Chemistry

Gradual nitration of toluene


Fig.1

Trinitrotoluene

Trinitrotoluene (TNT), according to IUPAC nomenclature 2-methyl-1,3,5-trinitrobenzene is an explosive. The structural formula of the compound shows a benzene ring with a methyl group (-CH3) and three ortho or para nitro groups (-NO2) as substituents. The compound is formed by nitrating toluene using nitrating acid, a mixture of nitric and sulfuric acids.

  • 2,4,6-trinitrotoluene
  • Trinitrotoluene
  • 2-methyl-1,3,5-trinitrobenzene
  • 2-methyl-1,3,5-trinitrobenzene (IUPAC)
  • 1-methyl-2,4,6-trinitrobenzene
  • TNT
  • Trotyl
  • AT
  • Tol
  • Tolite
  • Tritol
  • Tutol

colorless to yellowish rhombohedral crystals or needles [1]

very bad in water (140 mg l −1, 20 ° C) [1]

Switzerland: 0.01 ml · m −3 or 0.1 mg · m −3 [4]

First synthesized in 1863 by Julius Wilbrand (1839–1906), [5] the large-scale production of TNT began in 1901 in the German Empire. The TNT equivalent serves as a measure of the energy released in an explosion.


1. History and use

In the 19th century, organic chemistry developed faster and faster and the various classes of substances were processed more systematically. Among other things, it was found that the reactions of concentrated nitric acid with organic substances formed explosive substances. The first chemist in this field was the French Henri Braconnot, who in 1832 investigated the effects of nitric acid on starch, sugar, wood fibers and other substances and obtained highly flammable products. This type of reaction was then referred to as “nitration”. In order to bind the resulting water, concentrated sulfuric acid was added to the nitric acid. This mixture of acids, with which the nitration was now even easier, was soon referred to as “nitrating acid”. From then on, nitration primarily served the purpose of producing various explosives.

1.1. Nitroglycerin

In 1846, the Italian professor Ascanio Sobrero succeeded in producing what he called “pyroglycerin” - nitroglycerin by allowing nitrating acid to act on glycerin. Strictly speaking, however, this name is incorrect as it is a nitric acid ester and not a nitro compound. A major disadvantage of this oil was that its ignition was almost unpredictable, which led to many accidents. It was only when Alfred Nobel invented the ignition principle in May 1862 that nitroglycerin could be reliably detonated. His principle was based on the fact that a small detonator made of black powder, which was later replaced by fumed mercury, was placed in the oil. The explosion of this capsule then triggered the detonation of the surrounding nitroglycerin. Nevertheless, many accidents still happened during the production and transport of nitroglycerin, as it was very sensitive to impact. Finally, Alfred Nobel succeeded in finding a suitable carrier for the oil in the naturally occurring kieselguhr. In 1867 he applied for a patent for this product, known as “dynamite”. This discovery eventually allowed nitroglycerin to be used as a commercial explosive.

1.2. Gun cotton

Another important explosive was discovered almost at the same time as nitroglycerin. In 1846, the Basel professor Christian Friedrich Schönbein allowed the nitrating acid described above to act on cotton and then received the gun cotton. After washing them several times with water, he let them dry and then ignited them. In doing so, it evaporated very quickly and free of residues, which is probably one of the most famous features of gun cotton. However, initially it was very difficult to produce consistent results and prevent accidental detonation. Despite great efforts, such as washing for 14 days or boiling in an alkaline solution, there were repeated unwanted explosions due to residual acid. Only the English chemist Sir Frederik Abel succeeded in completely removing these residues from the gun cotton. In his process, the still moist fibers were crushed into tiny particles, boiled out and then pressed into the end product. Abel's assistant Edwin A. Brown found out that gun cotton with 20% water content was a very explosive explosive that could be detonated with the help of the initial detonation. However, the production of gun cotton is formally an esterification and not a nitration.

1.3. Picric acid

The first explosive to be produced by "proper" nitration was picric acid. As early as 1843 it was produced purely synthetically by nitrating phenol. At first it was considered harmless and was only used as a dye. It was not until the French Eugene Turpin demonstrated in 1885 that picric acid is a very powerful explosive if it is detonated with an initial ignition. From this point on it found multiple uses in the military, especially as a filler for hand grenades. However, the picric acid could form highly explosive heavy metal picrates with the metal bodies of the grenades, which could detonate the main charge, which led to many accidents.

1.4. Trinitrotoluene

For the reasons mentioned above, picric acid was quickly replaced by TNT (trinitrotoluene), which was first produced by Paul Hepp in 1880. From around 1900, the large-scale synthesis of toluene, which was obtained from the paint industry, was possible. The advantages of TNT over picric acid were also quickly recognized. TNT is much less sensitive to impact and has a melting point of around 80 ° C. Even today, TNT is a very important explosive.


Steric effect - influence on the course of the reaction Exercise

Assess the reactivity of different haloalkanes in a $ S_N2 $ reaction.

The adjacent molecule shows the transition state of the $ S_N2 $ reaction on methyl bromide.

The $ S_N2 $ reaction is a nucleophilic substitution, the rate-limiting step of which is of the second order (kinetics). The rate-determining step depends on the concentration of the haloalkane and the concentration of the attacking nucleophile (e.g. $ I ^ - $ or $ EtO ^ - $), because these form a five-coordinate transition state.

The largest rate constant is found with a primary haloalkane, e.g. methyl bromide and ethyl bromide, followed by the secondary (e.g. isopropyl bromide) and the tertiary haloalkane (e.g. tert-butyl bromide).

With increasing spatial expansion of the residues, the rate constant decreases, since the formation of the transition state is made more difficult. It's comparable to the comic. The more voluminous the gray cat (halogen alkane), the more difficult it is for the red cat (nucleophile) to push the gray cat out of the basket (molecule).

Determine the major product of the bromination of toluene.

The attacking electrophile is the ion opposite.

The methyl group on toluene has a + I effect.

The methyl group from toluene pushes electron density into the ring system.

The bromination of toluene with iron (III) bromide and bromine yields para-bromotoluene as the main product.

The methyl group from toluene directs due to the + I effect the electrophile in the ortho- and para-positions. Since the attacking electrophile $ left [FeBr_3 cdots Br right] ^ + $ is relatively bulky, it has little space for attack in the ortho position and therefore attacks in the 4 position (para). This product is 67% made. In a few cases the ortho-substituted product is also formed. 2,4-Dibromotoluene is also possible, but the stoichiometry must then be 2: 1 instead of 1: 1. The meta-product is hardly / not achievable due to the inductive effect of the methyl group.

Decide what mesomeric effect subsequent substituents have.

The nitro group pulls electrons out of the ring.

The mesomeric effect (M effect) describes the ability of a substituent or a group in a conjugated system to change the $ pi $ electron density. The formation of an increase or decrease in charge occurs through mesomerism. This refers to the formation of $ pi $ bonds through the uptake of $ pi $ electrons, through a substituent or through the incorporation of a free electron pair (mostly on the heteroatom such as nitrogen or oxygen.)

  • Amino- (- $ NH_2 $)
  • Hydroxy (- $ OH $)
  • Halogen (- $ Cl $)
  • Alkoxy- (- $ - O-R $) $ rightarrow $ see illustration
  • Aldehyde- (- $ CHO $),
  • Keto- (- $ R (CO) R '$),
  • Carboxy (- $ COOH $),
  • Nitrile- (- $ C equiv N $),
  • Sulfonic acid- (- $ SO_3H $)
  • Nitro group (- $ NO_2 $)

Show the importance of the steric effect in the synthesis of 2-nitroresorcinol.

Mark the positions into which the OH groups direct. Remember that the oxygen atom has two lone pairs of electrons.

The figure opposite shows the reactive particle during sulfonation.

The hydroxyl groups of resorcinol direct due to the + M effects in the ortho- and para position. The markings in the adjacent figure show that the 2-position double is preferred.

The nitration on the aromatic is carried out using the nitrating acid $ (HNO_3 / H_2SO_4) $. Nitric acid and sulfuric acid form the nitronium ion as a reactive electrophile:

In order to achieve a substitution between the hydroxyl groups, the resorcinol is sulfonated beforehand. In contrast to nitration, the 2-position would not react in a sulfonation because the steric hindrance from the two hydroxyl groups would be too great. This effect is used here. In the sulfonation, for steric reasons, only the 4- and 6-positions (each marked with just an asterisk next to an OH group) are substituted. After these two positions have been occupied by a sulfone group $ (SO_3) $, nitration can take place.

In order to subsequently remove the sulfonic acid groups, which are known as Protecting groups served to remove, the compound is reacted with hydrochloric acid, only this and not the nitro or hydroxyl groups react.

Explain the term: steric effect.

Consider why, for example, monobromomethane reacts faster in a $ S_N2 $ reaction than, for example, isopropyl bromide.

The steric effect occurs when molecules are very branched and require a large amount of space. The centers in the vicinity of such bulky substituents are protected, so they do not react as quickly as usual. This prevents reactions on the one hand, but also stabilizes reactive, unstable compounds on the other.

If bulky reagents react, they react preferentially at centers that provide enough space, i.e. have no substituents in the vicinity.

Despite the steric hindrance, syntheses can also be favored. The most common drivers of reactions are the formation of Aromaticity (electrophilic aromatic substitution) or that Splitting off small molecules (Condensation reaction) and the precipitation of salts. According to the principle of Le Chatelier (Principle of the smallest constraint) the precipitation of salts (e.g. LiCl, AgCl) shifts the equilibrium to the side of the product, since one of the products is always removed from the system. This principle also applies to the escape of gases (e.g. nitrogen oxides, carbon dioxide or hydrochloric acid) or the formation of liquid substances that form another phase (e.g. the release of water in toluene produces a mixture with two phases, which is used, for example, by the esterification water separator exploited).

Determine the functionality (s) that will be protected first in the following sugar derivative.

The reaction to introduce the protective group is a $ S_N2 $ reaction (Williamson ether synthesis).

Secondary alcohols react more slowly than primary alcohols in a $ S_N2 $ reaction.

Williamson's ether synthesis is one $ S_N2 $ response. Due to the five-fold coordinated transition state, this proceeds much faster when the alcohol to be attacked primary is, i.e. if the hydroxyl group in the vicinity bears a $ CH_2-R $ group.

In the methyl $ beta $ -D-glucopyranoside there is only one primary hydroxyl group (see fig. Blue), which is protected by the trityl chloride. the secondaryThe hydroxyl groups (see fig. green) have a $ CHR_2 $ group in the vicinity and are therefore sterically inhibited from attack by the sterically demanding trityl chloride.

The glycosidic hydroxyl group (see fig. Orange) is already protected as a fully acetal, as otherwise the ring could open.

The trityl chloride is very suitable as a protective group, as it can also be split off again very easily because the triphenylmethane system is stabilized by conjugation (similar to crystal violet).


History of Explosives - From Greek Fire to Modern Explosives

In 1832, Henri Braconnot (1780–1855) allowed a mixture of nitric acid and sulfuric acid (nitrating acid) to act on wood shavings. He received a material that is very easily inflammable. Braconnot called it "xyloidin". It is a starch nitrate with a nitrogen content of up to 13%. The speed of detonation is given as 4970 m / s.

The French chemist Théophile-Jules Pelouze (1807–1867) also studied this substance in 1838. However, he used paper. He called the resulting product pyroxyline.
Sources: [37], [38], [39], [40], [41]

6.1 Gun cotton

  • Manufacture: nitriding of cotton
  • Uses: propellant charge, celluloid, pyrotechnics
  • Detonation speed: 6300 m / s
  • Density: 1.3-1.7 g / cm 3
  • Impact sensitivity: 3 joules

In 1846 Christian Friedrich Schönbein (1799–1868) dealt with xyloidin and pyroxylin. Schönbein converted other organic substances with nitrating acid. When using cotton, Schönbein was given a cotton-wool-like fabric that was extremely flammable and called it gun cotton, which is also often referred to as nitrocellulose. However, this designation is wrong because it is not a nitro compound, but a nitrate compound. If gun cotton is set on fire, it deflagrates at a speed of 700–900 m / s. However, if it is pressed and initially ignited, a detonation speed of 6300 m / s is achieved. Nitrocellulose is the main component of many propellant powders for hunting and military ammunition. In order to influence the burning behavior, the propellant charge powder is shaped differently.
Sources: [42], [43], [44], [45]

6.2 Nitroglycerin and the not very noble dynamite

The name nitroglycerin is also wrong, as it is also a nitrate. So the correct name is glycerol trinitrate (propanetriol trinitrate).

  • Production: nitration of glycerine
  • Uses: explosives, dynamite production, medicine
  • Detonation speed: 7600 m / s
  • Density: 1,599 g / cm 3
  • Impact sensitivity: 0.2 joules

In 1783 Carl Wilhelm Scheele (1742–1786) saponified olive oil with lead oxide. This procedure gave him an oily and sweet-tasting liquid, which he called "oil sweet". The French chemist Michel Eugène Chevreul (1786–1889) later introduced the name glycerine for "oil sweet".

Pelouze reported to his former assistant, from 1840-1842, Ascanio Sobrero (1812-1888), a doctor and chemist from Turin, about the gun cotton and the other new fabrics. Sobrero initially experimented with gun cotton, but he also investigated other substances for reaction with nitrating acid. In 1847, for example, he added glycerine to a mixture of sulfuric acid and potassium nitrate. From it Sobrero obtained a slightly milky liquid that is immiscible with water. When he heated a few drops in a glass, there was a violent explosion. He then called his discovery "piroglicerina (pyrophoric glycerine)". A later nitroglycerin explosion disfigured Sobrero's face forever.

In 1850 the Swedish arms manufacturer Immanuel Nobel (1801–1872) sent his son Alfred on a three-year study trip through America and Europe. Alfred Nobel (1833-1896) worked briefly at Théophile-Jules Pelouze as an assistant. It was there that Alfred Nobel heard of nitroglycerin for the first time. What is certain, however, is that the Russian chemist Nikolaj Nikolaevic Zinin (1812–1880) and the pharmacologist Jurij Trapp reported on nitroglycerin in 1853. Then he and his father started working with nitroglycerin. A controlled ignition of the nitroglycerin was also denied to the Nobels for the time being.

