The basics of ph

Born in Havrebjerg, Denmark, Sørensen was a Danish chemist, famous for the introduction of the concept of pH, a scale for measuring acidity and basicity. From 1901 to 1938 he was head of the prestigious Carlsberg Laboratory, Copenhagen. While working at the Carlsberg Laboratory he studied the effect of ion concentration on proteins,and because the concentration of hydrogen ions was particularly important, he introduced the pH- scale as a simple way of expressing it in 1909.

The Danish scientist Sørensen defined the concept of pH as follows:

pH equals the inverse of the logarithm to the base 10 of the hydrogen ion concentration, as shown by the formula:

pH = -10log [H+] = paH1) (1)

  • Later Sørensen found this definition to be incorrect, since more concentrated solutions appeared to give deviations between calculated and measured values.
  • The definition therefore had to be modified to:
  • pH equals the inverse of the logarithm to the base 10 of the hydrogen ion activity2) as shown by the formula:

pH = -10log aH+ = pH3) (2)

The activity of the hydrogen ions is not always linear with the concentration, since this activity is not only affected
by the concentration of ions, but also by other factors, such as:

  • The activity of other ions present in the solution
  • The temperature of the solution
  • The character of the solution.

To facilitate the accurate measurement of pH, and its presentation as a scale, a range of “standard liquids” or “buffer solutions” are used.

These liquids, whose constituents are accurately defined, have known stable values.

Although in the preceding text the relationship to hydrogen ions has been made, research has shown, that the activity of hydroxonium ions (H30+) is more relevant. In aqueous solutions free H+ ions do not occur, but are always in combination with water molecules.

H+ + H20 ↔ H30+

Consequently, a more correct definition for pH is:

pH = -10log aH30+ (3)

For clarity, the notation H+ will be used in the book as the hydroxonium ion.

Note 1. The notation -10log …. can also be written p ….
Note 2. See Appendix 2: Definitions.
Note 3. See Chapter 2.8: Buffer solutions.

2.2 The pH Scale

The Basics of pH
Fig. 2.2a. pH value of pure water against temperature.

Your starting point for the pH scale is pure water which is said to be neutral. Water dissociates1) into:

H20 ↔ H+ + OH- (4)

Water has an equilibrium constant 2)3):

The Basics of pH


-log Kw = pKw = -log [H+] + -log [OH-]

= 14 log 10 (6)

Pure water divides to give equal numbers of H+ and OH- ions and consequently, the concentrations of ions are 10-7 so that:

pH = pOH = 7

The pH value of pure water is 7.

This statement is incomplete, since the equilibrium constant depends on the temperature. The definition should be: The pH value of pure water is 7 @ 25°C.

Fig. 2.2a. and the table show the pH variation of pure water with temperature.

If the concentration of H+ ions in a solution is increased (e.g. to 10-4), then the solution has an acid character. In this case the pH value is lower than 7.

Some examples of common solutions with an acid character are:

H2S04 ↔ S042– + 2H+
Sulphuric acid

HCl ↔ Cl– + H+
Hydrochloric acid

If the concentration of OH- ions in a solution is increased (e.g. to 10-10) then the solution is said to have a base character. In this case the pH value of the solution is a number greater than 7.

T(oC) pKw pH
14,94 7,47
18 14,22 7,11
25 14,00 7,00
50 13,22 6,61
100 12,24 6,12

Some more examples are:

NaOH ↔ Na+ + OH-
Caustic soda

NH3+ H2O ↔ NH4+ + OH-
Ammonia aqueous ammonia

Note 1. See Appendix 2: Definitions
Note 2. The equilibrium constant is the ratio between the rate of decomposition and the rate of composition.
Note 3. The concentration H2O is supposed to be 1.

pH Table

Some examples of the difference in pH value of various liquids, foods and fruit are shown in fig. 2.2b. These can be compared with the pH values of common chemical compounds dissolved in water.

