Hydroxide (OH–)
Table of Contents
The hydroxide ion is a negatively charged species consisting of one oxygen atom covalently bonded to one hydrogen atom. It is represented by the chemical formula OH– and carries a single negative charge. Hydroxide ions readily combine with positively charged cations to form ionic compounds, collectively known as hydroxides. Common examples include: [1-4]
- Sodium hydroxide (NaOH)
- Potassium hydroxide (KOH)
- Calcium hydroxide (Ca(OH)2)
- Magnesium hydroxide (Mg(OH)2)
Hydroxides are responsible for the basic or alkaline character of many solutions. For example, sodium hydroxide releases hydroxide ions in water, producing a strongly basic solution used in soap manufacturing, paper production, and water treatment.
Structure [4]
In the hydroxide ion (OH–), hydrogen and oxygen share a pair of electrons through a covalent bond. Because oxygen is more electronegative than hydrogen, it attracts the shared electrons more strongly, resulting in a highly polar bond. The extra electron that gives the ion its -1 charge is largely localized on the oxygen atom, which also possesses lone pairs of electrons. These features contribute to an uneven charge distribution and strong polarity.
Due to this polarity, hydroxide ions form strong hydrogen bonds with water molecules. As a result, they are highly soluble and extremely reactive in aqueous solutions. Their charge, lone pairs, and polarity collectively govern both their structure and chemical behavior.
The Lewis structure is shown below.
Physical Properties [1]
It is important to note that the hydroxide ion does not exist as a solid on its own. Instead, it forms solid metal hydroxides. Therefore, the following physical properties describe hydroxide compounds rather than the free ion.
| Property | Description |
|---|---|
| Molar Mass of OH– | 17.01 g mol-1 |
| State of Matter | Most metal hydroxides are solids at room temperature; alkali hydroxides are typically crystalline solids, while transition-metal hydroxides often occur as gelatinous precipitates. |
| Color | Usually white or colorless; some transition-metal hydroxides are colored (e.g., Cr(OH)3: green; Fe(OH)3: reddish-brown) |
| Solubility in Water | Alkali hydroxides (NaOH, KOH) are highly soluble; many transition-metal hydroxides are insoluble (e.g., Fe(OH)3) |
| Melting and Boiling Points | Alkali and alkaline-earth hydroxides generally have high melting and boiling points, though some decompose on heating (e.g. Ca(OH)2) |
Formation
1. Dissociation of Bases
Strong bases such as sodium hydroxide and potassium hydroxide completely dissociate in water to produce hydroxide ions:
i. NaOH (aq) → Na+ (aq) + OH– (aq)
ii. KOH (aq) → K+ (aq) + OH– (aq)
This increases the concentration of OH– ions, making the solution strongly basic. Weak bases like ammonia also generate hydroxide ions, but only partially and indirectly.
2. Reaction of Metals with Water
Reactive alkali metals such as sodium and potassium react vigorously with water to form metal hydroxides and hydrogen gas, thereby increasing alkalinity:
i. 2 Na (s) + 2 H2O (l) → 2 NaOH (aq) + H2 (g)
ii. 2 K (s) + 2 H2O (l) → 2 KOH (aq) + H2 (g)
Some alkaline-earth metals, such as calcium, react more moderately but still form hydroxide and hydrogen.
Chemical Reactions [2,3]
1. Precipitation with Metal Ions
When hydroxide ions are added to metal salt solutions, sparingly soluble metal hydroxides precipitate:
i. Cu2+ (aq) + 2 OH– (aq) → Cu(OH)2 (s)
ii. Fe3+ (aq) + 3 OH– (aq) → Fe(OH)3 (s)
2. Neutralization Reactions
In acid–base neutralization, hydroxide ions react with hydrogen ions to form water:
i. H+ (aq) + OH– (aq) → H2O (l)
This reaction is essential in controlling acidity.
3. Hydrolysis Reactions
i. Ester Hydrolysis (Saponification)
Hydroxide ions break ester bonds to produce a salt of the carboxylic acid and an alcohol:
CH3COOC2H5 (aq) + NaOH (aq) → CH3COONa (aq) + C2H5OH (aq)
This reaction forms the basis of soap manufacture.
ii. Amide Hydrolysis
Hydroxide ions convert amides into salts of the corresponding acids and ammonia (or amines):
CH3CONH2 (aq) + NaOH (aq) → CH3COONa (aq) + NH3 (g)
Redox and Complex Formation Reactions
In alkaline environments, hydroxide ions may support oxidation processes. For example:
i. 4 Fe(OH)2 (s) + O2 (g) + 2 H2O (l) → 4 Fe(OH)3 (s)
ii. 2 Mn(OH)2 (s) + O2 (g) → 2 MnO2 (s) + 2 H2O (l)
Additionally, some metal hydroxides dissolve in excess hydroxide to form complex ions:
i. Al(OH)3 (s) + OH– (aq) → [Al(OH)4]– (aq)
ii. Zn(OH)2 (s) + 2 OH– (aq) → [Zn(OH)4]2- (aq)
Applications [3]
- Manufacture of soaps, detergents, paper, and textiles
- Water treatment and pH control
- Laboratory acid-base titrations and analytical chemistry
- Industrial chemical synthesis and process control
- Food processing and refining
- Environmental management, including effluent and wastewater treatment
The hydroxide ion is essential because it plays a key role in many chemical processes that affect industry, laboratories, and the environment. Its widespread use and influence make it one of the most significant ions in chemistry.
