AQUEOUS SOLUTIONS
S OLUTION OF E LECTROLYTES
Solutions of acids, bases, and salts contain ions. Charged ions move when driven by an electric po- tential, and electrolyte solutions conduct electricity. These ion-containing substances are called elec-
trolytes. As mentioned earlier, the high dielectric
constant of water reduces the attraction of ions with- in ionic solids and dissolves them. Furthermore, the polar water molecules surround ions, forming hy-
drated ions. The concentration of all ions and
molecular substances in a solution contributes to the osmotic potential.
Water can also be an acid or a base, because H2O
molecules can receive or provide a proton (H).
Such an exchange by water molecules in pure water, forming hydrated protons (H3Oor (H2O)4H), is
called self-ionization. However, the extent of ion- ization is small, and pure water is a very poor con- ductor.
Self-Ionization of Water
The self-ionization of water is a dynamic equilibri- um,
H2O(l) ↔ H(aq) OH(aq),
Kw [H][OH] 10-14at 298 K and 1 atm
where [H] and [OH] represent the molar concen- trations of H(or H
3O) and OHions, respective-
ly, and Kwis called the ion product of water. Values
of Kwunder various conditions have been evaluated
theoretically (Marshall and Franck 1981, Tawa and Pratt 1995). Solutions in which [H] [OH] are
said to be neutral. Both pH and pOH, defined by the following equations, have a value of 7 at 298 K for a neutral solution.
pH log10[H] pOH log10[OH] 7
(at 298 K)
The Hrepresents a hydrated proton (H
3O), which
dynamically exchanges a proton with other water molecules. The self-ionization and equilibrium are present in water and all aqueous solutions.
Solutions of Acids and Bases
Strong acids HClO4, HClO3, HCl, HNO3, and
H2SO4 completely ionize in their solutions to give
H(H
3O) ions and anions ClO4, ClO3, Cl,
NO3, and HSO4, respectively. Strong bases
NaOH, KOH, and Ca(OH)2also completely ionize
to give OHions and Na, K, and Ca2ions, re-
spectively. In an acidic solution, [H] is greater than
[OH]. For example, in a 1.00 mol/L HCl solution
at 298 K, [H] 1.00 mol/L, pH 0.00, [OH]
10-14mol/L.
Weak acids such as formic acid (HCOOH), acetic acid (HCH3COO), ascorbic acid (H2C6H6O6), oxalic
acid (H2C2O4), carbonic acid (H2CO3), benzoic acid
(HC6H5COO), malic acid (H2C4H4O5), lactic acid
(HCH3CH(OH)COO), and phosphoric acid (H3PO4)
also ionize in their aqueous solutions, but not com- pletely. The ionization of acetic acid is represented by the equilibrium
HCH3COO(aq) ↔ H(aq) CH3COO(aq),
where Ka, as defined above, is the acid dissociation constant.
The solubility of CO2in water increases with par-
tial pressure of CO2, according to Henry’s law. The
chemical equilibrium for the dissolution is H2O(l) CO2(g) ↔ H2CO3(aq)
Of course, H2CO3dynamically exchanges Hand
H2O with other water molecules, and this weak diprotic acid ionizes in two stages with acid disso-
ciation constants Ka1and Ka2:
H2O CO2(aq) ↔ H(aq) HCO3(aq), Ka1
4.30 10-7at 298 K
HCO3(aq) ↔ H(aq) CO32-(aq), Ka2
5.61 10-11.
Constants Ka1and Ka2increase as temperature rises,
but the solubility of CO2decreases. At 298 K, the
pH of a solution containing 0.1 mol/L H2CO3is 3.7.
At this pH, acidophilic organisms survive and grow, but most pathogenic organisms are neutrophiles, and they cease growing. Soft drinks contain other acids— citric, malic, phosphoric, ascorbic, and others. They lower the pH further.
All three hydrogen ions in phosphoric acid (H3PO4) are ionizable, and it is a triprotic acid.
Acids having more than one dissociable Hare called polyprotic acids.
Ammonia and many nitrogen-containing com- pounds are weak bases. The ionization equilibrium of NH3in water and the base dissociation constant Kbare
NH3 H2O ↔ NH4O ↔ NH4(aq) OH(aq),
Other weak bases react with H2O and ionize in a
similar way. Kb + 4 H [OH NH OH at 298 K. =[ ] −]= × − [ ] 1 70 10. 5 Ka 3 H CH COO HCH COO at 298 K =[ +][ −]= × − [ ] . 3 1 75 10 5
The ionization or dissociation constants of inor- ganic and organic acids and bases are extensive, and they have been tabulated in various books (for example Perrin 1965, 1982; Kortüm et al. 1961).
Titration
Titration is a procedure for quantitative analysis of
a solute in a solution by measuring the quantity of a reagent used to completely react with it. This meth- od is particularly useful for the determination of acid or base concentrations. A solution with a known concentration of one reagent is added from a burette to a definite amount of the other. The end
point is reached when the latter substance is com-
pletely consumed by the reagent from the burette, and this is detected by the color change of an indica- tor or by pH measurements. This method has many applications in food analysis.
The titration of strong acids or bases utilizes the rapid reaction between Hand OH. The unknown
quantity of an acid or base may be calculated from the amount used to reach the end point of the titra- tion.
The variation of pH during the titration of a weak acid using a strong base or a weak base using a strong acid is usually monitored to determine the end point. The plot of pH against the amount of reagent added is a titration curve. There are a num-
ber of interesting features on a titration curve. Titration of a weak acid HA using a strong base NaOH is based on two rapid equilibria:
H OH H
2O, K 5.56 1015
at 298 K.