In the truest sense of the word, Alfred Nobel had a brilliant idea: He used a capsule filled with mercury fulminate to cause the nitroglycerin to explode in a controlled manner. The so-called initial spark was born. Nobel applied for a patent for this initiator on July 15, 1863. This is probably his greatest and most important invention. The nitroglycerin seemed to be "tamed".

On September 3rd, 1864 there was a violent explosion in an explosives factory owned by Alfred Nobel, in which five people died, including his younger brother Oscar-Emil Nobel.

This incident prompted Alfred Nobel to make the nitroglycerin safer. He tried a wide variety of binders, including Sawdust, wood flour, brick dust, etc.

But it was not until the use of diatomaceous earth that the desired success was achieved in 1866. Mixing nitroglycerine with kieselguhr creates a paste-like mass that can be portioned at will and is not sensitive to knocks, bumps or friction.

Because of its powerful explosive effect, Alfred Nobel called his explosive "dynamite". The discovery of blasting gelatine was made by Alfred Nobel in 1875. Blasting gelatine consists of nitroglycerin (92%) and gunned cotton
(8%) and is even stronger than the Gur dynamite. Today's dynamite is made entirely from these gelatinized nitroglycerin products.

Firing a firearm produced a lot of smoke through the use of black powder. For a long time a battlefield cannot be surveyed because of this smoke. Therefore Nobel developed the low-smoke powder "Ballistit" from the explosive gelatine, which he patented in 1887. Due to its slow combustion, ballistite is ideal as a propellant powder for ammunition. It made possible the development of machine guns.
Sources: [46], [47], [48], [49], [50], [51], [52]

6.3 Picric acid (2,4,6-trinitrophenol) and its salts (picrates)

  • Production (today): sulfonation and subsequent nitration of phenol
  • Detonation speed: 7100 m / s
  • Density: 1.76 g / cm 3
  • Impact sensitivity: 7.4 joules

Around 1650 Johann Rudolph Glauber (1604–1670) allowed concentrated nitric acid to act on sheep's wool. The result was a strongly yellow colored and bitter (piquant, Greek: bitter) tasting acidic solution. Glauber added potash to neutralize it, making it the first to produce a salt of picric acid, potassium picrate. An alcoholic potassium picrate solution ("Tinctura nitri Glauberi") was used in medicine until well into the 18th century.

In 1771 Peter Woulfe (1727? –1803) allowed nitric acid to act on indigo. He too received a yellow liquid that was able to dye silk and other textiles. Hausmann produced yellow picric acid for the first time in 1788.

The chemist Jean Joseph Welter (approx. 1763–1852) produced pure picric acid for the first time in 1799 by treating silk with nitric acid. Welter also discovered the potassium picrate fizzled out when exposed to heat. Potassium picrate's sensitivity to impact did not go unnoticed for long either. The potassium picrate was mixed with saltpeter and other substances. These mixtures were filled into hollow spheres that exploded on impact, even without a detonator. During the American Civil War (1862–1865), these projectiles came into frequent use.

Picric acid itself was considered a harmless dye; it was thought that picric acid lacked oxygen in order to have explosive properties. In 1885, the French chemist François Eugène Turpin (1848–1927) demonstrated the detonation ability of picric acid by using Nobel's detonators.

From 1886 picric acid was used in military technology under different names:

  • Melinite (France, 1886)
  • Lyddit (England, 1888)
  • Schimose (Japan)
  • Garnet filling 88 (Germany)
  • Ekrasit (Austria)
  • Pertit (Italy)

The picric acid-filled grenades fired on the battlefields did not detonate the picric acid completely. Unused picric acid colored the soldiers' skin bright yellow, which is why they were referred to as canaries or canaries.

Picric acid reacts with metals to form so-called picrates. Heavy metal picrates, such as lead or copper picrate, are very sensitive to impact, shock and friction. There were therefore very often gunshots, unwanted explosions and spontaneous detonations. These picrates had initially ignited the picric acid.

A large number of explosion accidents were recorded within a very short time:

  • 1887 Cornbrook near Manchester
  • 1901 Griesheim near Frankfurt a. M. in Germany
  • 1903 Woolwich (England)
  • 1917 Halifax (Canada) et al.

To prevent the picric acid from reacting with the metal walls, the walls of the grenades were tinned or painted with varnish. Picric acid was soon to be replaced by a non-acidic and manageable substance: TNT
Sources: [53], [54], [55], [56], [57]

6.4 TNT - 2,4,6-trinitrotoluene

  • Production: multiple nitration of toluene
  • Use: Explosives for hand grenades and bomb or grenade filling
  • Detonation speed: 6900 m / s
  • Density: 1.64 g / cm 3
  • Impact sensitivity: 15 joules

Julius Bernhard Friedrich Adolph Wilbrand (1839–1906) carried out a multiple nitration of toluene in 1863. He received light yellow, needle-shaped crystals that melt between 80–81 ° C. In 1891 the technical production takes place by Karl Häussermann (1853–1918) and the military research office in Berlin recognizes the suitability as an explosive. From 1902 TNT was used as a garnet filling in Germany until it was used in military ammunition worldwide during World War II. Because of its toxicity, TNT is increasingly being displaced by other explosives.
Sources: [58], [59], [60], [61]

6.5 Nitropenta (pentaerythritol granitrate)

  • Production: nitration of pentaerythritol with nitric acid (& gt 95%)
  • Use: detonating cords, plastic explosives, lightning cords, propellant powder, pyrotechnics
  • Detonation speed: 8400 m / s
  • Density: 1.77 g / cm 3
  • Impact sensitivity: 3 joules

Bernhard Tollens (1841–1918) discovered pentaerythritol in 1891. Nitropenta is obtained by nitrating pentaerythritol with highly concentrated nitric acid. The substance was proposed as an additive for low-smoke powders in 1894. Nitropenta is used as a component of almost all plastic explosives. It is also used as a filling for detonating cords and lightning cords in blasting technology or in pyrotechnics.
Sources: [62], [63], [64], [65]

Details Manfred Seidl Published: August 27, 2013 Last updated: August 03, 2017 Created: August 27, 2013 Hits: 10432

4 human experiences

Before and during World War I, the armaments industry in England, Germany and the USA processed large amounts of t-TNT, which was considered non-toxic [36]. The pure, recrystallized trinitrotoluene was still considered to be non-toxic in 1917 [37] and the toxic effects of a wide variety of impurities were blamed. Descriptions of the first t-TNT poisoning came from England. There, in the years 1911–1915, 279 deaths (101 ♂♂, 178 ♀♀) occurred in ammunition factories due to t-TNT [38]. These figures relate to the entire workforce in the ammunition factories and represent a relatively small percentage. As usual, exact workplace concentrations were not determined, only the workplaces and working conditions were described: for example, the amount of t-TNT during melting was 6 mg per m 3 air. From this, an uptake of at least 16 mg t-TNT in 7.5 hours was estimated. The floor was swept three times a day, which resulted in an additional exposure of 9.1 mg t-TNT per day. When blowing out the ignition channels of bombs and mines, one worker ingested 2–3 mg t-TNT per breath [9]. Subsequent air measurements made while the melting kettle was being charged at the workers' nose level showed 10.7 mg of t-TNT per m 3 of air [39].

In times of war, the first signs of poisoning were not taken seriously enough, which is evident from an investigation of workers in an ammunition factory, according to which 72% of the workers experienced subjective complaints (fatigue, paleness, cyanosis), but these were not considered to be real symptoms [25].

Observations from the years up to 1921 on at least 1800 persons exposed to different levels of t-TNT who had worked in ammunition factories for a month to three years can be summarized as follows [9, 12, 13, 21, 38, 40-46]:

Irritating symptoms of the respiratory tract: Epistaxis, runny nose, burning eyes, headache, tightness i.d. Chest, dry cough of the digestive tract: bitter taste in the mouth, increased or decreased appetite, nausea, vomiting, abdominal pain (liver, stomach, diaphragm area), constipation, later diarrhea (polyuria, only: [45]) of the skin: Dermatitis on contact surfaces, itchy rash (erythema consisting of tiny red spots), wheals, blisters, yellowing of the skin (hands, feet, face).

Toxic symptoms of the digestive tract: Anorexia, biliary colic, gastritis, epigastric pain, toxic jaundice of the blood: MetHb formation and sequelae (pallor, cyanosis, hunger for breath), anemia, aplastic anemia, changes in the blood count (reduction and damage to erythrocytes, leukopenia, leukocytosis, lymphocytosis, fragmented cells), subcutaneous bleeding of the cycle: Bradycardia, palpitations swelling d. Hands and feet of the nervous system: Somnolence, depression, apathy, (peripheral neuritis, only: [40]) other: Irregular, weak menstruation, dark urine, increased sweating.

Studies from ammunition factories often point to the role of individual predisposition [15, 47, 48]. Experience in England showed that workers who had tolerated t-TNT for 5-6 months were no longer expected to have dangerous consequences. There was talk of a certain immunity achieved by such workers [48, 49]:

In numerous studies [15, 37, 38, 50, 51] it is stated that the susceptibility to t-TNT, e.g. varies with gender and race. In contrast, Morton [52] found no differences between men, women, blacks, Caucasians or different age groups in his studies. An even distribution of the symptoms of poisoning over the various age groups in men was also confirmed by McGee [5].

Poisonings - including fatal ones - increased during the night shifts from May to September in warm weather [4]. Teisinger [53] also found a “clear change for the better” in investigations in October compared to those in a hot August (see also “Skin toxicity”). Alcohol also increased the toxicity [53, 54].

4.1 Acute toxicity

In the foreground was mostly the change in blood values, which was to be expected especially in workers with glucose-6-phosphate dehydrogenase deficiency [55] and was also found [56].

After 2–3 days of relatively high t-TNT exposure (1.8–2.95 mg / m 3), three men (18, 23 and 35 years old) who were exposed to glucose-6-phosphate dehydrogenase Suffering from deficiency, a hematolytic crisis: Hemoglobin: 40–82 g / l (normal value ♂: 140–180 g / l). Hematocrit: 17–24% (normal value ♂: 40–54%). Reticulocytes (% in red blood cells): 10-26.2%. All three recovered quickly after changing jobs [56].

At exposures below the TLV value of 1.5 mg / m 3 at that time, however, changes in liver values ​​were also observed. At 0.8 mg / m 3 SGOT (serum glutamate oxaloacetate transaminase) and LDH (lactate dehydrogenase) were significantly increased [52]. These observations led to a reduction in the TLV value to 0.5 mg / m 3.

The acute lethal dose for humans is assumed to be 1–2 g of 2,4,6-TNT [57]. Of 1,100 workers who were regularly monitored, some of whom were exposed to t-TNT through inhalation and some through skin contact (hands), 1% fell ill after inhalation exposure and 10% of workers in the corresponding group after direct skin contact [11]. Oily hands favored absorption - severe symptoms were observed in 17% of workers with oily hands compared to 11% with dry hands [42]. Perspiration also had an intensifying effect, which explains the increased number of diseases, especially dermatitis, in the warm season [4, 37, 45, 58].

4.2 Skin toxicity

Itchy, burning inflammations of the skin, usually papular [59], sometimes vesicles or wheals [36], are considered to be one of the first (after five or more days [60]) signs of intoxication [47].

Yellow coloring of the skin and orange coloring of the hair is often described as evidence of TNT contamination by, among others, [4, 6, 8, 12, 53, 60]. It occurred after t-TNT skin contact on uncovered parts of the body (hands, forearms, face) sometimes during the first month [23], but mostly after about 2 months [61]. "Trotyl scabies" [62], as the dermatitis was previously called in explosives factories, developed preferentially in people when t-TNT was mixed with hygroscopic ammonium nitrate, whereby increased temperature and perspiration had an intensifying effect [54].

Workers who were busy loading or unloading t-TNT or melting them down in hollow bodies (bombs, mines, grenades) were particularly prone to allergies and dermatitis [5, 63], or to mild dermatoses on the hands remained limited [48]. Dermatitis was observed in some workers at a daily average of only 0.35 mg 2,4,6-TNT / m 3 air (? / 43) [52].

52 students (44 ♀, 8 ♂, on average 20 years old) worked as temporary workers for 33 days in the filling room (0.3–0.6 mg / m 3) of an ammunition factory. Of them, 18 had mild, 5 moderate, and 4 severe dermatitis on their hands. In contrast, out of 10 ♂ students who worked in the enamel room (0.3–1.3 mg / m 3), not a single one developed a skin rash [64].

Of 495 mild t-TNT poisoning cases (examined within one year), 181 had cyanosis, 107 gastritis and 207 dermatitis [15].

4.3 Subchronic and chronic toxicity

In exposed workers, mild anemia (hypochromatic, normochromatic or hyperchromatic), venous dermatitis, hepatitis and toxic polyneuritis were observed most frequently. This is evident from the results of a study on men (18–65 years old) from several factories (no information on exposure duration and level) [5]:


Description Electrophilic substitution

This video explains electrophilic substitution to you. As a previous knowledge you should know what aromatic compounds are and what is meant by the term mesomerism. First, you will learn what electophilic substitution is and how the mechanism of this reaction works. This knowledge is explained to you in more detail using a few examples. At the end you will learn how the reaction can be accelerated and delayed and what multiple substitutions are.

Transcript Electrophilic substitution

Good day and welcome!

This video is about electrophilic substitution. The film belongs to the series "Reaction Mechanisms". As a previous knowledge you should know what benzene is. You know the structure of the benzene molecule, you know what aromatic compounds are and you know about mesomerism. My goal is to give an overview of electrophilic substitution on the nucleus.

Outline: 1. What is it? 2. A rule of thumb 3. The mechanism 4. Examples 5. Acceleration and deceleration 6. Multiple substitution and 7. Summary.

In principle, electrophilic substitution is possible with aromatics and with aliphatics. With aromatics, however, it works much better than with aliphatics. There are more examples of this as well. Therefore we will only consider electrophilic substitution for aromatics. In the electrophilic substitution of aromatics, hydrogen atoms are exchanged for other atoms or groups. We consider the benzene molecule as a representative of the aromatic compounds. The molecule has 6 freely movable π electrons. It has an electron sextet. Electrophiles can initiate an electrophilic attack on these electrons, which is the start of electrophilic substitution.

In connection with electrophilic substitution, one often hears the KKK rule. The 1st K means that the substitution takes place on the nucleus, i.e. on the aromatic ring. The 2nd K means catalyst. In other words, a catalyst is required for the reaction. Finally, the 3rd K stands for cold. In organic chemistry, cold does not necessarily mean freezing temperatures. Also 20 ° C and maybe a bit above that means cold here.

Let's consider an example: chlorine reacts with benzene. A chlorine molecule reacts with a benzene molecule. The first stage of the mechanism is the formation of the electrophile. The chlorine molecule reacts with the catalyst. Chlorine reacts with iron (III) chloride. Because that's the catalyst. And so the reaction takes place. The chlorine cation is created. That is the electrophile of the reaction. In addition, the ion FeCl (4-) is formed. In chemical notation this means: Cl2 + FeCl3 → Cl (+) + FeCl (4-). FeCl3 is iron (III) chloride, the catalyst. Cl (+) is the chlorine cation, the electrophile of the reaction.

    The formation of the π complex: The chlorine cation interacts with the benzene molecule. It is coordinated by this one. A so-called π-complex is created. This is shown in chemical formula notation. The third step of the mechanism is the formation of the σ-complex. In the π complex, the chlorine cation is coordinated with all carbon atoms at the same time. In the σ-complex it grabs a carbon atom and forms a proper chemical bond with it. Double bonds and positive charge are then in the core. That is the first order.But this arrangement is also possible, and ultimately this one. The superposition of the boundary structures gives this formula for the σ-complex.

eventually chlorobenzene is formed. The σ-complex forms chlorobenzene. It's an aromatic compound again. It can only arise by splitting off a hydrogen atom H +.

Benzene or other aromatic compounds can be derivatized in various ways by electrophilic substitutions. First of all, halogenation, of which we discussed chlorination. Only chlorination and bromination work well. During nitration, 1 hydrogen atom is exchanged for the nitro group NO2. During the sulfonation, the sulfonic acid group SO3H is introduced. In the acylation, 1 hydrogen atom is exchanged for an acyl radical. In the case of alkylation, the substitution is made by an alkyl radical. During the diazotization, 1 hydrogen atom is exchanged for a diazonium radical.

The rate of electrophilic substitution can be changed by substituents compared to benzene. Let us compare benzene with a compound in which there is a methyl group in the benzene residue instead of a hydrogen atom. In the second comparison structure, there is a nitro group instead of a hydrogen atom. The reaction rate of the electrophilic substitution for benzene is then black-V under certain conditions. In the case of methylbenzene, the reaction rate is red-V under the same conditions. In the case of nitrobenzene, the reaction rate is corresponding to Blue-V. Then Red-V & gtBlack-V. And Black-V is again & gtBlue-V.

Why is that? The methyl group pushes electrons into the π-electron sextet. The connection becomes more reactive. And vice versa: The nitro group withdraws electrons from the π-electron sextet. The molecule becomes less active for electrophilic substitution. At which points do the substitutions take place? In the case of methylbenzene, initially directly next to the methyl groups. This is the so-called ortho position. In addition, substitution takes place exactly opposite the methyl group, in the so-called para position. In the case of nitrobenzene, the substitution takes place preferably in two positions next to the nitro group. This is the so-called meta position. Also take a look at the video "Second Substitution".

Let us imagine that we want to substitute several times with one and the same electrophile. Let's start with the alkylation. We introduce a methyl group into the benzene ring. The methyl group pushes electrons into the π-electron sextet. This makes the molecule more active. A second methyl group can be introduced. The molecule becomes even more active and a 3rd can be introduced much faster. And now it is happening in rapid succession. The 4th comes immediately and this makes the system even more active. And already we have the 5th substituent. And now the exchange is taking place at gigantic speed. All 6 hydrogen atoms are substituted by methyl groups. Did you notice something? Right, the reaction was practically unstoppable. A so-called polyalkylation takes place, which is sometimes useful, but mostly disadvantageous.

Now let's look at nitriding. A simple nitriding can be accomplished without major problems. But now it comes: The electron-withdrawing nitro group withdraws electrons from the π-electron sextet. The beautiful responsiveness of the benzene ring is spoiled. The reaction stops, only mononitration occurs. If you want to achieve a 2nd or even 3rd nitration, as we have heard in the meta positions, the reaction temperature must be increased, at least to 100 ° C and sometimes even higher.

The electrophilic substitution takes place on the core when it is cold and requires a catalyst. As an example, we have considered the mechanism of chlorination. Chlorine reacts with iron (III) chloride, the catalyst. The chlorine cation, the electrophile, is created. This reacts with the aromatic, the benzene molecule. A π complex is created. A σ complex arises from the π complex. The product chlorobenzene is formed by splitting off a proton. Examples of electrophilic substitution are halogenation, nitration, sulfonation, acylation, alkylation and diazotization. Is alkylated e.g. B. when introducing methyl groups, it comes to polyalkylation. In the case of nitration, only mononitration occurs under mild conditions. If a methyl group is already present, the reaction rate increases. It is conducted in ortho and para positions. In contrast, the electron-withdrawing nitro group slows the reaction rate. It directs the new substituents in the meta position.

I thank you for your attention. I wish you all the best - Goodbye!


Old exams

In the case of the reactor accident in Fukushima, inter alia the radioactive isotope 137 Cs leaked. Describe the structure of this atom from its elementary particles in as much detail as possible (how many particles of which type, with which mass and charge are where in the atom?) (2 P)

Give the names of the following compounds: a) H2SO3, b) NH4 + NO3 - ,
c) CH3COOH, d) H.2O2 (2 P)

In experiments, hydrochloric acid is poured over a piece of a) lime, b) zinc and c) copper. Describe what happens and explain the chemical relationships with the help of reaction equations. (2 P)

In a tank there are 100 liters of hydrochloric acid with a pH value of 1. Calculate how many grams of calcium hydroxide have to be added for complete neutralization. (2 P)

Chlorine gas can be produced in the laboratory by reacting hydrochloric acid with potassium dichromate (K2Cr2O7), whereby i.a. Cr 3+ ions arise as a by-product. Set up the partial and total reaction equations using the relevant oxidation numbers. (2 P)

Describe the similarities and differences between diamond and graphite in terms of their electronic (hybridization) and molecular structure and the resulting material properties. (2 P)

Give the name and structural formula of the simplest aromatic alcohol and explain how the pH changes when you put it in water. (2 P)

A: Natural product / food / biochemistry:

9A. Enter the names and formulas of three common preservatives. (2 P)

10A. What do you know about the biochemically important compounds AMP / ADP / ATP? (2 P)

B: Physical chemistry:

9B. Explain in as much detail as possible the conditions under which endothermic reactions can occur spontaneously. (2 P)

10B. Describe the general molecular structure as well as the thermal and mechanical properties of a) thermoplastics, b) thermosets and c) elastomers (2 P)

Exam
For lecture 62-082.1: "Basics of Chemistry"
6.2.2017 9: 15-10: 45 h

Calculate in how many liters of water a mass of m = 1 g of magnesium chloride has to be dissolved so that a concentration of c = 0.01 mol / L results. (2 P)

Explain the octet rule using the structural formulas of the compounds CH4, CH2O, H2S and AlCl3. (2 P)

Explain the terms hydrophilic and lipophilic using the structural formulas of the solvents water, tert-butanol, hexane, methanol and diethyl ether. (2 P)

Calculate the pH of a buffer solution that contains 10x as much acetic acid (pKS. = 4.75) how does acetate contain? (2 P)

Explain the technical lime cycle in as much detail as possible. (2 P)

In the first lighter, developed by Johann Wolfgang Döbereiner in 1823, zinc was brought into contact with sulfuric acid and the gas produced in the process was guided past a platinum sponge. Explain the chemical relationships and reactions that occur, also using reaction equations. (2 P)

Lactic acid has the systematic name 2-hydroxy propanoic acid. Draw the structural formula. (2 P)

A: Natural product / food / biochemistry:

9A. Explain the terms asymmetrically substituted carbon atom and Chirality based on the structural formula of a suitable natural product. (2 P)

10A. Show by a reaction equation that glucose is a reducing sugar. (2 P)

B: Physical chemistry:

9B. Describe the chemical processes in a lead battery in as much detail as possible.
10B. Describe the chemical processes involved in radical polymerization using an example of your own choice.

For lecture 62-082.1: "Basics of Chemistry"

Compare the two atoms 50 Ti and 50 V with regard to their structure from elementary particles (how many particles of which kind with which properties are each where in the atom). (2 P)

Calculate how many grams of magnesium chloride have to be dissolved in V = 250 ml of water in order to produce a concentration of c = 0.1 mol / l? (2 P)

Explain the terms mesomerism and boundary structure formulas - if possible using the example of the acetate ion. (2 P)

Set up a reaction equation and the law of mass action for the reaction of sulfur dioxide with oxygen to form sulfur trioxide. (2 P)

A tank contains 2400 liters of acid with a pH value of 1. Calculate how many liters of sodium hydroxide solution with a concentration of c = 6 mol / L you have to add for complete neutralization. (2 P)

The reaction of sulfuric acid with calcium hydroxide produces, among other things, Calcium sulfate. The reaction of sulfuric acid with hydrogen bromide produces, among other things. elemental bromine and sulfur dioxide. Set up the two reaction equations and compare the reaction mechanisms. (2 P)

Compare the following ions with regard to their summer formula and the oxidation number of chlorine: a) chloride, b) chlorate, c) hypochlorite, d) perchlorate. (2 P)

Give the names and structural formulas for three different compounds with the empirical formula C.3H8O on. (2 P)

Describe in as much detail as possible (reaction equations with structural formulas) which different reactions chlorine can enter into toluene (= methylbenzene).
(2 P)

Explain in as much detail as possible how and why solid fat can be made from cooking oil (fat hardening). (2 P)

For lecture 62-082.1: "Basics of Chemistry"

Compare the gases carbon monoxide and carbon dioxide in terms of their molar mass and molar volume. (2 P)

Calculate in what volume of water 7.9 g of potassium permanganate (KMnO4) have to be solved so that the concentration is c = 10 -3 mol / l?

Draw the valence line formulas (including free electron pairs!) For the following molecules and ions: a) HCN, b) hydrogen peroxide, c) hydrogen carbonate ion, d) nitrite ion (2 P)

If hydrogen and oxygen are mixed, the reaction does not take place immediately, although the reaction would be exothermic or exergonic. It needs a spark or the presence of platinum. Explain the chemical relationships. (2 P)

Use an example to explain how an acid-base buffer works. (2 P)

In the laboratory, chlorine gas is produced by oxidizing hydrochloric acid with potassium permanganate (KMnO4), whereby inter alia. Mn 2+ ions arise. Set up the reaction equation and give the relevant oxidation numbers. (2 P)

Explain the large-scale chlor-alkali electrolysis in as much detail as possible. (2 P)

Describe the chemical reaction of ethane with bromine in as much detail as possible!
(Reaction equation, reaction conditions, mechanism, products) (2 P)

Explain the terms "positional isomerism" and "mirror image isomerism" using the example of lactic acid (= L-2-hydroxypropanoic acid). To do this, draw at least one structural formula. (2 P)

Explain in as much detail as possible the difference between a- and b-D-glucose, also with regard to the polysaccharides formed from them. (2 P)

For lecture 62-082.1: "Basics of Chemistry"

on Monday, February 15, 2016

Name the classic three elementary particles and compare them with regard to their mass, charge and their space in the atom. (2 P)

Calculate how much water you have to add to 250 ml of a solution so that the concentration of c0= 0.2 mol / L on c1= 0.04 mol / L decreases? (2 P)

Explain the bond character and the stoichiometric ratios in compounds from a) Ca and F, b) Si and Cl, c) P and H, d) Cu and Zn. (2 P)

Explain in as much detail as possible, using examples, what catalysts are and how they work. (2 P)

Blood has a pH of 7.3. Can you calculate how many (pieces!) Of hydronium ions a person has in 6 liters of blood? (2 P)

Archae bacteria can convert carbon dioxide and hydrogen into methane and water in the absence of oxygen. Draw up a reaction equation and give the oxidation numbers of carbon and hydrogen. (2 P)

Compare gold, iron and aluminum with regard to their corrosion resistance, taking into account the respective chemical properties. (2 P)

The naturally occurring L-lactic acid has the systematic name L-2-hydroxy-propanoic acid. Draw the structural formula. (2 P)

Find an equation for the reaction of sodium ethanolate with iodomethane and name the product. What is the mechanism by which this reaction occurs? (2 P)

Explain in as much detail as possible and using a structural formula what is meant by unsaturated fats. (2 P)

For lecture 62-082.1: "Basics of Chemistry"

What electron configuration does sulfur have and how can this explain the bonds in hydrogen sulfide? (2 P)

Draw the valence line formulas of the following molecules and ions:
a) SO2, b) HCN, c) hydrogen peroxide, d) hydrogen carbonate ion (2 P)

Write an equation for the reaction of aluminum with iodine. (2 P)

Draw the phase diagram of water schematically and show why ice melts under pressure and how a pressure cooker works. (2 P)

What are the pH values ​​when you
a) 1 g perchloric acid (HClO4, strong acid)
b) 0.1 mol acetic acid (weak acid, pKS.=4,75)
dissolves each in 10 L water? (2 P)

Describe chlor-alkali electrolysis in as much detail as possible. (2 P)

Draw up a reaction equation for the dissolution of copper in nitric acid, where i.a. Forms nitric oxide and copper (II) nitrate. (2 P)

