MYP Chemistry 5.6 Neutral Solutions Notes

What Does It Mean for a Solution to Be Neutral?

Pure Water and Neutral Solutions

In pure water at $25^{\circ} \mathrm{C}(298 \mathrm{~K})$:

1. A few water molecules ionise very slightly into hydrogen ions and hydroxide ions: $$\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \rightleftharpoons \mathrm{H}^{+}(\mathrm{aq})+\mathrm{OH}^{-}(\mathrm{aq})$$
2. At this temperature, in a neutral solution: $$\left[\mathrm{H}^{+}\right]=\left[\mathrm{OH}^{-}\right]=1.0 \times 10^{-7} \mathrm{~mol} \mathrm{dm}^{-3}$$
3. The product of these concentrations is constant at $25^{\circ} \mathrm{C}$: $$\left[\mathrm{H}^{+}\right]\left[\mathrm{OH}^{-}\right]=1.0 \times 10^{-14}$$

This constant is called $K_w$, the ionic product of water.

Water as Both Acid and Base: Amphiprotic

1. As discussed in the previous article, in the Brønsted–Lowry model:
   1. An acid is a proton (H⁺) donor.
   2. A base is a proton (H⁺) acceptor.
2. Water can do both, depending on what it is reacting with:
   1. With a strong base like NaOH, water can donate a proton (acting as an acid).
   2. With a strong acid like HCl, water can accept a proton (acting as a base).

How Can We Test if a Solution Is Neutral in the Lab?

We never taste solutions to check if they are acidic or basic. Instead, we can:

1. Use universal indicator
   1. Neutral solution → green (around pH 7).
   2. Acidic solution → red/orange/yellow.
   3. Basic solution → blue/purple.
2. Use litmus paper
   1. Blue litmus in neutral solution → stays blue.
   2. Red litmus in neutral solution → stays red.
   3. If both colors stay the same, the solution is close to neutral.
3. Use a pH meter / pH probe
   1. Gives a numerical pH value.
   2. A reading close to 7.0 indicates a neutral solution.

How Do Acids and Bases Interact to Form a Neutral Solution?

From our previous discussion and the Brønsted–Lowry definitions, you know that, in an aqueous solution:

1. An acid increases the concentration of H⁺(aq).
2. A base increases the concentration of OH⁻(aq).

What Is Neutralization?

1. At the ionic level, the key reaction is: $$\mathrm{H}^{+}(\mathrm{aq})+\mathrm{OH}^{-}(\mathrm{aq}) \rightarrow \mathrm{H}_2 \mathrm{O}$$
2. If the acid and base are mixed in exactly the right amounts (stoichiometric ratio), the resulting solution will be neutral or close to neutral, depending on the strength of the acid and base.

Neutralization as a Type of Reaction

1. When an acid reacts with a base in water, the general pattern is: $$\text { acid + base → salt + water }$$
2. Salt → ionic compound formed from the positive ion from the base and the negative ion from the acid.
3. Water → formed from H⁺ and OH⁻.
4. If a strong acid and strong base are used in correct proportions, the final solution is close to pH 7.

Why Is Near-Neutral pH Important in Real Life?

Neutralization reactions are extremely important in industry, the environment, and the human body.

Environmental Examples

Soil pH and Crop Growth

1. Soil pH affects:
   1. Water-holding capacity
   2. Nutrient availability (how easily plants can absorb nutrients)
   3. Microbial activity (which controls decomposition and nutrient cycling)
   4. Soil structure and root growth
2. Most crops grow best in slightly acidic to neutral soil (roughly pH 6–7.5).
   1. If soil is too acidic, plant growth and yield decrease.
   2. Farmers often spread basic substances (like powdered limestone, CaCO₃) to neutralise excess acids and raise soil pH.

Rivers, Lakes and Buffers

1. In rivers and lakes, substances like hydrogencarbonates of sodium, potassium, calcium and magnesium help neutralize acids:
   1. These hydrogencarbonates act as natural buffers, reacting with added acid and helping keep pH close to neutral.
2. If the amount of these buffering ions is reduced:
   1. The water becomes more vulnerable to acid inputs (e.g. acid rain, industrial discharge).
   2. pH can drop, harming fish, invertebrates and plant life.
3. Maintaining a near-neutral pH in rivers is essential for healthy aquatic ecosystems.

Ocean Acidification

1. The oceans absorb large amounts of carbon dioxide (CO₂) from the atmosphere: $$\mathrm{CO}_2+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{H}_2 \mathrm{CO}_3$$
2. This forms carbonic acid (H₂CO₃), which can release H⁺: $$\mathrm{H}_2 \mathrm{CO}_3 \rightleftharpoons \mathrm{H}^{+}+\mathrm{HCO}_3^{-}$$
   1. As more CO₂ is absorbed, more H⁺ is produced → pH falls.
   2. The ocean becomes slightly more acidic – this is called ocean acidification.
3. Marine organisms such as corals, shellfish and some plankton use calcium carbonate (CaCO₃) to build shells and skeletons.
   1. Calcium carbonate can neutralise acids: (simplified idea) $$\mathrm{CaCO}_3+\mathrm{H}^{+} \rightarrow \mathrm{Ca}^{2+}+\mathrm{HCO}_3^{-}$$
   2. If too much acid (e.g. carbonic acid) is present, more CaCO₃ dissolves.
   3. This weakens shells and skeletons and reduces the availability of carbonate for new shell formation.