THE GEOCHEMISTRY OF NATURAL WATERS

1. 1 THE GEOCHEMISTRY OF NATURAL WATERS CHEMICAL PRINCIPLES CHAPTER 1a - Kehew (2001)

2. 2 LEARNING OBJECTIVES Review basic fundamentals of chemistry. Understand commonly used concentration units and unit conversions. Learn to calculate and understand the significance of water hardness and charge-balance error. Be introduced to some common ways of graphically displaying natural water compositions. Some of the material in this presentation is a review of very basic chemical fundamentals that most of you should have learned in Introductory College Chemistry courses. This includes the nature and structure of atoms and ions, ionic vs. covalent bonding, how to determine the oxidation states of elements in compounds, the structure of water, the nature and consequences of hydrogen bonding, and the hydration of aqueous ions. If any of these subjects do not sound familiar, please take the time to carefully review the slides and notes on these subjects in this Powerpoint presentation. You will be expected to have some basic understanding of these fundamental subjects when we discuss more advanced topics later in the course. If the review in this presentation is not sufficient for you, you should consult Chapters 5, 6, and 7 in the book Faure G. (1998) Principles and Applications of Geochemistry, 2nd ed., Prentice Hall, Upper Saddle River, New Jersey. Most of the material on concentration units and interconversions should also be review for most of you, although some points may be new. Finally, the remaining material in this presentation, e.g., the discussion of water hardness, charge-balance error, Stiff plots and Piper diagrams, may be new to most of you, but are relatively simple concepts that we introduce now as preliminary warm-ups for things to come. Some of the material in this presentation is a review of very basic chemical fundamentals that most of you should have learned in Introductory College Chemistry courses. This includes the nature and structure of atoms and ions, ionic vs. covalent bonding, how to determine the oxidation states of elements in compounds, the structure of water, the nature and consequences of hydrogen bonding, and the hydration of aqueous ions. If any of these subjects do not sound familiar, please take the time to carefully review the slides and notes on these subjects in this Powerpoint presentation. You will be expected to have some basic understanding of these fundamental subjects when we discuss more advanced topics later in the course. If the review in this presentation is not sufficient for you, you should consult Chapters 5, 6, and 7 in the book Faure G. (1998) Principles and Applications of Geochemistry, 2nd ed., Prentice Hall, Upper Saddle River, New Jersey. Most of the material on concentration units and interconversions should also be review for most of you, although some points may be new. Finally, the remaining material in this presentation, e.g., the discussion of water hardness, charge-balance error, Stiff plots and Piper diagrams, may be new to most of you, but are relatively simple concepts that we introduce now as preliminary warm-ups for things to come.

3. 3 ATOMIC STRUCTURE Each atom of an element is composed of a nucleus surrounded by electrons in various orbitals. Nucleus - a central concentration of mass consisting of protons and neutrons. Electrons - negatively charged particles of relatively low mass. Protons - positively charged particles of relatively high mass. Neutrons - particles with no charge but mass similar to that of protons. It is important to have at least a rough conception of the structure of atoms of the various elements because this structure dictates the chemical properties of elements, the way in which they bond to form compounds, and hence their chemical behavior in natural waters. The terms defined in this slide are illustrated pictorially in the next slide. An atom is composed of three types of particles: 1) relatively heavy, positively charged particles called a protons; 2) relatively heavy, neutral particles called the neutrons; and 3) relatively light, negatively charged particles called electrons. The protons and neutrons are bound tightly together into a central entity called the nucleus, whereas the electrons orbit about the nucleus. The hydrogen atom, atomic number (Z) = 1, consists of 1 proton and 1 electron. For each unit increase in the atomic number of the elements as one proceeds through the periodic table, one proton and one electron are added, together with a variable quantity of neutrons. Atoms with the same number of protons and electrons, but with different numbers of neutrons, are known as isotopes of the same element, e.g., 12C vs. 13C. It is important to have at least a rough conception of the structure of atoms of the various elements because this structure dictates the chemical properties of elements, the way in which they bond to form compounds, and hence their chemical behavior in natural waters. The terms defined in this slide are illustrated pictorially in the next slide. An atom is composed of three types of particles: 1) relatively heavy, positively charged particles called a protons; 2) relatively heavy, neutral particles called the neutrons; and 3) relatively light, negatively charged particles called electrons. The protons and neutrons are bound tightly together into a central entity called the nucleus, whereas the electrons orbit about the nucleus. The hydrogen atom, atomic number (Z) = 1, consists of 1 proton and 1 electron. For each unit increase in the atomic number of the elements as one proceeds through the periodic table, one proton and one electron are added, together with a variable quantity of neutrons. Atoms with the same number of protons and electrons, but with different numbers of neutrons, are known as isotopes of the same element, e.g., 12C vs. 13C.

