Discussion

Most of us have heard of or used the pH scale as a measure of the acidity of water, soil or shampoo. In fact, one of the most valuable uses of pH is as a measure of the acidity of a solution. The concept of pH was developed because the hydrogen ion concentration in solutions can vary by many orders of magnitude and may have values as low as 1x10^-14 M. To avoid the nuisance of writing these tiny numbers, a logarithmic or pH scale was established which more conveniently communicates acidity values. The letter p is defined as –log₁₀ thus pH means –log₁₀[H⁺] where [H⁺] is the concentration of hydrogen ions. The definition includes a negative sign to avoid negative signs in the pH values. Because of the negative sign, the lower the pH is, the higher the acidity.

Water dissociation:

Water dissociates into hydronium ions and hydroxide ions: H₂O + H₂O ⇄ H₃O⁺ + OH^1-

The reaction is characterized by the equilibrium expression K_w which has a value of 1x10^-14. The pH scale is based upon this reaction. When an acid or base is dissolved in water, the aqueous solution must maintain a relationship between the hydronium ion concentration [H₃O⁺] and the hydroxide ion concentration, [OH^1-].

K_w = 1 x 10^-14 = [H₃O⁺] [OH^1-]

The pH scale is also limited by this relationship. pH is equal to the - log [H₃O⁺]. The pOH and the pK can also be determined by taking the - log of the [OH^1-] or the K value. This leads to the relationship in which

- log K_w = - log (1 x 10^-14) = -log ([H₃O⁺] [OH^1-]) = - log [H₃O⁺] - log[OH^1-]

which simplifies to

pK_w = 14 = pH + pOH

[H₃O⁺] = 10^-pH or = [OH^1-] = 10^-pOH

The pH scale basically cannot be lower than 0 or higher than 14, where pH from 0 to just below 7 has a greater [H₃O⁺] and a lower [OH^1-] making them Acids. pH 7 is neutral because the [H₃O⁺] is equal to [OH^1-], while pH above 7 is basic because the [OH^1-] is greater than the [H₃O⁺].

The key to understanding acid-base chemistry is the ability to recognize the type of compound you are dealing with. The categories include strong acids, strong bases, weak acids, weak bases and salts. Strong acids are also strong electrolytes that when dissolved in water ionize almost completely into the hydrogen ion, H⁺, and the anion. Strong acids include: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₃, HClO₄, and H₃PO₄.

For a solution of a strong acid such as HCl which almost totally ionizes in water, a 0.010 M solution contains 0.010 M H⁺. For this solution, the pH is approximately 2.0.

Strong Acids

For a strong acid, the K_a indicates that the reaction goes almost to completion; therefore, the hydronium ion concentration, [H₃O⁺] at equilibrium, is determined to be the initial concentration of the acid.

[HCl]_{initial concentration} = [H₃O⁺]_eq; pH = - log [H₃O⁺]_eq

There are a limited number of strong acids, HCl, HBr, HI, HClO₃, HClO₄, HNO₃, H₂SO₄

Weak Acids

For acids in which the dissociation in an aqueous solution is impeded, there are significantly less ions in solution. These acids are considered weak electrolytes or weak acids. When a weak acid dissolves in water the reaction is

HA + H₂O ⇄ H₃O⁺ + A^1-

where HA represents a weak acid with a K_a less than 10^-3 and A^1- is the anion of the acid or its conjugate base. According to the Bronsted-Lowery definition, acids donate hydrogen ions to another species, generally water, while bases accept hydrogen ions which results in the conjugate pair.

The pH of a solution can be calculated if the initial concentration and the K_a are known.

