Biology A
Modified from http://www.visionlearning.com/library/module_viewer.php?mid=58
2
You have learned that water is polar, with the O portion of the molecule attracting most of the electrons and therefore acting negatively charged, while the H atoms are mostly left without nearby electrons and so act positively charged. The opposite poles attract one another through hydrogen bonds and create cohesion between water molecules. The picture at left represents the situation in ice, where all the molecules are locked into position. In liquid water, these hydrogen bonds are constantly breaking and forming with new partners.
Sometimes, the O does such a good job of attracting H’s that it actually pulls one H +
away from another water molecule like this:
This results in two ions. The hydronium with the extra H + is positively charged, while the hydroxide that lost the H
+
(but still has the electron from it) is negatively charged.
In a sample of 1 million water molecules, only a couple of molecules will do this at any one time and they will always be perfectly balanced. In fact, what we really care about is not so much the hydronium ions that can form, but the H + ions themselves. The reaction can be shown as
H
2
O → H + + OH
− and the ions can reform into H
2
O.
This is important because we define acids and bases by whether they release H
+
(acid) or OH
−
(base) when they are put in water.
Acids and Bases
Acids taste sour, are corrosive to metals, change litmus (a dye extracted from lichens) red, and become less acidic when mixed with bases.
Bases feel slippery, change litmus blue, and become less basic when mixed with acids.
In the late 1800s, the Swedish scientist Svante Arrhenius proposed that water can dissolve many compounds by separating them into their individual ions. Arrhenius suggested that acids are compounds that contain hydrogen and can dissolve in water to release hydrogen ions into solution. For example, hydrochloric acid (HCl) dissolves in water as follows:

HCl
H
2
O
H
+
( aq )
+ Cl
-
( aq )
Note that ( aq ) means this is happening in water
Arrhenius defined bases as substances that dissolve in water to release hydroxide ions (OH-) into solution. For example, sodium hydroxide (NaOH) is a typical base because it does this:
NaOH
H
2
O
Na
+
( aq )
+ OH
-
( aq )
The Arrhenius definition explains the observation that acids and bases counteract each other.
This idea, that a base can make an acid weaker, and vice versa, is called neutralization. This explanation didn’t explain how substances like sodium bicarbonate NaHCO
3 that don’t contain
OH- could neutralize acids, but they do it very well. Most stomach antacids contain sodium bicarbonate.
As you can see from the equations, acids release H
+
into solution and bases release OH
-
. If we were to mix an acid and base together, the H + ion would combine with the OH ion to make the molecule H
2
O, or plain water:
H
+
( aq )
+ OH
-
( aq )
H
2
O
But remember that an acid is a substance that releases the H
+ ion and a base is the substance that releases the OH
-
ion. This means that while we are happy that we have made a molecule of water, we also have to think about what happens to the rest of the acid and base substances. The answer is that those substances are also oppositely charged ions and will form an ionically bonded compound called a salt if the water is removed.
The neutralization reaction of an acid with a base will always produce water and a salt, as shown below:
Acid
HCl
Or H
+
Cl
-
HBr
Base Water
+ NaOH H
2
O +
+ Na
+
OH
-
Salt
NaCl
H
+
OH
-
+ Na
+
Cl
-
+ KOH H
2
O + KBr
It’s interesting to note that sodium chloride, aka NaCl, aka table salt is what you get if you mix hydrochloric acid and sodium hydroxide (lye). Rearranging the ions turns two nasty substances into two harmless, and even necessary ones!

In 1923, the Danish scientist Johannes Brønsted modified Arrhenius' theory of acids and bases. The definition of acids remained any substance that can donate a hydrogen ion is an acid, but the Brønsted definition of bases is, however, quite different. According to Brønsted, a base is defined as any substance that can accept a hydrogen ion. In essence, a base is the opposite of an acid. NaOH and KOH, as we saw above, would still be considered bases because they can accept an H + from an acid to form water. However, the Brønsted definition also explains why substances that do not contain OH
-
can act like bases. Baking soda (NaHCO
3
), for example, acts like a base by accepting a hydrogen ion from an acid as illustrated below:
Acid Base Salt
HCl + NaHCO
3
H
2
CO
3
+ NaCl
In this example, the carbonic acid formed (H
2
CO
3
) rapidly breaks down to water and gaseous carbon dioxide, and so the solution bubbles as CO
2
gas is released. Bicarbonate is used in many parts of the human body to accept hydrogen ions and keep acids in check.
Great. So why should we care how the base is defined? Under the Brønsted definition, both acids and bases are related to the concentration of hydrogen ions present. Acids increase the concentration of hydrogen ions, while bases decrease the concentration of hydrogen ions (by accepting them). The acidity or basicity of something, therefore, can be measured by its hydrogen ion concentration. Now we have a way to label how strong and acid or base is.
In 1909, the Danish biochemist Sören Sörensen invented the pH scale for measuring acidity.
The pH scale is described by a terrible formula that you don’t need to worry about, but it does make some sense if we consider the concentration of H + ions in a solution. For example, a solution with an H
+
concentration = 1 x 10
-7
moles/liter has a pH equal to 7. A simpler way to think about pH is that it equals the exponent of the H
+
concentration, ignoring the minus sign. A solution with a pH of 6 would have an H
+
concentration of 1 x 10
-6 moles/liter, making it 10 times more concentrated than the solution with a pH of 7. So pH stands for p ower (as in exponent) of H ydrogen. The pH scale ranges from 0 (just H
+
with nothing to accept them) to 14
(no H
+
at all).

