A | B | C | D | E | F | G | H | CH | I | J | K | L | M | N | O | P | Q | R | S | T | U | V | W | X | Y | Z | 0 | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9
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| Names | |
|---|---|
| IUPAC name
Sodium hydroxide[3]
| |
| Other names | |
| Identifiers | |
3D model (JSmol)
|
|
| ChEBI | |
| ChemSpider | |
| ECHA InfoCard | 100.013.805 |
| EC Number |
|
| E number | E524 (acidity regulators, ...) |
| 68430 | |
| KEGG | |
| MeSH | Sodium+Hydroxide |
PubChem CID
|
|
| RTECS number |
|
| UNII | |
| UN number | 1823 (solid) 1824 (solution) |
CompTox Dashboard (EPA)
|
|
| |
| |
| Properties | |
| NaOH | |
| Molar mass | 39.9971 g/mol |
| Appearance | White, opaque crystals |
| Odor | odorless |
| Density | 2.13 g/cm3[4] |
| Melting point | 323 °C (613 °F; 596 K)[4] |
| Boiling point | 1,388 °C (2,530 °F; 1,661 K)[4] |
| 418 g/L (0 °C) 1000 g/L (25 °C)[4] 3370 g/L (100 °C) | |
| Solubility | soluble in glycerol, negligible in ammonia, insoluble in ether, slowly soluble in propylene glycol |
| Solubility in methanol | 238 g/L |
| Solubility in ethanol | <<139 g/L |
| Vapor pressure | <2.4 kPa (20 °C) 0.1 kPa (700 °C) |
| Acidity (pKa) | 15.7 |
| −15.8·10−6 cm3/mol (aq.)[5] | |
Refractive index (nD)
|
1.3576 |
| Structure[6] | |
| Orthorhombic, oS8 | |
| Cmcm, No. 63 | |
a = 0.34013 nm, b = 1.1378 nm, c = 0.33984 nm
| |
Formula units (Z)
|
4 |
| Thermochemistry[7] | |
Heat capacity (C)
|
59.5 J/(mol·K) |
Std molar
entropy (S⦵298) |
64.4 J/(mol·K) |
Std enthalpy of
formation (ΔfH⦵298) |
−425.8 kJ/mol |
Gibbs free energy (ΔfG⦵)
|
-379.7 kJ/mol |
| Hazards | |
| GHS labelling: | |
 
| |
| Danger | |
| H290, H302, H314 | |
| P280, P305+P351+P338, P310 | |
| NFPA 704 (fire diamond) | |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose)
|
40 mg/kg (mouse, intraperitoneal)[9] 140 - 340 mg/kg (rat, oral) 1350 mg/kg (rabbit, dermal) |
LDLo (lowest published)
|
500 mg/kg (rabbit, oral)[10] |
| NIOSH (US health exposure limits): | |
PEL (Permissible)
|
TWA 2 mg/m3[8] |
REL (Recommended)
|
C 2 mg/m3[8] |
IDLH (Immediate danger)
|
10 mg/m3[8] |
| Safety data sheet (SDS) | External SDS |
| Related compounds | |
Other anions
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Other cations
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Related compounds
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Except where otherwise noted, data are given for materials in their standard state (at 25 °C , 100 kPa).
| |
Sodium hydroxide, also known as lye and caustic soda,[1][2] is an inorganic compound with the formula NaOH. It is a white solid ionic compound consisting of sodium cations Na+ and hydroxide anions OH−.
Sodium hydroxide is a highly corrosive base and alkali that decomposes lipids and proteins at ambient temperatures and may cause severe chemical burns. It is highly soluble in water, and readily absorbs moisture and carbon dioxide from the air. It forms a series of hydrates NaOH·nH2O.[11] The monohydrate NaOH·H2O crystallizes from water solutions between 12.3 and 61.8 °C. The commercially available "sodium hydroxide" is often this monohydrate, and published data may refer to it instead of the anhydrous compound.
As one of the simplest hydroxides, sodium hydroxide is frequently used alongside neutral water and acidic hydrochloric acid to demonstrate the pH scale to chemistry students.[12]
Sodium hydroxide is used in many industries: in the making of wood pulp and paper, textiles, drinking water, soaps and detergents, and as a drain cleaner. Worldwide production in 2004 was approximately 60 million tons, while demand was 51 million tons.[13]
Properties
Physical properties
Pure sodium hydroxide is a colorless crystalline solid that melts at 318 °C (604 °F) without decomposition and boils at 1,388 °C (2,530 °F). It is highly soluble in water, with a lower solubility in polar solvents such as ethanol and methanol.[14] Sodium hydroxide is insoluble in ether and other non-polar solvents.