4. 4 PRIMITIVE MODEL OF THE ATOM We know that the primitive picture of the atom shown in this slide is not quite correct. For a rigorous description of the structure of atoms, we must turn to quantum mechanics. In the quantum mechanical view, electrons in an atom occupy orbitals which represent portions of space where we are most likely to find a given electron. However, there is a finite probability that a given electron can be anywhere. The model shown here is a more classical model, in which the atom is depicted as a kind of mini solar system, with electrons in definite orbits about the nucleus. For our purposes, this understanding of the atom is sufficient. The most important electrons in an atom, with respect to the chemical properties of the atom, are the outermost electrons or valence electrons. It is the outermost orbitals where electrons can be added or removed to form ions. We know that the primitive picture of the atom shown in this slide is not quite correct. For a rigorous description of the structure of atoms, we must turn to quantum mechanics. In the quantum mechanical view, electrons in an atom occupy orbitals which represent portions of space where we are most likely to find a given electron. However, there is a finite probability that a given electron can be anywhere. The model shown here is a more classical model, in which the atom is depicted as a kind of mini solar system, with electrons in definite orbits about the nucleus. For our purposes, this understanding of the atom is sufficient. The most important electrons in an atom, with respect to the chemical properties of the atom, are the outermost electrons or valence electrons. It is the outermost orbitals where electrons can be added or removed to form ions.

5. 5 IONS When electrons are removed from or added to the outer orbitals of atoms, charged particles called ions are formed. If electrons are removed, a positively charged particle called a cation is formed. If electrons are added, a negatively charged particle called a anion is formed. The number of electrons that can be removed or added is dependent on the electronic structure of the neutral atom. For more details on the electronic structure of atoms, see Faure (1998). The number of electrons that can be removed or added is dependent on the electronic structure of the neutral atom. For more details on the electronic structure of atoms, see Faure (1998).

6. 6 METALS VS. NON-METALS Ionization potential (IP) - The energy required to remove an electron from a neutral atom to an infinite distance from the nucleus. Electron affinity (EA) - The tendency of a neutral atom to attract electrons. Metal - An element with a low IP and a low EA; it tends to lose electrons and form a cation. Nonmetal - An element with a high IP and high EA; it tends to gain electrons and form an anion.

7. 7 IONIC BONDS Ionic bonds - bonds formed via the electrostatic attraction between oppositely charged ions. Coulomb’s law ?0 = permittivity of free space = 8.84x10-9 farads m-1 1 farad = 1 coulomb volt-1 F = force of attraction; Z+ and Z- are the charges of the cation and anion respectively; r is the cation-anion distance. If, in a compound, one element completely loses its outermost electrons to form a cation, and another element gains electrons to form an anion, then these two elements may be held together by the strong mutual Coulombic attraction that results. This is known as an ionic or electrostatic bond. According to Coulomb’s Law, the force of attraction increases with: a) increasing charge of the ions; and b) decreasing size of the ions, and therefore, decreasing distance r between them. Thus, the strongest ionic bonds would be formed between elements with the greatest charge-to-radius ratios (Z/r). If, in a compound, one element completely loses its outermost electrons to form a cation, and another element gains electrons to form an anion, then these two elements may be held together by the strong mutual Coulombic attraction that results. This is known as an ionic or electrostatic bond. According to Coulomb’s Law, the force of attraction increases with: a) increasing charge of the ions; and b) decreasing size of the ions, and therefore, decreasing distance r between them. Thus, the strongest ionic bonds would be formed between elements with the greatest charge-to-radius ratios (Z/r).

