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Sulfuric acid
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Not to be confused with Sulfurous acid.
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Sulfuric acid
Space-filling model
Ball-and-stick model
Names
IUPAC name
Sulfuric acid
Other names
Oil of vitriol
Identifiers
CAS Number
7664-93-9 Y
3D model (JSmol)
Interactive image
ChEBI
CHEBI:26836 Y
ChemSpider
1086 Y
ECHA InfoCard 100.028.763
EC Number 231-639-5
E number E513 (acidity regulators, ...)
Gmelin Reference
2122
KEGG
D05963 Y
PubChem CID
1118
RTECS number WS5600000
UNII
O40UQP6WCF Y
UN number 1830
InChI[show]
InChI=1S/H2O4S/c1-5(2,3)4/h(H2,1,2,3,4) Y
Key: QAOWNCQODCNURD-UHFFFAOYSA-N Y
InChI=1/H2O4S/c1-5(2,3)4/h(H2,1,2,3,4)
Key: QAOWNCQODCNURD-UHFFFAOYAC
SMILES[show]
OS(=O)(=O)O
Properties
Chemical formula
H
2SO
4
Molar mass 98.079 g/mol
Appearance Clear, colorless liquid
Odor odorless
Density 1.84 g/cm3, liquid
Melting point 10 °C (50 °F; 283 K)
Boiling point 337 °C (639 °F; 610 K) When sulfuric acid is above 300 °C (572 °F), it will decompose slowly
Solubility in water
miscible
Vapor pressure 0.001 mmHg (20 °C)[1]
Acidity (pKa) −3, 1.99
Viscosity 26.7 cP (20 °C)
Thermochemistry
Std molar
entropy (So298)
157 J·mol−1·K−1[2]
Std enthalpy of
formation (ΔfHo298)
−814 kJ·mol−1[2]
Hazards
Safety data sheet External MSDS
GHS pictograms
GHS signal word Danger
GHS hazard statements
H314
GHS precautionary statements
P260, P264, P280, P301+330+331, P303+361+353, P363, P304+340, P305+351+338, P310, P321, P310, P405, P501
NFPA 704
0
3
2
W
Flash point Non-flammable
Threshold limit value
15 mg/m3 (IDLH), 1 mg/m3 (TWA), 2 mg/m3 (STEL)
Lethal dose or concentration (LD, LC):
LD50 (median dose)
2140 mg/kg (rat, oral)[3]
LC50 (median concentration)
50 mg/m3 (guinea pig, 8 hr)
510 mg/m3 (rat, 2 hr)
320 mg/m3 (mouse, 2 hr)
18 mg/m3 (guinea pig)[3]
LCLo (lowest published)
87 mg/m3 (guinea pig, 2.75 hr)[3]
US health exposure limits (NIOSH):
PEL (Permissible)
TWA 1 mg/m3[1]
REL (Recommended)
TWA 1 mg/m3[1]
IDLH (Immediate danger)
15 mg/m3[1]
Related compounds
Related strong acids
Selenic acid
Hydrochloric acid
Nitric acid
Chromic acid
Related compounds
Sulfurous acid
Peroxymonosulfuric acid
Sulfur trioxide
Oleum
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is Y N ?)
Infobox references
Sulfuric acid (alternative spelling sulphuric acid) is a highly corrosive strong mineral acid with the molecular formula H2SO4 and molecular weight 98.079 g/mol. It is a pungent-ethereal, colorless to slightly yellow viscous liquid that is soluble in water at all concentrations.[4] Sometimes, it is dyed dark brown during production to alert people to its hazards.[5] The historical name of this acid is oil of vitriol.[6]
Sulfuric acid is a diprotic acid and shows different properties depending upon its concentration. Its corrosiveness on other materials, like metals, living tissues or even stones, can be mainly ascribed to its strong acidic nature and, if concentrated, strong dehydrating and oxidizing properties. It is also hygroscopic, readily absorbing water vapour from the air.[4] Sulfuric acid at a high concentration can cause very serious damage upon contact, since not only does it cause chemical burns via hydrolysis, but also secondary thermal burns through dehydration.[7][8] It can lead to permanent blindness if splashed onto eyes and irreversible damage if swallowed.[7]
Sulfuric acid has a wide range of applications including in domestic acidic drain cleaners,[9] as an electrolyte in lead-acid batteries and in various cleaning agents. It is also a central substance in the chemical industry. Principal uses include mineral processing, fertilizer manufacturing, oil refining, wastewater processing, and chemical synthesis. It is widely produced with different methods, such as contact process, wet sulfuric acid process, lead chamber process and some other methods.
Contents
[hide]
1 History
2 Physical properties
2.1 Grades of sulfuric acid
2.2 Polarity and conductivity
3 Chemical properties
3.1 Reaction with water and dehydrating property
3.2 Acid-base properties
3.3 Reactions with metals and strong oxidizing property
3.4 Reactions with non-metals
3.5 Reaction with sodium chloride
3.6 Electrophilic aromatic substitution
4 Occurrence
4.1 Stratospheric aerosol
4.2 Extraterrestrial sulfuric acid
4.2.1 Venus
4.2.2 Europa
5 Manufacture
5.1 Contact process
5.2 Wet sulfuric acid process
5.3 Other methods
6 Uses
6.1 Industrial production of chemicals
6.2 Sulfur–iodine cycle
6.3 Industrial cleaning agent
6.4 Catalyst
6.5 Electrolyte
6.6 Domestic uses
7 Safety
7.1 Laboratory hazards
7.2 Dilution hazards
7.3 Industrial hazards
8 Legal restrictions
9 See also
10 References
11 External links
History[edit]
John Dalton's 1808 sulfuric acid molecule shows a central sulfur atom bonded to three oxygen atoms, or sulfur trioxide, the anhydride of sulfuric acid.
The study of vitriol, a category of glassy minerals from which the acid can be derived, began in ancient times. Sumerians had a list of types of vitriol that they classified according to the substances' color. Some of the earliest discussions on the origin and properties of vitriol is in the works of the Greek physician Dioscorides (first century AD) and the Roman naturalist Pliny the Elder (23–79 AD). Galen also discussed its medical use. Metallurgical uses for vitriolic substances were recorded in the Hellenistic alchemical works of Zosimos of Panopolis, in the treatise Phisica et Mystica, and the Leyden papyrus X.[10]
Medieval Islamic era alchemists, Jābir ibn Hayyān (c. 721 – c. 815 AD, also known as Geber), Razi (865 – 925 AD), and Jamal Din al-Watwat (d. 1318, wrote the book Mabāhij al-fikar wa-manāhij al-'ibar), included vitriol in their mineral classification lists. Ibn Sina focused on its medical uses and different varieties of vitriol.[10]
Sulfuric acid was called "oil of vitriol" by medieval European alchemists because it was prepared by roasting "green vitriol" (iron(II) sulfate) in an iron retort. There are references to it in the works of Vincent of Beauvais and in the Compositum de Compositis ascribed to Saint Albertus Magnus. A passage from Pseudo-Geber´s Summa Perfectionis was long considered to be the first recipe for sulfuric acid, but this was a misinterpretation.[10]
In the seventeenth century, the German-Dutch chemist Johann Glauber prepared sulfuric acid by burning sulfur together with saltpeter (potassium nitrate, KNO
3), in the presence of steam. As saltpeter decomposes, it oxidizes the sulfur to SO
3, which combines with water to produce sulfuric acid. In 1736, Joshua Ward, a London pharmacist, used this method to begin the first large-scale production of sulfuric acid.