The Basics of pH

Fig. 2.2b.

2.3 Measuring the PH Scale

The pH value can be measured by different methods, e.g.:

  1. Colorimetric pH measurement
  2. Potentiometric pH measurement

2.3.1 Colorometric pH measurement

pH and alkalinity

Water quality and pH are often mentioned in the same sentence. The pH is a very important factor, because certain chemical processes can only take place when water has a certain pH. For instance, chlorine reactions only take place when the pH has a value of between 6,5 and 8. The pH is an indication for the acidity of a substance. It is determined by the number of free hydrogen ions (H+) in a substance. Acidity is one of the most important properties of water. Water is a solvent for nearly all ions. The pH serves as an indicator that compares some of the most water-soluble ions. The outcome of a pH-measurement is determined by a consideration between the number of H+ ions and the number of hydroxide (OH-) ions. When the number of H+ ions equals the number of OH- ions, the water is neutral. It will than have a pH of about 7. The pH of water can vary between 0 and 14. When the pH of a substance is above 7, it is a basic substance. When the pH of a substance is below 7, it is an acid substance. The further the pH lies above or below 7, the more basic or acid a solution is. The pH is a logarithmic factor; when a solution becomes ten times more acidic, the pH will fall by one unit. When a solution becomes a hundred times more acidic the pH will fall by two units.The common term for pH is alkalinity.

The word pH is short for “pondus Hydrogenium”. This literally means the weight of hydrogen. De pH is an indication for the number of hydrogen ions. It consisted when we discovered that water consists of hydrogen ions (H+) and hydroxide ions (OH-). The pH does not have a unit; it is merely expressed as a number. When a solution is neutral, the number of hydrogen ions equals the number of hydroxide ions. When the number of hydroxide ions is higher, the solution is basic. When the number of hydrogen ions is higher, the solution is acid. The Basics of pH Did you know that the pH of Coca-Cola is about 2? And did you know that it is useless to measure the pH of RO-water or demiwater? Both demiwater and RO-water do not contain any buffer ions. This means that the pH can be as low as four, but it can also be as high as 12. Both kinds of water are not readily usable in their natural form. They are always mixed before application!

Methods to determine the pH

There are several different methods to measure the pH. One of these is using a piece of pH indicator paper. When the paper is pushed into a solution it will change colour. Each different colour indicates a different pH-value.

This method is not very accurate and it is not suitable to determine more exact pH values. That is why there are now test-strings available, which are able to determine smaller pH-values, such as 3.5 or 8.5.

The most accurate method to determine the pH is measuring a colour change in a chemical lab experiment. With this method one can determine pH values, such as 5.07 and 2.03.

All of these methods are not suitable to determine a pH development in time.

The pH-electrode

A pH electrode is a tube that is small enough to put it in sample jars. It is tied to a pH-meter by means of a cable. A special type of fluid is located within the electrode; this is usually “3M Kalium Chlorine”.

Some electrodes contain a gel that has the same properties as the 3M-fluid. In the fluid there are silver and platinum wires. The system is quite fragile, because it contains a small membrane. The H+ and OH- ions will enter the electrode through this membrane.

The ions will create a slightly positive charge and a slightly negative charge in each end of the electrode. The potential of the charges determines the number of H+ and OH- ions and when this is determined the pH will appear digitally on the pH-meter.

The potential is co-dependent on the temperature of the solution. That is why the temperature is also presented on the pH-meter.

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Acids and bases

When acids enter the water, the ions will separate. For instance, hydrogen chloride will separate into hydrogen and chlorine ions (HCL à H+ + CL-). Bases also undergo separation of their ions when enter the water. When sodium hydroxide enters the water it will separate into sodium and hydroxide ions (NaOH à Na+ + OH-).

The Basics of pH The Basics of pH

When an acid substance ends up in water, it will give up a hydrogen ion to the water. The water will than become acid. The number of hydrogen ions that the water will receive determines the pH.

