laboratory preparation of hydrochloric acid
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Perchloric acid is the inorganic compound with the formula HClO4. Usually encountered as an aqueous solution, this colourless compound is a strong acid comparable in strength to sulfuric and nitric acids, as well as a powerful oxidizing agent. It is useful for preparing perchlorate salts, especially ammonium perchlorate, an important rocket fuel. Perchloric acid is also dangerously corrosive and readily forms explosive mixtures.
Perchloric acid is produced industrially by two routes. The traditional method exploits the very high aqueous solubility of sodium perchlorate (209 g/100 mL of water at room temperature). Treatment of such solutions with hydrochloric acid gives perchloric acid, precipitating solid sodium chloride:
- NaClO4 + HCl â†’ NaCl + HClO4
The concentrated acid can be purified by distillation. The alternative route, which is more direct and involves no salts, entails anodic oxidation of aqueous chlorine at a platinum electrode.
Treatment of barium perchlorate with sulfuric acid precipitates barium sulfate, leaving perchloric acid. It also can be made by mixing nitric acid with ammonium perchlorate. The reaction gives nitrous oxide and perchloric acid due to a concurrent reaction involving the ammonium ion.
Anhydrous perchloric acid is an oily liquid at room temperature. It forms at least five hydrates, several of which have been characterized crystallographically. These solids consist of the perchlorate anion linked via hydrogen bonds to H2O and H3O+ centers Perchloric acid forms an azeotrope with water, consisting of about 72.5% perchloric acid. This form of the acid is stable indefinitely and is commercially available. Such solutions are hygroscopic. Thus, if left open to the air, concentrated perchloric acid dilutes itself by absorbing water from the air.
Dehydration of perchloric acid gives the anhydride dichlorine heptoxide, which is even more dangerous:
- 2 HClO4 + P4O10â†’ Cl2O7 + "H2P4O11"
Perchloric acid is mainly produced as a precursor to ammonium perchlorate, which is used as rocket fuel. The growth in rocketry has led to increased production of perchloric acid. Several million kilograms are produced annually.
As an acid
Perchloric acid, a superacid, is one of the strongest BrÃ¸nsted-Lowry acids. Its pKa is âˆ’10. It provides strong acidity without interference from potential nucleophiles such as sulfate or chloride that complicate the use of sulfuric and hydrochloric acids. Other acids of noncoordinating anions, such as fluoroboric acid and hexafluorophosphoric acid are susceptible to hydrolysis, whereas perchloric acid is not. Despite hazards associated with the explosiveness of its salts, the acid is often preferred in certain syntheses. For similar reasons, it is a useful eluent in ion-exchange chromatography.
It is also used for electropolishing/etching of aluminum, molybdenum, and other metals.
Anhydrous and monohydrated perchloric acid are explosive, but the usual aqueous solutions are stable in the absence of organic compounds. It is very corrosive to skin and eyes. Upon contact with concentrated perchloric acid, organic materials such as cloth and wood ignite. Salts of perchloric acid are also powerful oxidizers that can be explosive. Perchlorate salts tend to be more stable than their chlorate counterparts, which has led to their increased use in pyrotechnic compositions due to safety concerns.
Because of these hazards, perchloric acid is usually handled under fume hoods with wash-down and air scrubbing capabilities, which are not available on standard laboratory fume hoods. The crystalline form of the acid, which is explosive and shock sensitive, can precipitate on hood surfaces; washing down the hood interior solves this problem.
O'Connor Electro-Plating Company Disaster
On February 20, 1947, in Los Angeles California, 17 people were killed and 150 injured when a bath, consisting of over 1000 litres of 75% Perchloric Acid and 25% Acetic Anhydride by volume, exploded. The plant, 25 other buildings and 40 automobiles were obliterated and 250 nearby homes were damaged.
The bath was being used to electro-polish aluminum furniture, and despite knowing the explosive dangers of the bath, the plant chemist Robert M. McGee allowed production to continue after the refrigeration system, to keep the batch cool, failed. In addition, organic compounds were added to the overheating bath when an iron rack was replaced with one coated with cellulose acetobutyrate (Tenit-2 plastic). A few minutes later the bath exploded.