Similar to the hydration of sulfuric acid, dissolution of solid sodium hydroxide in water is a highly exothermic reaction[15] where a large amount of heat is liberated, posing a threat to safety through the possibility of splashing. The resulting solution is usually colorless and odorless. As with other alkaline solutions, it feels slippery with skin contact due to the process of saponification that occurs between NaOH and natural skin oils.
Viscosity
Concentrated (50%) aqueous solutions of sodium hydroxide have a characteristic viscosity, 78 mPa·s, that is much greater than that of water (1.0 mPa·s) and near that of olive oil (85 mPa·s) at room temperature. The viscosity of aqueous NaOH, as with any liquid chemical, is inversely related to its temperature, i.e., its viscosity decreases as temperature increases, and vice versa. The viscosity of sodium hydroxide solutions plays a direct role in its application as well as its storage.[14]
Hydrates
Sodium hydroxide can form several hydrates NaOH·nH2O, which result in a complex solubility diagram that was described in detail by Spencer Umfreville Pickering in 1893.[16] The known hydrates and the approximate ranges of temperature and concentration (mass percent of NaOH) of their saturated water solutions are:[11]
- Heptahydrate, NaOH·7H2O: from −28 °C (18.8%) to −24 °C (22.2%).[16]
- Pentahydrate, NaOH·5H2O: from −24 °C (22.2%) to −17.7 °C (24.8%).[16]
- Tetrahydrate, NaOH·4H2O, α form: from −17.7 °C (24.8%) to 5.4 °C (32.5%).[16][17]
- Tetrahydrate, NaOH·4H2O, β form: metastable.[16][17]
- Trihemihydrate, NaOH·3.5H2O: from 5.4 °C (32.5%) to 15.38 °C (38.8%) and then to 5.0 °C (45.7%).[16][11]
- Trihydrate, NaOH·3H2O: metastable.[16]
- Dihydrate, NaOH·2H2O: from 5.0 °C (45.7%) to 12.3 °C (51%).[16][11]
- Monohydrate, NaOH·H2O: from 12.3 °C (51%) to 65.10 °C (69%) then to 62.63 °C (73.1%).[16][18]
Early reports refer to hydrates with n = 0.5 or n = 2/3, but later careful investigations failed to confirm their existence.[18]
The only hydrates with stable melting points are NaOH·H2O (65.10 °C) and NaOH·3.5H2O (15.38 °C). The other hydrates, except the metastable ones NaOH·3H2O and NaOH·4H2O (β) can be crystallized from solutions of the proper composition, as listed above. However, solutions of NaOH can be easily supercooled by many degrees, which allows the formation of hydrates (including the metastable ones) from solutions with different concentrations.[11][18]
For example, when a solution of NaOH and water with 1:2 mole ratio (52.6% NaOH by mass) is cooled, the monohydrate normally starts to crystallize (at about 22 °C) before the dihydrate. However, the solution can easily be supercooled down to −15 °C, at which point it may quickly crystallize as the dihydrate. When heated, the solid dihydrate might melt directly into a solution at 13.35 °C; however, once the temperature exceeds 12.58 °C it often decomposes into solid monohydrate and a liquid solution. Even the n = 3.5 hydrate is difficult to crystallize, because the solution supercools so much that other hydrates become more stable.[11]
A hot water solution containing 73.1% (mass) of NaOH is a eutectic that solidifies at about 62.63 °C as an intimate mix of anhydrous and monohydrate crystals.[19][18]
A second stable eutectic composition is 45.4% (mass) of NaOH, that solidifies at about 4.9 °C into a mixture of crystals of the dihydrate and of the 3.5-hydrate.[11]
The third stable eutectic has 18.4% (mass) of NaOH. It solidifies at about −28.7 °C as a mixture of water ice and the heptahydrate NaOH·7H2O.[16][20]
When solutions with less than 18.4% NaOH are cooled, water ice crystallizes first, leaving the NaOH in solution.[16]
The α form of the tetrahydrate has density 1.33 g/cm3. It melts congruously at 7.55 °C into a liquid with 35.7% NaOH and density 1.392 g/cm3, and therefore floats on it like ice on water. However, at about 4.9 °C it may instead melt incongruously into a mixture of solid NaOH·3.5H2O and a liquid solution.[17]
The β form of the tetrahydrate is metastable, and often transforms spontaneously to the α form when cooled below −20 °C.[17] Once initiated, the exothermic transformation is complete in a few minutes, with a 6.5% increase in volume of the solid. The β form can be crystallized from supercooled solutions at −26 °C, and melts partially at −1.83 °C.[17]
The "sodium hydroxide" of commerce is often the monohydrate (density 1.829 g/cm3). Physical data in technical literature may refer to this form, rather than the anhydrous compound.