The link below has the common name Pinguinon (you can see why). What is the systematic (IUPAC) name? (2 P)

Using an example of your own choice, explain what a nucleophilic substitution is.
(2 P)

Describe the structure of the DNA schematically. (2 P)

For lecture 62-082.1: "Basics of Chemistry"

Which ions have the same electron configuration as argon? (2 P)

Use the respective structural formulas to explain why all gaseous elements that are not noble gases occur as diatomic molecules. (2 P)

Explain in as much detail as possible what amalgam is (what elements it consists of, what stoichiometric ratio, what type of chemical bond is present).
(2 P)

Establish the reaction equation and the law of mass action for the formation reaction of ammonia from the elements. (2 P)

Calculate in how many liters of water m = 2 g perchloric acid (HClO4, strong acid) must be dissolved so that a pH value of 3 arises. (2 P)

14 mg calcium carbonate are soluble per liter of water. Calculate the solubility product. (2 P)

Describe in as much detail as possible how a car battery (= lead battery) works. (2 P)

Explain what chirality means using the structural formula of lactic acid
(= L-2-hydroxy propanoic acid). (2 P)

Compare esters and amides for their structure and manufacture (2 P)

Draw the structural formula of glucose and explain how the ring closure reaction (= formation of hemiacetal) occurs. (2 P)

For lecture 62-082.1: "Basics of Chemistry"

Compare the two atoms 50 Ti and 50 V with regard to their structure from the elementary particles. (2 P)

A camping gas stove contains 176 g propane. Calculate how many grams of carbon dioxide are produced if you burn it completely. To do this, set up a reaction equation. (2 P)

Which compounds (state the molecular formula and the type of bond) are formed from the following pairs of elements? (2 P)
a) C / Cl b) H / S c) Ca / F d) Li / O

Describe in as much detail as possible how a catalyst works in chemical reactions - if possible using the example of platinum. (2 P)

Calculate a) how many moles of perchloric acid (HClO4, strong acid) and b) how many moles of acetic acid (CH3COOH, weak acid pKS. = 4.75) you have to dissolve in V = 1 L of water so that a pH value of 4 is established in each case? (2 P)

Use the relevant oxidation numbers in the reaction equation to show that photosynthesis is a redox reaction. (2 P)

Name as many different substances as possible that are a) very reactive and b) unreactive and explain the difference with the help of the term Standard enthalpy of formation.
(2 P)

Draw the structural formulas of a) formic acid b) ethyl ethanoate
c) trimethylamine d) diethyl ether (2 P)

Compare ethane, ethene and ethyne in as much detail as possible with regard to their reaction with bromine (reaction conditions, mechanism, reaction equation with structural formulas, name of the products). (2 P)

Describe the saponification of fat molecules in as much detail as possible (reaction equation with structural formulas). (2 P)

For lecture 62-082.1: "Basics of Chemistry"

Which atom has twice as many electrons as argon? (2 P)

What is potassium nitrite and in how many liters of water do 17 g of it have to be dissolved so that the solution has a concentration of c = 10 -2 mol / L? (2 P)

Can you calculate how many grams of oxygen it takes to completely burn 4.4 g of propane? Write up an equation for the reaction. (2 P)

Compare the states of aggregation of H.2O, H2S, CO2 and SiO2 and explain the relationships based on the electronic and geometric structure of the molecules. (2 P)

Describe - if possible with an example - what an acid-base indicator is. (2 P)

State the partial and total reaction equations for the electrolysis of water and explain how the pH value changes in the cathode and anode compartment. (2 P)

Using a suitable example, describe what a coordinative bond is. (2 P)

Draw the structural formulas of
a) dimethyl ether
b) dimethylamine
c) ortho-dimethylbenzene (newer: 1,2-dimethylbenzene) and
d) 2,2-dimethyl-propan-1-ol. (2 P)

Describe the nitration of benzene (= benzene) in as much detail as possible. (2 P)

Explain in as much detail as possible (including structural formula) what is meant by unsaturated fatty acids? (2 P)

For lecture 62-082.1: "Basics of Chemistry"

    In the case of the reactor accident in Fukushima, inter alia the radioactive isotope 137 Cs leaked. Describe the structure of this atom from its elementary particles in as much detail as possible (how many particles of which type, with which properties (mass, charge) are located where in the atom?) (2 P)

The limit value for nitrate ions in drinking water is 25 mg / l. What concentration does that correspond to? (2 P)

Can you calculate how many kilograms of quicklime (CaO) can be obtained from 2t of lime (calcium carbonate)? (2 P)

Use a schematic phase diagram to explain two methods of gas liquefaction. (2 P)

A buffer solution contains acetic acid (pKS.= 4.75) and sodium acetate in a molar ratio of 10: 1. Calculate the pH. (2 P)

In the first lighter, developed by Johann Wolfgang Döbereiner in 1823, zinc was brought into contact with sulfuric acid and the gas produced in the process was guided past a platinum sponge. Explain the chemical relationships and reactions that occur, also using reaction equations. (2 P)

What is oil made of, how is it refined and what products are made? (2 P)

Draw the structural formula of L-lactic acid (= L-2-hydroxy-propanoic acid) (2 P)

Describe the reaction of methane with chlorine in as much detail as possible (reaction conditions, mechanism, products). (2 P)

Describe the general structure of fats (triglycerides) using a structural formula. (2 P)

For lecture 62-082.1: "Basics of Chemistry"

  1. Explain in as much detail as possible why nitrogen or argon is often used as a protective gas in chemical syntheses. (2 P)
  2. Calculate how many grams 1 liter of carbon dioxide weighs under normal conditions. (2 P)
  3. Draw the structural formulas (= valence line formula) of the respective compounds of hydrogen with a) sulfur, b) bromine, c) phosphorus and d) silicon. (2 P)
  4. Explain (using a schematic drawing, among other things) how the distillation of alcohol ("schnapps distilling") works. (2 P)
  5. Calculate the pH value that results when you add 12 g of acetic acid (weak acid, pKS. = 4.75) dissolves in 2 L water? (2 P)
  6. Name four large-scale chemical processes based on redox reactions with the corresponding reaction equations. (2 P)

7. Explain the different physical properties of graphite and diamond by their respective molecular structure. What is the hybridization of the carbon atoms? (2 P)

  1. Write down a reaction equation with structural formulas for the reaction of trans-2-pentene with chlorine and name the product. (2 P)

9. When is a chemical compound aromatic? Enter the name and structural formula for a pure hydrocarbon aromatic as well as for a heteroaromatic. (2 P)

For lecture 62-082.1: "Basics of Chemistry"

  1. Compare the atoms 50 Ti and 50 V with regard to their structure from elementary particles (similarities, differences). (2 P)
  2. Calculate how many grams n = 4 mol calcium hydroxide weigh? (2 P)
  3. Explain the octet rule using the structural formulas of water, ammonia, methane and boron trichloride. (2 P)
  4. Draw up a reaction equation for the exothermic chlorine oxyhydrogen reaction, in which hydrogen and chlorine are converted into hydrogen chloride gas, and explain how changes in temperature and pressure affect the equilibrium. (2 P)
  5. Cleopatra allegedly dissolved pearls in vinegar in order to ingest them.
    Draw up a reaction equation for this process, assuming that pearls are mainly made of lime. (2 P)
  6. Which oxidation numbers does sulfur have in the following compounds: (2 P)
    H2S, S8, H2SO4, HSO3 -, S.2O4 2-, p2O8 2- .
  7. Explain the different corrosion resistance of aluminum and iron.
    (2 P)
  8. Use an example of your choice to explain what is meant by chirality. (2 P)
  9. Under which conditions, to which products and by which mechanism does bromine react with a) cyclohexane, b) cyclohexene and c) benzene (outdated: benzene)? (2 P)
  10. Explain in as much detail as possible why fats do not have a specific melting temperature but a melting range? (2 P)

for lecture 62-082.1: "Basics of Chemistry"

  1. Calculate how many atoms (pieces!) There are in 1 g of copper. (2 P)
  2. Describe the physical and chemical properties of bromine. (2 P)
  3. In one experiment, hydrochloric acid is poured over a piece of a) lime, b) zinc and c) copper. Describe what happens and explain the chemical relationships with the help of reaction equations. (2 P)
  4. Compare the physical states of CO2, H2O and H2S and explain the relationships based on the electronic structure of the molecules. (2 P)
  5. Calculate how many grams of perchloric acid (HClO4, strong acid) must be dissolved in 200 ml of water so that a pH value of 2 arises. (2 P)
  6. Complete the following reaction equation with stoichiometric factors and give the respective oxidation numbers of Mn and C: (2 P)
    MnO4 - + C2O4 2- + H3O + --- & gt Mn 2+ + CO2 + H2O
  7. Describe the large-scale chlor-alkali electrolysis with the help of text, schematic drawing and reaction equation. (2 P)
  8. Draw the structural formulas of a) dimethyl ether b) dimethylamine
    c) 1,2-dimethylbenzene and d) 2,2-dimethyl-1-propanol. (2 P)
  9. Under which conditions, to which products and according to which mechanisms does bromine react with a) cyclohexane, b) cyclohexene and c) benzene? (2 P)
  10. Use the example of the a-L-amino acids to explain what is meant by chirality.
    (2 P)

For lecture 62-082.1: "Basics of Chemistry"

  1. Name and characterize the three types of radioactive radiation. (2 P)
  2. Calculate in which volume of solvent you have to dissolve a substance amount of n = 5 mol, so that a concentration of c = 0.25 mol / L results? (2 P)
  3. Name the elements that occur as two-atom molecules and explain the fact using the respective structural formulas. (2 P)
  4. Write up an equation of reaction for the combustion of ethene (C.2H4) and calculate the enthalpy of reaction DHR. from the following enthalpies of formation: (2 P)
    ΔHf in kJ / mol: C2H4: 52.3 CO2: -393.5 H.2O: -241.8
  5. Describe the principle of an acid-base buffer using an example. (2 P)
  6. Calculate what voltage a galvanic element delivers,
    in which a copper electrode is immersed in a Cu 2+ solution with c = 10 mol / L
    and immerses a zinc electrode in a Zn 2+ solution with c = 0.1 mol / L. (2 P)
    Standard redox potentials: DE0(Zn) = -0.76V
  7. What is sodium hypochlorite and why is it bleaching and disinfecting? (2 P)
  8. Draw the structural formulas of a) 2-methyl-1,3-butadiene, b) trimethylamine, c) hexachlorocyclohexane, d) trifluoroacetic acid: (2 P)
  9. Describe a detection reaction to differentiate between aldehydes and ketones (observation and reaction equation). (2 P)
  10. How do amylose and cellulose differ in terms of their chemical structure, their molecular geometry, their biochemical significance and their detectability by iodine? (2 P)

For lecture 62-082.1: "Basics of Chemistry"

  1. Explain what halogens are and how they are “salt formers”. (2 P)
  2. 1 g of an alkali hydroxide dissolved in 250 ml of water gives a concentration of
    c = 0.1 mol / L. Calculate which connection it is. (2 P)
  3. Write down an equation for the reaction of aluminum with bromine. (2 P)
  4. Give the names and the molecular formulas of five common acids and their corresponding bases. (2 P)
  5. A maximum of 7 g of lime can be dissolved in 1000 L of water. Calculate the solubility product of calcium carbonate. (2 P)
  6. Iodine reacts easily with itself in aqueous solution. Iodide and iodate ions (IO3 - ) develop. Write up an equation for the reaction. Which special case of a redox reaction is it? (2 P)
  7. Compare metal, graphite, and diamond for conductivity based on their molecular and electronic structure. (2 P)
  8. Explain what cis / trans (= Z / E) isomerism means using a self-selected example (name + structural formula). (2 P)
  9. Use suitable examples to explain how and why the melting and boiling points of alkanes change with increasing chain length? (2 P)

For lecture 62-082.1: "Basics of Chemistry"

  1. Calculate how much heavier a liter of helium is than a liter of hydrogen. (2 P)
  2. Which ions do the elements Al, Ba, S, N, Rb, F, Mg, O form when they reach the noble gas configuration? (2 P)
  3. Complete the following reaction equations: (2 P)
    + Ca (OH)2 --- & gt Ca (NO3)2 + 2 H.2O
    N2H4 + O2 --- & gt N2 + H2O
    Pb + HClO4 --- & gt Pb (ClO4)2 +
    Al + Fe2O3 --- & gt Fe +
  4. Use examples from inorganic and organic chemistry to explain what catalysts are and how they work. (2 P)
  5. Calculate how many grams of caustic soda (NaOH) it takes to completely deprotonate 9.8 g of phosphoric acid. (2 P)
  6. Describe the differences in the molecular structure and the material properties of diamond and graphite with the help of hybridization. (2 P)
  7. Describe the production of aluminum by fused-salt electrolysis in as much detail as possible. (2 P)
  8. Use the structural formulas of aminobutanoic acid to explain what is meant by positional isomerism. (2 P)
  9. Describe in as much detail as possible (reaction conditions, mechanisms, products) which different reactions bromine can enter into with toluene (methylbenzene). (2 P)
  10. Describe in as much detail as possible what the difference between alpha- and beta- (D) -glucose is and what influence it has on the structure and properties of cellulose and amylose. (2 P)

For lecture 62-082.1: "Basics of Chemistry"

  1. Calculate how many grams of calcium chloride have to be dissolved in 300 ml of water so that the solution has a concentration of c = 0.2 mol / L? (2 P)
  2. Which ions have the same electron configuration as neon? (2 P)
  3. Use common examples to explain what radicals are, how they arise, what properties they have and in which reactions they play an essential role. (2 P)
  4. Explain the terms a) emulsion, b) sublimation and c) suspension using typical examples! (2 P)
  5. Set up neutralization reactions in which as a product a) calcium sulfate,
    b) lithium hydrogen carbonate, c) ammonium nitrate and d) sodium hydrogen phosphate is formed. (2 P)
  6. Write up a reaction equation for the dissolution of silver by nitric acid, where i.a. Ag + ions and nitrogen monoxide (NO) are formed. Use the relevant oxidation numbers and set up partial oxidation and reduction equations. (2 P)
  7. What is oil made of, how is it refined and what products are made? (2 P)
  8. Metals themselves have neither smell nor taste, but they catalyze reactions on the skin that produce substances that cause the typical “metallic” smell or taste, e.g. B. Oct-1-en-3-one. Draw the structural formula.
    (2 P)
  9. Under which conditions, to which products and by which mechanism does bromine react with a) cyclohexane, b) cyclohexene and c) benzene? (2 P)
  10. Explain in as much detail as possible what is meant by the primary and secondary structure of proteins. (2 P)

for lecture 62-082.1: "Basics of Chemistry"

  1. Calculate the proportions of the two isotopes of silver 107 Ag and 109 Ag, so that an average atomic mass of mA.= 107.868 results. (2 P)
  2. Coca Cola contains 700 mg of phosphoric acid per liter. What is the concentration of phosphoric acid? (2 P)
  3. Explain the bond character and the stoichiometric ratios in compounds from a) Al and F, b) P and I, c) Cu and Zn. (2 P)
  4. Use a phase diagram to explain two methods by which gases can be liquefied.
  5. Explain in as much detail as possible and with equations what the difference between pH and pKS.-Is worth?

