Chemistry · Chemistry of the elements · Grade 8

Lime and classes of compounds

How do quicklime and slaked lime on a wall turn back into limestone?

The situation

In a village, to whitewash the walls of his house, a craftsman bought white lumps of quicklime at the bazaar. When he put the lime into a bucket of water, the bucket heated up and began to hiss; the wall was whitewashed with the resulting “milk of lime”, and within a few weeks it set hard and strong. Lime itself is made by burning limestone from the mountains in a kiln. Limestone → lime → milk of lime → limestone again: through which classes does the substance pass in this chain, and why does the wall harden in air?

Topic

Classes of inorganic compounds: genetic relationships using lime as the example. Burning limestone, slaking and whitewashing a wall: CaCO₃ → CaO → Ca(OH)₂ → CaCO₃ — the wall hardens with CO₂ from the air

Key ideas

Oxide, base, acid, salt
Genetic relationships
Amphoteric behaviour

Formulas

CaCO_{3} \overset{\overset{}{t}}{\to} CaO + CO_{2} ↑

Burning limestone: the temperature at which CO₂ escapes freely at 1 atm is 898 °C; ΔH ≈ +178 kJ/mol (heat is absorbed).

CaO + H_{2} O \overset{}{\to} Ca (OH)_{2}

Slaking: ΔH ≈ −65 kJ/mol — the mixture heats up.

Ca (OH)_{2} + CO_{2} \overset{}{\to} CaCO_{3} ↓ + H_{2} O

Limewater turning cloudy and a whitewashed wall setting.

Zn (OH)_{2} + 2 HCl \overset{}{\to} ZnCl_{2} + 2 H_{2} O Zn (OH)_{2} + 2 NaOH \overset{}{\to} Na_{2} [Zn (OH)_{4}]

An amphoteric hydroxide dissolves in both acid and alkali.

m (CaO) = m (C a C O_{3}) \cdot \frac{M (CaO)}{M (C a C O_{3})} = m \cdot \frac{56.08}{100.09}

Calculation from the reaction equation: 100 g of limestone gives 56 g of lime and 44 g of CO₂.

Theory

Main classes of inorganic compounds

Oxides consist of two elements, one of which is oxygen: basic (CaO, CuO — they correspond to bases), acidic (CO₂, P₂O₅ — they correspond to acids) and amphoteric (ZnO, Al₂O₃). Bases consist of a metal atom and hydroxide groups OH: those soluble in water are alkalis (NaOH, Ca(OH)₂), the insoluble ones are Cu(OH)₂, Mg(OH)₂. Acids consist of hydrogen atoms and an acid residue: oxygen-containing (H₂SO₄) and oxygen-free (HCl); basicity is the number of replaceable H atoms. Salts consist of a metal cation and an acid residue: normal (CaCO₃), acid (NaHCO₃) and basic ((CuOH)₂CO₃). Crystalline hydrates are salts with water in their crystals (gypsum, blue vitriol).

Genetic relationships

One class of substance can be made from another — this is the genetic relationship. The metal series: metal → basic oxide → base → salt ( $Ca \overset{}{\to} CaO \overset{}{\to} Ca (OH)_{2} \overset{}{\to} CaCO_{3}$ ). The non-metal series: non-metal → acidic oxide → acid → salt ( $C \overset{}{\to} CO_{2} \overset{}{\to} H_{2} CO_{3} \overset{}{\to} CaCO_{3}$ ). Both series meet at the salt: a base and an acid (or their oxides) form a salt. Lime is a natural link in this relationship: limestone is burnt to lime, lime is slaked to a base, and the base absorbs CO₂ from the air and returns to limestone. The active link is highlighted on the board.

Lime: burning, slaking, setting

When limestone (chalk and marble are also CaCO₃) is heated, it decomposes. This is a reversible process: at every temperature CO₂ has an equilibrium pressure — ≈ 24 kPa at 800 °C, 101 kPa (1 atm) at 898 °C. Above that the gas escapes freely, so lime kilns are kept at about 900–1000 °C; at 700 °C the pressure is only ≈ 3 kPa and decomposition is very slow. Slaking is strongly exothermic: 1 mol of CaO releases 65 kJ of heat. Ca(OH)₂ is sparingly soluble in water (1.73 g/L at 20 °C): the excess floats as “milk of lime”, and the clear filtrate is limewater (pH ≈ 12.4). Lime spread on a wall absorbs CO₂ from the air and slowly turns into CaCO₃, which is how it sets; you can test it with phenolphthalein: a fresh layer turns pink, a set one stays colourless. If CO₂ is in excess, the precipitate turns into soluble Ca(HCO₃)₂ — this is the origin of the hardness of natural waters.