When a basic substance enters the water it will take up hydrogen ions. This will raise the pH of the water. When a substance is strongly acidic it will give up more H+ ions to the water.

Strong bases will give up more OH-.

Here we have summed up a list of products and their pH:

pH product
14 sodium hydroxide
13 lye
12.4 lyme
11 ammonia
10.5 manganese
8.3 backing powder
7.4 human blood
7.0 pure water
6.6 milk
4.5 tomatoes
4.0 win
3.0 apples
2.0 lemon juice
hydrochloric acid

What Are The Basics of pH? Do You Know The Importance?

What are the Basics of pH? The nature of the solution (acidic, alkaline (or) neutral) can be represented in terms of either hydrogen ion concentration (or) hydroxyl ion concentration.

In 1909, Sorenson used a logarithmic scale for expressing the H+ concentration. This scale was called “pH”, Where “P” stands for ‘Power’ and ‘H’ for hydrogen ion concentration.

  • Sorenson defines the pH of a solution as
  • “ The negative logarithm of hydrogen ions concentration (in moles/liter)”
  • Thus,
  •                            1
  • PH = log ————– = – log [H+] [H+]
  • The symbol “P” denotes “negative logarithm of ”.
  • For a precisely neutral solution at 250C in which the concentration of hydrogen ions is 1.0 X 10-7 M, the pH can be calculated as follows:
  • 1
  • PH =  log———————= log (1.0 X 107)
  •                           1.0 X 10-7
  •                          =log 1.0 + log 107
  •                          = 0 + 7.0
  •                          = 7.0

The value of 7.0 for the pH of a precisely neutral solution is not an arbitrarily chosen figure. It is derived from the absolute value of the ion product of water at 250 C, which by convenient coincidence is a round number.

The Basics of pH

Importance of pH

  • Soils of specific pH are required for optimum group growth and better yields of crops.
  • Specific pH values are to be maintained for the biological process and industrial process to occur.
  • Specific pH is also to be maintained by the blood.
  • PH plays an important role in chemical analysis.

Development reason for pH

  • The pH scale was developed taking water as the standard.
  • It is an experimental fact that only 1 mole in 5,50,000,000 moles of water ionizes into an H+ and OH–.
  • This is the same proportion as one-gram hydrogen ion in 10,000,000 liters of water.
  • Hence, one liter of water contains 1/10,000,000 (or)1/107 of a gram of H+.
  • For everyday use, only the ‘Power’ figure was used and the symbol pH placed before it.

The Basics of pH

Biological Significance of pH

A) Tautomeric forms of purines and Pyrimidines

Tautomerization is a special type of isomerism where a proton migrates in one direction and a covalent bond shifts in the opposite direction within the molecule.  Purine and pyrimidine bases exist different tautomeric forms according to pH.

Their specific tautomeric at the body PH of nearly 7.4 is essential for the hydrogen bonding of complementary base pairs in the DNA double helices and RNA strands. So pH maintains the natural three-dimensional forms of nucleic acid molecules.

B) Isoelectric pH

PH influences the ionization of ionizable polar groups of amino acids, proteins, nucleic acids, Phospholipids, and mucopolysaccharides.

At a specific pH called the isoelectric pH of the molecule, each such molecule exists as dipolar zwitterions bearing both anionic acid and cationic groups and minimum net charge.

Zwitter ions do not migrate in electric fields and precipitate easily by aggregation due to minimum electrostatic repulsion.

C) Isolation of Proteins and Amino acids

The pH depends on the charged forms of proteins and amino acids are utilized in separating and isolating them from biological materials by methods such as “ion-exchange chromatography, Paper electrophoresis, and Isoelectrophoresis”.

D) Optimum pH

By influencing ionized states of proteins, pH affects the ionic and hydrogen bonds which stabilize the three-dimensional structure of proteins.