Strong Bases:

Strong bases generally contain the polyatomic ion, hydroxide, OH^1-. Common strong bases are NaOH, KOH, Ca(OH)₂. Like strong acids, strong bases are also strong electrolytes which ionize almost completely when dissolved in water. The hydroxide ion concentration is therefore determined from the concentration of the strong base:

[NaOH] = [OH^1-] and the

pOH = - log [OH^1-] and using the relationship of K_w

pH = 14 - pOH

Weak Bases:

When a weak base is dissolved in water the reaction is: B + H₂O ⇄ HB⁺ + OH^1-

A good example of a common weak base is NH₃, ammonia which has a K_b = 1.8 x 10^-5; however, true bases are not common, most bases are actually the conjugate base of a weak acid which is dissolved as a sodium or potassium salt, for example, NaHSO₄, sodium bisulfate where bisulfate ion is the conjugate base of sulfuric acid, H₂SO₄ or KCN, potassium cyanide, where CN^1- is the conjugate base of HCN, hydrocyanic acid. The conjugate base of an acid can be represented as the anion, A^1- of the weak acid HA.

The conjugate base of the acid, HA, when part of a salt can act as a base in water with the following reaction

A^1- + H₂O ⇄ HA + OH^1-

The equilibrium expression for the above reactions is:

There are very few reported K_b values because most of the known weak bases are actually conjugate bases of weak acids. The few exceptions are ammonia, organic amines and amides, which contain -NH_x groups.

Salts or ionic compounds:

Many ions react with water in what is called a hydrolysis reaction. The most commonly encountered compounds that undergo hydrolysis are the salts of weak acids such as sodium acetate. Sodium acetate completely ionizes when dissolved in water but then a small fraction of the acetate ions react with water as in the reaction below with A^- representing the acetate.

Salts or ionic compounds can be neutral, acidic or basic depending upon the ions present in the compound. An ionic compound or salt is composed of a cation and an anion. The cation, HB⁺ or Mⁿ⁺, is the conjugate acid of a base, while the anion is the conjugate base of an acid, A^1-. Therefore, a salt, always has two competing reactions occurring when the ions are dissolved in water. However, the relative strength of the ions determines whether the solution is basic, neutral or acidic. The strength of the conjugate is itself dependent on the strength of the weak acid or base where K_a x K_b = K_w. As K_a increase, K_b has to decrease, so as the weak acid becomes stronger, its corresponding conjugate base becomes weaker. The stronger the acid, the more likely it is to undergo the acid dissociation reaction, while the weaker an acid becomes, the less likely the reaction will occur.

Buffers

Buffers are useful because they can maintain a constant pH in the solution even when small amounts of acids or bases are added to the buffer solution. A buffer has to be able to react with added acid, therefore, it must contain a basic ion, but it must also contain an acidic ion to be able to reaction with added base. This can be achieved not by have a strong acid such as HCl with its extremely weak conjugate base, chloride ion, but with a weak acid such as acetic acid and its conjugate base, acetate, in the solution as a sodium salt. The weak acid can react with added base in the following reaction

HC₂H₃O₂ + OH^1- = H₂O + C₂H₃O₂^1-

producing a greater amount of conjugate base. While the conjugate base can react with added acid to produce the weak acid

C₂H₃O₂^1- + H⁺ = H₂O + HC₂H₃O₂

The weak acid reaction can be used to determine the pH of a buffer solution. Stoichiometry is needed to determine the initial quantities of the weak acid and conjugate base remaining once the added acid or base interacts to change the concentrations.

Procedure

Part III. pH of acids, bases and salts.

Use the videos and simulation below to record the pH of the 0.1 M acid, base or salt solutions.

IV. Buffers.

Preparation of Buffer solution

pH of water with added NaOH and HCl,

We will add 1 M NaOH and 1 M HCl to 25 mL of water by drops and measure the pH of the mixture, then look at how the pH changed with the addition of the NaOH versus the addition of the the HCl.

pH of buffer solution with added NaOH and HCl

We will add 1 M NaOH and 1 M HCl to 25 mL of an acetic buffer solution by 1 mL increments, measuring the pH. We will look at how the pH changes in the buffer solutions versus how it changed in the water.

Report and Conclusion Paragraph

Complete the data and result tables, graphs, calculations and answer all required questions for Parts I and II for this part of the combined experiment.

Complete a conclusion paragraph using the RERUNS method for the combined experiment 30/31/32. A discussion of how to write a conclusion paragraph is given in Appendix D: How to Write a Formal Laboratory Report