Caffeine: How much is too much?
By Mayo Clinic staff
If you rely on caffeine to wake you up and keep you going, you aren't alone. Caffeine stimulates the central nervous system, alleviating fatigue and increasing wakefulness.
When to consider cutting back
For most people, moderate doses of caffeine — 200 to 300 milligrams (mg), or about two to four cups of brewed coffee a day — aren't harmful. But some circumstances may warrant limiting or even ending your caffeine routine. Read on to see if any of these apply to you.
'Grande' is your middle name
Though moderate caffeine intake isn't likely to cause harm, too much can noticeably affect your health. Heavy daily caffeine use — more than 500 to 600 mg a day, or about four to seven cups of coffee — can cause:
* Insomnia
* Nervousness
* Restlessness
* Irritability
* Nausea or other gastrointestinal problems
* Fast or irregular heartbeat
* Muscle tremors
* Headaches
* Anxiety
Even a little makes you jittery
Some people are more sensitive to caffeine than are others. If you're susceptible to the effects of caffeine, just small amounts — even one cup of coffee or tea — may prompt unwanted effects, such as anxiety, restlessness, irritability and sleep problems. How you react to caffeine may be determined in part by how much caffeine you're used to drinking. So, people who don't regularly consume caffeine tend to be more sensitive to its negative effects. Other factors may include body mass, age, smoking habits, drug or hormone use, stress and health conditions such as anxiety disorders. Sex may even play a role: Research suggests that men are more susceptible to caffeine than are women.
Caffeine: How much is too much?
You're not getting enough sleep
Most adults need seven to eight hours of sleep each night. But caffeine can interfere with this much-needed sleep. Chronically losing sleep — whether it's from work, travel, stress or too much caffeine — results in sleep deprivation. Sleep loss is cumulative, and even small nightly decreases can add up and disturb your daytime alertness and performance.
Using caffeine to mask sleep deprivation can create an unwelcome cycle. For example, you drink caffeinated beverages because you have trouble staying awake during the day. But the caffeine keeps you from falling asleep at night, shortening the length of time you sleep. Caffeine can also increase the number of times you wake up during the night and interfere with deep sleep, making your night less restful. You wake up tired the next day and reach for your morning jolt of Java.
The best way to break this cycle is to reduce the caffeine and add more hours of quality sleep each day. Try to avoid caffeinated beverages eight hours before your desired bedtime. Your body doesn't store caffeine, but it takes many hours for it to eliminate the stimulant and its effects.
You're taking certain medications and supplements
Certain medications and herbal supplements negatively interact with caffeine. Here are some examples.
* Some antibiotics. Ciprofloxacin (Cipro) and norfloxacin (Noroxin) — types of antibacterial medications — can interfere with the breakdown of caffeine. This may increase the length of time caffeine remains in your body and amplify its unwanted effects.
* Theophylline (Theo-24, Uniphyl, others). This medication — which opens up bronchial airways by relaxing the surrounding muscles (a bronchodilator) — tends to have some caffeine-like effects. Taking it along with caffeinated foods and beverages may increase the concentration of theophylline in your blood. This can cause ill effects, such as nausea, vomiting and heart palpitations. If you take theophylline, your doctor may advise you to avoid caffeine.
* Ephedra (ma-huang). This herbal dietary supplement increases your risk of heart attack, stroke, seizures and death. Combined with caffeine, it becomes especially risky. The Food and Drug Administration has banned ephedra because of health concerns. The ban applies to dietary supplements but not herbal teas, which may still contain this herb.
Talk to your doctor or pharmacist about whether caffeine might affect your prescription. He or she can say whether you need to reduce or eliminate caffeine from your diet.
How to curb your caffeine habit
Whether it's for one of the reasons above — or because you want to trim your spending on pricey coffee drinks — cutting back on caffeine can be challenging. Too abrupt a decrease in caffeine can cause caffeine withdrawal with signs and symptoms such as headaches, fatigue, irritability and nervousness. Fortunately, these symptoms usually resolve after several days.
To change your caffeine habit more gradually, try these tips:
* Keep tabs. Start paying attention to how much caffeine you're getting from foods and beverages. It may be more than you think. Read labels carefully. Even then, your estimate may be a little low because not all foods list caffeine. Chocolate, which has a small amount, doesn't.
* Cut back. But do it gradually. For example, drink one less can of soda or drink a smaller cup of coffee each day. This will help your body get used to the lower levels of caffeine and thereby lessen the withdrawal effects.
* Go decaf. Most decaffeinated beverages look and taste the same as their caffeinated counterparts.
* Make it quick or herbal. When making tea, brew it for less time. This cuts down on its caffeine content. Or choose herbal teas, which don't contain the stimulant.
* Check the bottle. Some over-the-counter pain relievers contain caffeine — as much as 130 mg of caffeine in one dose. Look for caffeine-free pain relievers instead.
Wednesday, August 26, 2009
Sunday, August 9, 2009
Sodium metal
Characteristics
At room temperature, sodium metal is soft enough that it can be cut with a knife. In air, the bright silvery luster of freshly exposed sodium will rapidly tarnish. The density of alkali metals generally increases with increasing atomic number, but sodium is denser than potassium.