The pH Scale

Most people are familiar with the words acid and acidic—whether it’s because of acid rain or acidic foods like lemon juice. However, fewer people are aware of acid’s opposite: base (also called alkaline).

Basic substances include things like baking soda, soap, and bleach. Distilled water is a neutral substance. The pH scale, which measures from 0 to 14, provides an indication of just how acidic or basic a substance is.

Most parts of our body (excluding things like stomach acid) measure around 7.2 and 7.6 on the pH scale (a 7 is neutral on the scale). If foreign strong substances dramatically change this pH, our bodies can no longer function properly.

In this outcome, we’ll learn about acids and bases, and what impact they can have on living systems.

Learning Outcomes

  • Identify the characteristics of acids
  • Identify the characteristics of bases
  • Define buffers and discuss the role they play in human biology

The pH scale ranges from 0 to 14. The pH of a solution is a measure of its acidity or alkalinity (base).

You have probably used litmus paper, paper that has been treated with a natural water-soluble dye so it can be used as a pH indicator, to test how much acid or base (alkalinity) exists in a solution.

You might have even used some to make sure the water in an outdoor swimming pool is properly treated.

The Basics of pH

Figure 1. The pH scale measures the amount of hydrogen ions (H+) in a substance. (credit: modification of work by Edward Stevens)

This pH test measures the amount of hydrogen ions that exists in a given solution. High concentrations of hydrogen ions yield a low pH (acidic substances), whereas low levels of hydrogen ions result in a high pH (basic substances). The overall concentration of hydrogen ions is inversely related to its pH and can be measured on the pH scale (Figure 1).

Therefore, the more hydrogen ions present, the lower the pH; conversely, the fewer hydrogen ions, the higher the pH. A change of one unit on the pH scale represents a change in the concentration of hydrogen ions by a factor of 10, a change in two units represents a change in the concentration of hydrogen ions by a factor of 100.

Thus, small changes in pH represent large changes in the concentrations of hydrogen ions. Pure water is neutral. It is neither acidic nor basic, and has a pH of 7.0. Anything below 7.0 (ranging from 0.0 to 6.9) is acidic, and anything above 7.0 (from 7.1 to 14.0) is alkaline. The blood in your veins is slightly alkaline (pH = 7.4).

The environment in your stomach is highly acidic (pH = 1 to 2). Orange juice is mildly acidic (pH = approximately 3.5), whereas baking soda is basic (pH = 9.0).

Acids are substances that provide hydrogen ions (H+) and lower pH, whereas bases provide hydroxide ions (OH–) and raise pH. The stronger the acid, the more readily it donates H+.

For example, hydrochloric acid and lemon juice are very acidic and readily give up H+ when added to water. Conversely, bases are those substances that readily donate OH–. The OH– ions combine with H+ to produce water, which raises a substance’s pH.

Sodium hydroxide and many household cleaners are very alkaline and give up OH– rapidly when placed in water, thereby raising the pH.


Most cells in our bodies operate within a very narrow window of the pH scale, typically ranging only from 7.2 to 7.6. If the pH of the body is outside of this range, the respiratory system malfunctions, as do other organs in the body. Cells no longer function properly, and proteins will break down. Deviation outside of the pH range can induce coma or even cause death.

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So how is it that we can ingest or inhale acidic or basic substances and not die? Buffers are the key. Buffers readily absorb excess H+ or OH–, keeping the pH of the body carefully maintained in the aforementioned narrow range.

Carbon dioxide is part of a prominent buffer system in the human body; it keeps the pH within the proper range. This buffer system involves carbonic acid (H2CO3) and bicarbonate (HCO3–) anion.

If too much H+ enters the body, bicarbonate will combine with the H+ to create carbonic acid and limit the decrease in pH.

Likewise, if too much OH– is introduced into the system, carbonic acid will rapidly dissociate into bicarbonate and H+ ions. The H+ ions can combine with the OH– ions, limiting the increase in pH.