Substances with a pH between 0 and less than 7 are acids; pH and H
+
concentration are inversely related so lower pH means higher H + .

Substances with a pH greater than 7 and up to 14 are bases; higher pH means lower H
+ concentration.

Right in the middle, at pH = 7, are neutral substances, for example, pure water.
Remember that we started this explanation by telling you about the dissociation of H
2
O into H
+ and OH
-
that doesn’t happen very often? It turns out that 1 x 10
-7
moles/liter is exactly the concentration of H
+
in a sample of pure water. That’s why pure water sits right in the middle of the pH scale with a pH of 7.

The relationship between [H
+
] and pH is shown in the table below alongside some common examples of acids and bases in everyday life.
Acids
[H + ] pH Example
1
X
10
0
0 HCl
1 x 10 -1 1 Stomach acid
1 x 10
-2
2 Lemon juice
1 x 10
-3
3 Vinegar
1 x 10
-4
4 Soda
1 x 10
-5
5 Rainwater
1 x 10
-6
6 Milk
Neutral 1 x 10
-7
7 Pure water
Bases
1 x 10 -8 8 Egg whites
1 x 10
-9
9 Baking soda
1 x 10 -10 10 Tums
®
antacid
1 x 10
-11
11 Ammonia
1 x 10 -12 12 Mineral lime - Ca(OH)
2
1 x 10
-13
13 Drano
®
1 x 10
-14
14 NaOH
For an interactive pH scale and more explanation of the basis for the pH scale, go to http://www.johnkyrk.com/pH.html
Biology?
You should recall that living things must be able to maintain homeostasis, and the pH of various bodily fluids and tissues is critical. Most body pH is maintained between about 6.5 and 7.5. The blood must be maintained at a pH of between 7.35 and 7.45, but stomach acid must be at a low

pH (generally between 1 and 3). This acidic environment allows digestive enzymes from the stomach to work properly. After leaving the stomach, partially digested food must be brought to a pH of 8 or higher in order for the digestive enzymes of the small intestines to work. Saliva typically has a pH of between 6.35 and 7.4.
Buffers
In all of these areas, we maintain the H
+
concentration between these various desired ranges by using buffers. A buffer is a substance like a weak acid or base that can either absorb or release
H
+
as needed to maintain a steady pH, and bicarbonates play a big role here. Blood uses a bicarbonate to interact with carbonic acid to maintain pH. Sodium bicarbonate is the active ingredient in many antacids. Also note that a buffer does not necessarily neutralize a solution, instead it keeps the pH stable at whatever level is needed.

Name
Biology A
What is pH ?
: Hydroxide Ion (OH
-
) : Hydrogen Ion (H
+
)
A B C D
1.
Count the Hydrogen and Hydroxide ions for each beaker. Rank them in order of their approximate pH from most acidic to the most basic.
2.
Which beaker contains only pure water? Explain how you know.
3.
Consider a solution with a pH of 4 and a solution with a pH of 5. Explain what the numbers tell us about the relative numbers of hydrogen ions in these solutions.
4.
a. What is the pH range of most areas of your body? b. Name two places in your body where the pH is NOT in this range. Include what the pH is in each place and why is it necessary for the pH to be different in order for each part to function properly.

5.
What is a buffer?
6.
Explain why it is incorrect to say “buffers exist in the blood to make sure the pH of blood remains neutral.”
8.
What happens to the concentration of hydrogen ions when an acid is added to a beaker containing water? What happens to the pH?
9.
What happens to the concentration of hydroxide ions when a base is added to a beaker containing water? What happens to the pH?
10.
Solution A (pH 4) and Solution B (pH 9) are mixed. a.
What type of reaction happens? b.
What are the two products that will form? c.
Approximately where on the pH scale will the resulting mixture fall?