considered to be a relatively weak acid.