For lecture 62-082.1: “Fundamentals of Chemistry” on Thursday, August 23, 2012 from 10.15 am to 11.45 am

  1. What electron configuration does chlorine have and how can this explain the bond in hydrogen chloride (draw the structural formula)? (2 P)
  2. Draw up a reaction equation for the combustion of propane gas to carbon dioxide and water (2 P)
  3. Draw the structural formulas of a) carbonic acid, b) sulfite ion c) ammonia)
    d) Hydrocyanic acid (2 P)
  4. Use examples to explain what catalysts are and how they work. (2 P)
  5. Draw the titration curves of hydrochloric acid, sulfuric acid and acetic acid schematically in a diagram and compare them. (2 P)
  6. Which of the metals Cu, Fe, Ag, Zn, Ca dissolve in 1 molar hydrochloric acid? Give reasons for your answer and set up an exemplary reaction equation. (2 P)
  7. Describe the large-scale chlor-alkali electrolysis with the help of a schematic drawing and reaction equations. (2 P)
  8. Describe the chemical reaction of an alkane with a halogen in as much detail as possible. (2 P)
  9. Explain the terms ortho, meta and para position using an example of an aromatic that you have chosen yourself. (2 P)
  10. Explain the similarities and differences between glucose and fructose using their structural formulas. (2 P)

For lecture 62-082.1: "Basics of Chemistry" on Monday, July 23, 2012

  1. In the case of the Fukushima reactor disaster, inter alia the radioactive isotope cesium-137 has been released. How many protons and neutrons does its nucleus consist of and what does it mean that it is a b-emitter? (2 P)
  2. What is the molar amount of 2 g of aluminum hydroxide? (2 P)
  3. Calculate how much water you have to add to 25 ml of a solution so that the concentration of c0= 0.2 mol / L on c1= 0.04 mol / L decreases? (2 P)
  4. Set up the reaction equation and the formula for the law of mass action for the combustion of carbon monoxide to carbon dioxide. (2 P)
  5. A maximum of 7 g calcium carbonate can be dissolved in 1000 liters of water. Do you calculate the solubility product? (2 P)
  6. Give the names and formulas of a) two common oxidizing agents and b) two common reducing agents. (2 P)
  7. Draw the structural formulas of a) glycerin, b) acetone, c) butyric acid and
    d) formaldehyde. (2 P)
  8. Draw up a reaction equation for the detection of ethanol in test tubes using potassium dichromate (K2Cr2O7) whereby acetaldehyde (= ethanal) and green chromium (III) ions are formed. (2 P)
  9. Use the a-L-amino acids as an example to explain what is meant by chirality. (2 P)
  10. What do you know about polystyrene? (2 P)

for lecture 62-082.1: “Basics of Chemistry” on Thursday, May 24th. 6:15 pm-7:00pm

  1. Calculate in how many liters of water 4.2 g of sodium hydrogen carbonate have to be dissolved in order to produce a concentration of 0.4 mol / L?
  2. Write up an equation for the reaction of aluminum with sulfur (p8) to aluminum sulfide.
  3. * (V): Assign the following substances to their groups:
    Basic substance homogeneous mixture Heterogeneous mixture
    Bronze r rr
    Milk r r r
    Citric acid r r r
    Granite r r
    Calcium carbonate r r r
    Vodka r r r
    White gold r r r
    Tap water r r r
  1. Explain - also with reaction equations - why carbonated mineral water bubbles more when you add lemon.
  2. Blood has a pH of 7.3. Can you calculate how many hydronium ions a person has in 6 liters of blood?

for lecture 62-082.1: "Basics of Chemistry"

1. What percentage of the isotopes of copper 63 Cu and 65 Cu occur when the mean atomic mass is 63.546 u? (2 P)

2. Name the 10 most common elements of the earth's crust and what does the earth's core essentially consist of? (2 P)

3. When 149.1 g of an alkali halide is dissolved in 10 liters of water a solution of the concentration
c = 0.2 mol / L, which alkali halide is it? (The calculation method must be comprehensible!) (2 P)

4. Draw up a reaction equation for the production of ammonia from hydrogen and nitrogen and explain why a higher yield of product is obtained at high pressure? (2 P)

5. Calculate the pH levels that adjust when one
(a) 0.1 mole sodium hydroxide
(b) 0.1 mol of nitric acid
(c) 0.1 mol acetic acid (weak acid, pKS.=4,75)
dissolves each in 10 L water? (2 P)

6thThe reaction of sulfuric acid with hydrogen iodide produces, among other things. elemental iodine and hydrogen sulfide gas. Set up the reaction equations for the oxidation, the reduction step and the overall reaction, taking into account the relevant oxidation numbers. (2 P)

7. Describe the mode of operation of a galvanic element using the example of the Daniell element (Cu / Zn element) (drawing, text, reaction equations) and calculate the voltage it delivers under standard conditions.
DE0(Cu) = 0.35 V DE0(Zn) = -0.76 V (2 P)

8. Give the names and structural formulas of the compounds that are formed when (a) 1-butanol and (b) 2-butanol are oxidized as much as possible (without destroying the C-C structure)? (2 P)

9. Describe in as much detail as possible (reaction conditions, mechanism, products) which different reactions bromine can enter into with toluene (= methylbenzene).
(2 P)

10. How do amylose and cellulose differ in terms of their chemical structure, their molecular geometry, their biochemical significance and their detectability by iodine? (2 P)

For lecture 62-082.1: "Basics of Chemistry"

  1. How many grams does 2 mol of phosphoric acid weigh? (2 P)
  2. Write up an equation for the reaction of aluminum with sulfur (p8) to aluminum sulfide. (2 P)
  3. What are the two main components of air and how can they be separated? (2 P)
  4. Acetic acid (CH3-COOH) has a pKS.Value of 4.76, trifluoroacetic acid (CF3-COOH) on the other hand pKS.= 0.23. Explain where this difference comes from and how it affects acid properties. (2 P)
  5. A maximum of 7 g calcium carbonate can be dissolved in 1000 liters of water. How big is the solubility product? (2P)

for lecture 62-082.1: "Basics of Chemistry"

on Friday, December 11th, 2010 6:15 pm-7:00pm

  1. An isotonic 0.9% saline solution contains 0.9 g sodium chloride per liter of water. Which substance concentration does this correspond to? (2 P)
  2. Explain the bond character and the stoichiometric ratios in compounds from a) Mg and F, b) C and H, c) Cu and Zn. (2 P)
  3. Write down a reaction equation for the reaction of aluminum with oxygen (2 P)
  4. Explain what ampholytes are using the structural formulas of at least 2 examples. (2 P)
  5. Battery cells have leaked in a submarine and 100 L acid with a pH value of 0 must be neutralized. How many grams of calcium hydroxide do you need for this? (2 P)

for lecture 62-082.1: "Basics of Chemistry"

on Thursday, August 19th, 2010 from 2.15pm - 3.45pm

1. Explain in as much detail as possible what the similarities and differences between 12 C and 13 C atoms are. (2 P)

2. How many grams of oxygen do you need to completely burn 1.6 g of methane?
Write up an equation for the reaction. (2 P)

3. Explain the terms emulsion, suspension and sublimation using examples. (2 P)

4. Describe the technical lime cycle. (2 P)

5. Battery cells have leaked in a submarine and 100 liters of acid have to be used
pH value 1 can be neutralized. How many grams of caustic soda (NaOH) do you need for this?

6. Write up a reaction equation for the dissolution of copper in nitric acid, where i.a. Nitric oxide and Cu 2+ ions are formed. (2 P)

7. Explain the properties of graphite by its molecular structure. What is the hybridization of the carbon atoms? (2 P)

8. An organic compound reacts positively to Fehling and Tollens reactions (silver mirror test) and has a molar mass of 44 g / mol. Enter the name and structural formula of the compound and explain your decision (2 P)

9. When is a chemical compound aromatic (4 criteria)? Enter the name and structural formula for a pure hydrocarbon aromatic as well as for a heteroaromatic. (2 P)

10. Describe the general structure of those important for protein synthesis

a-L-amino acids. Why are amino acids usually soluble in water? (2 P)

For lecture 62-082.1: "Basics of Chemistry"

on Thursday, July 22nd, 2010 from 2.15pm - 3.45pm

  1. How many grams of sodium hydrogen carbonate have to be dissolved in 100 ml of water so that the solution has a concentration of c = 10 -2 mol / l? (2 P)
  2. Explain similarities and differences between CO2 and SiO2 based on their ties. (2 P)
  3. Explain what the law of mass action is and describe an application in inorganic chemistry. (2 P)
  4. 23 mg of sodium are added to 100 ml of water. Describe what happens, set up a reaction equation and calculate the pH value after the reaction. (2 P)
  5. Name a) two common oxidizing agents and b) two common reducing agents (2P)
  6. Explain why the blue salt, copper sulfate pentahydrate, turns white when heated, but turns blue when it is subsequently dissolved in water. (2 P)

7. Explain what electrolysis is using a large-scale process (2 P)

  1. Describe a detection reaction to differentiate between aldehydes and ketones (observation and reaction equation). (2 P)
  2. Draw the structural formulas of formaldehyde, aniline, ethyl acetate and pyridine (2 P)
  3. Describe the structure of the DNS in as much detail as possible. (2 P)
Mid-term exam
For lecture 62-082.1: "Basics of Chemistry"
On Fri. 04.06. 17: 15-18: 00
  1. Which ions have the same electron configuration as argon? (2 P)
  2. Use typical examples to describe a) the metal bond, b) the ionic bond, c) the non-polar atomic bond and d) the polar atomic bond. When does which type of bond arise and how does the character of the bond affect the properties of the fabric? (2 P)
  3. Use a schematic phase diagram to explain two methods of gas liquefaction. (2 P)
  4. Calculate how many grams of sodium hydroxide it takes to convert 1 mole of phosphoric acid into sodium phosphate.
    Write down the reaction equation. (2 P)
  5. Set up a reaction equation for the reaction of sulfuric acid with hydrogen iodide, in which elemental iodine and hydrogen sulfide are formed, and state the relevant oxidation numbers. (2 P)
Written exam (rewriting date)
For lecture 62-082.1: "Basics of Chemistry"
On Monday, February 22nd, 2010 from 9.15 a.m. to 10.45 a.m.
  1. How many grams of magnesium iodide have to be dissolved in 200 ml of water so that the concentration is c = 10 -2 mol / L? (2 P)
  2. Explain with the help of hybridization the geometry of a carbon dioxide molecule and its influence on the dipole character of the molecule and the properties of the compound. (2 P)
  3. Draw the valence line formulas of a) carbonic acid b) hydrogen peroxide
    c) nitrate ion d) ammonium ion. e) acetate ion. (2 P)
  4. Explain the terms hydrophilic and lipophilic using the solvents water, diethyl ether, methanol and hexane. (2 P)
  5. In one experiment, hydrochloric acid is poured over a piece of a) lime, b) zinc and c) copper. Describe what happens and explain the chemical relationships with the help of reaction equations. (2 P)
  6. 1 ml of hydrochloric acid with a concentration of c = 0.1 mol / l is diluted with 99 ml of water. What is the pH of the resulting solution? (2P)
  7. Describe the large-scale aluminum production in as much detail as possible. (2 P)
  8. Use an example chosen by you to describe what an esterification is (reaction equation with structural formulas). (2 P)
  9. When is a chemical compound aromatic? Enter the name and structural formula for a pure hydrocarbon aromatic as well as for a heteroaromatic. (2 P)
  10. Describe the structure of a fat using a structural formula and explain what polyunsaturated fatty acids (w-3,6,9-) are and why they are more common in liquid fats (oils). (2 P)
Exam
For lecture 62-082.1: "Basics of Chemistry"
On Wed. 02/10/2010 from 9.15 a.m. - 10.45 a.m.
  • Describe in as much detail as possible the structure of an atom of the element fluorine from its elementary particles (number, location, mass, charge). (2 P)
  • With how many grams of chlorine do you have to react with 4 g of calcium in order for complete conversion to take place, i.e. neither calcium nor chlorine remains?
    Draw up a reaction equation, the calculation must be understandable. Pay attention to complete equations and correct units. (2 P)
  • Explain the terms mesomerism and boundary structure formulas - if possible using the example of the carbonate ion. (2 P)
  • Draw the phase diagram of water schematically and show why ice melts under pressure and how does a pressure cooker work? (2 P)
  1. Which pH values ​​arise when you add 1 mol of a) nitric acid (strong acid) and b) nitrous acid (weak acid pKS.= 3.29) in each 1 liter of water. Explain the different properties of the acids using the structural formulas. (2P)
  • Describe as precisely as possible what sodium hypochlorite is and why it disinfects and bleaches? (2 P)
  • Describe in as much detail as possible how you can clean copper from foreign metals (such as Fe, Zn, Ag, Au) by means of electrolysis. (2 P)
  • Give the names and structural formulas of all isomers of butene. (2 P)
  • Give examples of a primary, sec. and tert. Alcohol (structural formula)
    and use reaction equations to show how they differ in their oxidizability. (2 P)
  1. Describe what the difference between a- and b- (D) -glucose is and what influence it has on the structure and properties of cellulose and amylose. (2 P)

for lecture 62-082.1: "Basics of Chemistry"

on Thursday, 08/20/2009 from 2:15 p.m. - 3:45 p.m.