While carbonic acid is an important product in this reaction, its presence is fleeting because the carbonic acid is released from the body as carbon dioxide gas each time we breathe.

Without this buffer system, the pH in our bodies would fluctuate too much and we would fail to survive.

The pH of a solution is a measure of the concentration of hydrogen ions in the solution. A solution with a high number of hydrogen ions is acidic and has a low pH value. A solution with a high number of hydroxide ions is basic and has a high pH value.

The pH scale ranges from 0 to 14, with a pH of 7 being neutral. Buffers are solutions that moderate pH changes when an acid or base is added to the buffer system.

Buffers are important in biological systems because of their ability to maintain constant pH conditions.

Using a pH meter, you find the pH of an unknown solution to be 8.0. How would you describe this solution?

  1. weakly acidic
  2. strongly acidic
  3. weakly basic
  4. strongly basic

The pH of lemon juice is about 2.0, whereas tomato juice’s pH is about 4.0. Approximately how much of an increase in hydrogen ion concentration is there between tomato juice and lemon juice?

  1. 2 times
  2. 10 times
  3. 100 times
  4. 1000 times

Answer the question(s) below to see how well you understand the topics covered in the previous section. This short quiz does not count toward your grade in the class, and you can retake it an unlimited number of times.

Use this quiz to check your understanding and decide whether to (1) study the previous section further or (2) move on to the next section.

Acids, Bases, & the pH Scale

It all has to do with hydrogen ions (abbreviated with the chemical symbol H+). In water (H2O), a small number of the molecules dissociate (split up). Some of the water molecules lose a hydrogen and become hydroxide ions (OH−).

The “lost” hydrogen ions join up with water molecules to form hydronium ions (H3O+). For simplicity, hydronium ions are referred to as hydrogen ions H+. In pure water, there are an equal number of hydrogen ions and hydroxide ions.

The solution is neither acidic or basic.

An acid is a substance that donates hydrogen ions. Because of this, when an acid is dissolved in water, the balance between hydrogen ions and hydroxide ions is shifted. Now there are more hydrogen ions than hydroxide ions in the solution. This kind of solution is acidic.

A base is a substance that accepts hydrogen ions. When a base is dissolved in water, the balance between hydrogen ions and hydroxide ions shifts the opposite way. Because the base “soaks up” hydrogen ions, the result is a solution with more hydroxide ions than hydrogen ions. This kind of solution is alkaline.

Acidity and alkalinity are measured with a logarithmic scale called pH.

Here is why: a strongly acidic solution can have one hundred million million, or one hundred trillion (100,000,000,000,000) times more hydrogen ions than a strongly basic solution! The flip side, of course, is that a strongly basic solution can have 100,000,000,000,000 times more hydroxide ions than a strongly acidic solution. Moreover, the hydrogen ion and hydroxide ion concentrations in everyday solutions can vary over that entire range.

In order to deal with these large numbers more easily, scientists use a logarithmic scale, the pH scale. Each one-unit change in the pH scale corresponds to a ten-fold change in hydrogen ion concentration.

The pH scale is theoretically open-ended but most pH values are in the range from 0 to 14.

It's a lot easier to use a logarithmic scale instead of always having to write down all those zeros! By the way, notice how one hundred million million is a one with fourteen zeros after it? It is not coincidence, it is logarithms!

To be more precise, pH is the negative logarithm of the hydrogen ion concentration:

pH = −log [H+]

The square brackets around the H+ automatically mean “concentration” to a chemist. What the equation means is just what we said before: for each 1-unit change in pH, the hydrogen ion concentration changes ten-fold.

Pure water has a neutral pH of 7. pH values lower than 7 are acidic, and pH values higher than 7 are alkaline (basic).