In 1746 in Birmingham, John Roebuck adapted this method to produce sulfuric acid in lead-lined chambers, which were stronger, less expensive, and could be made larger than the previously used glass containers. This process allowed the effective industrialization of sulfuric acid production. After several refinements, this method, called the lead chamber process or "chamber process", remained the standard for sulfuric acid production for almost two centuries.[2]
Sulfuric acid created by John Roebuck's process approached a 65% concentration. Later refinements to the lead chamber process by French chemist Joseph Louis Gay-Lussac and British chemist John Glover improved concentration to 78%. However, the manufacture of some dyes and other chemical processes require a more concentrated product. Throughout the 18th century, this could only be made by dry distilling minerals in a technique similar to the original alchemical processes. Pyrite (iron disulfide, FeS
2) was heated in air to yield iron(II) sulfate, FeSO
4, which was oxidized by further heating in air to form iron(III) sulfate, Fe2(SO4)3, which, when heated to 480 °C, decomposed to iron(III) oxide and sulfur trioxide, which could be passed through water to yield sulfuric acid in any concentration. However, the expense of this process prevented the large-scale use of concentrated sulfuric acid.[2]
In 1831, British vinegar merchant Peregrine Phillips patented the contact process, which was a far more economical process for producing sulfur trioxide and concentrated sulfuric acid. Today, nearly all of the world's sulfuric acid is produced using this method.[11]
Physical properties[edit]
Grades of sulfuric acid[edit]
Although nearly 99% sulfuric acid can be made, the subsequent loss of SO
3 at the boiling point brings the concentration to 98.3% acid. The 98% grade is more stable in storage, and is the usual form of what is described as "concentrated sulfuric acid". Other concentrations are used for different purposes. Some common concentrations are:[12][13]
Mass fraction
H2SO4
Density
(kg/L)
Concentration
(mol/L)
Common name
10% 1.07 ≈1 dilute sulfuric acid
29–32% 1.25–1.28 4.2–5 battery acid
(used in lead–acid batteries)
62–70% 1.52–1.60 9.6–11.5 chamber acid
fertilizer acid
78–80% 1.70–1.73 13.5–14 tower acid
Glover acid
98% 1.84 ≈18 concentrated sulfuric acid
"Chamber acid" and "tower acid" were the two concentrations of sulfuric acid produced by the lead chamber process, chamber acid being the acid produced in lead chamber itself (<70% to avoid contamination with nitrosylsulfuric acid) and tower acid being the acid recovered from the bottom of the Glover tower.[12][13] They are now obsolete as commercial concentrations of sulfuric acid, although they may be prepared in the laboratory from concentrated sulfuric acid if needed. In particular, "10M" sulfuric acid (the modern equivalent of chamber acid, used in many titrations) is prepared by slowly adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80 °C (176 °F) or higher.[13]
Sulfuric acid reacts with its anhydride, SO
3, to form H
2S
2O
7, called pyrosulfuric acid, fuming sulfuric acid, Disulfuric acid or oleum or, less commonly, Nordhausen acid. Concentrations of oleum are either expressed in terms of% SO
3 (called% oleum) or as% H
2SO
4 (the amount made if H
2O were added); common concentrations are 40% oleum (109% H
2SO
4) and 65% oleum (114.6% H
2SO
4). Pure H
2S
2O
7 is a solid with melting point of 36 °C.
Pure sulfuric acid has a vapor pressure of <0.001 mmHg at 25 °C and 1 mmHg at 145.8 °C,[14] and 98% sulfuric acid has a <1 mmHg vapor pressure at 40 °C.[15]
Pure sulfuric acid is a viscous clear liquid, like oil, and this explains the old name of the acid ('oil of vitriol').
Commercial sulfuric acid is sold in several different purity grades. Technical grade H
2SO
4 is impure and often colored, but is suitable for making fertilizer. Pure grades, such as United States Pharmacopeia (USP) grade, are used for making pharmaceuticals and dyestuffs. Analytical grades are also available.
There are nine hydrates known, but three of them were confirmed to be tetrahydrate (H2SO4·4H2O), hemihexahydrate (H2SO4· 6 1⁄2H2O) and octahydrate (H2SO4·8H2O).
Polarity and conductivity[edit]
Equilibrium of anhydrous sulfuric acid[16]
Species
mMol/kg
HSO4- 15.0
H3SO4+ 11.3
H3O+ 8.0
HS2O7- 4.4
H2S2O7 3.6
H2O 0.1
Anhydrous H2SO4 is a very polar liquid, having a dielectric constant of around 100. It has a high electrical conductivity, caused by dissociation through protonating itself, a process known as autoprotolysis.[16]
2 H2SO4 ⇌ H3SO4+ + HSO4-
The equilibrium constant for the autoprotolysis is[16]
Kap(25 °C)= [H3SO4+] [HSO4-] = 6996270000000000000♠2.7×10−4
The comparable equilibrium constant for water, Kw is 10−14, a factor of 1010 (10 billion) smaller.
In spite of the viscosity of the acid, the effective conductivities of the H3SO4+ and HSO4- ions are high due to an intra-molecular proton-switch mechanism (analogous to the Grotthuss mechanism in water), making sulfuric acid a good conductor of electricity. It is also an excellent solvent for many reactions.
Chemical properties[edit]
Reaction with water and dehydrating property[edit]
Drops of concentrated sulfuric acid rapidly decompose a piece of cotton towel by dehydration.
Play media
An experiment that demonstrates the dehydration properties of concentrated sulfuric acid. When concentrated sulfuric acid comes into contact with sucrose, slow carbonification of the sucrose takes place. The reaction is accompanied by the evolution of gaseous products that contribute to the formation of the foamy carbon pillar that rises above the beaker.
Because the hydration reaction of sulfuric acid is highly exothermic, dilution should always be performed by adding the acid to the water rather than the water to the acid.[17] Because the reaction is in an equilibrium that favors the rapid protonation of water, addition of acid to the water ensures that the acid is the limiting reagent. This reaction is best thought of as the formation of hydronium ions:
H
2SO
4 + H
2O → H
3O+
+ HSO4− K1 = 2.4×106 (strong acid)
HSO−
4 + H
2O → H
3O+
+ SO2−
4 K
2 = 1.0×10−2 [18]
HSO−
4 is the bisulfate anion and SO2−
4 is the sulfate anion. K1 and K2 are the acid dissociation constants.