[edit] Chemical properties
Sodium metal (approx 10g) under oil
Compared with other alkali metals, sodium is generally less reactive than potassium and more reactive than lithium,[2] in accordance with "periodic law": for example, their reaction in water, chlorine gas, etc.;
Sodium reacts exothermically with water: small pea-sized pieces will bounce across the surface of the water until they are consumed by it, whereas large pieces will explode. While sodium reacts with water at room temperature, the sodium piece melts with the heat of the reaction to form a sphere, if the reacting sodium piece is large enough. The reaction with water produces very caustic sodium hydroxide (lye) and highly flammable hydrogen gas. These are extreme hazards (see Precautions section below). When burned in air, sodium forms sodium peroxide Na2O2, or with limited oxygen, the oxide Na2O (unlike lithium, the nitride is not formed). If burned in oxygen under pressure, sodium superoxide NaO2 will be produced. In chemistry, most sodium compounds are considered soluble but nature provides examples of many insoluble sodium compounds such as the feldspars. There are other insoluble sodium salts such as sodium bismuthate NaBiO3, sodium octamolybdate Na2Mo8O25• 4H2O, sodium thioplatinate Na4Pt3S6, sodium uranate Na2UO4. Sodium meta-antimonate's 2NaSbO3•7H2O solubility is 0.3g/L as is the pyro form Na2H2Sb2O7• H2O of this salt. Sodium metaphosphate NaPO3 has a soluble and an insoluble form.[3]
[edit] Compounds
See also: Category:Sodium compounds
Sodium compounds are important to the chemical, glass, metal, paper, petroleum, soap, and textile industries. Hard soaps are generally sodium salt of certain fatty acids (potassium produces softer or liquid soaps).[4]
The sodium compounds that are the most important to industries are common salt (NaCl), soda ash (Na2CO3), baking soda (NaHCO3), caustic soda (NaOH), sodium nitrate (NaNO3), di- and tri-sodium phosphates, sodium thiosulfate (hypo, Na2S2O3 • 5H2O), and borax (Na2B4O7 • 10H2O).[4]
Activated carbon
Activated carbon, also called activated charcoal or activated coal, is a form of carbon that has been processed to make it extremely porous and thus to have a very large surface area available for adsorption or chemical reactions. [1] The word activated in the name is sometimes substituted by active. Due to its high degree of microporosity, just one gram of activated carbon has a surface area of approximately 500 m² (or about 2.17 tennis courts), as determined typically by nitrogen gas adsorption. Sufficient activation for useful applications may come solely from the high surface area, though further chemical treatment often enhances the adsorbing properties of the material. Activated carbon is usually derived from charcoal.
Production
Activated carbon is produced from carbonaceous source materials like nutshells, wood and coal. It can be produced by one of the following processes:
1. Physical reactivation: The precursor is developed into activated carbons using gases. This is generally done by using one or a combination of the following processes:
o Carbonization: Material with carbon content is pyrolyzed at temperatures in the range 600-900 °C, in absence of air (usually in inert atmosphere with gases like argon or nitrogen)
o Activation/Oxidation: Raw material or carbonised material is exposed to oxidizing atmospheres (carbon dioxide, oxygen, or steam) at temperatures above 250 °C, usually in the temperature range of 600-1200 °C.
2. Chemical activation: Prior to carbonization, the raw material is impregnated with certain chemicals. The chemical is typically an acid, strong base, or a salt (phosphoric acid, potassium hydroxide, sodium hydroxide, zinc chloride, respectively). Then, the raw material is carbonized at lower temperatures (450-900 °C). It is believed that the carbonization / activation step proceeds simultaneously with the chemical activation. This technique can be problematic in some cases, because, for example, zinc trace residues may remain in the end product. However, chemical activation is preferred over physical activation owing to the lower temperatures and shorter time needed for activating material.
Properties
A gram of activated carbon can have a surface area in excess of 500 m², with 1500 m² being readily achievable[2]. Carbon aerogels, while more expensive, have even higher surface areas, and are used in special applications.
Under an electron microscope, the high surface-area structures of activated carbon are revealed. Individual particles are intensely convoluted and display various kinds of porosity; there may be many areas where flat surfaces of graphite-like material run parallel to each other, separated by only a few nanometers or so. These micropores provide superb conditions for adsorption to occur, since adsorbing material can interact with many surfaces simultaneously. Tests of adsorption behaviour are usually done with nitrogen gas at 77 K under high vacuum, but in everyday terms activated carbon is perfectly capable of producing the equivalent, by adsorption from its environment, liquid water from steam at 100 °C and a pressure of 1/10,000 of an atmosphere.
Physically, activated carbon binds materials by Van der Waals force or London dispersion force.
Activated carbon does not bind well to certain chemicals, including alcohols, glycols, ammonia, strong acids and bases, metals and most inorganics, such as lithium, sodium, iron, lead, arsenic, fluorine, and boric acid. Activated carbon does adsorb iodine very well and in fact the iodine number, mg/g, (ASTM D28 Standard Method test) is used as an indication of total surface area.
Activated carbon can be used as a substrate for the application of various chemicals to improve the adsorptive capacity for some inorganic (and problematic organic) compounds such as hydrogen sulfide (H2S), ammonia (NH3), formaldehyde (HCOH), radioisotopes iodine-131 (131I) and mercury (Hg). This property is known as chemisorption.
Soda lime is a mixture of chemicals, used in granular form in closed breathing environments, such as general anaesthesia, submarines, rebreathers and recompression chambers, to remove carbon dioxide from breathing gases to prevent CO2 retention and carbon dioxide poisoning.[1][2]
It is made by treating slaked lime with concentrated sodium hydroxide solution.
[edit] Chemical components
The main components of soda lime are
• Calcium hydroxide, Ca(OH)2 (about 75%),
• Water, H2O (about 20%),
• Sodium hydroxide, NaOH (about 3%), and
• Potassium hydroxide, KOH (about 1%).
[edit] Anesthesia use
While administering general anesthesia, the patient's expired gases which contain carbon dioxide, are passed through an Anaesthetic machine breathing circuit filled with soda lime granules.[1] Medical grade soda lime has indicating dye which changes color when the soda lime loses its carbon dioxide absorbing capacity.