Crystal structure
NaOH and its monohydrate form orthorhombic crystals with the space groups Cmcm (oS8) and Pbca (oP24), respectively. The monohydrate cell dimensions are a = 1.1825, b = 0.6213, c = 0.6069 nm. The atoms are arranged in a hydrargillite-like layer structure, with each sodium atom surrounded by six oxygen atoms, three each from hydroxide ions and three from water molecules. The hydrogen atoms of the hydroxyls form strong bonds with oxygen atoms within each O layer. Adjacent O layers are held together by hydrogen bonds between water molecules.[21]
Chemical properties
Reaction with acids
Sodium hydroxide reacts with protic acids to produce water and the corresponding salts. For example, when sodium hydroxide reacts with hydrochloric acid, sodium chloride is formed:
- NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)
In general, such neutralization reactions are represented by one simple net ionic equation:
- OH−(aq) + H+(aq) → H2O(l)
This type of reaction with a strong acid releases heat, and hence is exothermic. Such acid–base reactions can also be used for titrations. However, sodium hydroxide is not used as a primary standard because it is hygroscopic and absorbs carbon dioxide from air.
Reaction with acidic oxides
Sodium hydroxide also reacts with acidic oxides, such as sulfur dioxide. Such reactions are often used to "scrub" harmful acidic gases (like SO2 and H2S) produced in the burning of coal and thus prevent their release into the atmosphere. For example,
- 2 NaOH + SO2 → Na2SO3 + H2O
Reaction with metals and oxides
Glass reacts slowly with aqueous sodium hydroxide solutions at ambient temperatures to form soluble silicates. Because of this, glass joints and stopcocks exposed to sodium hydroxide have a tendency to "freeze". Flasks and glass-lined chemical reactors are damaged by long exposure to hot sodium hydroxide, which also frosts the glass. Sodium hydroxide does not attack iron at room temperature, since iron does not have amphoteric properties (i.e., it only dissolves in acid, not base). Nevertheless, at high temperatures (e.g. above 500 °C), iron can react endothermically with sodium hydroxide to form iron(III) oxide, sodium metal, and hydrogen gas.[22] This is due to the lower enthalpy of formation of iron(III) oxide (−824.2 kJ/mol) compared to sodium hydroxide (−500 kJ/mol) and positive entropy change of the reaction, which implies spontaneity at high temperatures (ΔST > ΔH, ΔG < 0) and non-spontaneity at low temperatures (ΔST < ΔH, ΔG > 0). Consider the following reaction between molten sodium hydroxide and finely divided iron filings:
- 4 Fe + 6 NaOH → 2 Fe2O3 + 6 Na + 3 H2
A few transition metals, however, may react quite vigorously with sodium hydroxide under milder conditions.
In 1986, an aluminium road tanker in the UK was mistakenly used to transport 25% sodium hydroxide solution,[23] causing pressurization of the contents and damage to tankers. The pressurization is due to the hydrogen gas which is produced in the reaction between sodium hydroxide and aluminium:
- 2 Al + 2 NaOH + 6 H2O → 2 Na[Al(OH)4 + 3 H2
Precipitant
Unlike sodium hydroxide, which is soluble, the hydroxides of most transition metals are insoluble, and therefore sodium hydroxide can be used to precipitate transition metal hydroxides. The following colours are observed:
- Copper - blue
- Iron(II) - green
- Iron(III) - yellow / brown
Zinc and lead salts dissolve in excess sodium hydroxide to give a clear solution of Na2ZnO2 or Na2PbO2.
Aluminium hydroxide is used as a gelatinous flocculant to filter out particulate matter in water treatment. Aluminium hydroxide is prepared at the treatment plant from aluminium sulfate by reacting it with sodium hydroxide or bicarbonate.
- Al2(SO4)3 + 6 NaOH → 2 Al(OH)3 + 3 Na2SO4
- Al2(SO4)3 + 6 NaHCO3 → 2 Al(OH)3 + 3 Na2SO4 + 6 CO2
Saponification
Sodium hydroxide can be used for the base-driven hydrolysis of esters (also called saponification), amides and alkyl halides.[14] However, the limited solubility of sodium hydroxide in organic solvents means that the more soluble potassium hydroxide (KOH) is often preferred. Touching a sodium hydroxide solution with bare hands, while not recommended, produces a slippery feeling. This happens because oils on the skin such as sebum are converted to soap. Despite solubility in propylene glycol it is unlikely to replace water in saponification due to propylene glycol's primary reaction with fat before reaction between sodium hydroxide and fat.