PH
Understanding pH
pH is a scale used to specify the acidity or basicity of an aqueous solution. It is an essential concept in chemistry, particularly in the study of ionic equilibrium. The pH scale ranges from 0 to 14, with 7 being neutral. A pH less than 7 indicates acidity, while a pH greater than 7 indicates basicity.
Definition
The pH of a solution is defined as the negative logarithm (base 10) of the hydrogen ion concentration:
$$ pH = -\log[H^+] $$
where $[H^+]$ is the concentration of hydrogen ions in moles per liter (M).
Calculating pH
For a strong acid or base, which dissociates completely in water, the pH can be calculated directly from its concentration. For example, if the concentration of hydrochloric acid (HCl) is 0.01 M, the pH is calculated as follows:
$$ pH = -\log[0.01] = -(-2) = 2 $$
For weak acids and bases, which do not dissociate completely, the calculation is more complex and often involves the use of equilibrium constants.
The pH Scale
The pH scale is logarithmic, which means each whole pH value below 7 is ten times more acidic than the next higher value. For example, a pH of 3 is ten times more acidic than a pH of 4 and a hundred times more acidic than a pH of 5.
Here is a table summarizing the pH scale:
| pH Value | Acidity/Basicity | Example |
|---|---|---|
| < 7 | Acidic | Lemon juice |
| 7 | Neutral | Pure water |
| > 7 | Basic (alkaline) | Baking soda |
Importance of pH
pH is important in many areas, including:
- Biology: Enzyme activity and metabolic processes are pH-dependent.
- Medicine: Blood pH is tightly regulated, as deviations can be harmful.
- Agriculture: Soil pH affects nutrient availability and crop growth.
- Industry: pH control is crucial in many manufacturing processes.
Examples
Example 1: Strong Acid
Calculate the pH of a 0.005 M HCl solution.
$$ pH = -\log[0.005] = -\log[5 \times 10^{-3}] = -(-2.3) = 2.3 $$
Example 2: Weak Acid
Calculate the pH of a 0.1 M acetic acid (CH3COOH) solution, given that its dissociation constant (Ka) is $1.8 \times 10^{-5}$.
First, write the dissociation equation and the expression for Ka:
$$ CH3COOH \rightleftharpoons CH3COO^- + H^+ $$ $$ Ka = \frac{[CH3COO^-][H^+]}{[CH3COOH]} $$
Assuming $x$ is the concentration of $H^+$ that dissociates:
$$ Ka = \frac{x^2}{0.1 - x} \approx \frac{x^2}{0.1} $$
Solving for $x$:
$$ x^2 = Ka \times 0.1 $$ $$ x^2 = 1.8 \times 10^{-6} $$ $$ x = \sqrt{1.8 \times 10^{-6}} $$ $$ x = 1.34 \times 10^{-3} $$
Now, calculate the pH:
$$ pH = -\log[1.34 \times 10^{-3}] \approx 2.87 $$
Example 3: Basic Solution
Calculate the pH of a 0.01 M NaOH solution.
Since NaOH is a strong base, it dissociates completely:
$$ NaOH \rightarrow Na^+ + OH^- $$
The concentration of $OH^-$ is equal to the concentration of NaOH. To find the pH, we first calculate the pOH, which is the negative logarithm of the hydroxide ion concentration:
$$ pOH = -\log[0.01] = 2 $$
Since pH + pOH = 14 for aqueous solutions at 25°C, we can find the pH:
$$ pH = 14 - pOH = 14 - 2 = 12 $$
Conclusion
Understanding pH is crucial for studying chemical reactions, biological systems, and industrial processes. It helps us to predict the behavior of acids and bases in different environments and to maintain the necessary conditions for various applications.