[edit] Undersea use
Exhaled gas must be passed through a "carbon dioxide scrubber" where the carbon dioxide is absorbed before the gas is made available to be breathed again. In rebreathers the scrubber is a part of the breathing loop.[2][3] Color indicating dye was removed from US Navy fleet use in 1996 when it was suspected of releasing chemicals into the circuit.[4] In larger environments, such as recompression chambers or submarines, a fan is used to pass gas through the canister.[2]
[edit] Chemical reaction
The overall reaction is:
CO2 + Ca(OH)2 → CaCO3 + H2O + heat (in the presence of water)
The reaction can be considered as a strong base catalysed, water facilitated reaction.
steps:
1) CO2 + H2O → CO2 (aq) (CO2 dissolves in water - slow and rate determining)
2) CO2 (aq) + NaOH → NaHCO3 (bicarbonate formation at high pH)
3) NaHCO3 + Ca(OH)2 → CaCO3 + H2O + NaOH (NaOH recycled to step 2) - hence a catalyst)
each mole of CO2 (44g) reacted produces one mole of water (18g)
Iron(III) chloride, also called ferric chloride, is an industrial scale commodity chemical compound, with the formula FeCl3. The colour of iron(III) chloride crystals depends on the viewing angle: by reflected light the crystals appear dark green, but by transmitted light they appear purple-red. Anhydrous iron(III) chloride is deliquescent, forming hydrated hydrogen chloride mists in moist air. It is rarely observed in its natural form, mineral molysite, known mainly from some fumaroles.
When dissolved in water, iron(III) chloride undergoes hydrolysis and gives off heat in an exothermic reaction. The resulting brown, acidic, and corrosive solution is used as a coagulant in sewage treatment and drinking water production, and as an etchant for copper-based metals in printed circuit boards. Anhydrous iron(III) chloride is a fairly strong Lewis acid, and it is used as a catalyst in organic synthesis.
•
Chemical and physical properties
Iron(III) chloride has a relatively low melting point and boils at around 315 °C. The vapour consists of the dimer Fe2Cl6 (compare aluminium chloride) which increasingly dissociates into the monomeric FeCl3 (D3h point group molecular symmetry) at higher temperature, in competition with its reversible decomposition to give iron(II) chloride and chlorine gas.[2]
[edit] Reactions
Iron(III) chloride is a moderately strong Lewis acid, forming adducts with Lewis bases such as triphenylphosphine oxide, e.g. FeCl3(OPPh3)2 where Ph = phenyl.
Iron(III) chloride reacts with other chloride salts to give the yellow tetrahedral FeCl4− ion. Salts of FeCl4− in hydrochloric acid can be extracted into diethyl ether.
When heated with iron(III) oxide at 350 °C, iron(III) chloride gives iron oxychloride, a layered solid and intercalation host.[citation needed]
FeCl3 + Fe2O3 → 3 FeOCl
In the presence of base, alkali metal alkoxides react to give the dimeric complexes:
2 FeCl3 + 6 C2H5OH + 6 NH3 → (Fe(OC2H5)3)2 + 6 NH4Cl
Oxalates react rapidly with aqueous iron(III) chloride to give [Fe(C2O4)3]3−. Other carboxylate salts form complexes, e.g. citrate and tartrate.
Sodium hydroxide (NaOH), also known as lye and caustic soda, is a caustic metallic base. Sodium hydroxide forms a strong alkaline solution when dissolved in a solvent such as water. However, only the hydroxide ion is basic. It is used in many industries, mostly as a strong chemical base in the manufacture of pulp and paper, textiles, drinking water, soaps and detergents and as a drain cleaner. Worldwide production in 1998 was around 45 million tonnes. Sodium hydroxide is a common base in chemical laboratories.
Pure sodium hydroxide is a white solid; available in pellets, flakes, granules and as a 50% saturated solution. It is hygroscopic and readily absorbs water from the air, so it should be stored in an airtight container. It is very soluble in water with liberation of heat. It also dissolves in ethanol and methanol, though it exhibits lower solubility in these solvents than potassium hydroxide. It is insoluble in ether and other non-polar solvents. A sodium hydroxide solution will leave a yellow stain on fabric and paper.
Physical properties
Δ H° dissolution for diluted aqueous -44.45 kJ / mol;
From aqueous solutions at 12.3-61.8°C, it crystallizes in monohydrate, with a melting point 65.1 °C and density of 1.829 g/cm 3;
Δ H° form -734.96 kJ / mol;
Monohydrate from -28 to -24°C;
Heptahydrate from -24 to -17.7°C;
Pentahydrate from -17.7 to -5.4°C;
Tetrahydrate (α- changed), at -5 , 4 - 12.3°C Also know metastable β- NaOH 4* H2O. Which above 61.8°C are crystallized.
[edit] Chemical properties
Sodium hydroxide is completely ionic, containing sodium cations and hydroxide anions. The hydroxide anion makes sodium hydroxide a strong base which reacts with acids to form water and the corresponding salts, e.g., 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) + H3O+(aq) → 2H2O
This type of reaction with a strong acid, releases heat, and hence is referred to as exothermic. Such acid-base reactions can also be used for titrations, which is a common method to determine the concentration of acids. Another type of reaction that sodium hydroxide is involved in is with acidic oxides. The reaction of carbon dioxide has already been mentioned, but other acidic oxides such as sulfur dioxide (SO2) also react completely. Such reactions are often used to "scrub" harmful acidic gases (like SO2 and H2S) and prevent their release into the atmosphere.
2NaOH + CO2 → Na2CO3 + H2O
Sodium hydroxide slowly reacts with glass to form sodium silicate, so glass joints and stopcocks exposed to NaOH have a tendency to "freeze". Flasks and glass-lined chemical reactors are damaged by long exposure to hot sodium hydroxide, and the glass becomes frosted. Sodium hydroxide does not attack iron since iron does not have amphoteric properties. A few transition metals, however, may react with sodium hydroxide in a vigorous way.