Because the hydration of sulfuric acid is thermodynamically favorable and the affinity of it for water is sufficiently strong, sulfuric acid is an excellent dehydrating agent. Concentrated sulfuric acid has a very powerful dehydrating property, removing water (H2O) from other compounds including sugar and other carbohydrates and producing carbon, heat, steam.
In the laboratory, this is often demonstrated by mixing table sugar (sucrose) into sulfuric acid. The sugar changes from white to dark brown and then to black as carbon is formed. A rigid column of black, porous carbon will emerge as well. The carbon will smell strongly of caramel due to the heat generated.[19]
C 12 H 22 O 11 ⏞ sucrose → H 2 SO 4 12 C ( black graphitic foam ) + 11 H 2 O ( g , l ) {\displaystyle {\ce {\overbrace {C12H22O11} ^{sucrose}->[{\ce {H2SO4}}]{\underset {(black~graphitic~foam)}{12C}}+{11H2O}_{(g,l)}}}}
Similarly, mixing starch into concentrated sulfuric acid will give elemental carbon and water as absorbed by the sulfuric acid (which becomes slightly diluted). The effect of this can be seen when concentrated sulfuric acid is spilled on paper which is composed of cellulose; the cellulose reacts to give a burnt appearance, the carbon appears much as soot would in a fire. Although less dramatic, the action of the acid on cotton, even in diluted form, will destroy the fabric.
( C 6 H 10 O 5 ) n ⏞ polysaccharide → H 2 SO 4 6 n C + 5 n H 2 O {\displaystyle {\ce {\overbrace {(C6H10O5)_{\mathit {n}}} ^{polysaccharide}->[{\ce {H2SO4}}]}}\ 6n{\ce {C}}+5n{\ce {H2O}}}
The reaction with copper(II) sulfate can also demonstrate the dehydration property of sulfuric acid. The blue crystal is changed into white powder as water is removed.
CuSO 4 ⋅ 5 H 2 O ( blue crystal ) ⏞ copper ( II ) sulfate hydrate → H 2 SO 4 CuSO 4 ( white powder ) ⏞ Anhydrous copper ( II ) sulfate + 5 H 2 O {\displaystyle {\ce {\overbrace {\underset {(blue~crystal)}{CuSO4.5H2O}} ^{copper(II)~sulfate \atop hydrate}->[{\ce {H2SO4}}]\overbrace {\underset {(white~powder)}{CuSO4}} ^{Anhydrous \atop copper(II)~sulfate}+5H2O}}}
Acid-base properties[edit]
As an acid, sulfuric acid reacts with most bases to give the corresponding sulfate. For example, the blue copper salt copper(II) sulfate, commonly used for electroplating and as a fungicide, is prepared by the reaction of copper(II) oxide with sulfuric acid:
CuO (s) + H
2SO
4 (aq) → CuSO
4 (aq) + H
2O (l)
Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with sodium acetate, for example, displaces acetic acid, CH
3COOH, and forms sodium bisulfate:
H
2SO
4 + CH
3COONa → NaHSO
4 + CH
3COOH
Similarly, reacting sulfuric acid with potassium nitrate can be used to produce nitric acid and a precipitate of potassium bisulfate. When combined with nitric acid, sulfuric acid acts both as an acid and a dehydrating agent, forming the nitronium ion NO+
2, which is important in nitration reactions involving electrophilic aromatic substitution. This type of reaction, where protonation occurs on an oxygen atom, is important in many organic chemistry reactions, such as Fischer esterification and dehydration of alcohols.
Solid state structure of the [D3SO4]+ ion present in [D3SO4]+[SbF6]−, synthesized by using DF in place of HF. (see text)
When allowed to react with superacids, sulfuric acid can act as a base and be protonated, forming the [H3SO4]+ ion. Salt of [H3SO4]+ have been prepared using the following reaction in liquid HF:
((CH3)3SiO)2SO2 + 3 HF + SbF5 → [H3SO4]+[SbF6]− + 2 (CH3)3SiF
The above reaction is thermodynamically favored due to the high bond enthalpy of the Si–F bond in the side product. Protonation using simply HF/SbF5, however, have met with failure, as pure sulfuric acid undergoes self-ionization to give [H3O]+ ions, which prevents the conversion of H2SO4 to [H3SO4]+ by the HF/SbF5 system:[20]
2 H2SO4 ⇌ [H3O]+ + [HS2O7]−
Reactions with metals and strong oxidizing property[edit]
Dilute sulfuric acid reacts with metals via a single displacement reaction as with other typical acids, producing hydrogen gas and salts (the metal sulfate). It attacks reactive metals (metals at positions above copper in the reactivity series) such as iron, aluminium, zinc, manganese, magnesium and nickel.
Fe (s) + H
2SO
4 (aq) → H
2 (g) + FeSO
4 (aq)
However, concentrated sulfuric acid is a strong oxidizing agent[7] and does not react with metals in the same way as other typical acids. Sulfur dioxide, water and SO42− ions are evolved instead of the hydrogen and salts.
2 H2SO4 + 2 e− → SO2 + 2 H2O + SO42−
It can oxidize non-active metals such as tin and copper, depending upon the temperature.
Cu + 2 H2SO4 → SO2 + 2 H2O + SO42− + Cu2+
Lead and tungsten, however, are resistant to sulfuric acid.
Reactions with non-metals[edit]
Hot concentrated sulfuric acid oxidizes non-metals such as carbon[21] (as bituminous coal) and sulfur.
C + 2 H2SO4 → CO2 + 2 SO2 + 2 H2O
S + 2 H2SO4 → 3 SO2 + 2 H2O
Reaction with sodium chloride[edit]
It reacts with sodium chloride, and gives hydrogen chloride gas and sodium bisulfate:
NaCl + H2SO4 → NaHSO4 + HCl
Electrophilic aromatic substitution[edit]
Benzene undergoes electrophilic aromatic substitution with sulfuric acid to give the corresponding sulfonic acids:[22]
Occurrence[edit]
Rio Tinto with its highly acidic water
Pure sulfuric acid is not encountered naturally on Earth in anhydrous form, due to its great affinity for water. Dilute sulfuric acid is a constituent of acid rain, which is formed by atmospheric oxidation of sulfur dioxide in the presence of water – i.e., oxidation of sulfurous acid. Sulfur dioxide is the main byproduct produced when sulfur-containing fuels such as coal or oil are burned.
