Food grade sodium bicarbonate
| IUPAC name
Sodium hydrogen carbonate
| Other names
Bicarbonate of soda
|Molar mass||84.0066 g/mol|
|Density|| 2.20 g/cm3 (solid)|
1.1-1.3 g/cm3 (powder)
|Melting point||50 °C (122 °F; 323 K) (decomposes to Na2CO3)|
| 6.9 g/100 ml (0 °C)|
8.2 g/100 ml (10 °C)
9.6 g/100 ml (20 °C)
10 g/100 ml (25 °C)
11.1 g/100 ml (30 °C)
12.7 g/100 ml (40 °C)
14.5 g/100 ml (50 °C)
16.5 g/100 ml (60 °C)
19.5 g/100 ml (80 °C)
23.6 g/100 ml (100 °C)
|Solubility||Insoluble in ethanol, halocarbons, hydrocarbons|
|Solubility in acetone||0.02 wt% (22 °C)|
|Solubility in methanol||2.13 wt% (22 °C)|
|Vapor pressure||~0 mmHg|
Std enthalpy of
|Safety data sheet||Sigma-Aldrich|
|Lethal dose or concentration (LD, LC):|
LD50 (Median dose)
|4,220 mg/kg (rat, oral)|
| Potassium bicarbonate|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Sodium bicarbonate, sodium hydrogen carbonate, bicarbonate of soda, Bicarb, or baking soda, is a commonly used reagent with the chemical formula NaHCO3. Generally appearing as a fine white powder or fine crystals, sodium bicarbonate is an amphoteric compound, fulfilling roles as both an acid and a base in different chemical reactions, though it is more often thought of as a base.
It exists in nature as the rare mineral nahcolite (thermokalite).
Sodium bicarbonate will react with acids to release carbon dioxide.
- NaHCO3 + HX → H2O + CO2 + NaX
Sodium bicarbonate reacts with bases such as sodium hydroxide to form carbonates:
- NaHCO3 + NaOH → Na2CO3 + H2O
This property allows sodium bicarbonate to be useful as a safe neutralizing agent for both acids and bases.
Sodium bicarbonate slowly decomposes to form sodium carbonate above 50 °C, though the rate of decomposition is much higher at or above 200 °C.
- 2 NaHCO3 → Na2CO3 + H2O + CO2
Sodium bicarbonate forms white crystals and is soluble in water, but insoluble in organic solvents. It has a slightly salty taste, similar to that of sodium carbonate.
Relatively pure sodium bicarbonate can be purchased in supermarkets, cheaply and in large quantities, as baking soda. It is essential to buy unscented baking soda, as scented baking soda may form side products with other reactants or give a nauseating odor in many reactions.
Sodium bicarbonate can be prepared by bubbling carbon dioxide in a concentrated solution of sodium carbonate. Cool the solution until sodium bicarbonate precipitates, then filter and air dry the bicarbonate precipitate. Do not boil the bicarbonate solution, as sodium bicarbonate decomposes above 50 °C.
Sodium bicarbonate is one of the most versatile and commonly encountered compounds in amateur chemistry. Neutralization of sodium bicarbonate with acids yields a sodium salt and large amounts of carbon dioxide; reaction vessels should have ample headroom to prevent spills. Due to its ease of access and relative cheapness, sodium bicarbonate can be used to neutralize and clean acid spills as well; it is convenient, because the reaction of neutralization in its case is endothermic and results in cooling, unlike with hydroxides, which give very exothermic reactions that may lead to boiling and sprinkling of acids. As many carbonates are not water-soluble, sodium bicarbonate can be used in double replacement reactions to produce the carbonates of many metals.
Sodium bicarbonate, like sodium carbonate, can be used in acid-base titrations.
Sodium bicarbonate is non-toxic, but if swallowed it will neutralize the gastric acid and release carbon dioxide, causing burps or if too much is ingested, vomit. It will also increase the sodium level in blood if large quantities are consumed.
Sodium bicarbonate should be stored in closed bottles away from any acidic vapors.
Sodium bicarbonate can be poured down the drain or dumped in trash. Waste sodium bicarbonate can also be used as neutralizing agent for acids and other corrosive solutions.