Table 1 has examples of substances with different pH values (Decelles, 2002; Environment Canada, 2002; EPA, date unknown).

pH Value H+ Concentration Relative to Pure Water Example
10 000 000 battery acid
1 1 000 000 gastric acid
2 100 000 lemon juice, vinegar
3 10 000 orange juice, soda
4 1 000 tomato juice, acid rain
5 100 black coffee, bananas
6 10 urine, milk
7 1 pure water
8 0.1 sea water, eggs
9 0.01 baking soda
10 0.001 Great Salt Lake, milk of magnesia
11 0.000 1 ammonia solution
12 0.000 01 soapy water
13 0.000 001 bleach, oven cleaner
14 0.000 000 1 liquid drain cleaner

Table 1. The pH Scale: Some Examples


For other uses, see PH (disambiguation).

pH values of some common substances

Acids and bases Acid types Base types
  • Acid
  • Acid–base reaction
  • Acid strength
  • Acidity function
  • Amphoterism
  • Base
  • Buffer solutions
  • Dissociation constant
  • Equilibrium chemistry
  • Extraction
  • Hammett acidity function
  • pH
  • Proton affinity
  • Self-ionization of water
  • Titration
  • Lewis acid catalysis
  • Frustrated Lewis pair
  • Chiral Lewis acid
  • Brønsted–Lowry
  • Lewis
  • Acceptor
  • Mineral
  • Organic
  • Oxide
  • Strong
  • Superacids
  • Weak
  • Solid
  • Brønsted–Lowry
  • Lewis
  • Donor
  • Organic
  • Oxide
  • Strong
  • Superbases
  • Non-nucleophilic
  • Weak
  • v
  • t
  • e

In chemistry, pH (/piːˈeɪtʃ/) (abbr. power of hydrogen[1] or potential for hydrogen[2]) is a scale used to specify how acidic or basic a water-based solution is. Acidic solutions have a lower pH, while basic solutions have a higher pH. At room temperature (25°C or 77°F), pure water is neither acidic nor basic and has a pH of 7.

The pH scale is logarithmic and inversely indicates the concentration of hydrogen ions in the solution (a lower pH indicates a higher concentration of hydrogen ions).

This is because the formula used to calculate pH approximates the negative of the base 10 logarithm of the molar concentration[a] of hydrogen ions in the solution.

More precisely, pH is the negative of the base 10 logarithm of the activity of the hydrogen ion.[3]

At 25 °C, solutions with a pH less than 7 are acidic, and solutions with a pH greater than 7 are basic. The neutral value of the pH depends on the temperature, being lower than 7 if the temperature increases. The pH value can be less than 0 for very strong acids, or greater than 14 for very strong bases.[4]

The pH scale is traceable to a set of standard solutions whose pH is established by international agreement.

[5] Primary pH standard values are determined using a concentration cell with transference, by measuring the potential difference between a hydrogen electrode and a standard electrode such as the silver chloride electrode.

The pH of aqueous solutions can be measured with a glass electrode and a pH meter, or a color-changing indicator. Measurements of pH are important in chemistry, agronomy, medicine, water treatment, and many other applications.


The concept of pH was first introduced by the Danish chemist Søren Peder Lauritz Sørensen at the Carlsberg Laboratory in 1909[6] and revised to the modern pH in 1924 to accommodate definitions and measurements in terms of electrochemical cells. In the first papers, the notation had the H as a subscript to the lowercase p, as so: pH.

See also:  ‘so’ and ‘so that’: coordinating or subordinating conjunction?

The exact meaning of the p in pH is disputed, as Sørensen did not explain why he used it.[7] He describes a way of measuring it using potential differences, and it represents the negative power of 10 in the concentration of hydrogen ions.

All the words for these start with p in French, German and Danish, all languages Sørensen published in: Carlsberg Laboratory was French-speaking, German was the dominant language of scientific publishing, and Sørensen was Danish. He also used “q” in much the same way elsewhere in the paper.

