A pH scale is a measure of how acidic or basic a substance is. While the pH scale formally measures the activity of hydrogen ions in a substance or solution, it is typically approximated as the concentration of hydrogen ions. However, this approximation is applicable only under low concentrations.
The pH scale is logarithmic, meaning that an increase or decrease of an integer value changes the concentration by a tenfold. For example, a pH of 3 is ten times more acidic than a pH of 4. Likewise, a pH of 3 is one hundred times more acidic than a pH of 5. Similarly a pH of 11 is ten times more basic than a pH of 10. Because of the amphoteric nature of water (i.e., acting as both an acid or a base), water does not always remain as \(H_2O\). In fact, two water molecules react to form hydronium and hydroxide ions:
\[ 2\, H_2O \;(l) \rightleftharpoons H_3O^+ \;(aq) + OH^− \; (aq) \]
This is also called the self-ionization of water. The concentration of H_{3}O^{+} and OH^{-} are equal in pure water because of the stoichiometric ratio. The molarity of H_{3}O^{+} and OH^{-} in water are also both \(1.0 \times 10^{-7} \,M\) at 25° C. Therefore, a constant of water (\(K_w\)) is created to show the equilibrium condition for the self-ionization of water. The product of the molarity of hydronium and hydroxide ion is always \(1.0 \times 10^{-14}\).
\[K_w= [H_3O^+][OH^-] = 1.0 \times 10^{-14}\]
This equations also applies to all aqueous solutions. However, \(K_w\) does change at different temperatures, which affects the pH range discussed below. Note: H^{+} and H_{3}O^{+} is often used interchangeably. The equation for water equilibrium is:
\[ H_2O \rightleftharpoons H^+ + OH^- \]
Because the constant of water, K_{w} is always 1.0 X 10^{-14}, the pK_{w} is 14, the constant of water determines the range of the pH scale. To understand what the pK_{w} is, it is important to understand first what the "p" means in pOH, and pH. The danish biochemist Soren Sorenson proposed the term pH to refer to the "potential of hydrogen ion." He defined the "p" as the negative of the logarithm, -log, of [H^{+}]. Therefore the pH is the negative logarithm of the molarity of H. The pOH is the negative logarithm of the molarity of OH^{-} and the pK_{w}_{ }is the negative logarithm of the constant of water. These definitions give the following equations:
pH= -log [H^{+}]
pOH= -log [OH^{-}]
pK_{w}= -log [K_{w}]
A Logarithm, used in the above equations, of a number is how much a power is raised to a particular base in order to produce that number. To simplify this, look at the equation: log_{b}a=x. This correlates to b^{x}=a. A simple example of this would be log_{10}100=2, or 10^{2}=100. It is assumed that the base of Logarithms is ten if it is not stated. So for the sake of pH and pOH problems it will always be ten. When x is a negative number that means you are dividing it by the power. So, if log_{10}0.01=-2 which can be written 10^{-2}=0.01. 10^{-2} also means 1/10^{2}. The log function can be found on your scientific calculator. Now if we apply this to pH and pOH we can better understand how we calculate the values.
The constant of water is always 1.0 X 10^{-}^{14}. So pK_{w}=-log [1.0 X 10^{-14}]. Using what we know about Logarithms, we can write this as 10^{-pK}^{w}=10^{-14}. By substituting we see that pK_{w} is 14. The equation also shows that each increasing unit on the scale decreases by the factor of ten on the molarity. For example, a pH of 1 has a molarity ten times more concentrated than a solution of pH 2. Also, the pK_{w} of water is 14 and the addition of pH and pOH is always 14 at 25° Celsius.
\[pK_w= pH + pOH = 14\]
The pH scale is often referred to as ranging from 0-14 or perhaps 1-14. Neither is correct. The pH range does not have an upper nor lower bound, since as defined above, the pH is an indication of concentration of H^{+}. For example, at a pH of zero the hydronium ion concentration is one molar, while at pH 14 the hydroxide ion concentration is one molar. Typically the concentrations of H^{+} in water in most solutions fall between a range of 1 M (pH=0) and 10^{-14} M (pH=14). Hence a range of 0 to 14 provides sensible (but not absolute) "bookends" for the scale. However, in principle, one can go somewhat below zero and somewhat above 14 in water, because the concentrations of hydronium ions or hydroxide ions can exceed one molar. Figure 1 depicts the pH scale with common solutions and where they are on the scale.
Figure 1: Solutions and the placement of them on pH scale
In 1909 S.P.L. Sorensen published a paper in Biochem Z in which he discussed the effect of H^{+} ions on the activity of enzymes. In the paper he invented the term pH to describe this effect and defined it as the -log[H^{+} ]. In 1924 Sorensen realized that the pH of a solution is a function of the "activity" of the H^{+} ion not the concentration and published a second paper on the subject. A better definition would be pH=-log[aH^{+} ], where aH^{+} denotes the activity of the H^{+} ion. The activity of an ion is a function of many variables of which concentration is one. It is unfortunate that chemistry texts use a definition for pH that has been obsolete for over 50 years.