8. 8 COVALENT BONDS Covalent bonds - bonds formed by sharing of electrons. Many gaseous species and organic compounds contain covalent bonds, e.g., N2. A few minerals, such as diamond and graphite, contain covalent bonds. Sharing may not be equal. That is, covalent bonds may have some ionic character, and vice versa. A purely covalent bond is one in which there is equal sharing of the electrons between the two atoms comprising the bond. In order to form a covalent bond, orbitals on the two atoms involved must overlap. Very few bonds, if any, can be considered purely ionic or purely covalent. Most bonds are a mixture of these two endmembers. However, homopolar bonds, i.e., bonds between two atoms of the same element, such as in N2, O2, graphite, metallic gold, etc., can be considered 100% covalent. A purely covalent bond is one in which there is equal sharing of the electrons between the two atoms comprising the bond. In order to form a covalent bond, orbitals on the two atoms involved must overlap. Very few bonds, if any, can be considered purely ionic or purely covalent. Most bonds are a mixture of these two endmembers. However, homopolar bonds, i.e., bonds between two atoms of the same element, such as in N2, O2, graphite, metallic gold, etc., can be considered 100% covalent.

9. 9 MEASURE OF IONIC/COVALENT CHARACTER Electronegativity (?) - measures the ability of an atom in a compound to attract electrons to itself. Metals have low ? values and non-metals have high ? values (see Table 1-2 in Kehew for electronegativity values of the elements). We use differences in electronegativity to determine degree of ionic character of a bond. Example: CsF ?Cs = 0.7; ?F = 4.0 ?F - ?Cs = 3.3 > 92% ionic; < 8% covalent Electronegativity is a fundamental property of elements that measures their tendency to attract electrons. The percentage of ionic/covalent character of any bond is dependent the difference in the electronegativities of the two atoms involved. If the two atoms have nearly identical electronegativities, then these two atoms will share their valence electrons and form a covalent bond. If one of the atoms has a much higher electronegativity than the other, then the atom with the relatively low electronegativity will give up its electrons to form a cation and the atom with the relatively high electronegativity will accept the electrons given up by the cation to form an anion. We can actually determine the percentage ionic/covalent character of a bond by calculating the value of the difference in electronegativities of the high- and low-electronegativity atoms and referring to the table in the next slide. In the case shown in this slide, we see that CsF is almost, but not quite, a nearly purely ionic bond. Electronegativity is a fundamental property of elements that measures their tendency to attract electrons. The percentage of ionic/covalent character of any bond is dependent the difference in the electronegativities of the two atoms involved. If the two atoms have nearly identical electronegativities, then these two atoms will share their valence electrons and form a covalent bond. If one of the atoms has a much higher electronegativity than the other, then the atom with the relatively low electronegativity will give up its electrons to form a cation and the atom with the relatively high electronegativity will accept the electrons given up by the cation to form an anion. We can actually determine the percentage ionic/covalent character of a bond by calculating the value of the difference in electronegativities of the high- and low-electronegativity atoms and referring to the table in the next slide. In the case shown in this slide, we see that CsF is almost, but not quite, a nearly purely ionic bond.

10. 10 PERCENT IONIC CHARACTER We can use this table to convert electronegativity differences to percent ionic character. Strictly speaking, the values provided here apply to the case where the anion is oxygen (O2-). However, the Table will yield a rough indication of the degree of covalency of a bond between any two elements. We can use this table to convert electronegativity differences to percent ionic character. Strictly speaking, the values provided here apply to the case where the anion is oxygen (O2-). However, the Table will yield a rough indication of the degree of covalency of a bond between any two elements.

11. 11 OTHER EXAMPLES ZnS ?Zn = 1.6; ?S = 2.5 ?S - ?Zn = 2.5 - 1.6 = 0.9 19% ionic; 81% covalent H2 ?H = 2.1 ?H - ?H = 0 0% ionic; 100% covalent CCl4 ?C = 2.5; ?Cl = 3.0 ?Cl - ?C = 3.0 - 2.5 = 0.5 6% ionic; 94% covalent The bonds shown in this slide all have relatively high degrees of covalency. Of course, the homopolar H-H bond in hydrogen gas is as perfect a covalent bond as is possible, whereas the C-Cl and Zn-S bonds are somewhat less covalent. The bonds shown in this slide all have relatively high degrees of covalency. Of course, the homopolar H-H bond in hydrogen gas is as perfect a covalent bond as is possible, whereas the C-Cl and Zn-S bonds are somewhat less covalent.

12. 12 WHY DO WE CARE IF A BOND IS IONIC OR COVALENT? All physical and chemical properties of a compound depend on the character of the bonds. Example: Solubility in water A general rule is that, like dissolves like. As we will see, water is a polar covalent solvent. Ionically bonded compounds have high aqueous solubilities e.g., NaCl ?Cl - ?Na = 3.0 - 0.9 = 2.1 67% ionic, relatively soluble in water C(diamond) ?C - ?C = 0 0% ionic, nearly insoluble in water The nature of chemical bonding is very useful in determining the physical and chemical properties of minerals and dissolved constituents. Properties of minerals such as cleavage, index of refraction of light, hardness and others depend on the type of bonding. In general, dominantly ionically bonded solids tend to have bonds that are equally strong in every direction. This results in a high degree of isotropy in cleavage, hardness and index of refraction. In other words, these properties are the same in all directions in the mineral. Also, ionically bonded solids tend to have relatively high boiling and melting points. On the other hand, dominantly covalently bonded solids tend to have very directional bonds. This means that, bonding is strong in certain directions, particularly bonds between atoms in a molecule. However, bonds can be quite weak in other directions, such as bonds among individual molecules. Thus, covalently bonded solids tend to be very anisotropic in their physical properties, and also tend to have low melting and boiling points. For our purposes, solubility is the most important property affected by the nature of chemical bonding. Water, a polar substance, dissolves ionic solids (e.g., NaCl, K2SO4) very well. On the other hand, polar water is a very poor solvent for many organic compounds, e.g., toxic benzene (C6H6). These organic molecules are best dissolved by organic solvents that are also bonded covalently, e.g., chloroform (CH3Cl) or acetone ((CH3)2CO). The nature of chemical bonding is very useful in determining the physical and chemical properties of minerals and dissolved constituents. Properties of minerals such as cleavage, index of refraction of light, hardness and others depend on the type of bonding. In general, dominantly ionically bonded solids tend to have bonds that are equally strong in every direction. This results in a high degree of isotropy in cleavage, hardness and index of refraction. In other words, these properties are the same in all directions in the mineral. Also, ionically bonded solids tend to have relatively high boiling and melting points. On the other hand, dominantly covalently bonded solids tend to have very directional bonds. This means that, bonding is strong in certain directions, particularly bonds between atoms in a molecule. However, bonds can be quite weak in other directions, such as bonds among individual molecules. Thus, covalently bonded solids tend to be very anisotropic in their physical properties, and also tend to have low melting and boiling points. For our purposes, solubility is the most important property affected by the nature of chemical bonding. Water, a polar substance, dissolves ionic solids (e.g., NaCl, K2SO4) very well. On the other hand, polar water is a very poor solvent for many organic compounds, e.g., toxic benzene (C6H6). These organic molecules are best dissolved by organic solvents that are also bonded covalently, e.g., chloroform (CH3Cl) or acetone ((CH3)2CO).

