Wednesday, January 24, 2018

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Chemistry - pH

Hydrogen ion activity in aqueous (water-solvent) solutions is a very important parameter for a wide variety of industrial processes. A number of reactions important to chemical processing are inhibited or significantly slowed if the hydrogen ion activity of a solution does not fall within a narrow range. Some additives used in water treatment processes (e.g. flocculants) will fail to function efficiently if the hydrogen ion activity in the water is not kept within a certain range. Alcohol and other fermentation processes strongly depend on tight control of hydrogen ion activity, as an incorrect level of ion activity will not only slow production, but may also spoil the product. Hydrogen ions are always measured on a logarithmic scale, and referred to as pH.

Free hydrogen ions (H+) are rare in a liquid solution, and are more often found attached to whole water molecules to form a positive ion called hydronium (H3O+). However, process control professionals usually refer to these positive ions simply as “hydrogen” even though the truth is a bit more complicated.

pH is mathematically defined as the negative common logarithm of hydrogen ion activity in a solution. Hydrogen ion activity is expressed as a molarity (number of moles of active ions per liter of solution), with “pH” being the unit of measurement for the logarithmic result:

pH = log[H+]

For example, an aqueous solution with an active hydrogen concentration of 0.00044 M has a pH value of 3.36 pH.

Water is a covalent compound, and so there is little separation of water molecules in liquid form. Most of the molecules in a sample of pure water remain as whole molecules (H2O) while a very small percentage ionize into positive hydrogen ions (H+) and negative hydroxyl ions (OH). The mathematical product of hydrogen and hydroxyl ion molarity in water is known as the ionization constant (Kw), and its value varies with temperature:

Kw = [H+] × [OH]

At 25 degrees Celsius (room temperature), the value of Kw is very nearly equal to 1.0 × 10-14. Since each one of the water molecules that does ionize in this absolutely pure water sample separates into exactly one hydrogen ion (H+) and one hydroxyl ion (OH), the molarities of hydrogen and hydroxyl ions must be equal to each other. The equality between hydrogen and hydroxyl ions in a pure water sample means that pure water is neutral, and that the molarity of hydrogen ions is equal to the square root of Kw:

Since we know pH is defined as the negative logarithm of hydrogen ion activity, and we can be assured all hydrogen ions present in the solution will be “active” since there are no other positive ions to interfere with them, the pH value for water at 25 degrees Celsius is:

pH of pure water at 25oC = log(1.0 × 107 M) = 7.0 pH

As the temperature of a pure water sample changes, the ionization constant changes as well. Increasing temperature causes more of the water molecules to ionize into H+ and OHions, resulting in a larger Kw value. The following table shows Kw values for pure water at different temperatures:

0 oC   1.139 × 10-15 
5 oC   1.846 × 10-15 
10 oC   2.920 × 10-15 
15 oC   4.505 × 10-15 
20 oC   6.809 × 10-15 
25 oC   1.008 × 10-14 
30 oC   1.469 × 10-14 
35 oC   2.089 × 10-14 
40 oC   2.919 × 10-14 
45 oC   4.018 × 10-14 
50 oC   5.474 × 10-14 
55 oC   7.296 × 10-14 
60 oC   9.614 × 10-14 

This means that while any pure water sample is neutral (an equal number of positive hydrogen ions and negative hydroxyl ions) at any temperature, the pH value of pure water actually changes with temperature, and is only equal to 7.0 pH at one particular (“standard”) temperature: 25 oC. Based on the Kw values shown in the table, pure water will be 6.51 pH at 60 oC and 7.47 pH at freezing.

If we add an electrolyte to a sample of pure water, (at least some of) the molecules of that electrolyte will separate into positive and negative ions. If the positive ion of the electrolyte happens to be a hydrogen ion (H+), we call that electrolyte an acid. If the negative ion of the electrolyte happens to be a hydroxyl ion (OH), we call that electrolyte a caustic, or alkaline, or base. Some common acidic and alkaline substances are listed here, showing their respective positive and negative ions in solution:


Sulfuric acid is an acid (produces H+ in solution)

H2SO4 2H+ + SO42

Nitric acid is an acid (produces H+ in solution)

HNO3 H+ + NO3

Hydrocyanic acid is an acid (produces H+ in solution)