Because of the difficulty in accurately measuring the activity of the H^{+} ion for most solutions the International Union of Pure and Applied Chemistry (IUPAC) and the National Bureau of Standards (NBS) has defined pH as the reading on a pH meter that has been standardized against standard buffers. The following equation is used to calculate the pH of all solutions:
The activity of the H^{+} ion is determined as accurately as possible for the standard solutions used. The identity of these solutions vary from one authority to another, but all give the same values of pH to ± 0.005 pH unit. The historical definition of pH is correct for those solutions that are so dilute and so pure the H^{+} ions are not influenced by anything but the solvent molecules (usually water). In most solutions the pH differs from the -log[H^{+} ] in the first decimal point.
Above, the pH was approximated as the measure of H^{+} concentration:
pH= -log [H^{+}]
Note: concentration is abbreviated by using square brackets, thus [H^{+}] = hydrogen ion concentration. When measuring pH, [H^{+}] is in units of moles of H^{+} per liter of solution. This is a reasonably accurate definition at low concentrations (the dilute limit) of H^{+}. At very high concentrations (10 M hydrochloric acid or sodium hydroxide, for example,) a significant fraction of the ions will be associated into neutral pairs such as H^{+}Cl^{–}, thus reducing the concentration of “available” ions to a smaller value which we will call the effective concentration. It is the effective concentration of H^{+} and OH^{–} that determines the pH and pOH. The pH scale as shown above is called sometimes "concentration pH scale" as opposed to the "thermodynamic pH scale". The main difference between both scales is that in thermodynamic pH scale one is interested not in H^{+}concentration, but in H^{+}activity. What a person measures in the solution is just activity, not the concentration. Thus it is thermodynamic pH scale that describes real solutions, not the concentration one.
For solutions in which ion concentrations don't exceed 0.1 M, the formulas pH = –log [H^{+}] and pOH = –log[OH^{–}] are generally reliable, but don't expect a 10.0 M solution of a strong acid to have a pH of exactly –1.00! However, this definition is only an approximation (albeit very good under most situations) of the proper definition of pH, which depends on the activity of the hydrogen ion:
pH= -log a{H+}
The activity is a measure of the "effective concentration" of a substance, is often related to the true concentration via an activity coefficient, γ:
a{H^{+}}=γ[H^{+}]
Calculating the activity coefficient requires detailed theories of how charged species interact in solution at high concentrations (e.g., the Debye-Hückel Theory). The following table gives experimentally determined pH values for a series of HCl solutions of increasing concentration at 25 °C.
Table 3: HCl Solutions with Corresponding pH values
Molar Concentration of HCl
| pH defined as Concentration | Experimentally Determined pH |
0.00050 | 3.30 | 3.31 |
0.0100 | 2 | 2.04 |
0.100 | 1 | 1.10 |
0.40 | 0.39 | 0.52 |
7.6 | -0.88 | -1.85 |
Christopher G. McCarty and Ed Vitz, Journal of Chemical Education, 83(5), 752 (2006) and G.N. Lewis, M. Randall, K. Pitzer, D.F. Brewer, Thermodynamics (McGraw-Hill: New York, 1961; pp. 233-34).
Example 1 |
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If the concentration of NaOH in a solution is 2.5 X 10^{-4} M, what is the concentration of H_{3}O^{+}? SOLUTION Because 1.0 X 10^{-14} = [H_{3}O^{+}][OH^{-}], (1.0 X 10^{-14})/[OH^{-}] = [H_{3}O^{+}] |
Example 2 |
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1) Find the pH of a solution of 0.002M of HCl. SOLUTION [H^{+}]= 2.0 X 10^{-3} M 2. The equation for pOH is -log [OH^{-}] [OH^{-}]= 5.0 X 10^{-5} M |
Example 3 |
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If soil has a pH of 7.84, what is the H^{+} concentration of the soil solution? SOLUTION pH = -log [H^{+}] (Hint: place -7.84 in your calculator and take the antilog (often inverse log or 10^{x}) = 1.45 x 10^{-8}M |
Molecules that make up or are produced by living organisms usually function within a narrow pH range (near neutral) and a narrow temperature range (body temperature). Many biological solutions, such as blood, have a pH near neutral. pH influences the structure and the function of many enzymes (protein catalysts) in living systems. Many of these enzymes have narrow ranges of pH activity.
Cellular pH is so important that death may occur within hours if a person becomes acidotic (having increased acidity in the blood). As one can see pH is critical to life, biochemistry, and important chemical reactions. Common examples of how pH plays a very important role in our daily lives are given below:
Table 2: pH in Living Systems
Compartment | pH |
Gastric Acid | 1 |
Lysosomes | 4.5 |
Granules of Chromaffin Cells | 5.5 |
Human Skin | 5.5 |
Urine | 6 |
Neutral H_{2}O at 37 degrees Celsius | 6.81 |
Cytosol | 7.2 |
Cerebrospinal Fluid | 7.3 |
Blood | 7.43-7.45 |
Mitochondrial Matrix | 7.5 |
Pancreas Secretions | 8.1 |
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