13. 13 RULES TO DETERMINE OXIDATION STATES 1) The oxidation state of all pure elements is zero. 2) The oxidation state of H is +1, except in hydrides (e.g., LiH, PdH2), where it is -1. 3) The oxidation state of O is -2, except in peroxides (e.g., H2O2), where it is -1. 4) The algebraic sum of oxidation state must equal zero for a neutral molecule or the charge on a complex ion. The oxidation state may be thought of as the charge an atom would acquire if it were dissolved in water as a simple ion. It is important to be able to determine the oxidation state of atoms in compounds for a number of reasons. First and foremost, we must make this determination when dealing with oxidation-reduction reactions. Another reason is that the chemical properties of elements vary according to their oxidation state. Many elements can exist in variety of oxidation states. However, the oxidation states of oxygen and hydrogen exhibit relatively little variation. As pure elements, O2 and H2, their oxidation state is 0. In almost all other compounds, O exists as -2 and H as +1. The main exceptions for oxygen are the relatively rare peroxides, such as hydrogen peroxide (H2O2) and sodium peroxide (Na2O2). In the case of hydrogen, the exception is the class of compounds known as hydrides, in which the oxidation state of H is -1. Such compounds generally do not occur in nature. Knowing the possible oxidation states of O and H, and using rule number 4, we can determine the oxidation states of many other elements. Example 1: H2S(g). We know that each H atom has an oxidation state of +1 and there are 2 of them. We also know that the overall H2O molecule is neutral. Thus, if we let x = the oxidation state of S, we can write: 2?(+1) + x = 0. Solving this, we get x = -2 as the oxidation state of S in H2S(g). Example 2: CrO42-. We know each O atom has an oxidation state of -2. The overall charge on CrO42- is -2. If we let y = the oxidation state of Cr, we can write: 4?(-2) + y = -2, or solving we get y = +6 as the oxidation state of Cr in CrO42-. The oxidation state may be thought of as the charge an atom would acquire if it were dissolved in water as a simple ion. It is important to be able to determine the oxidation state of atoms in compounds for a number of reasons. First and foremost, we must make this determination when dealing with oxidation-reduction reactions. Another reason is that the chemical properties of elements vary according to their oxidation state. Many elements can exist in variety of oxidation states. However, the oxidation states of oxygen and hydrogen exhibit relatively little variation. As pure elements, O2 and H2, their oxidation state is 0. In almost all other compounds, O exists as -2 and H as +1. The main exceptions for oxygen are the relatively rare peroxides, such as hydrogen peroxide (H2O2) and sodium peroxide (Na2O2). In the case of hydrogen, the exception is the class of compounds known as hydrides, in which the oxidation state of H is -1. Such compounds generally do not occur in nature. Knowing the possible oxidation states of O and H, and using rule number 4, we can determine the oxidation states of many other elements. Example 1: H2S(g). We know that each H atom has an oxidation state of +1 and there are 2 of them. We also know that the overall H2O molecule is neutral. Thus, if we let x = the oxidation state of S, we can write: 2?(+1) + x = 0. Solving this, we get x = -2 as the oxidation state of S in H2S(g). Example 2: CrO42-. We know each O atom has an oxidation state of -2. The overall charge on CrO42- is -2. If we let y = the oxidation state of Cr, we can write: 4?(-2) + y = -2, or solving we get y = +6 as the oxidation state of Cr in CrO42-.