Amphoterism (lab 1)

An amphoteric hydroxide shows the properties of both a base and an acid. If NaOH is added dropwise to a ZnSO₄ solution in a test tube, a white gelatinous precipitate of Zn(OH)₂ forms: 2 mol of OH⁻ per 1 mol of Zn²⁺. The precipitate is split in two: to one part HCl is added and it dissolves ( $Zn (OH)_{2} + 2 HCl \overset{}{\to} ZnCl_{2} + 2 H_{2} O$ ); the other dissolves in excess NaOH too, forming sodium tetrahydroxozincate ( $Zn (OH)_{2} + 2 NaOH \overset{}{\to} Na_{2} [Zn (OH)_{4}]$ ), which uses up another 2 mol of OH⁻. Al(OH)₃ is amphoteric as well (it precipitates at 3 mol of OH⁻ and dissolves at 4 mol). Mg(OH)₂ and Cu(OH)₂ are basic: they do not dissolve in excess alkali. The graph shows the amount of precipitate against the OH⁻ added: a “triangle” for zinc, a “staircase” for magnesium.

Calcium and magnesium (Grade 9)

Both are in Group 2 (IIA) — alkaline-earth metals, with 2 electrons in the outer shell and +2 in compounds. Calcium is more reactive: it reacts vigorously with cold water ( $Ca + 2 H_{2} O \overset{}{\to} Ca (OH)_{2} + H_{2} ↑$ ), whereas magnesium hardly reacts with cold water and releases hydrogen slowly in hot water. Magnesium burns in air with a dazzling white flame: $2 Mg + O_{2} \overset{}{\to} 2 MgO$ — the mass of the oxide formed is 1.66 times that of the metal (oxygen is added). When gypsum $CaSO_{4} \cdot 2 H_{2} O$ is heated to about 150 °C, it loses three quarters of its water and becomes plaster of Paris $CaSO_{4} \cdot \frac{1}{2} H_{2} O$ (mass falls by 15.7 %); when plaster of Paris is mixed with water it turns back into gypsum and sets quickly — this is the basis of plaster casts, sculptures and ganch ornaments. Ca²⁺ and Mg²⁺ ions cause the hardness of water (details in the “Water purification” experiment).

Safety and the model

Safety: quicklime and alkalis burn the skin and eyes — wear safety goggles and gloves; if any gets into the eyes, rinse with plenty of water and see a doctor. Pour acid into water, never the other way round. Hold a hot crucible or test tube with tongs, and do not point the mouth of a test tube at yourself or others. Do not look directly at a burning magnesium flame. Model: the burning rate is simplified — it is taken to be proportional to the difference between the equilibrium pressure of CO₂ and the pressure of CO₂ in the air, and a larger sample is taken to heat up more slowly; time is sped up for the kiln and for gypsum (1 s ≈ 1 min), while the balance reading follows exact stoichiometry. CO₂ passed through limewater is taken to be absorbed completely. The speed of the reaction of Ca and Mg with water is shown qualitatively, while the volume of H₂ released follows the equation (STP). The thermochemical values come from the enthalpies of formation in the CRC Handbook.

Did you know?

Carved ganch ornaments in Uzbekistan are made from a mixture of gypsum and clay: the gypsum in burnt ganch, like plaster of Paris, sets when mixed with water. The craftsman applies two layers to the wall — the lower sets quickly, the upper stays soft longer, and the pattern is carved into it.
At the end of the 19th century photographers took pictures in a dark room by burning magnesium powder: a magnesium flame gives off very bright white light — that is how the first camera flashes worked.

Inside the experiment

The experiment starts from a real-life situation: you predict the result, change the controls, run the experiment and draw a conclusion. Inside are missions, theory, a quiz and a guided tour in which the camera moves and explains the key points.

More experiments in this section

Air: composition and pollution — Grade 7
Oxygen and ozone — Grade 7
Hydrogen — the lightest gas — Grade 7
Water: composition and properties — Grade 7
Mineral resources — Grade 7
Non-metals and halogens — Grade 8
Chlorine and hydrochloric acid — Grade 8
The oxygen group and sulfur — Grade 8
Sulfuric acid and the contact process — Grade 8
Nitrogen, ammonia and nitric acid — Grade 8