Sulfuric acid is formed naturally by the oxidation of sulfide minerals, such as iron sulfide. The resulting water can be highly acidic and is called acid mine drainage (AMD) or acid rock drainage (ARD). This acidic water is capable of dissolving metals present in sulfide ores, which results in brightly colored, toxic streams. The oxidation of pyrite (iron sulfide) by molecular oxygen produces iron(II), or Fe2+
:
2 FeS
2 (s) + 7 O
2 + 2 H2O → 2 Fe2+
+ 4 SO2−
4 + 4 H+
The Fe2+
can be further oxidized to Fe3+
:
4 Fe2+
+ O
2 + 4 H+
→ 4 Fe3+
+ 2 H2O
The Fe3+
produced can be precipitated as the hydroxide or hydrous oxide:
Fe3+
+ 3 H2O → Fe(OH)
3↓ + 3 H+
The iron(III) ion ("ferric iron") can also oxidize pyrite:
FeS
2(s) + 14 Fe3+
+ 8 H2O → 15 Fe2+
+ 2 SO2−
4 + 16 H+
When iron(III) oxidation of pyrite occurs, the process can become rapid. pH values below zero have been measured in ARD produced by this process.
ARD can also produce sulfuric acid at a slower rate, so that the acid neutralizing capacity (ANC) of the aquifer can neutralize the produced acid. In such cases, the total dissolved solids (TDS) concentration of the water can be increased from the dissolution of minerals from the acid-neutralization reaction with the minerals.
Sulfuric acid is used as a defence by certain marine species, for example, the phaeophyte alga Desmarestia munda (order Desmarestiales) concentrates sulfuric acid in cell vacuoles.[23]
Stratospheric aerosol[edit]
In the stratosphere, the atmosphere's second layer that is generally between 10 and 50 km above Earth's surface, sulfuric acid is formed by the oxidation of volcanic sulfur dioxide by the hydroxyl radical:[24]
SO
2 + HO• → HSO
3
HSO
3 + O
2 → SO
3 + HO
2
SO
3 + H2O → H
2SO
4
Because sulfuric acid reaches supersaturation in the stratosphere, it can nucleate aerosol particles and provide a surface for aerosol growth via condensation and coagulation with other water-sulfuric acid aerosols. This results in the stratospheric aerosol layer.[24]
Extraterrestrial sulfuric acid[edit]
Venus[edit]
Sulfuric acid is produced in the upper atmosphere of Venus by the Sun's photochemical action on carbon dioxide, sulfur dioxide, and water vapor. Ultraviolet photons of wavelengths less than 169 nm can photodissociate carbon dioxide into carbon monoxide and atomic oxygen. Atomic oxygen is highly reactive. When it reacts with sulfur dioxide, a trace component of the Venusian atmosphere, the result is sulfur trioxide, which can combine with water vapor, another trace component of Venus's atmosphere, to yield sulfuric acid. In the upper, cooler portions of Venus's atmosphere, sulfuric acid exists as a liquid, and thick sulfuric acid clouds completely obscure the planet's surface when viewed from above. The main cloud layer extends from 45–70 km above the planet's surface, with thinner hazes extending as low as 30 km and as high as 90 km above the surface. The permanent Venusian clouds produce a concentrated acid rain, as the clouds in the atmosphere of Earth produce water rain.
The atmosphere exhibits a sulfuric acid cycle. As sulfuric acid rain droplets fall down through the hotter layers of the atmosphere's temperature gradient, they are heated up and release water vapor, becoming more and more concentrated. When they reach temperatures above 300 °C, sulfuric acid begins to decompose into sulfur trioxide and water, both in the gas phase. Sulfur trioxide is highly reactive and dissociates into sulfur dioxide and atomic oxygen, which oxidizes traces of carbon monoxide to form carbon dioxide. Sulfur dioxide and water vapor rise on convection currents from the mid-level atmospheric layers to higher altitudes, where they will be transformed again into sulfuric acid, and the cycle repeats.
Europa[edit]
Infrared spectra taken by NASA's Galileo spacecraft show distinct absorptions on Jupiter's moon Europa that have been attributed to one or more sulfuric acid hydrates. Sulfuric acid in solution with water causes significant freezing-point depression of water's melting point, down to 210 K (−63 °C), and this would make the existence of liquid solutions beneath Europa's icy crust more likely. The interpretation of the spectra is somewhat controversial. Some planetary scientists prefer to assign the spectral features to the sulfate ion, perhaps as part of one or more minerals on Europa's surface.[25]
Manufacture[edit]
Main articles: Contact process, Wet sulfuric acid process, and Lead chamber process
Sulfuric acid is produced from sulfur, oxygen and water via the conventional contact process (DCDA) or the wet sulfuric acid process (WSA).
Contact process[edit]
Main article: Contact process
In the first step, sulfur is burned to produce sulfur dioxide.
S (s) + O
2 (g) → SO
2 (g)
This is then oxidized to sulfur trioxide using oxygen in the presence of a vanadium(V) oxide catalyst. This reaction is reversible and the formation of the sulfur trioxide is exothermic.
2 SO
2 (g) + O
2 (g) ⇌ 2 SO
3 (g) (in presence of V
2O
5)
The sulfur trioxide is absorbed into 97–98% H
2SO
4 to form oleum (H
2S
2O
7), also known as fuming sulfuric acid. The oleum is then diluted with water to form concentrated sulfuric acid.
H
2SO
4 (l) + SO
3 (g)→ H
2S
2O
7 (l)
H
2S
2O
7 (l) + H
2O (l) → 2 H
2SO
4 (l)
Note that directly dissolving SO
3 in water is not practical due to the highly exothermic nature of the reaction between sulfur trioxide and water. The reaction forms a corrosive aerosol that is very difficult to separate, instead of a liquid.
SO
3 (g) + H
2O (l) → H
2SO
4 (l)
Wet sulfuric acid process[edit]
Main article: Wet sulfuric acid process
In the first step, sulfur is burned to produce sulfur dioxide:
S(s) + O
2(g) → SO
2(g)
or, alternatively, hydrogen sulfide (H
2S) gas is incinerated to SO
2 gas:
2 H
2S + 3 O
2 → 2 H
2O + 2 SO
2 (−518 kJ/mol)
This is then oxidized to sulfur trioxide using oxygen with vanadium(V) oxide as catalyst.
2 SO
2 + O
2 → 2 SO
3 (−99 kJ/mol) (reaction is reversible)
The sulfur trioxide is hydrated into sulfuric acid H
2SO
4:
SO
3 + H
2O → H
2SO
4(g) (−101 kJ/mol)
The last step is the condensation of the sulfuric acid to liquid 97–98% H
2SO
4:
H
2SO
4(g) → H
2SO
4(l) (−69 kJ/mol)
Other methods[edit]
Another method is the less well-known metabisulfite method, in which metabisulfite is placed at the bottom of a beaker, and 12.6 molar concentration hydrochloric acid is added. The resulting gas is bubbled through nitric acid, which will release brown/red vapors. The completion of the reaction is indicated by the ceasing of the fumes. This method does not produce an inseparable mist, which is quite convenient.
