Explain why (a) the pH decreases when lactic acid is added to a sodium lactate solution. (b) the \(\mathrm{pH}\) of \(0.1 \mathrm{M} \mathrm{NH}_{3}\) is less than \(13.0\). (c) a buffer resists changes in pH caused by the addition of \(\mathrm{H}^{+}\) or \(\mathrm{OH}^{-} .\) (d) a solution with a low \(\mathrm{pH}\) is not necessarily a strong acid solution.

Understanding the pH level is central to grasping acid-base reactions. The pH is a scale used to specify how acidic or basic (alkaline) an aqueous solution is. It's defined as the negative logarithm of the hydrogen ion concentration, written as \(\text{pH} = -\log[H^+]\). The scale ranges from 0 to 14, with 7 being neutral. Values below 7 indicate acidity, while values above 7 indicate alkalinity.

When we talk about the pH changing, like in the case of lactic acid added to a sodium lactate solution, we're discussing an increase in the hydrogen ion concentration. This higher concentration translates directly to a lower pH value. Similarly, when an ammonia solution doesn't reach a pH of 13, it's because the weak base only partially ionizes in water, producing fewer hydroxide ions (OH-) and therefore having a lesser effect on the pH. Understanding the relationship between ion concentration and pH is key to predicting the behavior of solutions in acid-base reactions.

A buffer solution is a special chemical solution that resists changes to its pH when acids or bases are added or when dilution occurs. It's usually made up of a weak acid and its conjugate base, or vice versa. The principle behind its operation is beautifully illustrated when explaining why a buffer resists change upon the addition of H+ or OH-.

According to Le Chatelier's Principle, any change in the conditions of a reaction in equilibrium causes the system to adjust to counteract the change. In the case of buffers, when extra hydrogen ions (H+) are introduced, they're consumed by the conjugate base in the buffer. This minimizes the increase in the hydrogen ion concentration, thereby preventing a substantial change in pH. Conversely, the addition of hydroxide ions (OH-) leads to their combination with H+ to form water, with the buffer producing more H+ to maintain equilibrium. This dynamic equilibrium is what enables buffers to maintain a stable pH in biological systems and industrial processes.

Acids are substances that can donate hydrogen ions (H+), and their strength is defined by their ability to do so. A strong acid, such as hydrochloric acid (HCl), dissociates completely in solution, releasing a large number of hydrogen ions and thus has a high acid strength. This complete dissociation leads to a low pH, indicating a very acidic solution.

On the other hand, weak acids like acetic acid (CH3COOH) only partially dissociate in water, resulting in a higher pH compared to a strong acid at the same concentration. This partial dissociation is why explaining that a solution with a low pH is not always a strong acid solution is important. A solution can have a low pH due to the high concentration of a weak acid, not necessarily the presence of a strong acid. Moreover, this concept clarifies why the pH of an acid solution isn't solely dependent on acid strength but also on the concentration of the acid in the solution.

Le Chatelier's Principle provides a predictive tool in chemical reactions, including those involving acids and bases. It states that if a dynamic equilibrium is disturbed by changing the conditions (concentration, pressure, temperature), the position of equilibrium moves to counteract the change.

Applying this principle to acid-base equilibrium, when an acid (like the added lactic acid) is introduced to a solution in equilibrium (such as the sodium lactate solution), the equilibrium will shift to reduce the effect of the added acid. In this scenario, the system shifts to consume the extra H+ ions, leading to the formation of more weak acid molecules. By the same token, the addition of a base to a solution results in the equilibrium shifting to counteract the increase in OH- ions, typically by generating more H+ ions. Le Chatelier's Principle is crucial for understanding how buffer solutions work, how changing conditions affect chemical equilibria, and how systems respond to maintain balance.