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AMD Basics

AMD Chemistry

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Acids, bases, pH

For a more complete treatment on this subject, see the chapter Introduction to acid-base theory (PDF 22 pages) by S.K. Lower.  It covers the fundamental concepts of acids and bases. Except for some stoichiometry and a discussion on pH, this section is largely qualitative.

Understanding AMD water chemistry requires an appreciation of acids, bases, and pH.  When we talk about acids, bases, and pH we are referring to a single property of water and its solutions.  That property is the concentration of hydrogen ion [H+].*  (When a chemical symbol is placed in square brackets, i.e. [H+], it indicates the its concentration in moles per liter of solution.) The pH of water is a measure of its [H+], but it does it in a way that may seem odd.  pH is defined as -log [H+], where log is the base 10 logarithm function.  The following two equations show the relationship of [H+] and pH.

pH= -log [H+]  (1)

[H+] = 10-pH  (2)

In studying water-based systems, we are typically interested in a pH range of 0 to 14 where most water-based chemistry occurs.  By examining the table below, you will notice that every unit increase in pH corresponds to a 10 fold decrease in [H+].  You will also notice a very large dynamic range of [H+]  (An [H+] of 1 has a pH of 0; an [H+] of 0.00000000000001 has a pH of 14).  pH lets us express this broad dynamic [H+] range conveniently.  (The Richter Scale for expressing the intensity of earthquakes works in a similar way.)   Note that pure water, having a pH of 7, is defined as neutral.

pH
 

[H+]
decimal form
 

[H+]
scientific notation

[OH-]
decimal form
 

[OH-]
scientific notation

 

0

1

1 x 100

0.00000000000001

1 x 10-14

very acidic

1

0.1

1 x 10-1

0.0000000000001

1 x 10-13

quite acidic

2

0.01

1 x 10-2

0.000000000001

1 x 10-12

quite acidic

3

0.001

1 x 10-3

0.00000000001

1 x 10-11

moderately acidic

4

0.0001

1 x 10-4

0.0000000001

1 x 10-10

moderately acidic

5

0.00001

1 x 10-5

0.000000001

1 x 10-9

mildly acidic

6

0.000001

1 x 10-6

0.00000001

1 x 10-8

mildly acidic

7

0.0000001

1 x 10-7

0.0000001

1 x 10-7

pure water, neutral

8

0.00000001

1 x 10-8

0.000001

1 x 10-6

mildly basic

9

0.000000001

1 x 10-9

0.00001

1 x 10-5

mildly basic

10

0.0000000001

1 x 10-10

0.0001

1 x 10-4

moderately basic

11

0.00000000001

1 x 10-11

0.001

1 x 10-3

moderately basic

12

0.000000000001

1 x 10-12

0.01

1 x 10-2

quite basic

13

0.0000000000001

1 x 10-13

0.1

1 x 10-1

quite basic

14

0.00000000000001

1 x 10-14

1

1 x 100

very basic

The presence of other chemicals in water may very well change the concentration of H+ and hence its pH.  A solution whose pH is less than 7 is said to be acidic, whereas a solution whose pH is greater than 7 is said to be basic or alkaline. Chemicals that cause the number of H+  ions in water to increase are acids; chemicals that cause the number of OH- ions to increase are bases.  As we shall see, [H+] and [OH-] in water are intimately related: when one increases, the other decreases, and we calculate one if we know the other.

Why is pH Important?

Dissociation of Water

Water molecules have a natural tendency to split apart as symbolized in the expression:

H2O H+  + OH-  (3)

This process, known as dissociation or ionization, involves a water molecule splitting into a hydrogen ion (H+ ) and a hydroxide ion (OH-).  In a liter of absolutely pure water (55.6 moles) at 25C,  1 x 10-7 moles of water molecules will always dissociate into 1 x 10-7 moles of hydrogen ions and an equal number of moles of hydroxide ions.  Since there are still about 55.6 moles of undissociated water molecules in the liter of water, this doesn't amount to much on a percentage basis.  Yet a little goes a long way as far as hydrogen ions are concerned.  Dissociation is the reason why the pH of pure water is 7.  (See the row in the table for pH=7)

Now notice that [H+] x [OH-] = (1 x 10-7) x (1 x 10-7) = 1 x 10-14 for pure water.  The quantity of [H+] x [OH-] is known as the dissociation constant, Kw. 

Kw = [H+][OH-] = 1 x 10-14  (4)

The really surprising thing is that if you were to dump in some acid that raised the [H+], the [OH-] would automatically decrease so that the product of [H+] x [OH-] would remain at 1 x 10-14! If you measure both the [H+] and the [OH-] of any water-based system, the product [H+] x [OH-] (at 25C) will be 1 x 10-14. Because of this amazing property, we can always figure what the [OH-] is if we know the [H+], and vice versa. 

[H+] = Kw / [OH-]  (5a)
[OH-] = Kw / [H+]  (5b)

Look again at the table and verify that for any pH that  [H+] x [OH-] = 1 x 10-14

Acids

An acid may be thought of as a molecule that splits apart (dissociates) in water yielding a hydrogen ion and a negative anion, symbolized as A-

HA H+  + A-  (6)

A- is also referred to as the conjugate base of the acid HA.

Acids are generally classified as either strong or weak.  A strong acid is one where virtually all the original acid molecules HA dissociate into ions, H+ and A-. This can be represented by equation 6a where the reaction goes to completion.

HA --> H+  + A-  (6a)  (Strong acid)

Hydrochloric acid (HCl) is an example of a strong acid.  Dissolved in water, hydrochloric acid exists only as hydrogen ions and chloride ions.  Sulfuric acid H2SO4 is also a strong acid that exists in aqueous solutions only as H+ and SO42-.

Weak acids however have only a fraction of the molecules dissociating into ions.  In equation 6, a large number of HA molecules remain unreacted.  The degree to which a weak acid with the general formula HA dissociates is defined by the equation

Ka= [H+][A-] / [HA]  (7)

where Ka is known as the acid constant.  Each weak acid has its own unique Ka.

See  pH is a measure of the [H3O +] in solution for supplemental info.

As stated earlier, the presence of other chemicals in the water may change its [H+] and hence its pH.  The addition of H+ or OH- directly (as in the case of adding a strong acid (e.g. hydrochloric acid HCl) or a strong base (e.g. sodium hydroxide NaOH)) will likely alter the pH.  Chemical reactions may modify the pH by producing or consuming either H+ or OH-.

See Visionlearning's Introductory Lesson on Acids and Bases for another treatment of the same subject.

See John L. Park's more rigorous treatment of Acids and Bases.

See Acids and Bases pH Tutorial for another reasonably complete treatment.

From the standpoint of a stream, the healthiest streams (as measured by the biodiversity within the streams) generally have a pH that is close (within 1 pH unit) to neutral or pH=7.  Acidity tends to be the more common problem.  As the pH drops below 6, the aquatic life supported becomes less diverse as acid intolerant species either die off or the creatures that feed on them loose their food supply.


* In this treatment, we refer to the hydrogen ion, H+ as a chemical species.  In actuality, this is not accurate.  The true species we are interested in is H3O+ , an H+ bonded to a water molecule.  This is a notational shortcut that works out, but we should be mindful that in water solutions the species we are really talking about is H3O+ .