  1. Explain how Ernest Rutherford inferred his atomic model from the experimental results of the scattering of a-particles on gold foil. (2 P)
  2. What is the molar mass of ammonium nitrate? (2 P)

3. Which connections are formed from the following pairs of elements:
Li / O Al / F C / S N / H? Indicate whether each bond is ionic or covalent. (2 P)

  1. Name compounds that do not meet the octet rule and explain their behavior in aqueous solution and their function in electrophilic substitution on aromatic compounds. (2 P)
  2. One tank contains 1000 liters of an aqueous solution with pH = 2. How many grams of caustic soda (NaOH) do you have to add to neutralize? (2 P)

6. What is electrolysis? For the electrolysis of copper iodide, give the partial reaction equations at the anode and the cathode as well as the overall reaction equation. (2 P)

  1. Use hybridization to describe the molecular structure of diamond and graphite and explain why diamond is hard and graphite is soft. (2 P)

8. Which structural isomers are there of hexane (structural formula + name)? (2 P)

  1. Use examples to describe what polycyclic aromatic hydrocarbons (PAHs) are.
  2. Describe in as much detail as possible what is meant by the primary and secondary structure of proteins. (2 P)

for lecture 62-082.1: "Basics of Chemistry"

on Thursday, July 23, 2009 from 2.15pm - 3.45pm

  1. Draw up an equation for the reaction of butane gas combustion. (2 P)
  2. In how much water do you have to dissolve 9 g of fructose (fruit sugar) so that the concentration is 0.5 mol / L? (2 P)
  3. Draw the valence line formulas of a) nitrous acid b) hydrogen carbonate ion c) hydrogen sulfide d) phosphoric acid. (2 P)
  4. Name the 10 most common elements of the earth's crust and what does the earth's core essentially consist of? (2 P)
  5. What are the pH values ​​when you
    a) 0.1 mol strong, mono-protonic acid and
    b) 0.1 mol weak, monoprotonic acid (pKS. = 4,0)
    dissolve in 10 liters of water?
    ------------------------------------------------------------------------------------------------------
  6. Write up an equation for the reaction of potassium dichromate (K2Cr2O7) with hydrochloric acid to form chromium (III) ions and chlorine. (2 P)
  7. Describe the structure and function of a car battery (lead battery) using text, drawings and reaction equations. (2 P)
  8. Give the name and the structural formula of the compound that is formed when 2-hexene is allowed to react with chlorine. (2 P)
  9. Use the structural formulas of aminopropanoic acid to explain what is meant by positional isomerism. (2 P)
  10. Draw the structural formula of D-glucose and describe its reaction with alkaline copper sulfate solution (Fehling's sample) (observation + reaction equation). (2 P)

for lecture 62-082.1: "Basics of Chemistry"

1. The new EU guideline value stipulates that a car per 1 km has a maximum carbon dioxide emission of 120 g. How many liters of CO2 does this correspond under normal conditions? (2 P)

2. Which ions have the same electron configuration as argon? (2P)

3. Draw schematically the phase diagram of water and show why ice melts under pressure and how one
Does the pressure cooker work? (2P)

4. The solubility product of calcium carbonate is 5 × 10 -9 mol 2 / L 2. How many milligrams of calcium carbonate dissolve in one liter of water? (2 P)

5. Explain in words and reaction equations how acid rain occurs and what reaction causes it to dissolve limestone?

for lecture 62-082.1: "Basics of Chemistry"

on Wed. February 25th, 2009 from 2.15pm - 3.45pm

1. What are alkaline earth metals? Name at least three representatives of this element group. What is their valence electron configuration? Write down an equation for the reaction of an alkaline earth metal with a halogen. (2 P)

2. How many grams of ammonium nitrate do you have to dissolve in 100 ml of water in order to obtain a solution with a molar concentration of c = 10 -2 mol / l? (2 P)

3. The oxidation reaction of sulfur dioxide gas to sulfur trioxide gas is exothermic and reversible. Draw up a reaction equation and explain how and why an increase in temperature or pressure affects the position of equilibrium. (2 P)

4. Explain - also using examples - what radicals are, how they arise and how they react. (2 P)

5. Give the reaction equations for the reaction of phosphoric acid with sodium hydroxide solution and draw the titration curve. (2 P)

6. Describe in as much detail as possible what a normal hydrogen electrode is and calculate its electrochemical potential in neutral water. (2 P)

7. Describe in as much detail as possible (sketch, reaction equations, text) how iron is produced from ore in a blast furnace.

8. Give the name and structural formula of the simplest alkane with a chiral center. (2 P)

9. When is a chemical compound aromatic? Enter the name and structural formula for a pure hydrocarbon aromatic as well as for a heteroaromatic. (2 P)

10. Explain in as much detail as possible what the difference is between mineral oil (= petroleum) and fatty oil. (2 P)

on Wed. 02/11/2009 from 2.15 p.m. - 3.45 p.m.

  1. What electron configuration does nitrogen have and how can this explain the bonds in ammonia? (2 P)
  2. How many grams of the salt silver nitrate have to be dissolved in 250 ml of water in order to obtain a solution with a molar concentration of c = 0.1 mol / L? (2 P)
  3. Name inorganic and organic compounds with a triple bond! (2 P)
  4. Use examples to explain what catalysts are and how they work. (2 P)
  5. What are the pH values ​​when you
    (a) Dissolve 4 g sodium hydroxide in 10 L water?
    (b) 4 mol acetic acid (weak acid, pKS.= 4.75) dissolve in 100 ml water?
    (2 P)
  6. Some of the Pb 2+ ions in an aqueous solution can be used as poorly soluble lead sulfide PbS (solubility product Lp(PbS)= 10 -28 mol 2 / L 2) precipitate by introducing hydrogen sulfide gas.
    a) Set up the reaction equation for this.
    b) What is the minimum concentration of sulfide ions in the solution so that c (Pb 2+) <10 -25 mol / L? (2 P)
  7. Write up an equation for the reaction of oxalic acid (C.2H2O4) with potassium permanganate (KMnO4) in an acidic environment to carbon dioxide and Mn 2+ ions. (2 P)
  8. Draw the structural formulas of acetone, formaldehyde, ethyl acetate and pyridine. (2 P)
  9. Under what conditions, to what products and by what mechanism does chlorine react with a) cyclohexane, b) cyclohexene and c) benzene? (2 P)
  10. How do starch and cellulose differ in terms of their chemical structure, their molecular geometry, their biochemical significance and their detectability by iodine? (2 P)

for lecture 62-082.1: "Basics of Chemistry"
on Fri. 5.12. 2008 from 5 p.m. to 6 p.m.

1. Which connections are formed from the following pairs of elements:
Mg / Br C / H K / O Al / F C / S?
Indicate in each case whether the bond is ionic or covalent. (2 P)

2. How many grams of calcium chloride do you have to dissolve in 100 ml of water in order to achieve the concentration of the substance
c = 10 -2 mol / l?

  1. Can you explain the terms emulsion, suspension and sublimation using examples? (2 P)
  2. Write up an equation for the reaction of aluminum with oxygen. (2 P)

5. There are 300 liters of an aqueous solution with pH = 2 in one tank. How many liters of sodium hydroxide solution with a concentration of c = 0.5 mol / L do you have to add for complete neutralization? The calculation must contain complete equations with units! (2 P)

to lecture 13.820: "Basics of Chemistry"

on Thursday, August 28th, 2008 from 2.15pm - 3.45pm

  1. Describe why and to what extent (example reaction equation) lithium, sodium and potassium are similar in their chemical behavior? (2 P)
  2. Name a salt in which the cations and anions occur in a ratio of 2: 3. (2P)
  3. What is an ampholyte? Give at least two examples. (2P)
  4. One tank contains 1000 liters of an aqueous solution with pH = 3.How many liters of 0.5 molar sodium hydroxide solution do you have to add to neutralize? (2P)
  5. Write down the equation for the reaction of nitrogen dioxide with ammonia to form nitrogen and water. (2P)
  6. Do you describe the electrolysis of water in as much detail as possible? How does the pH value at the anode and cathode change? (2P)
  7. Draw the structural formula of 2,2,4-trimethylpentane, the main component of gasoline. (2P)
  8. Give the name and the structural formula that results when 2-butyne is completely reacted with bromine. (2P)
  9. How many grams of grape sugar do you have to dissolve in 100 ml of water to get a 1 molar solution? (2P)
  10. Can you describe what a disaccharide is in as much detail as possible and with an example? (2P)

About lecture 13.820: "Basics of Chemistry"

on Thursday, July 17th, 2008 from 2.15pm - 3.45pm

1. What electron configuration does oxygen have and how can it be used to explain the bonds in water. (2 P)

    Draw the valence line formulas of the following molecules and ions:
    Ammonia, nitric acid, carbonate ion, sulfite ion. (2 P)

    How much water do you need to dilute 50 ml of 1 molar hydrochloric acid with so that the resulting solution has a pH of 2? (2 P)

    Describe the large-scale chlor-alkali electrolysis with the help of text, a schematic drawing and reaction equations. (2 P)

8. Enter the name and structural formula for the simplest alkane with a quaternary carbon atom. (2 P)

    Do you describe a detection reaction to differentiate between aldehydes and ketones (observation and reaction equation)? (2 P)

to lecture 13.820: "Basics of Chemistry"

on Thursday, May 29 2008 from 4.15pm to 5pm

  1. Describe the structure of a fluorine atom from its elementary particles in as much detail as possible. (2 P)
  2. Which connections are formed from the following pairs of elements:
    a) Mg / Br b) C / H c) K / O d) Al / S e) S / H? Indicate whether each bond is ionic or covalent. (2 P)
  3. How does a material separation by distillation work? Draw an apparatus schematically. (2 P)
  4. What is acid rain, how does it arise and what reaction (equation!) Does it dissolve limestone? (2 P)
  5. In how many liters of water do you have to add 2 g caustic soda (NaOH) to create a solution with pH = 12? Complete equations with units! (2 P)

About lecture 13.820: "Basics of Chemistry"

on Thursday, February 28th, 2008 from 2.15pm - 3.45pm

    What are Alkali Metals? Name at least three representatives of this element group. What is their valence electron configuration? Why do they not occur in nature in elemental form? Give the reaction equation of an alkali metal with chlorine.

About lecture 13.820: "Basics of Chemistry"

WS 07/08 on February 14th, 2008 from 2.15pm - 3.45pm

1. Name at least two noble gases. Why are noble gases so unreactive? (2 P)

2. What is electronegativity and how does it influence the bonds and thus the properties of methane, ammonia, water and in hydrogen fluoride? (2 P)

3. How do the oxides of sodium, magnesium and aluminum differ in their composition? (2 P)

4. Describe the dissolution of calcium carbonate in hydrochloric acid using a reaction equation. (2 P)

5. What are the pH values ​​when you
(a) Dissolve 0.01 mol of sodium hydroxide in 100 ml of water?
(b) 1 mol acetic acid (weak acid, pKS.= 4.75) dissolve in 1 liter of water? (2 P)

6. Describe the mode of operation of a galvanic element using the example of the Daniell element (Cu / Zn element). What voltage does it deliver under standard conditions? DE0(Cu) = 0.35 V DE0(Zn) = -0.76 V (2 P)

7. Explain the properties of graphite by its molecular structure. What is the hybridization of the carbon atoms? (2 P)

8. Give the structural formula and the name for the (a) simplest trihydric alcohol and for the (b) simplest tertiary alcohol. (2P)

9. When is a chemical compound aromatic? Enter the name and structural formula for a pure hydrocarbon aromatic as well as for a heteroaromatic. (2 P)

About lecture 13.820: "Basics of Chemistry"

on Thursday, August 23, 2007 from 2.15pm - 3.45pm

1. Explain what isotopes are - if possible using the example of carbon? (2 P)

2. Why do metals have good electrical and thermal conductivity? (2 P)

3. Give the reaction equations for the step-by-step protolysis of phosphoric acid and draw the titration curve schematically. (3 P)

4. How many mg of silver chloride dissolve in 10 liters of water? (2 P)
KL.(AgCl) = 10 -10 mol 2 / l 2 M (AgCl) = 143.5 g / mol

5. Do you give two common oxidizing agents and two common reducing agents? (2 P)

6. Describe the passivation (corrosion resistance) of aluminum with words and a reaction equation. (2 P)

7. Describe how metal complexes are built up, if possible using hexaqua calcium dichloride. What is a coordinative bond? (3 P)

8. Enter the structural formula for nitroglycerin. What alcohol and acid is it made from? (2 P)

9. Describe the fat saponification with the help of reaction equations. Why do the soap molecules dissolve fat? (3 P)

10. Describe the structure of the DNA schematically. (2 P)

Exam
About lecture 13.820: "Basics of Chemistry"
SS07 on Thursday, July 19, 2007 from 2:15 p.m. - 3:45 p.m.

