Question

For solutions of the same concentration, as acid strength increases, indicate what happens to each of the following (increases, decreases, or doesn't change). $$ \begin{array}{ll}{\text { a. }\left[\mathrm{H}^{+}\right]} & {\text { d. pOH }} \\ {\text { b. pH }} & {\text { e. } K_{\mathrm{a}}} \\ {\text { c. }\left[\mathrm{OH}^{-}\right]}\end{array} $$

Short answer

As acid strength increases in solutions of the same concentration, the following changes are observed: a. \([H^+]\) ion concentration increases. b. pH decreases. c. \([OH^-]\) ion concentration decreases. d. pOH increases. e. \(K_a\) (Acid dissociation constant) increases.

Step-by-step solution

a. H+ ion concentration

When the acid strength increases, the concentration of H+ ions will also increase as the acid becomes more effective at releasing its protons. So, H+ ion concentration increases.

b. pH

The pH scale is defined as pH = -log10[H+]. Since the H+ ion concentration increases as the acid strength increases, the pH of the solution will become lower (more negative logarithm). Therefore, the pH decreases.

c. OH- ion concentration

As the concentration of H+ ions increases, the concentration of OH- ions decreases to maintain the balance and keep the product of [H+] and [OH-] equal to the ion product of water (Kw = 1.0 × 10^(-14) at 25°C). So, OH- ion concentration decreases.

d. pOH

pOH is defined as pOH = -log10[OH-]. Since the OH- ion concentration decreases as the acid strength increases, the pOH of the solution will become higher (more negative logarithm). Therefore, pOH increases.

e. Ka (Acid dissociation constant)

Ka is defined as the ratio of the concentration of the dissociated form of an acid to the concentration of the undissociated form at equilibrium: Ka = ([H+][A-]) / [HA]. As the acid strength increases, more molecules dissociate into H+ ions and A- ions, and the Ka value representing the degree of ionization also increases. So, Ka increases.

Key concepts

pH Scale

The pH scale is an integral part of understanding acid strength and its effects in solutions. It is used to measure the acidity or basicity of a solution. The scale ranges from 0 to 14, where a pH less than 7 indicates an acidic solution, a pH of 7 indicates a neutral solution, and a pH greater than 7 indicates a basic solution.

The formula to calculate pH is given by:

• \( \text{pH} = -\log_{10}[\text{H}^+] \)

As acid strength increases, more hydrogen ions \([\text{H}^+]\) are released into the solution, which leads to a decrease in pH. A lower pH indicates a higher concentration of hydrogen ions and a stronger acid. The use of the logarithm in the pH scale means that each whole number on the scale represents a tenfold difference in ion concentration. This scale is crucial for chemists and biologists in assessing the acidity in various environments and reactions.

Acid Dissociation Constant (Ka)

The Acid Dissociation Constant, represented by \( K_a \), is a crucial concept to understand when assessing the strength of an acid. It provides a quantitative measure of the extent of ionization of an acid in a solution, which is essentially how well an acid donates protons to the solution.

The formula for \( K_a \) is:

• \( K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]} \)

In this expression, \([\text{HA}]\) is the concentration of the undissociated acid, while \([\text{H}^+]\) and \([\text{A}^-] \) are the concentrations of the dissociated ions. The higher the value of \( K_a \), the stronger the acid, as it indicates a greater degree of ionization.

Increasing the acid strength increases the value of \( K_a \), reflecting an enhanced ability to donate protons. Understanding \( K_a \) helps in predicting the strength of acids in various chemical reactions, and is important for fields like organic chemistry and biochemistry.

Ion Concentration

Ion concentration plays an essential role in the chemistry of acids and bases. For an acidic solution, it generally refers to the concentration of hydrogen ions \([\text{H}^+]\).

When an acid is dissolved in water, it dissociates to release \([\text{H}^+]\) ions. The concentration of these ions is directly tied to the strength of the acid. A stronger acid releases more \([\text{H}^+]\) ions, resulting in a higher ion concentration.

It's important to balance this with the concentration of hydroxide ions \([\text{OH}^-]\) in the solution:

• The product of \([\text{H}^+]\) and \([\text{OH}^-]\) must equal the ion product of water \( (K_w = 1.0 \times 10^{-14} \) at 25°C).

This means that an increase in \([\text{H}^+]\) leads to a decrease in \([\text{OH}^-]\). This relationship is crucial in maintaining the chemical equilibrium in solutions.

pOH

pOH is another scale used to measure the concentration of hydroxide ions \([\text{OH}^-]\) in a solution. It is closely related to the pH scale and is particularly handy when dealing with bases or when specific hydroxide ion concentration details are needed.

The formula for pOH is:

• \( \text{pOH} = -\log_{10}[\text{OH}^-] \)

In acidic solutions, as the concentration of \([\text{H}^+]\) increases due to stronger acid, \([\text{OH}^-]\) will decrease due to the ion product relationship with water. Because of this relationship, if \([\text{OH}^-]\) goes down, pOH will go up, indicating lesser basicity. The sum of pH and pOH is always equal to 14 at 25°C.

Understanding pOH is vital for figuring out the basicity of a solution and gives a complete picture when combined with pH, especially in chemical titration and buffer solutions.