14. 14 ELEMENTS WITH VARIABLE OXIDATION STATES Sulfur: SO42-(+6), SO32-(+4), S(0), FeS2(-1), H2S(-2) Carbon: CO2(+4), C(0), CH4(-4) Nitrogen: NO3-(+5), NO2-(+3), NO(+2), N2O(+1), N2(0), NH3(-3) Iron: Fe2O3(+3), FeO(+2), Fe(0) Manganese: MnO4-(+7), MnO2(+4), Mn2O3(+3), MnO(+2), Mn(0) Copper: CuO(+2), Cu2O(+1), Cu(0) Tin: SnO2(+4), Sn2+(+2), Sn(0) Uranium: UO22+(+6), UO2(+4), U(0) Arsenic: H3AsO40(+5), H3AsO30(+3), As(0), AsH3(-1) Chromium: CrO42-(+6), Cr2O3(+3), Cr(0) Gold: AuCl4-(+3), Au(CN)2-(+1), Au(0) Most elements have many possible oxidation states. The oxidation states of elements are very important in determining their environmental behavior. Several examples are given below: Example 1: Cr(+6) vs. Cr(+3). Cr(+6) is highly soluble as HCrO4- or CrO42- and hence is free to move (mobile) in the environment. It is also highly toxic and carcinogenic. In contrast, Cr(+3) is far less toxic, and it is also far less soluble. It tends to be immobilized as solids such as Cr(OH)3. Clearly, Cr(+6) is much more hazardous in the environment than Cr(+3). Example 2: U(+6) vs. U(+4). Uranium is most soluble and mobile as U(+6) species such as UO22+. On the other hand, U(+4) is quite insoluble and will be immobilized as UO2 (uraninite). Example 3: Fe(+2) vs. Fe(+3). Iron tends to be most soluble as Fe(+2), i.e., the aqueous Fe2+ ion, whereas Fe(+3) usually will be immobilized as Fe(OH)3, FeOOH, or Fe2O3. Most elements have many possible oxidation states. The oxidation states of elements are very important in determining their environmental behavior. Several examples are given below: Example 1: Cr(+6) vs. Cr(+3). Cr(+6) is highly soluble as HCrO4- or CrO42- and hence is free to move (mobile) in the environment. It is also highly toxic and carcinogenic. In contrast, Cr(+3) is far less toxic, and it is also far less soluble. It tends to be immobilized as solids such as Cr(OH)3. Clearly, Cr(+6) is much more hazardous in the environment than Cr(+3). Example 2: U(+6) vs. U(+4). Uranium is most soluble and mobile as U(+6) species such as UO22+. On the other hand, U(+4) is quite insoluble and will be immobilized as UO2 (uraninite). Example 3: Fe(+2) vs. Fe(+3). Iron tends to be most soluble as Fe(+2), i.e., the aqueous Fe2+ ion, whereas Fe(+3) usually will be immobilized as Fe(OH)3, FeOOH, or Fe2O3.

15. 15 STRUCTURE OF WATER Ionic character: ?O - ?H = 3.5 - 2.1 = 1.4 39% ionic; 61% covalent - water is polar covalent. Water is a polar substance, meaning there is a positive and a negative pole to the molecule. The water molecule has the shape of the letter “V”. Because each O-H bond is polar (i.e., is a dipole), the overall V-shaped water molecule is also polar. This means that there is a positive and a negative end to the molecule. The negative end is the oxygen atom at the vertex of the “V”, and the positive end are the two hydrogen atoms on the “wingtips” of the “V”. This contrasts with a linear molecule such as CO2 (O=C=O) which is not polar. Because it is polar, the positive end of one water molecule can attract the negative end of a second water molecule, leading to hydrogen bonding. The water molecule has the shape of the letter “V”. Because each O-H bond is polar (i.e., is a dipole), the overall V-shaped water molecule is also polar. This means that there is a positive and a negative end to the molecule. The negative end is the oxygen atom at the vertex of the “V”, and the positive end are the two hydrogen atoms on the “wingtips” of the “V”. This contrasts with a linear molecule such as CO2 (O=C=O) which is not polar. Because it is polar, the positive end of one water molecule can attract the negative end of a second water molecule, leading to hydrogen bonding.