SO2 + HNO3 + H2O→ H2SO4 + NO
Sulfuric acid can be produced in the laboratory by burning sulfur in air and dissolving the gas produced in a hydrogen peroxide solution.
SO2 + H2O2 → H2SO4
Prior to 1900, most sulfuric acid was manufactured by the lead chamber process.[26] As late as 1940, up to 50% of sulfuric acid manufactured in the United States was produced by chamber process plants.
In early to mid nineteenth century "vitriol" plants existed, among other places, in Prestonpans in Scotland, Shropshire and the Lagan Valley in County Antrim Ireland where it was used as a bleach for linen. Early bleaching of linen was done using milk but this was a slow process and the use of vitriol sped up the bleaching process.[citation needed]
Uses[edit]
Sulfuric acid production in 2000
Sulfuric acid is a very important commodity chemical, and indeed, a nation's sulfuric acid production is a good indicator of its industrial strength.[27] World production in 2004 was about 180 million tonnes, with the following geographic distribution: Asia 35%, North America (including Mexico) 24%, Africa 11%, Western Europe 10%, Eastern Europe and Russia 10%, Australia and Oceania 7%, South America 7%.[28] Most of this amount (≈60%) is consumed for fertilizers, particularly superphosphates, ammonium phosphate and ammonium sulfates. About 20% is used in chemical industry for production of detergents, synthetic resins, dyestuffs, pharmaceuticals, petroleum catalysts, insecticides and antifreeze, as well as in various processes such as oil well acidicizing, aluminium reduction, paper sizing, water treatment. About 6% of uses are related to pigments and include paints, enamels, printing inks, coated fabrics and paper, and the rest is dispersed into a multitude of applications such as production of explosives, cellophane, acetate and viscose textiles, lubricants, non-ferrous metals and batteries.[29]
Industrial production of chemicals[edit]
The major use for sulfuric acid is in the "wet method" for the production of phosphoric acid, used for manufacture of phosphate fertilizers. In this method, phosphate rock is used, and more than 100 million tonnes are processed annually. This raw material is shown below as fluorapatite, though the exact composition may vary. This is treated with 93% sulfuric acid to produce calcium sulfate, hydrogen fluoride (HF) and phosphoric acid. The HF is removed as hydrofluoric acid. The overall process can be represented as:
Ca
5F(PO
4)
3 + 5 H
2SO
4 + 10 H
2O → 5 CaSO
4·2 H
2O + HF + 3 H
3PO
4
Ammonium sulfate, an important nitrogen fertilizer, is most commonly produced as a byproduct from coking plants supplying the iron and steel making plants. Reacting the ammonia produced in the thermal decomposition of coal with waste sulfuric acid allows the ammonia to be crystallized out as a salt (often brown because of iron contamination) and sold into the agro-chemicals industry.
Another important use for sulfuric acid is for the manufacture of aluminium sulfate, also known as paper maker's alum. This can react with small amounts of soap on paper pulp fibers to give gelatinous aluminium carboxylates, which help to coagulate the pulp fibers into a hard paper surface. It is also used for making aluminium hydroxide, which is used at water treatment plants to filter out impurities, as well as to improve the taste of the water. Aluminium sulfate is made by reacting bauxite with sulfuric acid:
2 AlO(OH) + 3 H
2SO
4 → Al
2(SO
4)
3 + 4 H
2O
Sulfuric acid is also important in the manufacture of dyestuffs solutions.
Sulfur–iodine cycle[edit]
The sulfur-iodine cycle is a series of thermo-chemical processes used to obtain hydrogen. It consists of three chemical reactions whose net reactant is water and whose net products are hydrogen and oxygen. Step one of cycle is the Bunsen reaction.
2 H
2SO
4 → 2 SO
2 + 2 H
2O + O
2 (830 °C)
I
2 + SO
2 + 2 H
2O → 2 HI + H
2SO
4 (120 °C)
2 HI → I
2 + H
2 (320 °C)
The sulfur and iodine compounds are recovered and reused, hence the consideration of the process as a cycle. This process is endothermic and must occur at high temperatures, so energy in the form of heat has to be supplied.
The sulfur-iodine cycle has been proposed as a way to supply hydrogen for a hydrogen-based economy. It does not require hydrocarbons like current methods of steam reforming. But note that all of the available energy in the hydrogen so produced is supplied by the heat used to make it.
The sulfur-iodine cycle is currently being researched as a feasible method of obtaining hydrogen, but the concentrated, corrosive acid at high temperatures poses currently insurmountable safety hazards if the process were built on a large scale.[30]
Industrial cleaning agent[edit]
Main article: Cleaning agent
Sulfuric acid is used in large quantities by the iron and steelmaking industry to remove oxidation, rust and scaling from rolled sheet and billets prior to sale to the automobile and major appliances industry.[citation needed] Used acid is often recycled using a spent acid regeneration (SAR) plant. These plants combust spent acid[clarification needed] with natural gas, refinery gas, fuel oil or other fuel sources. This combustion process produces gaseous sulfur dioxide (SO
2) and sulfur trioxide (SO
3) which are then used to manufacture "new" sulfuric acid. SAR plants are common additions to metal smelting plants, oil refineries, and other industries where sulfuric acid is consumed in bulk, as operating a SAR plant is much cheaper than the recurring costs of spent acid disposal and new acid purchases.
Hydrogen peroxide (H
2O
2) can be added to sulfuric acid to produce piranha solution, a powerful but very toxic cleaning solution with which substrate surfaces can be cleaned. Piranha solution is typically used the microelectronics industry, and also in laboratory settings to clean glassware.
Catalyst[edit]
Sulfuric acid is used for a variety of other purposes in the chemical industry. For example, it is the usual acid catalyst for the conversion of cyclohexanone oxime to caprolactam, used for making nylon. It is used for making hydrochloric acid from salt via the Mannheim process. Much H
2SO
4 is used in petroleum refining, for example as a catalyst for the reaction of isobutane with isobutylene to give isooctane, a compound that raises the octane rating of gasoline (petrol). Sulfuric acid is also often used as a dehydrating or oxidising agent in industrial reactions, such as the dehydration of various sugars to form solid carbon.
Electrolyte[edit]
Acidic drain cleaners usually contain sulfuric acid at a high concentration which turns a piece of pH paper red and chars it instantly, demonstrating both the strong acidic nature and dehydrating property.
Sulfuric acid acts as the electrolyte in lead–acid batteries (lead-acid accumulator):
At anode:
Pb + SO
42− ⇌ PbSO
4 + 2 e−
At cathode:
PbO
2 + 4 H+ + SO
42− + 2 e− ⇌ PbSO
4 + 2 H2O
An acidic drain cleaner can be used to dissolve grease, hair and even tissue paper inside water pipes.