In 1986 an aluminium road tanker in the UK was mistakenly used to transport 25% sodium hydroxide solution, causing pressurization of the contents and damage to the tanker. The pressurization was due to the hydrogen gas which is produced in the reaction between sodium hydroxide and aluminium:
2Al(s) + 6NaOH(aq) → 3H2(g) + 2Na3AlO3(aq)
Unlike NaOH, the hydroxides of most metals are insoluble, and therefore sodium hydroxide can be used to precipitate metal hydroxides. One such hydroxide is aluminium hydroxide, used as a gelatinous floc to filter out particulate matter in water treatment. Aluminium hydroxide is prepared at the treatment plant from aluminium sulfate by reacting with NaOH. This reaction is highly profitable, and is hence an important synthesis reaction.
Sodium hydroxide reacts readily with carboxylic acids to form their salts and is even a strong enough base to form salts with phenols. NaOH can be used for the base-driven hydrolysis of esters (as in saponification), amides and alkyl halides. However, the limited solubility of NaOH in organic solvents means that the more soluble KOH is often preferred.
Basic hydrolysis of an ester
Iron(III) chloride is a mild oxidising agent, for example capable of oxidising copper(I) chloride to copper(II) chloride. Reducing agents such as hydrazine convert iron(III) chloride to complexes of iron(II).
[edit] Structure
Iron(III) chloride adopts the BiI3 structure, which features octahedral Fe(III) centres interconnected by two-coordinate chloride ligands.
[edit] Preparation and production
Anhydrous iron(III) chloride may be prepared by union of the elements:[3]
2 Fe(s) + 3 Cl2(g) → 2 FeCl3(s)
Solutions of iron(III) chloride are produced industrially both from iron and from ore, in a closed-loop process.
1. Dissolving pure iron in a solution of iron(III) chloride
Fe(s) + 2 FeCl3(aq) → 3 FeCl2(aq)
2. Dissolving iron ore in hydrochloric acid
Fe3O4(s) + 8 HCl(aq) → FeCl2(aq) + 2 FeCl3(aq) + 4 H2O
3. Upgrading the iron(II) chloride with chlorine
2 FeCl2(aq) + Cl2(g) → 2 FeCl3(aq)
Alternatively, iron(II) chloride can be oxidised with sulfur dioxide:
32 FeCl2 + 8 SO2 + 32 HCl → 32 FeCl3 + S8 + 16 H2O
Like many other hydrated metal chlorides, hydrated iron(III) chloride can be converted to the anhydrous salt by refluxing with thionyl chloride.[4] The hydrate cannot be converted to anhydrous iron(III) chloride by only heat, as instead HCl is evolved and iron oxychloride forms.
Silver nitrate, also known as lunar caustic, is a soluble chemical compound with chemical formula AgNO3. This compound is a versatile precursor to many other silver compounds, such as those used in photography. Comparatively, it is far less sensitive to light than the halides. It is called lunar caustic because silver was called luna by the ancient alchemists.[1]
In solid silver nitrate, the silver ions are three-coordinated in a trigonal planar arrangement.[2]
Preparation
Silver nitrate can be prepared by simply reacting silver, for example a silver buillion or silver foil, with nitric acid. The equation is: Ag + 2HNO3 → AgNO3 + NO2 + H2O
This must be performed under a fume hood[3]
[edit] Uses
[edit] Precursor to other silver compounds
Silver nitrate is the least expensive salt of silver; it offers several other advantages as well. It is non-hygroscopic, in contrast to silver fluoroborate and silver perchlorate. It is relatively stable to light. Finally it dissolves in numerous solvents, including water. The nitrate can be easily replaced by other ligands, rendering AgNO3 versatile. Treatment with solutions of halide ions gives a precipitate of AgX (X = Cl, Br, I). When making photographic film, silver nitrate is treated with halide salts of sodium or potassium to form insoluble silver halide in situ in photographic gelatin, which is then applied to strips of tri-acetate or polyester. Similarly, silver nitrate is used to prepare some silver-based explosives, such as the fulminate, azide, or acetylide, through a precipitation reaction.
Treatment of silver nitrate with base gives dark grey silver oxide:[4]
2 AgNO3 + 2 NaOH → Ag2O + 2 NaNO3 + H2O
[edit] Halide abstraction
The silver cation quickly and effectively irreversibly reacts with halide anions to produce the insoluble silver halide, which is usually a creamy precipitate. This reaction is commonly used in inorganic chemistry to abstract the halide as the insoluble silver salt:
Ag+ (aq) + X- (aq) → AgX (s) (X = Cl, Br, I)
Other silver salts with non-coordinating anions, namely silver tetrafluoroborate and silver hexafluorophosphate are used for more demanding applications.
Similarly, this reaction is used in analytical chemistry to confirm the presence of chloride, bromide, or iodide ions can be tested by adding silver nitrate solution. Samples are typically acidifed with dilute nitric acid to remove interfering ions, e.g. carbonate ions and sulfide ions. This step avoids confusion of silver sulfide or silver carbonate precipitates with that of silver halides. The color of precipitate varies with the halide: white (silver chloride), pale yellow/cream (silver bromide), yellow (silver iodide). AgBr and especially AgI photo-decompose to the metal, as evidence by a grayish color on exposed samples.
[edit] Organic synthesis
Silver nitrate is used in many ways in organic synthesis, e.g. for deprotection and oxidations. Ag+ binds alkenes reversibly, and silver nitrate has been used to separate mixtures of alkenes by selective absorption. The resulting adduct can be decomposed with ammonia to release the free alkene.[5]
Iron(II) sulfate (Iron(II) sulphate) or ferrous sulfate (ferrous sulphate) is the chemical compound with the formula (FeSO4), known since ancient times as copperas. It is most commonly encountered as the blue-green heptahydrate.
edit] Hydrates
Iron(II) sulfate can be found in various states of hydration, and several of these forms exist in nature.
