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Chemistry of Calcium

Calcium is the 20th element in the periodic table. It is a group 2 metal, also known as an alkaline-earth metal, and no populated d-orbital electrons. Calcium is the fifth most abundant element by mass (3.4%) in both the Earth's crust and in seawater. All living organisms require calcium for survival. Calcium is a silver-gray metal which takes its name from the Latin word calx, which means lime. It is the fifth most abundant element in the earth's crust and is widely distributed as limestone (CaCO3), quicklime (CaO) and calcium fluoride.

General Properties of Calcium

Symbol Ca
Color dull gray or silver
Atomic Number 20
Category alkaline earth metal
Atomic Weight 40.078 g•mol−1
Group,Period,Block 2,4,s
Electron Configuration [Ar]4s2
Valence Electrons 2
Phase (room temperature) solid
Melting Point 1115 K, 842°C
Boiling Point 1757 K, 1484 °C
Atomic Radius 197 pm
Oxidation States 2
Density at room temp 1.55 g•cm−3
Electronegativity 1.00 (Pauling)
First ionization energy 589.8 kJ•mol−1
Number of stable isotopes 4 (2 more are fairly stable)

Electronic Structure


Fig. 1: Calcium Atom

Discovery and Properties of Calcium

 In 1808, British chemist Sir Humphry Davy first isolated elemental calcium using electrolysis. Calcium is the lightest of all metals, with a density 1.55 g/cm3. It reacts with both air and water, usually in reactions involving calcium carbonate, but this reaction is quite slow because calcium hydroxide, Ca(OH)2, is not very soluble in water. 


Reaction of Calcium with Halides 

Calcium forms salts with halides, such as \(CaCl_2\) or \(CaF_2\). They have a variety of uses, but the most usage most familiar to chemistry students is the use of calcium chloride as a desiccant (drying agent). 

\[ CaCl_2+ 2 H_2O(l) \rightarrow CaCl_2 \cdot 2H_2O \]

Reaction of Calcium with Carbonates 

Calcium carbonate is important in the formation of cave stalactites and stalagmites. This reaction allows calcium carbonate to be dissolved into solution as calcium bicarbonate.

\[ CaCO_3(s)+ CO_2(l)+ H_2O(l) \rightarrow Ca(HCO_3)_2 \]

The reverse reaction then allows the solution to become solid calcium carbonate once again, forming spikes of limestone in caves as the calcium bicarbonate solution drips vertically for several millennia. 

Reaction of Calcium with Water

Calcium metal is fairly reactive and combines with water at room temperature of produce hydrogen gas and calcium hydroxide

\[ Ca(s) + 2H_2O(g) \rightarrow Ca(OH)_2(aq) + H_2(g)\]

Product will reveal hydrogen bubbles on calcium metal's surface.

Reaction of Calcium with Acid

Calcium dissolves in acid to form dissociated ions of Ca and Cl along with Hydrogen gas.

\[ Ca(s) + 2HCl(aq) \rightarrow Ca^{2+}(aq) + 2Cl^-(aq) + H_2 (g)\]

Reaction of Calcium with Oxygen

Calcium metal slowly oxidizes in air, becoming encrusted with white \(CaO\) and \(CaCO_3\), which protect from attack by air. When ignited, Calcium burns to give calcium xxide.

\[ 2Ca \,(s)+ O_2 \,(g)\rightarrow 2CaO\,(s) \]

Calcium in living organisms

Because calcium is essential for life, it can be found in all organisms, living or dead. Shells of aquatic organisms, snail shells, and egg shells are all composed of mostly calcium carbonate, which can be dissolved in acid. Besides skeletal functions, the Ca2+ ion in animals and many organisms also plays an essential role in signal transduction pathways, neurotransmission, muscle function, fertilization, and enzymatic function. In plants, calcium is also important in the cell wall, membrane, and vacuole.

One of the most important calcium deposits is in coral reefs, which are comprised of mostly calcium carbonate. Coral secrete calcium carbonate over the period of their life, then die to allow new coral to build on top of their calcium carbonate structure. Over massive amounts of time, these calcium deposits grow into gigantic reefs, some of which can be seen from space (like the Great Barrier Reef in Australia). With the waters rich in sunlight and minerals like calcium, photosynthesis in sea plants is highly favored, allowing fish and other marine life to flourish in these regions.