1. What electron configuration does nitrogen have and how can this explain the bonds in ammonia? (2 P)

2. Draw the phase diagram of water schematically and show why ice melts under pressure? (2 P)

3. Give the names and the empirical formulas of five different acids and their corresponding bases. (3 P)

4. If you dissolve 1.8 g of glucose in 100 ml of water, you get a 0.1 molar solution.
From this, calculate the molar mass of glucose. The invoice must be understandable. (2 P)

5. Describe what happens if you add 23 mg of sodium (MN / A= 23 g / mol) in 100 ml of water? Write down the reaction equation. Calculate how the pH changes. (3 P)

6. What is electrolysis? Give the partial reaction equations at the anode and the cathode as well as the overall reaction equation for the electrolysis of water. (2 P)

7. Describe the three steps of nitric acid production using the Ostwald process
with the help of reaction equations. (2 P)

8. Enter the names and structural formulas for three different compounds with the empirical formula C.3H8O on. (3 P)

9. Explain the terms ortho-, meta- and para-position using a self-chosen example of an aromatic. (2 P)

10. Describe the general structure of the α-L-amino acids that are important for protein synthesis. Why are amino acids usually soluble in water? (2 P)
Examination exam SS06
For chemistry lecture 13.820
On 07/13/2006

1. Draw the electron energy scheme for the ground state of the carbon atom. Which hybridizations of carbon are present in benzene, methane, Erhan, ethene and ethanol?

2. How is the chemical concentration defined?

3. How do Lewis acids differ from Brönstedt acids?

4. What is the pH value of a solution of 4 g sodium hydroxide in 100 L water? (Atomic mass Na = 27)

5. The solubility product for silver chloride is 10 -10 mol 2 / L 2. What is the concentration of silver ions in a silver chloride / water suspension?

6. What chemical reactions take place in a car battery?

7. How is sulfuric acid made? Brief description of the process with equations!

8. What is an electrophilic substitution (example - if possible, also diazotization)?

9. What is the difference between starch and cellulose?

10. Name the monomers from which the following plastics are built:
Polyethylene, polystyrene, PVC, polyester-PET?

About lecture 13.820: "Chemistry for students of physics, computer science, geology, mineralogy, wood management, business administration (merchandise), high school teaching (biology), GTW"

  • Give the valence electron configuration of the atoms of oxygen, neon and magnesium. Which typical reactions can be expected in each case?
  • Draw the valence formulas of the following molecules:
    a) HCCl3 b) SO4 2- c) H.2O2 d) HCN
  • How many mg of potassium permanganate (KMnO4) have to be weighed in to produce 100 ml of a solution with a molar concentration of 0.1 mol / l?
    Molar masses: M (O) = 16 g / mol, M (K) = 39.1 g / mol, M (Mn) = 54.9 g / mol
  • The solubility product of lead sulfate is Lp (PbSO4) = 10 -8 mol 2 / l 2.
    How many mg of lead (in the form of Pb 2+) are there in 1 liter of a saturated solution?
    Molar mass: M (Pb) = 207 g / mol
  • Complete the following reaction equation with stoichiometric factors and give the respective oxidation numbers of Mn and C:
    MnO4 - + C2O4 2- + H + ® Mn 2+ + CO2 + H2O
  • What is the electrochemical potential of a copper electrode in 0.1 M CuSO?4-Solution? E.0= 0.35V
  • Which structural isomers are there of pentane (structural formula + name)?
  • Why is a catalyst necessary for the bromination of benzene, but not for that of aniline (aminobenzene)?
  • How are fats structured and why are some solid and others liquid?
  1. Describe the term mutarotation using the example of D-glucose using structural formulas.

1.) Name the 10 most common elements of the earth's crust!
2.) How is the ionization energy defined?

3.) Which ions have the same electron configuration as neon?

4.) Describe (formulate) the technical production of sulfuric acid!

5.) Draw the titration curve of the titration of acetic acid with caustic soda!


Description Functionalized aromatics and their reaction behavior

This video shows you what influence functional groups have on the reaction behavior of aromatic compounds. At the beginning you will be shown important functional groups and you will be explained what properties they have and how they can affect the reaction behavior of molecules. In the further course you will then be explained at which positions on the ring the second substituents preferentially attack and how the positions are named.

Transcript Functionalized aromatics and their reaction behavior

Reaction behavior of functionalized aromatics

Hello! Have you ever wondered how headache pills are made?

In order to be able to manufacture such active ingredients, it is important to understand the chemistry behind them. In headache tablets, the active ingredient is often acetylsalicylic acid, or ASA for short. This is a compound with an aromatic ring. To understand the chemistry of such compounds, I'll first show you what an aromatic ring is. Then you will learn about different functional groups and how they influence the reaction behavior of the ring when they are attached to the ring.

So let's start with benzene: Benzene is an aromatic hydrocarbon that contains 6 pi electrons. These are represented in the formula by the double bonds. In reality, however, the pi electrons are freely distributed over the entire ring, so they are delocalized. Therefore two mesomeric structures can be formulated. The hydrogen atoms on this ring can now be exchanged for functional groups. The ring is thus functionalized.

Frequently occurring functional groups are the hydroxyl group, the carboxy group, the cyano group, the amino group, the nitro group and the methyl group. Such groups now have different effects on the ring. Some pull electrons out of the system, so they have a negative inductive effect. This is abbreviated with the minus I effect. These include the hydroxy, carboxy, cyano, amino and nitro groups. Other groups, such as the methyl group, push electrons into the system and thus have a positive inductive effect. This is abbreviated as the plus-I effect.

The mesomeric effect also acts on the ring. If the substituent takes part in the mesomerism of the ring, it can either increase the electron density in the ring, like the amino, the methyl and the hydroxyl group, then one speaks of the plus-M effect, or it can decrease it, like the carboxy or the cyano - and the nitro group, then we speak of the minus-M effect.

If a functional group is already on the ring, these effects of the group influence the further substitution on the ring. In order to name the position of the second group in relation to the first, these are titled with Greek prefixes. The position exactly opposite the substituent is called para, the position directly next to the substituent ortho and the position in between is called meta.

An example: in the case of ortho-nitrotulol, for example, the methyl group substitutes exactly in the position adjacent to the nitro group. To understand how the effects of the functional group work, let's look at nitrobenzene as a first example, i.e. a compound with a nitro group on the ring: It has an electron-withdrawing effect and thus pulls electrons out of the system. This is made possible by the oxygen atoms on the nitrogen, which have an electron-withdrawing effect and are able to form double bonds. Thus, mesomeric limit formulas can develop in the substituent with the inclusion of the benzene ring. This should be shown using the following structural formulas.

Here you can see that the nitro group pulls electrons out of the aromatic system and thereby positive charges arise in the ortho position and in the para position.

This explains why another substituent that attacks this structure electrophilically would not be directed in the ortho / or para position. Since it is electrophilic, it prefers attack at the point with the highest electron density: this is the meta position. The nitro group directs a second substituent that attacks electrophilically, i.e. into the meta position.

Second, I will explain the plus-M effect to you using aniline as an example. Aniline is a benzene ring attached to an amino group. The amino group has a lone pair of electrons on nitrogen and can push electrons into the system through the formation of mesomeric boundary structures. Since hydrogen atoms are bonded to nitrogen here, there is no electron-withdrawing substituent.

The electrons can only be pushed into the ring. Due to the high electron density that now prevails in the ring, negative charges develop within the ring. These are in the ortho and para positions. If a second substituent attacks electrophilically, the positions with the highest electron density are preferred. In this case, these are the ortho positions and the para position.

So you see how different substituents and functional groups can influence the reaction behavior of molecules and how they can be described. The meta-directing groups include: the nitro group, the cyano group, and the carboxy group. The methyl, amino and hydroxyl groups are ortho and para-directing.

So now you know that functional groups can influence the distribution of electrons on the aromatic ring. If a group pushes the electrons into the ring, a second substituent attacks in the ortho or para position. When the group pulls the electrons out of the ring, the second substituent in the meta position attacks. Such effects must also be taken into account when synthesizing headache tablets. A precursor of ASA is salicylic acid, i.e. a compound in which the carboxy and hydroxyl group must be in the ortho position. So it's always important to understand the chemistry behind it.


Subjects

1 Speech and presentation methods put to the test

In school books, but also in university textbooks, there are various formulas for the sulfate ion and for sulfuric acid. In the classic case, the Lewis formula is drawn with two S = O double bonds and with two S-O single bonds (formula I). On the other hand, more modern university textbooks contain more and more formulas in which the sulfate ion is drawn with four SO single bonds and therefore four single negatively charged oxygen atoms and one double positively charged sulfur atom (Formula II) (Mortimer, Müller, Beck, 2015). In this notation, the sulfur atom obeys the octet rule (Fig. 1).

Both of these formulas have a historical origin. It was none other than Gilbert Newton Lewis (1875-1946) who, when he introduced his electron pair notation, established that in the molecules and ions of the main group elements there are always 8 electrons around an atom (Jensen, 1984)

The Lewis formula representation seems to reach its limits if more than four atoms or four atoms plus lone pairs are arranged in the coordination sphere, especially the heavier elements of the third and higher periods. In order to be able to reasonably explain such “hypervalent” cases, as it were, Linus Pauling suggested sp 3 d hybridization for trigonal bipyramidal molecules, as in the case of PF5 and the sp 3 d 2 hybridization for octahedral molecular ions as in PF6 - before. To Pauling, the charge accumulation appeared to be a less favorable condition than the expansion of the octet by the inclusion of d orbitals (Pauling, 1960).

Potential for confusion in the above modes of speech and representation

If we look at the “traditional” hypervalent notation (Formula I), students repeatedly find it difficult to decide whether and when an octet expansion is “allowed” and when not.In the case of nitrogen-oxygen compounds in particular, it is very tempting to illegally exceed the octet, for example to draw two double bonds in the case of the nitrate ion.

Another difficulty with this representation arises when one takes a closer look at the structure of the sulfate ion or when the electron pair repulsion model is used to predict the spatial structure. In fact, the sulfate molecule has a symmetrical tetrahedral structure, with S-O distances of the same length and regular angles (Mortimer, Müller & Beck, 2015). Here, the mesomerism concept must also be used, according to which the formula only represents one of six boundary structures, but the actual state lies in between. This is particularly problematic because in chemistry lessons the Lewis notation and the electron pair repulsion model are usually introduced in secondary school, mostly in 9/10, but delocalization and mesomeric boundary structures are only discussed in secondary school. As a teacher, you have to either “reach around” such examples or accept “incorrect” representations or inconsistencies.

On the other hand, this problem cannot be completely avoided anyway, for example in the formulas of the SO2- or the SO3-Molecule, if the octet rule is strictly valid, resonance structures can also occur.

The validity of the octet rule usually brings more formal charges into the formula representation. This creates difficulties for many students. However, these are unavoidable, because with the Lewis notation to determine the formal charges by counting the electrons around an atom and forming the difference to the number of valence electrons according to the position in the periodic table, is part of the basic tools of the formula notation. Since more formal charges are usually added to the formulas when the octet rule is strictly valid, there are additional exercise options that can help the learner to deal with this issue competently.

2 suggestions to minimize confusion and understanding

Since the 1980s, quantum mechanical calculations have shown that the bonding relationships in “hypervalent” molecules or ions can be described as well as without the involvement of d orbitals (Kutzelnigg, 1984). At the beginning of this century, electron density determinations by means of X-ray structure analyzes were able to prove the validity of the octet rule using many examples. Perhaps the clearest example of this was the high-resolution determination of the electron density of potassium sulfate K.2SO4 (Schmøkel et al., 2012) summarized in (Irmer, Stalke, 2017). These investigations show very convincingly for the sulfate ion that all S-O bonds can be described as pure σ bonds with polar components (this also causes the shortening compared to a pure single bond). On each of the oxygen atoms (formally simply negatively charged) three electron density maxima can be observed, which can be interpreted as lone pairs of electrons. The sulfur atom has an electron density that is in good agreement with a twofold positive charge.

The formula (II) in Fig. 1, taking into account the octet rule, reproduces the bonding relationships in the sulfate ion much better than the formula (I). The same applies to hypervalent molecules such as sulfur dioxide, sulfuric acid and phosphoric acid or other ions such as phosphate.

At this point the question arises less how one can make it “easier” for pupils, but rather the question of how close one is to the “actual” state of the molecule or ion with one's spelling, so it is primarily about that technical correctness. If the result is an easier approach for the students, so much the better.

The formula representation, taking into account the octet rule, reflects the bond relationships in "hypervalent" molecules or ions of the main group elements such as sulfuric acid, sulfate, phosphoric acid or phosphate better than a notation with an extension of the octet. For schoolchildren, this technically more correct representation also makes it easier to write formulas. The distinction between compliance with the octet rule in the second period and the extension in the third and higher period is no longer necessary.

4 impulses to think ahead

A few pitfalls should be pointed out.

Molecules like PF5 or SF6 must then consequently be written divided into ions (Fig. 3). On the other hand, however, these compounds hardly play a role in chemistry classes in schools.

A problem lurks in a completely different place. One also draws the sulfonic acid groups

(-SO3H) on aromatic systems consequently according to the octet rule with four single bonds, it is no longer possible to formulate an -M effect which, for example, can explain the meta-directing effect of the sulfonic acid group in the electrophilic second substitution (e.g. nitration of sulfonic acids). However, you don't necessarily need an M-Effect to describe the directing effect. If the σ-complexes are formulated, for example for the nitration of the sulfonic acid at the different ring positions, one recognizes that when the nitrosyl cation attacks in the ortho or para position on the ring carbon atom that carries the sulfonyl group, a boundary structure is created with a positive charge (Fig. 4). This is energetically unfavorable because the sulfur atom also carries a positive charge. Only with substitution in the meta position does not exist a boundary structure with such a positive charge on the carbon atom immediately adjacent to the sulfo group. Some organic textbooks therefore no longer work with M-effects at all (SYKES, 1996).

Irmer, E. Stalke, D. (2017). Teaching chemical bonds - a plea for the octet rule. MNU Journal 70(4), 227–234.

Jensen, W. B. (1984). Abegg, Lewis, Langmuir and the Octet Rule. Journal of Chemical Education 61(3), 191–200.

Kutzelnigg, W. (1984). The chemical bond with the higher main group elements. applied Chemistry 96(4), 262–286.

Mortimer, C. E., Müller, U., Beck, J. (2015). Chemistry. The basic knowledge of chemistry, 12th edition. Thieme, Stuttgart.