• FeSO4•H2O (mineral: szomolnokite, relatively rare)
• FeSO4•4H2O (mineral: rozenite, white, relatively common, may be dehydratation product of melanterite)
• FeSO4•5H2O (mineral: siderotil, relatively rare)
• FeSO4•6H2O (mineral: ferrohexahydrite, relatively rare)
• FeSO4•7H2O (mineral: melanterite, blue, relatively common)
At 90°C, the heptahydrate, also called green vitriol or copperas, loses water to form the colorless monohydrate, In its anhydrous, crystalline state, its standard enthalpy of formation is ΔfH°solid = -928.4 kJ.mol-1 and its standard molar entropy is S°solid = 107.5 J.K-1.mol-1. All mentioned mineral forms are connected with oxidation zones of Fe-bearing ore beds (pyrite, marcasite, chalcopyrite etc.) and related environments (like coal fire sites). Many undergo rapid dehydratation and sometimes oxidation.
[edit] Production and reactions
In the finishing of steel prior to plating or coating, the steel sheet or rod is passed through pickling baths of sulfuric acid. This treatment produces large quantities of iron(II) sulfate as a by-product.[1]
Fe + H2SO4 → FeSO4 + H2
Another source of large amounts results from the production of titanium dioxide from ilmenite via the sulfate process. Ferrous sulfate is also prepared commercially by oxidation of pyrite:
2 FeS2 + 7 O2 + 2 H2O → 2 FeSO4 + 2 H2SO4
[edit] Reactions
On heating, iron(II) sulfate first loses its water of crystallization and the original green crystals are converted into a dirty-yellow anhydrous solid. When further heated, the anhydrous material releases sulfur dioxide and white fumes of sulfur trioxide, leaving a reddish-brown iron(III) oxide. Decomposition of iron(II) sulfate begins at about 480 °C.
2 FeSO4 → Fe2O3 + SO2 + SO3
Like all iron(II) salts, iron(II) sulfate is reducing agent. For example, it reduces nitric acid to nitrogen oxide and chlorine to chloride:
6 Fe2+ + 6 H+ + 2 HNO3 → 6 Fe3+ + 4 H2O + 2 NO
2 Fe2+ + Cl2 → 2 Fe3+ + 2 Cl-
[edit] Uses
Characteristics
At room temperature, sodium metal is soft enough that it can be cut with a knife. In air, the bright silvery luster of freshly exposed sodium will rapidly tarnish. The density of alkali metals generally increases with increasing atomic number, but sodium is denser than potassium.
[edit] Chemical properties
Sodium metal (approx 10g) under oil
Compared with other alkali metals, sodium is generally less reactive than potassium and more reactive than lithium,[2] in accordance with "periodic law": for example, their reaction in water, chlorine gas, etc.;
Sodium reacts exothermically with water: small pea-sized pieces will bounce across the surface of the water until they are consumed by it, whereas large pieces will explode. While sodium reacts with water at room temperature, the sodium piece melts with the heat of the reaction to form a sphere, if the reacting sodium piece is large enough. The reaction with water produces very caustic sodium hydroxide (lye) and highly flammable hydrogen gas. These are extreme hazards (see Precautions section below). When burned in air, sodium forms sodium peroxide Na2O2, or with limited oxygen, the oxide Na2O (unlike lithium, the nitride is not formed). If burned in oxygen under pressure, sodium superoxide NaO2 will be produced. In chemistry, most sodium compounds are considered soluble but nature provides examples of many insoluble sodium compounds such as the feldspars. There are other insoluble sodium salts such as sodium bismuthate NaBiO3, sodium octamolybdate Na2Mo8O25• 4H2O, sodium thioplatinate Na4Pt3S6, sodium uranate Na2UO4. Sodium meta-antimonate's 2NaSbO3•7H2O solubility is 0.3g/L as is the pyro form Na2H2Sb2O7• H2O of this salt. Sodium metaphosphate NaPO3 has a soluble and an insoluble form.[3]
[edit] Compounds
See also: Category:Sodium compounds
Sodium compounds are important to the chemical, glass, metal, paper, petroleum, soap, and textile industries. Hard soaps are generally sodium salt of certain fatty acids (potassium produces softer or liquid soaps).[4]
The sodium compounds that are the most important to industries are common salt (NaCl), soda ash (Na2CO3), baking soda (NaHCO3), caustic soda (NaOH), sodium nitrate (NaNO3), di- and tri-sodium phosphates, sodium thiosulfate (hypo, Na2S2O3 • 5H2O), and borax (Na2B4O7 • 10H2O).[4]
Activated carbon
Activated carbon, also called activated charcoal or activated coal, is a form of carbon that has been processed to make it extremely porous and thus to have a very large surface area available for adsorption or chemical reactions. [1] The word activated in the name is sometimes substituted by active. Due to its high degree of microporosity, just one gram of activated carbon has a surface area of approximately 500 m² (or about 2.17 tennis courts), as determined typically by nitrogen gas adsorption. Sufficient activation for useful applications may come solely from the high surface area, though further chemical treatment often enhances the adsorbing properties of the material. Activated carbon is usually derived from charcoal.
Production
Activated carbon is produced from carbonaceous source materials like nutshells, wood and coal. It can be produced by one of the following processes:
1. Physical reactivation: The precursor is developed into activated carbons using gases. This is generally done by using one or a combination of the following processes:
o Carbonization: Material with carbon content is pyrolyzed at temperatures in the range 600-900 °C, in absence of air (usually in inert atmosphere with gases like argon or nitrogen)
o Activation/Oxidation: Raw material or carbonised material is exposed to oxidizing atmospheres (carbon dioxide, oxygen, or steam) at temperatures above 250 °C, usually in the temperature range of 600-1200 °C.