Human bones are made up of mostly calcium phosphate (\(Ca_3(PO_4)_2\)). Cow milk also contains a large amount of calcium phosphate, which is why human culture encourages children and those particularly susceptible to osteoporosis to drink milk.

Uses of Calcium

The first known uses of calcium were by the Romans in the first century to make calcium oxide. Other written documentation around 975 C.E. suggests that plaster of Paris and Calcium Sulfate were medically useful. Calcium is available in a wide variety of forms, from limestone and chalk (calcium carbonate) to marble (calcite) and pearls. It also has many mineral forms (see link provided in Outside Links section). Calcium carbonate, in moderate amounts, can also be used as an antacid or a calcium supplement. Calcium nitrate is also a common fertilizer. Pure \(CaCo_3\) can be extracted from limestone in a series of three reactions:

  • Calcination:   \[CaCO_3 \rightarrow CaO + CO_2\]
  • Slaking:   \[CaO + H_2O \rightarrow Ca(OH)_2 \]
  • Carbonation:  \[Ca(OH)_2 + CO_2 → CaCO_3 + H_2O\]


Fig. 2: A Blue Starfish resting on hard Acropora coral. Lighthouse, Ribbon Reefs, Great Barrier Reef. Image used with permission from Wikipedia.

Hard water  

Hard water, as opposed to soft water, has a high mineral content of calcium sulfate (\(CaSO_4\)) or calcium carbonate (\(CaCO_3\)). It also includes magnesium ions (Mg2+) and sometimes iron, aluminum, and manganese. When left to evaporate, white calcium minerals can be seen on sinks, showers, etc.

Calcium oxide (quicklime) 

Calcium oxide, often referred to as quicklime, has many commercial functions, some of which include making mortar and pottery, food, construction, agriculture, pollution control, and in medicine. It also heats quite readily with water in this reaction: 

\[CaO  + H_2O \rightarrow Ca(OH)_2 \;\;\;\; \Delta H = −63.7\; kJ/mol\]

Outside links


  1.  Pettrucci, Ralph H. General Chemistry: Principles and Modern Applications. 9th. Upper Saddle River: Pearson Prentice  Hall, 2007.
  2.  Van Dyke, Fred. Conservation Biology: Foundations, Concepts, Applications. 2nd. McGraw-Hill Publishing Co. 2008.
  3. Allotropic Modification of Calcium. J. F. Smith, O. N. Carlson, and R. W. Vest, J. Electrochem. Soc. 103, 409 (1956), DOI:10.1149/1.2430364
  4. Davy, Humphry (1808). "On some new Phenomena … of the fixed Alkalies…". Philosophical Transactions of the Royal Society of London 98: 1–45.
  5. Fig. 1:
  6. Fig. 2:
  7. Fig. 3:


  1. Chalk (calcium carbonate) is a common compound of calcium. When an acid such as acetic acid is added to chalk, carbon dioxide is formed. Write and balance this reaction. What are the products of this reaction? 
  2. How can pure calcium carbonate be produced? What are the reactions called? 
  3. Which reaction is responsible for the formation of stalactites and stalagmites? Under what conditions? 
  4. Which compound of calcium is found in cow's milk, and how does this relate to overall bone health? 
  5. A sample of hard water has 156.9ppm Ca2+. An ion exchange column removes all Ca2+ and replaces it with H3O+. What is the final pH of the water as it leaves the column, assuming it began at pH of 7 and no other ions were produced? 


1) 2CaCo3 + HC2H3O2  Ca(C2H3O2)2 + CO2 + H2O
     Calcium acetate, c
arbon dioxide, and water. 

2) Purification of limestone through calcination, slaking, and carbonation.
CaCO3  CaO + CO2 

CaO + H2O Ca(OH)2 

Ca(OH)2 + CO2  CaCO3 + H2

3) CaCO3 + CO2 + H2O Ca(HCO3)

          Under acidic conditions. 

4) Calcium phosphate (Ca3(PO4)2). Bones are made from a mineral of this called hydroxylapatite with the formula Ca10(PO4)6(OH)2. More calcium phosphate means more mineral for bone growth, ensuring that enough calcium is available for the body to both make bones and have enough Ca2+ ions for other important signaling processes. 

5) 156.9 mg Ca2+  x   (1 mol Ca2+)/(40080 mg Ca2+)   x   (2 mol H3O+)/(1 mol Ca2+)= 7.83 E-3 M H3O+      

          -log(7.83 E-3)= pH 2.11 


  • Anna Zhu (UCD)