So the “p” could stand for the French puissance, German Potenz, or Danish potens, meaning “power”, or it could mean “potential”. He might also just have labelled the test solution “p” and the reference solution “q” arbitrarily; these letters are often paired.

[8] There is little to support the suggestion that “pH” stands for the Latin terms pondus hydrogenii (quantity of hydrogen) or potentia hydrogenii (power of hydrogen).

Currently in chemistry, the p stands for “decimal cologarithm of”, and is also used in the term pKa, used for acid dissociation constants[9] and pOH, the equivalent for hydroxide ions.

Bacteriologist Alice C.

Evans, famed for her work's influence on dairying and food safety, credited William Mansfield Clark and colleagues (of whom she was one) with developing pH measuring methods in the 1910s, which had a wide influence on laboratory and industrial use thereafter. In her memoir, she does not mention how much, or how little, Clark and colleagues knew about Sørensen's work a few years prior.[10]:10 She said:

In these studies [of bacterial metabolism] Dr. Clark's attention was directed to the effect of acid on the growth of bacteria. He found that it is the intensity of the acid in terms of hydrogen-ion concentration that affects their growth. But existing methods of measuring acidity determined the quantity, not the intensity, of the acid.

Next, with his collaborators, Dr. Clark developed accurate methods for measuring hydrogen-ion concentration. These methods replaced the inaccurate titration method of determining acid content in use in biologic laboratories throughout the world.

Also they were found to be applicable in many industrial and other processes in which they came into wide usage.[10]:10

The first electronic method for measuring pH was invented by Arnold Orville Beckman, a professor at California Institute of Technology in 1934.[11] It was in response to local citrus grower Sunkist that wanted a better method for quickly testing the pH of lemons they were picking from their nearby orchards.[12]

Definition and measurement

pH is defined as the decimal logarithm of the reciprocal of the hydrogen ion activity, aH+, in a solution.[5]

pH Definition – basic (alkaline) and acidic

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pH can be viewed as an abbreviation for power of Hydrogen – or more completely, power of the concentration of the Hydrogen ionin a liquid.

The mathematical definition of pH is a bit less intuitive but in general more useful. It says that the pH is equal to to the negative logarithmic value of the Hydrogen ion (H+) concentration, or

pH = -log [H+]

pH can alternatively be defined mathematically as the negative logarithmic value of the Hydroxonium ion (H3O+) concentration. Using the Bronsted-Lowry approach

pH = -log [H3O+]

pH values are calculated in powers of 10. The hydrogen ion concentration in a solution with pH 1.0 is 10 times larger than the hydrogen concentration in a solution with pH 2.0. The larger the hydrogen ion concentration – the smaller the pH.

  • when the pH is above 7 the solution is basic (alkaline)
  • when the pH is below 7 the solution is acidic

In pure neutral water the concentration of hydrogen and hydroxide ions are both 10-7 equivalents per liter.

pHIon Concentration (gram equivalent per liter)Type of Solution
1.0 Acid Solution – Hydrogen ions – H+
1 0.1
2 0.01
3 0.001
4 0.0001
5 0.00001
6 0.000001
7 0.0000001 Neutral Solution (pure neutral water)
8 0.000001 Basic (alkaline) Solution – Hydroxide ions -OH-
9 0.00001
10 0.0001
11 0.001
12 0.01
13 0.1
14 1.0

Some common Products and their pH Values

pH values in some common products:

Battery Acid
HCl in stomach acid 1
Lemon juice, vinegar 2-3
Orange juice 3-4
Acid rain 4
Black coffee 5
Urine, salvia 6
Pure water 7
Sea water 8
Baking soda 9
Ammonia solution 10-11
Soapy water 12
Bleach 13
Oven cleaner 13-14
Drain cleaner 14