2. Chemical activation: Prior to carbonization, the raw material is impregnated with certain chemicals. The chemical is typically an acid, strong base, or a salt (phosphoric acid, potassium hydroxide, sodium hydroxide, zinc chloride, respectively). Then, the raw material is carbonized at lower temperatures (450-900 °C). It is believed that the carbonization / activation step proceeds simultaneously with the chemical activation. This technique can be problematic in some cases, because, for example, zinc trace residues may remain in the end product. However, chemical activation is preferred over physical activation owing to the lower temperatures and shorter time needed for activating material.
Properties
A gram of activated carbon can have a surface area in excess of 500 m², with 1500 m² being readily achievable[2]. Carbon aerogels, while more expensive, have even higher surface areas, and are used in special applications.
Under an electron microscope, the high surface-area structures of activated carbon are revealed. Individual particles are intensely convoluted and display various kinds of porosity; there may be many areas where flat surfaces of graphite-like material run parallel to each other, separated by only a few nanometers or so. These micropores provide superb conditions for adsorption to occur, since adsorbing material can interact with many surfaces simultaneously. Tests of adsorption behaviour are usually done with nitrogen gas at 77 K under high vacuum, but in everyday terms activated carbon is perfectly capable of producing the equivalent, by adsorption from its environment, liquid water from steam at 100 °C and a pressure of 1/10,000 of an atmosphere.
Physically, activated carbon binds materials by Van der Waals force or London dispersion force.
Activated carbon does not bind well to certain chemicals, including alcohols, glycols, ammonia, strong acids and bases, metals and most inorganics, such as lithium, sodium, iron, lead, arsenic, fluorine, and boric acid. Activated carbon does adsorb iodine very well and in fact the iodine number, mg/g, (ASTM D28 Standard Method test) is used as an indication of total surface area.
Activated carbon can be used as a substrate for the application of various chemicals to improve the adsorptive capacity for some inorganic (and problematic organic) compounds such as hydrogen sulfide (H2S), ammonia (NH3), formaldehyde (HCOH), radioisotopes iodine-131 (131I) and mercury (Hg). This property is known as chemisorption.
Soda lime is a mixture of chemicals, used in granular form in closed breathing environments, such as general anaesthesia, submarines, rebreathers and recompression chambers, to remove carbon dioxide from breathing gases to prevent CO2 retention and carbon dioxide poisoning.[1][2]
It is made by treating slaked lime with concentrated sodium hydroxide solution.
[edit] Chemical components
The main components of soda lime are
• Calcium hydroxide, Ca(OH)2 (about 75%),
• Water, H2O (about 20%),
• Sodium hydroxide, NaOH (about 3%), and
• Potassium hydroxide, KOH (about 1%).
[edit] Anesthesia use
While administering general anesthesia, the patient's expired gases which contain carbon dioxide, are passed through an Anaesthetic machine breathing circuit filled with soda lime granules.[1] Medical grade soda lime has indicating dye which changes color when the soda lime loses its carbon dioxide absorbing capacity.
[edit] Undersea use
Exhaled gas must be passed through a "carbon dioxide scrubber" where the carbon dioxide is absorbed before the gas is made available to be breathed again. In rebreathers the scrubber is a part of the breathing loop.[2][3] Color indicating dye was removed from US Navy fleet use in 1996 when it was suspected of releasing chemicals into the circuit.[4] In larger environments, such as recompression chambers or submarines, a fan is used to pass gas through the canister.[2]
[edit] Chemical reaction
The overall reaction is:
CO2 + Ca(OH)2 → CaCO3 + H2O + heat (in the presence of water)
The reaction can be considered as a strong base catalysed, water facilitated reaction.
steps:
1) CO2 + H2O → CO2 (aq) (CO2 dissolves in water - slow and rate determining)
2) CO2 (aq) + NaOH → NaHCO3 (bicarbonate formation at high pH)
3) NaHCO3 + Ca(OH)2 → CaCO3 + H2O + NaOH (NaOH recycled to step 2) - hence a catalyst)
each mole of CO2 (44g) reacted produces one mole of water (18g)
Iron(III) chloride, also called ferric chloride, is an industrial scale commodity chemical compound, with the formula FeCl3. The colour of iron(III) chloride crystals depends on the viewing angle: by reflected light the crystals appear dark green, but by transmitted light they appear purple-red. Anhydrous iron(III) chloride is deliquescent, forming hydrated hydrogen chloride mists in moist air. It is rarely observed in its natural form, mineral molysite, known mainly from some fumaroles.
When dissolved in water, iron(III) chloride undergoes hydrolysis and gives off heat in an exothermic reaction. The resulting brown, acidic, and corrosive solution is used as a coagulant in sewage treatment and drinking water production, and as an etchant for copper-based metals in printed circuit boards. Anhydrous iron(III) chloride is a fairly strong Lewis acid, and it is used as a catalyst in organic synthesis.
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Chemical and physical properties
Iron(III) chloride has a relatively low melting point and boils at around 315 °C. The vapour consists of the dimer Fe2Cl6 (compare aluminium chloride) which increasingly dissociates into the monomeric FeCl3 (D3h point group molecular symmetry) at higher temperature, in competition with its reversible decomposition to give iron(II) chloride and chlorine gas.[2]
[edit] Reactions
Iron(III) chloride is a moderately strong Lewis acid, forming adducts with Lewis bases such as triphenylphosphine oxide, e.g. FeCl3(OPPh3)2 where Ph = phenyl.
Iron(III) chloride reacts with other chloride salts to give the yellow tetrahedral FeCl4− ion. Salts of FeCl4− in hydrochloric acid can be extracted into diethyl ether.