Pauling, L. (1960). The nature of the chemical bond and the structure of molecules and crystals. An introduction to modern structural chemistry, 3rd ed. Cornell Univ. Press, Ithaca, New York.

Schmøkel, M. S., Cenedese, S., Overgaard, J., Jørgensen, M. R. V., Chen, Y.-S., Gatti, C., Stalke, D., Iversen, B. B. (2012). Testing the concept of hypervalency: charge density analysis of K2SO4. Inorganic chemistry 51(15), 8607–8616.

Sykes, P., Hopf, H. (Eds.) (1996). How do organic reactions work? Reaction mechanisms for beginners. VCH, Weinheim.

Erhard Irmer is a teacher at the Otto Hahn Gymnasium in Göttingen and is currently responsible, among other things, for the involvement of student teachers in the XLAB student laboratory at the University of Göttingen. In addition to the implementation of teacher training courses, he was co-editor of the magazine "Praxis der Naturwissenschaften - Chemie in der Schule" and works as a textbook author and editor for the magazine "CHEMKON".

Georg-August-Universität Göttingen, XLAB - Göttingen experimental laboratory for young people, Justus-von-Liebig-Weg 8, D-37077 Göttingen, Tel. 0551 39 25270

Comments

Prof. Dr. Dr. H. c. Günter Baars, Bern

Let's finally take our high school students seriously.

A reply to the essay by Erhard Irmer «Octet expansion or consistent application of the octet rule», MNU 06, 2020, put up for discussion

In my lectures, when it came to the octet rule, I often said that my students in Switzerland “can count to eight. You no longer have to practice this with electrons. " When will we finally stop teaching that atoms combine to make an octet on the outermost shell? An idea that unfortunately makes a strong and lasting impression on young people's minds. During the oral Abitur exams in the canton of Bern, as an expert in this regard, I recently heard the following answer from a high school graduate to the question why non-metal atoms combine with one another: H. eight electrons on the valence shell. ”After further inquiries, he was unable to name any other criteria.

Why don't we ask our students to argue with electrically charged particles and their interactions? In chemical reactions there is a shift of electrons, thus a change of Coulomb forces and consequently to new energy states, quantum-chemically speaking, one speaks of a change in electron density. As a result, the entropy changes in the system and in the environment. But one after the other, using the example of reactions between non-metal atoms.

The reaction between non-metal atoms

Single and double occupied electron clouds on the valence shell:

The valence electrons of the atoms can be described in terms of models with single and double-occupied electron clouds. A simple model that is based on the fundamentals of quantum chemistry and can be made plausible in the basic subject even in beginner lessons. No statement is made regarding the structure of the clouds. [H. Based on the Kimball model, R. Christen introduced the spherical cloud model (KWM), which provided for a structure of the valence shell. Christians and the author of this text then abolished the KWM in his / her / our textbooks in order to correct the misconception that the molecular structure can be derived from the structure of the valence shell of an atom. This can be easily obtained with the electron pair repulsion model (EPA) applied to molecules.]

Comparison of metal and non-metal atoms:

The experimental values ​​of the ionization energies show that the metal atoms have large atomic cores and a small core charge and thus bind electrons weakly. Non-metal atoms, on the other hand, have small atomic cores with high charges and thus strongly attract electrons. This leads to electron pair bonds (molecules) in the connection of non-metal atoms, in the reaction metal + non-metal to ions (ionic compounds, salts) and the metal atoms form lattices from the atomic cores with the easily movable binding electrons in between.

Derivation of the sum and Lewis formulas of molecules:

I.
Molecules are created by simply occupied electron clouds from two non-metal atoms superimposed to form common, binding electron pairs. More than two electrons per cloud are not possible (Pauli principle). This creates stronger attractive forces and consequently a lower energy state - the binding enthalpies all have a negative sign. This corresponds to the calculations of quantum chemistry, which result in a higher electron density between two atomic cores and thus stronger attractive forces. With this simple rule, a large part of all molecules used in high school chemistry lessons can be derived from the elements and the corresponding Lewis formulas can be drawn. So there is no need to discuss whether this rule is e.g. B. also applies to the third period of the periodic table.

The virial theorem shows that a lower-energy state is formed when an electron pair bond occurs: the decrease in potential energy when an electron pair bond is formed is twice as large as the increase in kinetic energy.

II
Of course, the students reach their limits with rule I when it comes to the Lewis formulas of more complex molecules and ions such as SO2, SO3, O3, H2SO4, SO42-, H3PO4, HClO2, HClO3, HClO4, etc. This helps here an extension of rule I: an atom can also provide two electrons for an electron pair bond. This is useful because z. B. Oxygen atoms have a higher electronegativity than sulfur atoms and can therefore attract a non-bonding pair of electrons from an S atom. To do this, the distribution of the valence electrons of an oxygen atom is modified by forming a doubly occupied cloud from two singly occupied clouds (this is a model concept, and studies show that no d electrons are involved in these bonds). This means that all of the above-mentioned molecules and ions can be easily understood. The fact that these are limit formulas does not initially play a role in the rule extension (see III). With very few exceptions, most of the school-relevant molecules are available.

No further explanations such as exceeding octets, hypervalent notation, etc. are necessary. Formal charges are also not allowed, as they are polar bonds.

III
Another difficulty arises with benzene. The bonds between the carbon atoms are shorter than single and longer than double bonds. In addition, benzene prefers substitution and not addition reactions (teacher's experiment with toluene). With this, the phenomenon of systems of delocalized electrons can be derived as an example and the molecules mentioned under II can be recognized as limit formulas because of the experimentally determined bond lengths. Their Lewis formulas are therefore also available in the classroom. Delocalized electrons can be expected when bond lengths are shorter than single bonds and longer than double bonds. This means that the students have practically all molecules “under control”.

Molecules are particles whose atoms are held together by Coulomb forces, not an octet. Their composition, and thus also the Lewis formulas, can be derived without much effort using the model of the single and double occupied electron clouds. According to II and as an extension of the model, a doubly occupied cloud can be made from two singly occupied electron clouds. An atom then contributes two electrons to the electron pair bond.

One more word about the so-called "formal charges". It is nonsense (!) To introduce formal charges. Why do you have to treat the students with charges that do not exist and that have to be differentiated from the ionic charges in salts? Bonds that arise when a non-metal atom provides two electrons (a non-binding pair of electrons) are simply polar and can therefore be identified by partial charges. After all, we are all talking about “common” electron pairs in an electron pair bond.

«These [formal charges note d. Author] cannot be avoided, however, because determining the formal charges in the Lewis notation by counting the electrons around an atom and calculating the difference to the number of valence electrons according to the position in the periodic table is one of the basic tools of the formula notation. Since more formal charges are usually included in the formula representation if the octet rule is strictly valid, there are additional exercise options that can help the learners to deal with this issue competently. " This is what my colleague Erhard Irmer writes. No (!), This is not part of the basic hand tool the students use to write Lewis formulas. What a waste of time when you have to practice something that has absolutely nothing to do with the essence of chemistry. In addition, according to rules I and II, it is clear to every student where the electrons come from that appear in a Lewis formula.

To conclude, I would like to express it once again: The universe consists of protons, electrons and neutrons (actually from d- and u-quarks and electrons), which make up all substances. This means that the Coulomb forces are at the center of understanding chemical processes, linked to the entropy due to the random movement of particles. My experience of fifty years of teaching at the Neufeld-Gymnasium Bern and thirty years in parallel at the University and PH Bern (chemistry didactics, research on basic chemistry and quantum chemistry at the SII) have confirmed this approach.

Since the reply to the essay by Erhard Irmer is short and z. For example, if it does not contain any information on the formation of salts or metal grids, some information on more extensive texts is given below.

What a lesson looks like in the sense outlined above can be read:
(1) Baars, G., Deuber, R. 2020: Chemistry for high school. hep-Verlag Bern. 2nd Edition

How rule I can be introduced as an example and experimentally can be found at:
(2) Zitt, J., Baars, G. 2020: Experiments on the book. Methane. www.hep-verlag.ch/chemie-zusatzmaterial-lehrdienstleistungen (free access)

If you are interested in the virial theorem and the corresponding calculation for the formation of an H atom from a proton and an electron, you will find this in the teaching units:
(3) Baars, G. 2010: Quantum Chemistry and Chemistry of Colored Substances Module 2, Quantum Chemistry and Chemical Bonding, Section 4.2
https://www.phbern.ch/dienstleistungen/unterrichtsmedien/ideenset-quantenchemie-und-chemie-farbiger-stoffe (free access)

Attachment theory with a difference. Five mnemonics and the doctrine of attachment "one has a grip":
(4) Baars, G. 2019: The doctrine of attachment - with a difference. Bond theory without octet rule, formal charges and hybrid orbitals. Chemistry lessons 1/19. Friedrich. Pp. 43 - 45

There is also a corresponding text for the procedure for introducing systems of delocalized electrons:
(5) Baars, G. 2015: Systems of delocalized electrons. An experimental approach. Practice of science. Chemistry in school. 3/64. Aulis publishing house. S.21-26

Answer to the reply from Prof. Günter Baars

I would like to thank my esteemed colleague Günter Baars for his reply, which gives me the opportunity to emphasize some aspects again.

First of all, I absolutely have to agree with Günter Baars: The octet rule does not explain why non-metal atoms combine with one another! Atoms do not combine "to get an octet on the outermost shell". The octet state of an isolated atom or ion is also not fundamentally the most energetically favorable state. This can easily be made clear by looking at the first ionization energy of the sodium atom: For the formation of the sodium cation from the sodium atom, an amount of energy of 496 kJ / mol has to be applied. The isolated sodium ion is therefore no more stable than the sodium atom. The technical error of the noble gas configuration as the most energetically favorable state can also be found in the literature and in fact very easily settles deeply as a misconception in the minds of students, but also of students and teachers, as well as empirical ones Studies confirm [1]. This misconception must be countered unconditionally and continuously.

Therefore, the introduction of the octet rule must not replace an occupation with the energetic fundamentals and the electrostatic forces in the formation of ions and molecules. I am very grateful to Günter Baars that he has shown time and again from his teaching experience and in his diverse publications that it is possible to work out sophisticated model concepts with younger students, which later enable them to be expanded to a quantum mechanical wave model.

The octet rule does not have an explanatory but a mnemonic and heuristic function for the rapid writing of Lewis formulas. The octet or noble gas rule provides a formal description of the bonding state of atoms and ions in many main group compounds. A corresponding formulation of the octet rule could be: “In many main group compounds, each atom or ion can be assigned as many electrons as a noble gas atom, so mostly 8 electrons.” In this formulation, the verb “to add” makes it clear that it is a formal and not an explanatory criterion. The use of the octet rule as an aid when writing electron pair formulas is therefore not to be understood as a contrast to the explanatory models proposed by Günter Baars, but as a supplement to the formulation of Lewis formulas.

In my contribution I wanted to show that the statements of the octet rule are also supported by modern scientific research. The octet rule does not represent a technically questionable simplification, but can enable a description of binding states that can be quickly grasped and implemented in an uncomplicated manner. That is why the octet rule is also widely used in universities, especially in organic chemistry. In my opinion, there is little to be said against using this simple tool for writing Lewis formulas in schools.

With the rules he presented, Günter Baars also comes up with formulas that are described by the octet rule. In this way, too, the molecule of sulfuric acid is represented with four single bonds on the sulfur atom. To what extent the extension of his rule that an atom can also provide two electrons for the electron pair bond, later in the coordinative bond cannot cause confusion, I do not like to judge. The same procedure could just as easily be used to arrive at formulas that exceed valence, for example for sulfuric acid (two S = O double bonds).

I would like to vehemently object to one of Günter Baars' arguments: "It is nonsense (!) To introduce formal charges" and "Even formal charges are not justified because they are polar bonds". It is correct: formal charges are not "real" charges as we know them from ion charges. The term formal charge already bears its characteristic in the name: It is a formal assignment of charges under the assumption that the electrons in the electron pair bond are assigned to the participating atoms equally (!).

The notation of Lewis formulas is also initially a purely formal representation: a line (or originally two points) is assigned to a pair of electrons. By comparing the number of electrons in the Lewis formula with the number of electrons in the isolated atom, a charge can be assigned to the atom in the molecule. To easily identify the charge of charged molecules, it has been proven in my teaching experience to use counting electrons. If the formal charges of all atoms in the molecule are determined (mostly this applies to individual atoms - if at all), then the ion charge of the entire molecular ion can easily be determined by adding up the formal charges. I find it confusing when an oxygen atom with a single bond and three free electron pairs without a (formal) charge is drawn in the sulfuric acid formula, while in the acetate ion an oxygen atom with the same form is supposed to receive an ionic charge.

Here, too, the following applies: It is a challenge for teachers and students to distinguish between the different types of cargo. However, I am convinced that this topic is very much something that has to do with the "essence of chemistry" and that dealing with these different formulas should very well be part of the "basic tools of the students". According to the educational standards [2], not only keeping the substance and particle level apart, but also the critical handling of the representative representation level is a central task in the area of ​​process-related competencies, knowledge acquisition, communication and evaluation (E7, E9, K7, B1). Competent handling of the model character, including the formulaic representations of the particle level, has to be discussed, applied and practiced again and again. I am convinced that we should not spare our students this, but can use it in a targeted manner to bring the learners closer to an important aspect that is particularly relevant when it comes to chemical content. The different charge types in chemical formulas (to which one can also add the formal oxidation numbers) are in my opinion particularly well suited.

literature
[1] Yayon, M., Mamlok-Naaman, R., Fortus, D. (2012). Characterizing and representing student's conceptual knowledge of chemical bonding. Chem. Educ. Res. Pract. 13/3, 248-267.
[2] Secretariat of the Conference of Ministers of Education (2020). Educational standards in chemistry for the general university entrance qualification. Resolution of the Conference of Ministers of Education on June 18, 2020. Secretariat of the Conference of Ministers of Education, Berlin.


Video: Q No. #60. Based on concept of Nitration of Toluene. IIT JEE. NEET (December 2021).