Formic acid (also called methanoic acid) is the simplest carboxylic acid. Its chemical formula is HCOOH or HCO2H. It is an important intermediate in chemical synthesis and occurs naturally, most notably in the venom of bee and ant stings. In fact, its name comes from the Latin word for ant, formica, referring to its early isolation by thedistillation of ant bodies. Esters, salts, and the anion derived from formic acid are referred to as formates.
Formic acid is miscible with water and most polar organicsolvents, and somewhat soluble in hydrocarbons. In hydrocarbons and in the vapor phase, it consists of hydrogen-bonded dimers rather than individual molecules. Owing to its tendency to hydrogen-bond, gaseous formic acid does not obey the ideal gas law. Solid formic acid (two polymorphs) consists of an effectively endless network of hydrogen-bonded formic acid molecules. This relatively complicated compound also forms a low-boiling azeotrope with water (22.4%) and liquid formic acid also tends to supercool.
In nature, it is found in the stings and bites of many insects of the order Hymenoptera, mainly ants. Because of its abundance in their diet, giant anteaters (unlike most mammals) do not produce hydrochloric acid for their gastric acid.
In 1988, the worldwide capacity for producing this compound was 330,000 tonnes/annum. It is commercially available in solutions of various concentrations between 85 and 99 w/w %.
From methyl formate and formamide
- CH3OH + CO â†’ HCO2CH3
In industry, this reaction is performed in the liquid phase at elevated pressure. Typical reaction conditions are 80 Â°C and 40 atm. The most widely-used base is sodium methoxide. Hydrolysis of the methyl formate produces formic acid:
- HCO2CH3 + H2O â†’ HCO2H + CH3OH
Efficient hydrolysis of methyl formate requires a large excess of water. Some routes proceed indirectly by first treating the methyl formate with ammonia to give formamide, which is then hydrolyzed with sulfuric acid:
- HCO2CH3 + NH3â†’ HC(O)NH2 + CH3OH
- 2 HC(O)NH2 + 2 H2O + H2SO4â†’ 2HCO2H + (NH4)2SO4
This approach suffers from the need to dispose of the ammonium sulfate byproduct. This problem has led some manufacturers to develop energy efficient means for separating formic acid from the large excess amount of water used in direct hydrolysis. In one of these processes (used by BASF) the formic acid is removed from the water via liquid-liquid extraction with an organic base.
By-product of acetic acid production
A significant amount of formic acid is produced as a byproduct in the manufacture of other chemicals. At one time, acetic acid was produced on a large scale by oxidation of alkanes, via a process that cogenerates significant formic acid. This oxidative route to acetic acid is declining in importance, so that the aforementioned dedicated routes to formic acid have become more important.
Hydrogenation of carbon dioxide
The catalytic hydrogenation of CO2 has long been studied. This reaction can be conducted homogeneously.
In the laboratory, formic acid can be obtained by heating oxalic acid in anhydrous glycerol and extraction by steam distillation. Another preparation (which must be performed under a fume hood) is the acid hydrolysis of ethyl isonitrile (C2H5NC) using HCl solution.
- C2H5NC + 2 H2O â†’ C2H5NH2 + HCO2H
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Answers:I will answer because I know, but you should submit your attempt for correction,if you really want to learn. a) 100 cm3 measuring cylinder, or for better accuracy 4 x 20 cm3 pipette. b) A gas syringe c) Burette d) test tube e) watch glass f) test tube. or watch glass depending on it's size. A pipette would be most accurate, however 25.5 cm3 is an unlikely number for a pipette So the next best alternative is the burette. Not available in your list is a gravimetric option, the most accurate.
Answers:Here's what I think the answer is: Acyl chloride's functional group is -COCl while carboxylic acid has -COOH Acyl chloride is very reactive because of the chlorine on the C=O (carbonyl carbon) which makes the structure very electronegative. Electronegativity is the ability of an atom to attract electron density towards itself (thus making it a very good electrophile). Carboxylic acids are polar and forms hydrogen bonding. If you compare the reactivity of this two, then acyl chloride is the better reagent because of the chlorine which in a way activates the the carbonyl group due to increased electronegativity. Hope this helps.. Ciao
Answers:not so difficult. First assume you need to prepare one liter. So you will need one mole of hydrochloric acid. From the molar weight Thats 36,5 grams of pure HCl. Now 36.5 divided by 1.18 is 30.93 mL of hypothetically pure HCl, but actually it comes as a 36% solution, so 30.93 * 100/36 = .... thats it!