When heated with iron(III) oxide at 350 °C, iron(III) chloride gives iron oxychloride, a layered solid and intercalation host.[citation needed]
FeCl3 + Fe2O3 → 3 FeOCl
In the presence of base, alkali metal alkoxides react to give the dimeric complexes:
2 FeCl3 + 6 C2H5OH + 6 NH3 → (Fe(OC2H5)3)2 + 6 NH4Cl
Oxalates react rapidly with aqueous iron(III) chloride to give [Fe(C2O4)3]3−. Other carboxylate salts form complexes, e.g. citrate and tartrate.
Sodium hydroxide (NaOH), also known as lye and caustic soda, is a caustic metallic base. Sodium hydroxide forms a strong alkaline solution when dissolved in a solvent such as water. However, only the hydroxide ion is basic. It is used in many industries, mostly as a strong chemical base in the manufacture of pulp and paper, textiles, drinking water, soaps and detergents and as a drain cleaner. Worldwide production in 1998 was around 45 million tonnes. Sodium hydroxide is a common base in chemical laboratories.
Pure sodium hydroxide is a white solid; available in pellets, flakes, granules and as a 50% saturated solution. It is hygroscopic and readily absorbs water from the air, so it should be stored in an airtight container. It is very soluble in water with liberation of heat. It also dissolves in ethanol and methanol, though it exhibits lower solubility in these solvents than potassium hydroxide. It is insoluble in ether and other non-polar solvents. A sodium hydroxide solution will leave a yellow stain on fabric and paper.
Physical properties
Δ H° dissolution for diluted aqueous -44.45 kJ / mol;
From aqueous solutions at 12.3-61.8°C, it crystallizes in monohydrate, with a melting point 65.1 °C and density of 1.829 g/cm 3;
Δ H° form -734.96 kJ / mol;
Monohydrate from -28 to -24°C;
Heptahydrate from -24 to -17.7°C;
Pentahydrate from -17.7 to -5.4°C;
Tetrahydrate (α- changed), at -5 , 4 - 12.3°C Also know metastable β- NaOH 4* H2O. Which above 61.8°C are crystallized.
[edit] Chemical properties
Sodium hydroxide is completely ionic, containing sodium cations and hydroxide anions. The hydroxide anion makes sodium hydroxide a strong base which reacts with acids to form water and the corresponding salts, e.g., 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) + H3O+(aq) → 2H2O
This type of reaction with a strong acid, releases heat, and hence is referred to as exothermic. Such acid-base reactions can also be used for titrations, which is a common method to determine the concentration of acids. Another type of reaction that sodium hydroxide is involved in is with acidic oxides. The reaction of carbon dioxide has already been mentioned, but other acidic oxides such as sulfur dioxide (SO2) also react completely. Such reactions are often used to "scrub" harmful acidic gases (like SO2 and H2S) and prevent their release into the atmosphere.
2NaOH + CO2 → Na2CO3 + H2O
Sodium hydroxide slowly reacts with glass to form sodium silicate, so glass joints and stopcocks exposed to NaOH have a tendency to "freeze". Flasks and glass-lined chemical reactors are damaged by long exposure to hot sodium hydroxide, and the glass becomes frosted. Sodium hydroxide does not attack iron since iron does not have amphoteric properties. A few transition metals, however, may react with sodium hydroxide in a vigorous way.
In 1986 an aluminium road tanker in the UK was mistakenly used to transport 25% sodium hydroxide solution, causing pressurization of the contents and damage to the tanker. The pressurization was due to the hydrogen gas which is produced in the reaction between sodium hydroxide and aluminium:
2Al(s) + 6NaOH(aq) → 3H2(g) + 2Na3AlO3(aq)
Unlike NaOH, the hydroxides of most metals are insoluble, and therefore sodium hydroxide can be used to precipitate metal hydroxides. One such hydroxide is aluminium hydroxide, used as a gelatinous floc to filter out particulate matter in water treatment. Aluminium hydroxide is prepared at the treatment plant from aluminium sulfate by reacting with NaOH. This reaction is highly profitable, and is hence an important synthesis reaction.
Sodium hydroxide reacts readily with carboxylic acids to form their salts and is even a strong enough base to form salts with phenols. NaOH can be used for the base-driven hydrolysis of esters (as in saponification), amides and alkyl halides. However, the limited solubility of NaOH in organic solvents means that the more soluble KOH is often preferred.
Basic hydrolysis of an ester
Iron(III) chloride is a mild oxidising agent, for example capable of oxidising copper(I) chloride to copper(II) chloride. Reducing agents such as hydrazine convert iron(III) chloride to complexes of iron(II).
[edit] Structure
Iron(III) chloride adopts the BiI3 structure, which features octahedral Fe(III) centres interconnected by two-coordinate chloride ligands.
[edit] Preparation and production
Anhydrous iron(III) chloride may be prepared by union of the elements:[3]
2 Fe(s) + 3 Cl2(g) → 2 FeCl3(s)
Solutions of iron(III) chloride are produced industrially both from iron and from ore, in a closed-loop process.
1. Dissolving pure iron in a solution of iron(III) chloride
Fe(s) + 2 FeCl3(aq) → 3 FeCl2(aq)
2. Dissolving iron ore in hydrochloric acid
Fe3O4(s) + 8 HCl(aq) → FeCl2(aq) + 2 FeCl3(aq) + 4 H2O
3. Upgrading the iron(II) chloride with chlorine
2 FeCl2(aq) + Cl2(g) → 2 FeCl3(aq)
Alternatively, iron(II) chloride can be oxidised with sulfur dioxide:
32 FeCl2 + 8 SO2 + 32 HCl → 32 FeCl3 + S8 + 16 H2O
Like many other hydrated metal chlorides, hydrated iron(III) chloride can be converted to the anhydrous salt by refluxing with thionyl chloride.[4] The hydrate cannot be converted to anhydrous iron(III) chloride by only heat, as instead HCl is evolved and iron oxychloride forms.