Overall:
Pb + PbO
2 + 4 H+ + 2 SO
42− ⇌ 2 PbSO
4 + 2 H2O
Domestic uses[edit]
Sulfuric acid at high concentrations is frequently the major ingredient in acidic drain cleaners[9] which are used to remove grease, hair, tissue paper, etc. Similar to their alkaline versions, such drain openers can dissolve fats and proteins via hydrolysis. Moreover, as concentrated sulfuric acid has a strong dehydrating property, it can remove tissue paper via dehydrating process as well. Since the acid may react with water vigorously, such acidic drain openers should be added slowly into the pipe to be cleaned.
Safety[edit]
Laboratory hazards[edit]
Drops of 98% sulfuric acid char a piece of tissue paper instantly. Carbon is left after the dehydration reaction staining the paper black.
Superficial chemical burn caused by two 98% sulfuric acid splashes (forearm skin).
Sulfuric acid is capable of causing very severe burns, especially when it is at high concentrations. In common with other corrosive acids and alkali, it readily decomposes proteins and lipids through amide and ester hydrolysis upon contact with living tissues, such as skin and flesh. In addition, it exhibits a strong dehydrating property on carbohydrates, liberating extra heat and causing secondary thermal burns.[7][8] Accordingly, it rapidly attacks the cornea and can induce permanent blindness if splashed onto eyes. If ingested, it damages internal organs irreversibly and may even be fatal.[4] Protective equipment should hence always be used when handling it. Moreover, its strong oxidizing property makes it highly corrosive to many metals and may extend its destruction on other materials.[7] Because of such reasons, damage posed by sulfuric acid is potentially more severe than that by other comparable strong acids, such as hydrochloric acid and nitric acid.
Sulfuric acid must be stored carefully in containers made of nonreactive material (such as glass). Solutions equal to or stronger than 1.5 M are labeled "CORROSIVE", while solutions greater than 0.5 M but less than 1.5 M are labeled "IRRITANT". However, even the normal laboratory "dilute" grade (approximately 1 M, 10%) will char paper if left in contact for a sufficient time.
The standard first aid treatment for acid spills on the skin is, as for other corrosive agents, irrigation with large quantities of water. Washing is continued for at least ten to fifteen minutes to cool the tissue surrounding the acid burn and to prevent secondary damage. Contaminated clothing is removed immediately and the underlying skin washed thoroughly.
Dilution hazards[edit]
Preparation of the diluted acid can be dangerous due to the heat released in the dilution process. To avoid splattering, the concentrated acid must be added to water and not the other way around. Consider of the physical properties of acid and water:
Comparison of sulfuric acid and water
Physical property
H2SO4
Water
Units
Density
1.84 1.0 kg/L
Volumetric heat capacity
2.54 4.18 kJ/L
Boiling point
337 100 °C
Water has a higher heat capacity than acid, and a vessel of cold water will absorb heat as acid is added.
Also, because acid is denser than water, it sinks to the bottom. Heat is generated at interface between acid and water, which is at the bottom of the vessel. Acid will not boil, because of its higher boiling point. Warm water near the interface rises due to convection, which cools the interface, and prevents boiling of either acid or water.
In contrast, addition of water to concentrated sulfuric acid results in a thin layer of water on top of the acid. Heat generated in this thin layer of water can boil, leading to the dispersal of a sulfuric acid aerosol or worse, an explosion.
Preparation of solutions greater than 6 M (35%) in concentration is most dangerous, because the heat produced may be sufficient to boil the diluted acid: efficient mechanical stirring and external cooling (such as an ice bath) are essential.
Reaction rates double for about every 10 degree Celsius increase in temperature.[31] Therefore, the reaction will become more violent as dilution proceeds, unless the mixture is given time to cool. Adding acid to warm water will cause a violent reaction.
On a laboratory scale, sulfuric acid can be diluted by pouring concentrated acid onto crushed ice made from de-ionized water. The ice melts in an endothermic process while dissolving the acid. The amount of heat needed to melt the ice in this process is greater than the amount of heat evolved by dissolving the acid so the solution remains cold. After all the ice has melted, further dilution can take place using water.
Industrial hazards[edit]
Although sulfuric acid is non-flammable, contact with metals in the event of a spillage can lead to the liberation of hydrogen gas. The dispersal of acid aerosols and gaseous sulfur dioxide is an additional hazard of fires involving sulfuric acid.
The main occupational risks posed by this acid are skin contact leading to burns (see above) and the inhalation of aerosols. Exposure to aerosols at high concentrations leads to immediate and severe irritation of the eyes, respiratory tract and mucous membranes: this ceases rapidly after exposure, although there is a risk of subsequent pulmonary edema if tissue damage has been more severe. At lower concentrations, the most commonly reported symptom of chronic exposure to sulfuric acid aerosols is erosion of the teeth, found in virtually all studies: indications of possible chronic damage to the respiratory tract are inconclusive as of 1997. Repeated occupational exposure to sulfuric acid mists may increase the chance of lung cancer by up to 64 percent.[32] In the United States, the permissible exposure limit (PEL) for sulfuric acid is fixed at 1 mg/m3: limits in other countries are similar. There have been reports of sulfuric acid ingestion leading to vitamin B12 deficiency with subacute combined degeneration. The spinal cord is most often affected in such cases, but the optic nerves may show demyelination, loss of axons and gliosis.
Legal restrictions[edit]
International commerce of sulfuric acid is controlled under the United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances, 1988, which lists sulfuric acid under Table II of the convention as a chemical frequently used in the illicit manufacture of narcotic drugs or psychotropic substances.[33]
See also[edit]
Diethyl ether – also known as "sweet oil of vitriol"
Aqua regia
Piranha solution
Sulfuric acid poisoning
Sulfur oxoacid
References[edit]
^ Jump up to: a b c d "NIOSH Pocket Guide to Chemical Hazards #0577". National Institute for Occupational Safety and Health (NIOSH).
^ Jump up to: a b c d Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A23. ISBN 0-618-94690-X.
^ Jump up to: a b c "Sulfuric acid". Immediately Dangerous to Life and Health. National Institute for Occupational Safety and Health (NIOSH).
^ Jump up to: a b c "Sulfuric acid safety data sheet" (PDF). arkema-inc.com. Clear to turbid oily odorless liquid, colorless to slightly yellow.
Jump up ^ "Sulfuric acid". chemicalland21.com. Colorless (pure) to dark brown, oily, dense liquid with acrid odor.
Jump up ^ sulfuric acid. Encyclopædia Britannica. 2010.
^ Jump up to: a b c d e "Sulfuric acid – uses". dynamicscience.com.au.
^ Jump up to: a b "BASF Chemical Emergency Medical Guidelines – Sulfuric acid (H2SO4)" (PDF). BASF Chemical Company. 2012. Retrieved December 18, 2014.
