The material known as “lime” is a fundamental chemical compound, not to be confused with the citrus fruit, and is one of the world’s most widely used industrial substances. In this context, lime refers to calcium oxide (\(text{CaO}\)) or calcium hydroxide (\(text{Ca}(text{OH})_2\)), derived from naturally occurring limestone. Used since ancient times in construction and agriculture, this inorganic material is essential in processes ranging from steel production and water purification to chemical manufacturing. The manufacturing process transforms this common rock into highly reactive chemical products through thermal chemistry.
Preparing the Limestone
The process begins with sourcing high-quality limestone, a sedimentary rock composed mainly of calcium carbonate (\(text{CaCO}_3\)). Limestone is typically extracted from quarries or mines, and its purity directly influences the quality of the final lime product. High-calcium limestone, containing over 90% \(text{CaCO}_3\), is preferred for producing high-grade lime.
The extracted stone is prepared for thermal transformation by mechanical crushing to reduce its size. This is followed by screening to achieve a uniform particle size, often specified between 10 and 50 millimeters. This specific sizing ensures that the subsequent heating process, known as calcination, occurs evenly, preventing a mix of under-burnt and over-burnt material that would lower product quality.
The Calcination Process
The core of lime manufacturing is calcination, a thermal decomposition reaction where prepared limestone is subjected to intense heat. This highly endothermic process requires substantial energy input to break the chemical bonds within the calcium carbonate molecule (\(text{CaCO}_3\)). The reaction drives off carbon dioxide (\(text{CO}_2\)), leaving behind calcium oxide (\(text{CaO}\)), and is represented by the formula: \(text{CaCO}_3\) + Heat \(rightarrow text{CaO} + text{CO}_2\).
The process temperature is precisely controlled, typically ranging between 900°C and 1200°C. Modern industrial production relies on specialized equipment, such as rotary or vertical shaft kilns, to facilitate this high-temperature reaction. Temperature control is critical: if the temperature is too low, the decomposition is incomplete, leading to under-burnt lime. If the temperature exceeds the optimal range, the calcium oxide becomes over-fired, a condition known as “dead-burned” lime, which is chemically unreactive.
Handling Quicklime
The direct product of calcination is quicklime, or calcium oxide (\(text{CaO}\)), a white, highly alkaline, and caustic material. Quicklime is known for its intense chemical reactivity, particularly its strong affinity for water, with which it reacts vigorously and exothermically.
Due to this high reactivity, quicklime is used extensively in industrial sectors such as steelmaking, where it acts as a flux to remove impurities. It is also employed in environmental applications, such as water treatment, to adjust pH levels and remove heavy metals. Quicklime must be kept in sealed environments to protect it from atmospheric moisture and carbon dioxide, which would diminish its purity and utility.
Creating Hydrated Lime
For many applications, quicklime is an intermediate step in creating hydrated lime, or calcium hydroxide (\(text{Ca}(text{OH})_2\)). This secondary manufacturing step is known as slaking, where quicklime is intentionally mixed with a controlled amount of water. The reaction is \(text{CaO} + text{H}_2text{O} rightarrow text{Ca}(text{OH})_2\) + Heat.
The slaking process is highly exothermic, releasing a significant amount of heat. Industrial slakers are engineered to manage this thermal energy, often maintaining the reaction temperature between 79°C and 91°C to ensure the resulting particles are highly reactive. The precise ratio of water to quicklime is managed to produce a dry, fine, white powder. This powdered hydrated lime is widely used in construction for mortars and plasters, and in agriculture to neutralize acidic soils.