Understanding pH and buffers
Objectives
After completing this exercise, you should be able to:
• Define and correctly use the following terms: dissociation, acid, base, pH, buffer, buffering range, and buffering c
...
Understanding pH and buffers
Objectives
After completing this exercise, you should be able to:
• Define and correctly use the following terms: dissociation, acid, base, pH, buffer, buffering range, and buffering capacity
• Explain how solutions with different pHs compare with respect to their H+ and OH- concentrations
• Calibrate and use a pH meter to measure the pH of a solution
• Plot and interpret a pH titration curve to determine the buffering range and buffering capacity of a buffered solution
Before Coming to Lab
• Read the Prelab Discussion and do all the Your Turn exercises.
• You will need to use safety glasses or goggles and closed toed shoes.
Prelab
Dissociation of water
The chemistry of life is based largely on the chemistry of water. When ionic or polar solutes are mixed with water, the polar water molecules will be attracted to them. Molecules held together by ionic bonds may be pulled apart to separate the oppositely charged ions. This process is called dissociation. In fact, water molecules may even pull other water molecules apart to produce oppositely charged hydrogen ions (H+) and hydroxide ions (OH-):
H2O ↔ H+ + OH-
The double-headed arrow in the equation above indicates that this is a reversible reaction – the ions can join back together to reform water. Because each water molecule that dissociates produces one H+ and one OH-, pure water always has equal concentrations of H+ and OH-, which turn out to be 1 × 10-7 M:
In pure water, [H+ ] = [OH- ] = 1 x 10-7 M
The square brackets means “concentration of”, so “[H+]” means “concentration of hydrogen ions”. H+ goes by several names. It is often called a “hydrogen ion”, but can also be called a “hydronium ion” or a “proton”.
Acids and bases
Although acids and bases have been defined in several different ways, in this class we will use the Bronsted-Lowry definition of an acid as an H+ donor, and a base as an H+ acceptor. Examine the following chemical equation involving hydrochloric acid (HCl) and ammonia (NH3):
HCl + NH3 Cl- + NH4+
For the reaction that occurs in the direction of the larger arrow, hydrochloric acid (HCl) is the acid because it dissociates to form H+ and Cl-, and is thus acting as an H+ donor. Ammonia (NH3) is the base because it binds the released H+ to form an ammonium ion (NH4+).
The reaction can occur at a slower rate in the opposite direction, as indicated by the smaller arrow. In this direction, NH4+ is the H+ donor, or conjugate acid, because it dissociates to form H+ and NH3. Cl- is the H+ acceptor, or conjugate base, because it binds the released H+ to form HCl.
Because acids release H+ when they dissociate in water, acidic solutions have a greater [H+] than [OH-].
Bases, on the other hand, decrease the [H+] in a solution in one of two possible ways:
• Some bases dissociate to form OH-. For example, sodium hydroxide (NaOH) dissociates to form Na+ and OH-. Thus, an aqueous solution of NaOH would have a greater [OH-] than [H+]. [H+] is further lessened because the extra OH- binds some of the H+ ions in the solution to form water molecules.
• Some bases do not dissociate and release OH-. Instead, these bases bind H+ ions, thus removing “free” H+ from the solution. As the supply of free H+ decreases, water molecules in the solution dissociate and increase the supply of OH- ions. The NH3 in the example from above is this type of base.
In summary, when an acid is added to an aqueous solution, it increases [H+] so that it is greater than [OH-]. When a base is added to an aqueous solution, [OH-] will be greater than [H+].
Acids and bases can vary in their “strengths”. Strong acids, such as hydrochloric acid (HCl) and sulfuric acid (H2SO4), almost completely dissociate in water. Therefore, they cause a relatively large increase in [H+] and a correspondingly large decrease in [OH-]. Weak acids, like acetic acid or citric acid, dissociate much more slowly. This results in a relatively small increase in [H+] and a correspondingly small decrease in [OH-].
Similarly, strong bases, such as NaOH and potassium hydroxide (KOH), dissociate easily and cause a large increase in [OH-] when dissolved in water, while weak bases such as ammonia (NH3) cause a much smaller increase in [OH-]. Many biomolecules act as weak acids and/or bases.
pH
The pH of a solution is a continuous random variable and is a positive value that ranges from 0 to 14. pH is defined as the negative base-10 logarithm of the hydrogen ion concentration of the solution:
pH = - log [H+ ]
Since it is a logarithm, pH is a unitless measurement.
In this equation, the base-10 logarithm (or “log”) is the power to which 10 must be raised to give the desired number. For example, the log of 100 is 2 (to get 100, we must raise 10 to the power of 2) and the log of 0.01 is -2 (to get 0.01, we must raise 10 to the power of -2).
Notice that the equation takes the negative log of the [H+]. This ensures that the pH will be a positive value. Example: A neutral solution has a hydrogen ion concentration of 1 × 10-7 M. The log of 1 × 10-7 is -7, so the negative log of 1 × 10-7 is 7. Therefore, the pH of a neutral solution is 7.
The [H+] of a solution is inversely related to its pH:
An acid has a [H+] greater than 1 x10-7 M and a pH lower than 7. A base has a [H+] less than 1 x10-7 M and a pH higher than 7.
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