16. 16 HYDROGEN BONDING Because each water molecule has a positive and a negative end, these can attract one another to form a hydrogen bond. The hydrogen bonding of water is responsible for many anomalous properties of water. For example, the boiling points of H2O, CH4 and CCl4 are 100°C, -161.5°C and 76.54°C, respectively. Clearly, water has an anomalously high boiling point because of H-bonding. The melting points of these three substances are 0°C, -182.5°C and -22.99°C, respectively. Once again, water has the highest melting point, reflecting the H-bonding. Water also has a higher heat of fusion, heat of vaporization, heat capacity and dielectric constant than CH4 or H2O. If we compare H2O, H2S, H2Se and H2Te, all of which are V-shaped molecules formed between hydrogen and elements in the same column in the periodic table, we see that water, the molecule with the lowest mass, is the only one which occurs as a liquid at standard temperature and pressure (STP: 25°C and 1 bar). Water is the only compound in which H-bonding is possible. The others are gases at STP. The behavior of water is anomalous because we would expect the least massive molecule to have the lowest boiling point. Because life as we know it requires the presence of liquid water, hydrogen-bonding is what makes water capable of supporting life on Earth! The hydrogen bonding of water is responsible for many anomalous properties of water. For example, the boiling points of H2O, CH4 and CCl4 are 100°C, -161.5°C and 76.54°C, respectively. Clearly, water has an anomalously high boiling point because of H-bonding. The melting points of these three substances are 0°C, -182.5°C and -22.99°C, respectively. Once again, water has the highest melting point, reflecting the H-bonding. Water also has a higher heat of fusion, heat of vaporization, heat capacity and dielectric constant than CH4 or H2O. If we compare H2O, H2S, H2Se and H2Te, all of which are V-shaped molecules formed between hydrogen and elements in the same column in the periodic table, we see that water, the molecule with the lowest mass, is the only one which occurs as a liquid at standard temperature and pressure (STP: 25°C and 1 bar). Water is the only compound in which H-bonding is possible. The others are gases at STP. The behavior of water is anomalous because we would expect the least massive molecule to have the lowest boiling point. Because life as we know it requires the presence of liquid water, hydrogen-bonding is what makes water capable of supporting life on Earth!

17. 17 ION HYDRATION Also because of the polar nature of water, ions will be surrounded by water dipoles (hydrated) in solution. Hydration isolates the ions from their neighbors and neutralizes the attractive forces that hold minerals together. In the drawing shown here, the big, light yellow circles represent oxygen atoms, and the smaller, orange circles represent hydrogen atoms. Cations tend to attract the negative ends (O) of the water molecule. Thus, cations are usually surrounded by water molecules with their oxygen pointing towards the cation. We say that the cation is hydrated. In a similar manner, anions tend to attract the positive ends (H) of water molecules, so anions are also surrounded by water molecules (hydrated). However, in this case, he hydrogens point toward the anion. The hydration of ions accounts in part for the excellent ability of water to dissolve ionically bound compounds. There is no such thing as a “bare” ion in an aqueous solution. Each ion, whether cation or anion, carries along some water baggage on its travels! In the drawing shown here, the big, light yellow circles represent oxygen atoms, and the smaller, orange circles represent hydrogen atoms. Cations tend to attract the negative ends (O) of the water molecule. Thus, cations are usually surrounded by water molecules with their oxygen pointing towards the cation. We say that the cation is hydrated. In a similar manner, anions tend to attract the positive ends (H) of water molecules, so anions are also surrounded by water molecules (hydrated). However, in this case, he hydrogens point toward the anion. The hydration of ions accounts in part for the excellent ability of water to dissolve ionically bound compounds. There is no such thing as a “bare” ion in an aqueous solution. Each ion, whether cation or anion, carries along some water baggage on its travels!