Acid-Base Indicators

IndicatorpH – RangeColor CangeAcidBase
Thymol blue 1.2 – 2.8 red yellow
Pentamethoxy red 1.2 – 2.3 red – violet colorless
Tropeolin 1.3 – 3.2 red yellow
2,4 – Dinitrophenol 2.4 – 4.0 colorless yellow
Methyl yellow 2.9 – 4.0 red yellow
Methyl orange 3.1 – 4.4 red orange
Congo red 3.0 – 4.2 blue-violet red-orange
Bromphenol blue 3.0 – 4.6 yellow blue – violet
Tetrabromphenol blue 3.0 – 4.6 yellow blue
Alizarin sodium sulfonate 3.7 – 5.2 yellow violet
α – Naphthyl red 3.7 – 5.0 red yellow
p – Ethoxychrysoidine 3.5 – 5.5 red yellow
Bromcresol green 4.0 – 5.6 yellow blue
Methyl red 4.4 – 6.2 red yellow
Bromcresol purple 5.2 – 6.8 yellow purple
Chlorphenol red 5.4 – 6.8 yellow red
Bromphenol blue 6.2 – 7.6 yellow blue
p – Nitrophenol 5.0 – 7.0 colorless yellow
Litmus 5.0 – 8.0 red blue
Azolitmin 5.0 – 8.0 red blue
Phenol red 6.4 – 8.0 yellow red
Neutral red 6.4 – 8.0 red yellow
Rosolic acid 6.8 – 8.0 yellow red
Cresol red 7.2 – 8.8 yellow red
α – Naphtholphthalein 7.3 – 8.7 rose green
Tropeolin 7.6 – 8.9 yellow rose – red
Thymol blue 8.0 – 9.6 yellow blue
Phenolphthalein 8.0 – 10.0 colorless red
α – Naphtholbenzein 9.0 – 11.0 yellow blue
Thymolphthalein 9.4 – 10.6 colorless blue
Nile blue 10.1 – 11.1 blue red
Alizarin yellow 10.0 – 12.0 yellow lilac
Salicyl yellow 10.0 – 12.0 yellow orange – brown
Diazo violet 10.1 – 12.0 yellow violet
Tropeolin 11.0 – 13.0 yellow orange – brown
Nitramine 11.0 – 13.0 colorless orange – brown
Poirrier's blue 11.0 – 13.0 blue violet – pink
Trinitrobenzoic acid 12.0 – 13.4 colorless orange – red

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BioMath: pH

Understanding pH is essential in chemistry and biology. pH is defined by the following equation,

pH = −log [H+] ,

where [H+] denotes the molar hydrogen ion concentration. Notice that we are required to take the common (base 10) logarithm of the hydrogen ion concentration in order to calculate pH.

Because pH is a measure of hydrogen ion concentration, it is used to quantitatively characterize solutions as acidic, neutral, or basic (alkaline). The typical pH scale runs from 0 – 14. A pH of 7 is neutral, a pH < 7 is called acidic while pH > 7 is called basic.

  • A note of caution while working with pH.
  • Remember that pH is calculated on a logarithmic scale, therefore small differences in pH represent much larger differences in hydrogen ion concentration.
  • For example,

a solution with pH 3 ( i.e. [H+] = 1 × 10 − 3 M )

is ten times more acidic than

a solution with pH 4  (i.e. [H+] = 1 × 10 − 4 M )

  1. In fact, for each unit increase in pH, there is a 10 fold increase in the hydrogen ion concentration.
  2. For more on pH see the Acids and Bases Problem Set of the Chemistry section.
  3. Now that you have been introduced to some pH basics, try the following 6 problems.
  4. Problem 1- Calculate the pH of lemon juice
  5. Problem 2 – Calculate the pH of an unknown solution
  6. Problem 3 – Find the hydrogen ion concentration of an unknown solution
  7. Problem 4 -Determine the hydrogen concentration of blood
  8. Problem 5 – Compare the acidity of two solutions
  9. Problem 6 – The pH range of blood
  10. Next Application: Drug Concentrations

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