As the Hions react with OH, more Hions are
produced due to the equilibrium HA H A, K
a(ionization constant of the acid).
Before any NaOH solution is added, HA is the dom- inant species; when half of HA is consumed, [HA] [A], which is called the half equivalence point.
At this point, the pH varies the least when a little H
or OHis added, and the solution at this point is the
most effective buffer solution, as we shall see later.
The pH at this point is the same as the pKa of the
weak acid. When an equivalent amount of OHhas
been added, the Aspecies dominates, and the solu- tion is equivalent to a salt solution of NaA. Of course, the salt is completely ionized.
Polyprotic acids such as ascorbic acid H2-
(H6C6O6) (Vitamin C; Ka1 7.9 10-5, Ka2 1.6
10-12) and phosphoric acid H
3PO4(Ka1 6.94
10-3, K
a2 6.2 10-8, Ka32.1 10-12) have more
than one mole of Hper mole of acid. A titration
curve of these acids will have two and three end points for ascorbic and phosphoric acids, respective- ly, partly due to the large differences in their dissoci- ation constants (Ka1, Ka2, etc.). In practice, the third
end point is difficult to observe in the titration of H3PO4. Vitamin C and phosphoric acids are often
used as food additives.
Many food components (e.g., amino acids, pro- teins, alkaloids, organic and inorganic stuff, vita- mins, fatty acids, oxidized carbohydrates, and com- pounds giving smell and flavor) are weak acids and bases. The pH affects their forms, stability, and reac- tions. When pH decreases by 1, the concentration of H, [H], increases 10-fold, accompanied by a 10-
fold decrease in [OH]. The Hand OHare very
active reagents for the esterification and hydrolysis reactions of proteins, carbohydrates, and lipids, as we shall see later. Thus, the acidity, or pH, not only affects the taste of food, it is an important parameter in food processing.
Solutions of Amino Acids
Amino acids have an amino group (NH3), a car-
boxyl group (COO), a H, and a side chain (R)
[H O2 ]
w
K
Figure 5.11. Titration curve of a 0.10 mol/L (or M) weak acid HA (Ka= 1 105) using a 0.10 mol/L strong base NaOH solution.
attached to the asymmetric alpha carbon. They are the building blocks of proteins, polymers of amino acids. At a pH called the isoelectric point, which depends on the amino acid in question, the dominant species is a zwitterion, RHC(NH3)(COO), which
has a positive and a negative site, but no net charge. For example, the isoelectric point for glycine is pH 6.00, and its dominant species is H2C(NH3)
COO. An amino acid exists in at least three forms
due to the following ionization or equilibria: RHC(NH3)(COOH) RHC(NH3)(COO)
H, K a1
RHC(NH3)(COO) RHC(NH2)(COO) H, Ka2.
Most amino acids behave like a diprotic acid with two dissociation constants, Ka1 and Ka2. A few
amino acids have a third ionizable group in their side chains.
Among the 20 common amino acids, the side chains of eight are nonpolar, and those of seven are polar, containing –OH, CuO, or –SH groups. Aspartic and glutamic acid contain acidic –COOH groups in their side chains, whereas arginine, histi- dine, and lysine contain basic –NH or –NH2groups.
These have four forms due to adding or losing pro- tons at different pH values of the solution, and they behave as triprotic acids. For example, aspartic acid [Asp (COOH)CH2C(NH3)(COO)] has these
forms:
AspH Asp H
Asp Asp H
Asp Asp2- H
Proteins, amino acid polymers, can accept or pro- vide several protons as the pH changes. At its iso-
electric point (a specific pH), the protein has no net
charge and is least soluble because electrostatic repulsion between its molecules is lowest, and the molecules coalesce or precipitate, forming a solid or gel.
Solutions of Salts
Salts consist of positive and negative ions, and these ions are hydrated in their solutions. Positive, hydrat- ed ions such as Na(H2O)6, Ca(H2O)82, and Al
(H2O)63have six to eight water molecules around
them. Figure 5.12 is a sketch of the interactions of water molecules with ions. The water molecules point the negative ends of their dipoles towards pos- itive ions, and their positive ends towards negative ions. Molecules in the hydration sphere constantly and dynamically exchange with those around them. The number and lifetimes of hydrated water mole- cules have been studied by various methods. These studies reveal that the hydration sphere is one layer deep, and the lifetimes of these hydrated water mol- ecules are in the order of picoseconds (1012 sec-
onds). The larger negative ions also interact with the polar water molecules, but not as strongly as do
Figure 5.12. The first hydration sphere of most cations M(H2O)6, and anions X(H2O)61. Small water molecules are below the plane containing the ions, and large water molecules are above the plane.
cations. The presence of ions in the solution changes the ordering of water molecules even if they are not in the first hydration sphere.
The hydration of ions releases energy, but break- ing up ions from a solid requires energy. The amount of energy needed depends on the substance, and for this reason, some substances are more soluble than others. Natural waters in oceans, streams, rivers, and lakes are in contact with minerals and salts. The concentrations of various ions depend on the solu- bility of salts (Moeller and O’Connor 1972) and the contact time.
All salts dissolved in water are completely ion- ized, even those formed in the reaction between weak acids and weak bases. For example, the com- mon food preservative sodium benzoate (NaC6H5
COO) is a salt formed between a strong base NaOH and weak benzoic acid (Ka 6.5 10-5). The ben-
zoate ions, C6H5COO, in the solution react with
water to produce OHions giving a slightly basic
solution: