Chemistry ch 16: pH, TITRATION, INDICATORS

Ionization of Water

Water ionizes only slightly according to the following equation: H₂O_(l) Ö H⁺¹ + OH^–1 . This produces equal numbers of H⁺¹ and OH^–1 . The H⁺¹ that is produced is very reactive and attaches itself to another water molecule to form the hydronium ion, H₃O⁺¹ . The self ionization of water can now be written as: H₂O_(l) + H₂O_(l) Ö H₃O⁺¹ + OH^–1.

In pure water [H₃O⁺¹] = [OH^–1] and the solution is neutral. At 25 ºC these concentrations have been measured to be [H₃O⁺¹] = [OH^–1] = 1.00 x 10^–7. Their product is called the ion product constant of water (or the ionization constant of water) and is given the symbol K_w and numerically at 25 ºC is 1.00 x 10^–14. Therefore: [H₃O⁺¹]^.[OH^–1] = 1.00 x 10^–14 . If you know either the [H₃O⁺¹] or the [OH^–1] you can determine the other:

and

If you add an acid to a neutral solution, the [H₃O⁺¹] increases and the [OH^–1] decreases so [H₃O⁺¹] > [OH^–1] and it is now considered to be an acidic solution. If you add a base to a neutral solution, the [OH^–1] increases and the [H₃O⁺¹] decreases so [H₃O⁺¹] < [OH^–1] and it is now considered to be a basic solution.

The pH Scale

The pH scale is a simple way to state what the [H₃O⁺¹] is in a solution. It identifies the relative acidity of the solution. pH is expressed as a power of ten, usually from 10⁰ to 10^-14 . If [H₃O⁺¹] = [OH^–1] = 1.00 x 10^–7 is considered neutral than the pH = 7. Definition: pH = –log₁₀ [H₃O⁺¹].

A known pH value can be used to find [H₃O⁺¹]. The [H₃O⁺¹] = 10^–pH.

The pH scale is usually expressed from 0 to 14. A neutral solution has a pH = 7. An acidic solution has a pH< 7 and a basic solution has a pH >7.

If there is a pH, there is also a pOH used to state what the [OH^–1] is in a solution.

Definition: pOH = –log₁₀ [OH^–1] . The [OH^–1] = 10^–pOH. Mathematically, pH + pOH = 14 .

Indicators

Indicators are used to determine the relative acidity of a solution. These can be an instrument or a chemical. The pH meter is an electrical devise used to determine the pH of a solution. But indicator usually implies a chemical that changes color with a changing pH. The range of pH over which the indicator changes is called the transition interval. Chemical indicators are weak organic acids or bases whose color is different from it's conjugate. An example is phenolphthalein: in an acidic solution phenolphthalein is colorless but in a basic solution (it's conjugate base) it is a bright fuchsia (or magenta or pink) color.

Chemical indicators have several drawbacks or limitations to their use: 1) they are usually only good over a narrow pH range, so you must know the approximate pH range of the substance you want to test; 2) since the pH is based on color differences, the solution you are testing must be colorless or near-colorless so it doesn't interfere with the pH color changes; and 3) since the pH is based on color differences, you must be able to distinguish the subtle color changes and the human eye is limited in this capacity (especially if someone is color-blind!).

Titration

In a chemistry laboratory it is sometimes necessary to experimentally determine the concentration of an acid solution or a base solution. A procedure for making this kind of determination is called an acid-base titration. It is a form of volumetric analysis. It is a quantitative process to determine the unknown concentration of a solution by comparing it with the known concentration of another solution. The equivalence point is the point where moles of [H₃O⁺¹] equal the moles of [OH^–1]. The pH of the equivalence point is NOT always 7, a neutral solution. A strong acid neutralized by a weak base will have a pH < 7, and a weak acid neutralized by a strong base will have a pH > 7.

In this procedure, a solution of known concentration, called the standard solution, is used to neutralize a precisely measured volume of the solution of unknown concentration to which one or two drops of an appropriate acid-base indicator have been added. If the solution of unknown concentration is an acid, a standard base solution is added to the acid solution until it is neutralized. If the solution of unknown concentration is a base, a standard acid solution is added to the base solution until it is neutralized.

When carrying out an acid-base titration, you must be able to recognize when to stop adding the standard solution, that is, when neutralization is reached. This is the purpose of the acid-base indicator mentioned above. A sudden change in color of the indicator signals that neutralization has occurred. At this point, the number of moles of the hydronium ions from the acid is equal to the number of moles of the hydroxide ions from the base. The point at which this occurs is called the endpoint of the titration. When the endpoint is reached, the volume of the standard solution used is carefully determined. The volume of the unknown solution is also carefully measured. Then, the measured volumes of the two solutions and the known concentration of the standard solution can be used to calculate the concentration of the other.

One of the following methods can be used to calculate unknown concentration or volumes in an acid-base titration:

1. Write a balanced equation for the reaction. From the coefficients, determine how many moles of acid react with one

mole of base or vice versa. Use coefficients to form a mole ratio.

2. The following relationship can be used to calculate the unknown concentration:M_{a *} n_{a *} V_a = M_{b *} n_{b *} V_b ,

where M_a is the molarity of the acid, M_b is the molarity of the base, V_a is the volume of the acid, V_b is the volume of the base, n_a is the number of H⁺¹ ionized from one acid and n_b is the number of OH^–1 ionized from one base. (This can also be used to determine the volume of an acid or a base needed to neutralize a solution.)