Hydrofluoric acid is an acid (produces H+ in solution)

HF H+ + F

Lithium hydroxide is a caustic (produces OHin solution)

LiOH Li+ + OH

Potassium hydroxide is a caustic (produces OHin solution)


Sodium hydroxide is a caustic (produces OHin solution)

NaOH Na+ + OH

Calcium hydroxide is a caustic (produces OHin solution)

Ca(OH)2 Ca2+ + 2OH


When an acid substance is added to water, some of the acid molecules dissociate into positive hydrogen ions (H+) and negative ions (the type of negative ions depending on what type of acid it is). This increases the molarity of hydrogen ions (the number of moles of H+ ions per liter of solution). The addition of hydrogen ions to the solution also decreases the molarity of hydroxyl ions (the number of moles of OHions per liter of solution) because some of the water’s OHions combine with the acid’s H+ ions to form deionized water molecules (H2O).

If an alkaline substance (otherwise known as a caustic, or a base) is added to water, some of the alkaline molecules dissociate into negative hydroxyl ions (OH) and positive ions (the type of positive ions depending on what type of alkaline it is). This increases the molarity of OHions in the solution, as well as decreases the molarity of hydrogen ions (again, because some of the caustic’s OHions combine with the water’s H+ ions to form deionized water molecules, H2O).

The result of this complementary effect (increasing one type of water ion, decreasing the other) keeps the overall ionization constant relatively constant, at least for dilute solutions. In other words, the addition of an acid or a caustic may change [H+], but it has little effect on Kw.

A simple way to envision this effect is to think of a laboratory balance scale, balancing the number of hydrogen ions in a solution against the number of hydroxyl ions in the same solution:

When the solution is pure water, this imaginary scale is balanced (neutral), with [H+] = [OH]. Adding an acid to the solution tips the scale one way, while adding a caustic to the solution tips it the other way1.

If an electrolyte has no effect on the hydrogen and hydroxyl ion activity of an aqueous solution, we call it a salt. The following is a list of some common salts, showing their respective ions in solution:


Potassium chloride is a salt (produces neither H+ nor OHnor O2in solution)

KCl K+ + Cl

Sodium chloride is a salt (produces neither H+ nor OHnor O2in solution)

NaCl Na+ + Cl

Zinc sulfate is a salt (produces neither H+ nor OHnor O2in solution)

ZnSO4 Zn+ + SO4


The addition of a salt to an aqueous solution should have no effect on pH, because the ions created neither add to nor take away from the hydrogen ion activity2.

Acids and caustics tend to neutralize one another, the hydrogen ions liberated by the acid combining (and canceling) with the hydroxyl ions liberated by the caustic. This process is called pH neutralization, and it is used extensively to adjust the pH value of solutions. If a solution is too acidic, just add caustic to raise its pH value. If a solution is too alkaline, just add acid to lower its pH value.

The result of a perfectly balanced mix of acid and caustic is deionized water (H2O) and a salt formed by the combining of the acid’s and caustic’s other ions. For instance, when hydrochloric acid (HCl) and potassium hydroxide (KOH) neutralize one another, the result is water (H2O) and potassium chloride (KCl), a salt. This production of salt is a necessary side-effect of pH neutralization, which may require addressing in later stages of solution processing. Such neutralizations are exothermic, owing to the decreased energy states of the hydrogen and hydroxyl ions after combination. Mixing of pure acids and caustics together without the presence of substantial quantities of water (as a solvent) is often violently exothermic, presenting a significant safety hazard to anyone near the reaction.


1It should be noted that the solution never becomes electrically imbalanced with the addition of an acid or caustic. It is merely the balance of hydrogen to hydroxyl ions we are referring to here. The net electrical charge for the solution should still be zero after the addition of an acid or caustic, because while the balance of hydrogen to hydroxyl ions does change, that electrical charge imbalance is made up by the other ions resulting from the addition of the electrolyte (anions for acids, cations for caustics). The end result is still one negative ion for every positive ion (equal and opposite charge numbers) in the solution no matter what substance(s) we dissolve into it.

2Exceptions do exist for strong concentrations, where hydrogen ions may be present in solution yet unable to react because of being “crowded out” by other ions in the solution.

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