^ Jump up to: a b "Sulphuric acid drain cleaner" (PDF). herchem.com.
^ Jump up to: a b c Karpenko, Vladimir and Norris, John A. (2001). Vitriol in the history of Chemistry, Charles University
Jump up ^ Philip J. Chenier (1 April 2002). Survey of industrial chemistry. Springer. pp. 28–. ISBN 978-0-306-47246-6. Retrieved 23 December 2011.
^ Jump up to: a b "sulfuric acid". The Columbia Encyclopedia (6th ed.). 2009. Retrieved 2010-03-16.
^ Jump up to: a b c "Sulphuric acid". Encyclopædia Britannica. 26 (11th ed.). 1910–1911. pp. 65–69.
Jump up ^ "Sulfuric acid". Determination of Noncancer Chronic Reference Exposure Levels Batch 2B December 2001 (PDF). 2001. Retrieved 2012-10-01.
Jump up ^ "Sulfuric Acid 98%" (PDF). rhodia.com. 2009. Retrieved 2014-07-02.
^ Jump up to: a b c Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9.
Jump up ^ Consortium of Local Education Authorities for the Provision of Science Equipment -STUDENT SAFETY SHEETS 22 Sulfuric(VI) acid
Jump up ^ "Ionization Constants of Inorganic Acids". .chemistry.msu.edu. Retrieved 2011-05-30.
Jump up ^ Sulphuric acid on sugar cubes chemistry experiment 8. Old Version. YouTube. Retrieved on 2011-07-18.
Jump up ^ Housecroft, Catherine E.; Sharpe, Alan G. (2008). "Chapter 16: The group 16 elements". Inorganic Chemistry, 3rd Edition. Pearson. p. 523. ISBN 978-0-13-175553-6.
Jump up ^ Kinney, Corliss Robert; Grey, V. E. (1959). Reactions of a Bituminous Coal with Sulfuric Acid (PDF). Pennsylvania State University.
Jump up ^ Carey, F. A. "Reactions of Arenes. Electrophilic Aromatic Substitution". On-Line Learning Center for Organic Chemistry. University of Calgary. Archived from the original on 6 July 2008. Retrieved 27 January 2008. CS1 maint: Unfit url (link)
Jump up ^ Pelletreau, K.; Muller-Parker, G. (2002). "Sulfuric acid in the phaeophyte alga Desmarestia munda deters feeding by the sea urchin Strongylocentrotus droebachiensis". Marine Biology. 141 (1): 1–9. doi:10.1007/s00227-002-0809-6.
^ Jump up to: a b Kremser, S.; Thomson, L.W. (2016). "Stratospheric aerosol—Observations, processes, and impact on climate". Review of Geophysics. 54 (2): 278–335. doi:10.1002/2015RG000511.
Jump up ^ Orlando, T. M.; McCord, T. B.; Grieves, G. A. (2005). "The chemical nature of Europa surface material and the relation to a subsurface ocean". Icarus. 177 (2): 528–533. Bibcode:2005Icar..177..528O. doi:10.1016/j.icarus.2005.05.009.
Jump up ^ Jones, Edward M. (1950). "Chamber Process Manufacture of Sulfuric Acid". Industrial and Engineering Chemistry. 42 (11): 2208–2210. doi:10.1021/ie50491a016.
Jump up ^ Chenier, Philip J. (1987). Survey of Industrial Chemistry. New York: John Wiley & Sons. pp. 45–57. ISBN 0-471-01077-4.
Jump up ^ Davenport, William George & King, Matthew J. (2006). Sulfuric acid manufacture: analysis, control and optimization. Elsevier. pp. 8, 13. ISBN 978-0-08-044428-4. Retrieved 23 December 2011.
Jump up ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 653. ISBN 0-08-037941-9.
Jump up ^ Ngo, Christian; Natowitz, Joseph (2016). Our Energy Future: Resources, Alternatives and the Environment. John Wiley & Sons. pp. 418–419. ISBN 9781119213369.
Jump up ^ Pauling, L.C. (1988) General Chemistry, Dover Publications
Jump up ^ Beaumont, JJ; Leveton, J; Knox, K; Bloom, T; McQuiston, T; Young, M; Goldsmith, R; Steenland, NK; Brown, DP; Halperin, WE (1987). "Lung cancer mortality in workers exposed to sulfuric acid mist and other acid mists". J Natl Cancer Inst. 79 (5): 911–21. PMID 3479642.
Jump up ^ Annex to Form D ("Red List"), 11th Edition, January 2007 (p. 4). International Narcotics Control Board. Vienna, Austria.
External links[edit]
Wikimedia Commons has media related to Sulfuric acid.
International Chemical Safety Card 0362
Sulfuric acid at The Periodic Table of Videos (University of Nottingham)
NIOSH Pocket Guide to Chemical Hazards
CDC – Sulfuric Acid – NIOSH Workplace Safety and Health Topic
External Material Safety Data Sheet
Calculators: surface tensions, and densities, molarities and molalities of aqueous sulphuric acid
Sulfuric acid analysis – titration freeware
Process flowsheet of sulfuric acid manufacturing by lead chamber process
[show]
v
t
e
Hydrogen compounds
H3AsO3
H3AsO4
HAt
HSO3F
HBF4
HBr
HBrO
HBrO2
HBrO3
HBrO4
HCl
HClO
HClO2
HClO3
HClO4
HCN
HCNO
H2CrO4/H2Cr2O7
H2CO3
H2CS3
HF
HFΟ
HI
HIO
HIO2
HIO3
HIO4
HMnO4
H2MoO4
HNC
NaHCO3
HNCO
HNO
HNO2
HNO3
H2N2O2
HNO5S
H3NSO3
H2O
H2O2
H2O3
H3PO2
H3PO3
H3PO4
H4P2O7
H5P3O10
H2PtCl6
H2S
H2S2
H2Se
H2SeO3
H2SeO4
H4SiO4
H2SiF6
HSCN
HNSC
H2SO3
H2SO4
H2SO5
H2S2O3
H2S2O6
H2S2O7
H2S2O8
CF3SO3H
H2Te
H2TeO3
H6TeO6
H4TiO4
H2Po
HCo(CO)4
[show]
v
t
e
Sulfur compounds
Al2S3
As2S2
As2S3
As5S2
As4S4
Au2S3
B2S3
BaS
BeS
Bi2S3
Br2S2
CS2
C3S2
CaS
CdS
CeS
SCl2
S2Cl2
CoS
Cr2S3
CuS
D2S
Dy2S3
Er2S3
EuS
SF4
SF6
FeS2
GaS
H2S
HfS2
HgS
InS
LaS
LiS
MgS
MoS3
NiS
SO2
SO3
P4S7
PbS
PbS2
PtS
ReS2
SrS
TlS
SV
SeS2
S2U
WS2
Sb2S3
Sb2S5
Sm2S3
Y2S3
Ag2SO4
SOBr2
CSTe
C2H4S
C2H6S3
C4H4S
CaSO4
C32H66S2
CuFeS2
H2SO4
H2SO3
F2OS
NaHS
K2SO3
O3S3Sb4
Yb2(SO4)3
AlKO8S2
CHCl3S
KSCN
CdSO3
PSCl3
SOCl2
Cs2O4S
Re2S7
Na2S
K2S
H2S2O7
H2SO5
NH5S
HgSO4
K2SO4
RaSO4
SnSO4
SrSO4
Zr(SO4)2
Ti(SO4)2
Tm2(SO4)3
AlNa(SO4)2
Er2(SO4)3
Eu2(SO4)3
CHNS
Co(SCN)2
C2H3SN