Silver nitrate, also known as lunar caustic, is a soluble chemical compound with chemical formula AgNO3. This compound is a versatile precursor to many other silver compounds, such as those used in photography. Comparatively, it is far less sensitive to light than the halides. It is called lunar caustic because silver was called luna by the ancient alchemists.[1]
In solid silver nitrate, the silver ions are three-coordinated in a trigonal planar arrangement.[2]
Preparation
Silver nitrate can be prepared by simply reacting silver, for example a silver buillion or silver foil, with nitric acid. The equation is: Ag + 2HNO3 → AgNO3 + NO2 + H2O
This must be performed under a fume hood[3]
[edit] Uses
[edit] Precursor to other silver compounds
Silver nitrate is the least expensive salt of silver; it offers several other advantages as well. It is non-hygroscopic, in contrast to silver fluoroborate and silver perchlorate. It is relatively stable to light. Finally it dissolves in numerous solvents, including water. The nitrate can be easily replaced by other ligands, rendering AgNO3 versatile. Treatment with solutions of halide ions gives a precipitate of AgX (X = Cl, Br, I). When making photographic film, silver nitrate is treated with halide salts of sodium or potassium to form insoluble silver halide in situ in photographic gelatin, which is then applied to strips of tri-acetate or polyester. Similarly, silver nitrate is used to prepare some silver-based explosives, such as the fulminate, azide, or acetylide, through a precipitation reaction.
Treatment of silver nitrate with base gives dark grey silver oxide:[4]
2 AgNO3 + 2 NaOH → Ag2O + 2 NaNO3 + H2O
[edit] Halide abstraction
The silver cation quickly and effectively irreversibly reacts with halide anions to produce the insoluble silver halide, which is usually a creamy precipitate. This reaction is commonly used in inorganic chemistry to abstract the halide as the insoluble silver salt:
Ag+ (aq) + X- (aq) → AgX (s) (X = Cl, Br, I)
Other silver salts with non-coordinating anions, namely silver tetrafluoroborate and silver hexafluorophosphate are used for more demanding applications.
Similarly, this reaction is used in analytical chemistry to confirm the presence of chloride, bromide, or iodide ions can be tested by adding silver nitrate solution. Samples are typically acidifed with dilute nitric acid to remove interfering ions, e.g. carbonate ions and sulfide ions. This step avoids confusion of silver sulfide or silver carbonate precipitates with that of silver halides. The color of precipitate varies with the halide: white (silver chloride), pale yellow/cream (silver bromide), yellow (silver iodide). AgBr and especially AgI photo-decompose to the metal, as evidence by a grayish color on exposed samples.
[edit] Organic synthesis
Silver nitrate is used in many ways in organic synthesis, e.g. for deprotection and oxidations. Ag+ binds alkenes reversibly, and silver nitrate has been used to separate mixtures of alkenes by selective absorption. The resulting adduct can be decomposed with ammonia to release the free alkene.[5]
Iron(II) sulfate (Iron(II) sulphate) or ferrous sulfate (ferrous sulphate) is the chemical compound with the formula (FeSO4), known since ancient times as copperas. It is most commonly encountered as the blue-green heptahydrate.
edit] Hydrates
Iron(II) sulfate can be found in various states of hydration, and several of these forms exist in nature.
• FeSO4•H2O (mineral: szomolnokite, relatively rare)
• FeSO4•4H2O (mineral: rozenite, white, relatively common, may be dehydratation product of melanterite)
• FeSO4•5H2O (mineral: siderotil, relatively rare)
• FeSO4•6H2O (mineral: ferrohexahydrite, relatively rare)
• FeSO4•7H2O (mineral: melanterite, blue, relatively common)
At 90°C, the heptahydrate, also called green vitriol or copperas, loses water to form the colorless monohydrate, In its anhydrous, crystalline state, its standard enthalpy of formation is ΔfH°solid = -928.4 kJ.mol-1 and its standard molar entropy is S°solid = 107.5 J.K-1.mol-1. All mentioned mineral forms are connected with oxidation zones of Fe-bearing ore beds (pyrite, marcasite, chalcopyrite etc.) and related environments (like coal fire sites). Many undergo rapid dehydratation and sometimes oxidation.
[edit] Production and reactions
In the finishing of steel prior to plating or coating, the steel sheet or rod is passed through pickling baths of sulfuric acid. This treatment produces large quantities of iron(II) sulfate as a by-product.[1]
Fe + H2SO4 → FeSO4 + H2
Another source of large amounts results from the production of titanium dioxide from ilmenite via the sulfate process. Ferrous sulfate is also prepared commercially by oxidation of pyrite:
2 FeS2 + 7 O2 + 2 H2O → 2 FeSO4 + 2 H2SO4
[edit] Reactions
On heating, iron(II) sulfate first loses its water of crystallization and the original green crystals are converted into a dirty-yellow anhydrous solid. When further heated, the anhydrous material releases sulfur dioxide and white fumes of sulfur trioxide, leaving a reddish-brown iron(III) oxide. Decomposition of iron(II) sulfate begins at about 480 °C.
2 FeSO4 → Fe2O3 + SO2 + SO3
Like all iron(II) salts, iron(II) sulfate is reducing agent. For example, it reduces nitric acid to nitrogen oxide and chlorine to chloride:
6 Fe2+ + 6 H+ + 2 HNO3 → 6 Fe3+ + 4 H2O + 2 NO
2 Fe2+ + Cl2 → 2 Fe3+ + 2 Cl-
[edit] Uses
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