While common table salt is sodium chloride, the chemical definition of a salt encompasses a vast class of compounds. These compounds are found everywhere, from the minerals in the Earth’s crust to the essential electrolytes within the human body. Understanding salts is necessary for grasping many basic principles of chemistry and their practical applications.
Defining Salts as Ionic Compounds
Salts are ionic compounds built from positively and negatively charged particles called ions. The positive ion, known as the cation, is typically derived from a metal or a polyatomic ion such as ammonium (\(\text{NH}_4^+\)). The negative ion, or anion, generally comes from a non-metal or a polyatomic group like sulfate (\(\text{SO}_4^{2-}\)). These oppositely charged ions are held together by strong electrostatic forces, known as ionic bonds, forming a crystal lattice structure.
The compound maintains electrical neutrality because the total positive charge from the cations balances the total negative charge from the anions. This balance is why, for instance, magnesium chloride (\(\text{MgCl}_2\)) requires two singly-charged chloride anions (\(\text{Cl}^-\)) to neutralize the double-positive charge of one magnesium cation (\(\text{Mg}^{2+}\)). This ionic composition gives salts characteristic properties, such as high melting points and the ability to conduct electricity when dissolved in water or melted, as the ions become mobile.
How Salts Are Formed
The most common method for forming a salt is through a neutralization reaction, which involves mixing an acid and a base. In this process, the acid contributes hydrogen ions (\(\text{H}^+\)) while the base contributes hydroxide ions (\(\text{OH}^-\)). These two ions react to form a neutral water molecule (\(\text{H}_2\text{O}\)).
The remaining ions combine to form the new salt compound. For example, when hydrochloric acid (\(\text{HCl}\)) reacts with sodium hydroxide (\(\text{NaOH}\)), the \(\text{H}^+\) and \(\text{OH}^-\) form water, leaving the sodium cation (\(\text{Na}^+\)) and the chloride anion (\(\text{Cl}^-\)) to bond, resulting in the salt sodium chloride (\(\text{NaCl}\)). This reaction is often summarized by the general formula: Acid + Base \(\rightarrow\) Salt + Water.
Classifying Salts Based on pH
Not all salts produce a neutral solution when dissolved in water, a phenomenon explained by a reaction called hydrolysis. The strength of the acid and base that originally formed the salt determines whether the resulting solution will be acidic, basic, or neutral. Ions from strong acids or bases are spectators and do not react with water, but ions from weak acids or bases will interact with water molecules. This interaction can alter the hydrogen or hydroxide ion concentration in the solution, shifting the pH.
A salt formed from a strong acid and a strong base, like \(\text{NaCl}\) (from \(\text{HCl}\) and \(\text{NaOH}\)), will result in a neutral solution with a \(\text{pH}\) of 7, since neither ion hydrolyzes water. Conversely, a salt from a strong acid and a weak base, such as ammonium chloride (\(\text{NH}_4\text{Cl}\)), creates an acidic solution because the cation reacts with water to produce \(\text{H}^+\) ions, lowering the \(\text{pH}\) below 7. When a weak acid and a strong base react, the resulting salt, like sodium bicarbonate (\(\text{NaHCO}_3\)), is basic because the anion reacts with water to generate \(\text{OH}^-\) ions, pushing the \(\text{pH}\) above 7.
Practical Examples of Salts
The diverse applications of salts extend far beyond the kitchen, demonstrating their varying chemical properties in everyday life and industry. Magnesium sulfate, commonly known as Epsom salt (\(\text{MgSO}_4 \cdot 7\text{H}_2\text{O}\)), is a hydrated salt used widely for therapeutic purposes in bathwater and is also a component in agricultural fertilizers. The compound acts as a source of both magnesium and sulfur, which are important plant nutrients.
Sodium Bicarbonate
Sodium bicarbonate (\(\text{NaHCO}_3\)), or baking soda, is a slightly basic salt often used as a leavening agent in baking because it releases carbon dioxide gas when heated or mixed with an acid. It is also used as a household deodorizer and an antacid to neutralize stomach acid.
Potassium Nitrate
Potassium nitrate (\(\text{KNO}_3\)) is primarily valued for its use in fertilizers, providing the necessary potassium and nitrogen elements for plant growth. It also has a historical role in the production of gunpowder and continues to be used in various industrial applications.