PSI3
ZrS2
SiS
CSSe
Chemical formulas
[hide]
v
t
e
Salts and esters of the sulfate ion
H2SO4 He
Li2SO4 BeSO4 B esters
ROSO3−
(RO)2SO2 (NH4)2SO4
[N2H5]HSO4
(NH3OH)2SO4
NOHSO4 HOSO4 F Ne
Na2SO4
NaHSO4 MgSO4 Al2(SO4)3
Al2SO4(OAc)4 Si P SO42−
HSO3HSO4
(HSO4)2 Cl Ar
K2SO4
KHSO4 CaSO4 Sc2(SO4)3 Ti(SO4)2
TiOSO4 VSO4
V2(SO4)3
VOSO4 CrSO4
Cr2(SO4)3 MnSO4
Mn2(SO4)3 FeSO4
Fe2(SO4)3 CoSO4
Co2(SO4)3 NiSO4 CuSO4
Cu2SO4
[Cu(NH3)4(H2O)]SO4 ZnSO4 Ga2(SO4)3 Ge As Se Br Kr
RbHSO4
Rb2SO4 SrSO4 Y2(SO4)3 Zr(SO4)2 Nb Mo Tc Ru Rh PdSO4 Ag2SO4 CdSO4 In2(SO4)3 SnSO4 Sb2(SO4)3 Te I Xe
Cs2SO4 BaSO4 Hf Ta W Re Os Ir Pt Au Hg2SO4
HgSO4 Tl2SO4
Tl2(SO4)3 PbSO4 Bi2(SO4)3 Po At Rn
Fr Ra Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
↓
La Ce2(SO4)3
Ce(SO4)2 Pr2(SO4)3 Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb2(SO4)3 Lu
Ac Th Pa U(SO4)2
UO2SO4 Np Pu Am Cm Bk Cf Es Fm Md No Lr

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Causes severe skin burns and eye damage

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text:
New Palestine High School
From Wikipedia, the free encyclopedia
Jump to navigation Jump to search
New Palestine High School
Address
4485 S Victory Dr
New Palestine, (Hancock County), Indiana 46163
United States
Coordinates
39°43′07″N 85°53′26″W / 39.718653°N 85.890571°W / 39.718653; -85.890571Coordinates: 39°43′07″N 85°53′26″W / 39.718653°N 85.890571°W / 39.718653; -85.890571
Information
Type
Public high school
Established
1919
School district
Southern Hancock County Community School Corporation
Principal
Keith Fessler
Faculty
61
Grades
9-12
Enrollment
1,137 [1] (2016-2017)
Student to teacher ratio
22.9:1
Color(s)
Athletics conference
Hoosier Heritage Conference
Mascot
Dragon
Newspaper
Crimson Messenger
Yearbook
Avalon
Vision statement
Achieving our personal best in pursuit of excellence
Website
NPHS
New Palestine High School is a rural public high school (grades 9-12) located in New Palestine, Indiana, managed by the Community School Corporation of Southern Hancock County. As of the 2016-2017 school year it serves 1,137 students and employed 61 faculty members.[2]
New Palestine High School has repeatedly been named a Four-Star School by the Indiana Department of Education.[3] In 2012, it was named a National Blue Ribbon School by the United States Department of Education. [4]
Contents
[hide]
1 Athletics and extracurricular activities
2 Notable graduates
3 References
4 External links
Athletics and extracurricular activities[edit]
New Palestine High School is part of the Hoosier Heritage Conference and offers soccer, tennis, cross country, American football, golf, volleyball, basketball, gymnastics, swimming, wrestling, baseball, softball and track and field. In addition, New Palestine High School has award-winning marching band, choir and journalism programs.
New Palestine has won six Indiana High School Athletic Association state championships: in baseball (2004), softball (2004, 2008, 2009, 2017), and football (2014). In addition, the Dragons have had State Finals appearances in baseball (2003, 2004), softball (2003, 2004, 2006, 2008, 2009, 2017) and football (1990, 2014, 2015), as well as team state finals appearances in golf and gymnastics. NPHS also has eight individual IHSAA state championships, won by Kyle Ulrey (1), Chad Red (4) and Alec White (1) in wrestling, as well as Lori Swegman (1, 880 yard run) and Samuel Voelz (1, 800 meter run) in track and field.[5]
A former football coach, Marvin Shepler, is a member of the Indiana Football Hall of Fame and a former basketball coach, Fred Keesling, is a member of the Indiana Basketball Hall of Fame, and the IHSAA's Mental Attitude Award for boys' golf is named for him. Former football player, Todd Yoder, was also inducted in the Indiana Football Hall of Fame in 2015.
Notable graduates[edit]
Angela Ahrendts, Senior Vice President, Retail and Online Stores, Apple Inc.; former CEO of Burberry
Todd Yoder, former professional football player
Kent Raymond, former professional basketball player
References[edit]
Jump up ^ http://www.ihsaa.org/Portals/0/ihsaa/documents/schools/201617SchoolDirectory.pdf
Jump up ^ "New Palestine High School". Indiana Department of Education. Retrieved March 1, 2017.
Jump up ^ http://www.doe.in.gov/student-services/accreditation/four-star-schools
Jump up ^ "Yorktown, New Palestine High Schools make Blue Ribbon list". 13 WTHR Indianapolis. 2012-09-07. Retrieved 2017-03-01.
Jump up ^ http://www.ihsaa.org/dnn/Sports/Girls/Track/IndividualStateChampions1/tabid/680/Default.aspx#2165177-880-yard-run
External links[edit]
New Palestine High School website
PublicSchoolReview
New Palestine Sports
List of high schools in Indiana
Hoosier Heritage Conference
hide
v
t
e
Hoosier Heritage Conference
Delta
Greenfield Central
Mount Vernon (Fortville)
New Castle
New Palestine
Pendleton Heights
Shelbyville
Yorktown
Retrieved from "https://en.wikipedia.org/w/index.php?title=New_Palestine_High_School&oldid=842133394"
Categories:
Public high schools in Indiana
Schools in Hancock County, Indiana
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This page was last edited on 20 May 2018, at 13:02.
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