Amine reaction with hcl

Amine reaction with hcl DEFAULT

Amines as Bases

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We are going to have to use two different definitions of the term "base" in this page. A base is

  • a substance which combines with hydrogen ions. This is the Bronsted-Lowry theory.
  • an electron pair donor. This is the Lewis theory.

The easiest way of looking at the basic properties of amines is to think of an amine as a modified ammonia molecule. In an amine, one or more of the hydrogen atoms in ammonia has been replaced by a hydrocarbon group. Replacing the hydrogens still leaves the lone pair on the nitrogen unchanged - and it is the lone pair on the nitrogen that gives ammonia its basic properties. Amines will therefore behave much the same as ammonia in all cases where the lone pair is involved.

The reactions of amines with acids

These are most easily considered using the Bronsted-Lowry theory of acids and bases - the base is a hydrogen ion acceptor. We'll do a straight comparison between amines and the familiar ammonia reactions. Ammonia reacts with acids to produce ammonium ions. The ammonia molecule picks up a hydrogen ion from the acid and attaches it to the lone pair on the nitrogen.

If the reaction is in solution in water (using a dilute acid), the ammonia takes a hydrogen ion (a proton) from a hydroxonium ion. (Remember that hydrogen ions present in solutions of acids in water are carried on water molecules as hydroxonium ions, H3O+.)

\[ NH_3 (aq) + H_3O^+ (aq) \rightarrow NH_4^+ + H_2O (l)\]

If the acid was hydrochloric acid, for example, you would end up with a solution containing ammonium chloride - the chloride ions, of course, coming from the hydrochloric acid. You could also write this last equation as:

\[ NH_3 (aq) + H^+ \rightarrow NH_4^+ (aq)\]

. . . but if you do it this way, you must include the state symbols. If you write H+ on its own, it implies an unattached hydrogen ion - a proton. Such things don't exist on their own in solution in water. If the reaction is happening in the gas state, the ammonia accepts a proton directly from the hydrogen chloride:

\[ NH_3 (aq) + HCl (g) \rightarrow NH_4^+ (s) + Cl^- (s)\]

This time you produce clouds of white solid ammonium chloride.

The nitrogen lone pair behaves exactly the same. The fact that one (or more) of the hydrogens in the ammonia has been replaced by a hydrocarbon group makes no difference.

Example 1: Ethylamine

If the reaction is done in solution, the amine takes a hydrogen ion from a hydroxonium ion and forms an ethylammonium ion.

Or:

The solution would contain ethylammonium chloride or sulfate or whatever. Alternatively, the amine will react with hydrogen chloride in the gas state to produce the same sort of white smoke as ammonia did - but this time of ethylammonium chloride.

\[ CH_3CH_2NH_2 (g) + HCl(g) \rightarrow CH_3CH_2NH_3^+ (s) + Cl^- (s)\]

These examples have involved a primary amine, but it makes no real difference if a secondary or tertiary amine were used.; the equations would just look more complicated. The product ions from diethylamine and triethylamine would be diethylammonium ions and triethylammonium ions respectively.

The Reactions of Amines with Water

Again, it is easiest to use the Bronsted-Lowry theory and, again, it is useful to do a straight comparison with ammonia. Ammonia is a weak base and takes a hydrogen ion from a water molecule to produce ammonium ions and hydroxide ions. However, the ammonia is only a weak base, and doesn't hang on to the hydrogen ion very successfully. The reaction is reversible, with the great majority of the ammonia at any one time present as free ammonia rather than ammonium ions.

\[ NH_3 (aq) + H_2O (l) \rightleftharpoons NH_4^+ (aq) + OH^- (aq)\]

The presence of the hydroxide ions from this reaction makes the solution alkaline. The amine still contains the nitrogen lone pair, and does exactly the same thing. For example, with ethylamine, you get ethylammonium ions and hydroxide ions produced.

\[ CH_3CH_2NH_2 (aq) + H_2O (l) \rightleftharpoons CH_3CH_2NH_3^+ (aq) + OH^- (aq)\]

There is, however, a difference in the position of equilibrium. Amines are usually stronger bases than ammonia (there are exceptions to this, though - particularly if the amine group is attached directly to a benzene ring).

The reactions of amines with copper(II) ions

Just like ammonia, amines react with copper(II) ions in two separate stages. In the first step, we can go on using the Bronsted-Lowry theory (that a base is a hydrogen ion acceptor). The second stage of the reaction can only be explained in terms of the Lewis theory (that a base is an electron pair donor).

Example 2: Reaction between Ammonia and Copper (II)

Copper(II) sulphate solution, for example, contains the blue hexaaquacopper(II) complex ion -\( [Cu(H_2O)_6]^{2+}\).

In the first stage of the reaction, the ammonia acts as a Bronsted-Lowry base. With a small amount of ammonia solution, hydrogen ions are pulled off two water molecules in the hexaaqua ion.

This produces a neutral complex - one carrying no charge. If you remove two positively charged hydrogen ions from a 2+ ion, then obviously there isn't going to be any charge left on the ion. Because of the lack of charge, the neutral complex isn't soluble in water, and so you get a pale blue precipitate.

\[ [Cu(H_2O)_6]^{2+} + 2NH_3 \rightleftharpoons [Cu(H_2O)_4(OH)_2] + 2NH_4^+\]

This precipitate is often written as Cu(OH)2 and called copper(II) hydroxide. The reaction is reversible because ammonia is only a weak base. That precipitate dissolves if you add an excess of ammonia solution, giving a deep blue solution.

The ammonia replaces four of the water molecules around the copper to give tetraamminediaquacopper(II) ions. The ammonia uses its lone pair to form a co-ordinate covalent bond (dative covalent bond) with the copper. It is acting as an electron pair donor - a Lewis base.

\[ [Cu(H_2O)_6]^{2+} + 4NH_3 \rightleftharpoons [Cu(NH_3)_4(H_2O)_2]^{2+} + 4H_2O\]

The color changes are:

The corresponding reaction with amines

The small primary amines behave in exactly the same way as ammonia. There will, however, be slight differences in the shades of blue that you get during the reactions.

Example 3: Reaction between Methylamine and Copper (II)

With a small amount of methylamine solution you will get a pale blue precipitate of the same neutral complex as with ammonia. All that is happening is that the methylamine is pulling hydrogen ions off the attached water molecules.

With more methylamine solution the precipitate redissolves to give a deep blue solution - just as in the ammonia case. The amine replaces four of the water molecules around the copper.

As the amines get bigger and more bulky, the formula of the final product may change - simply because it is impossible to fit four large amine molecules and two water molecules around the copper atom.

Sours: https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Supplemental_Modules_(Organic_Chemistry)/Amines/Reactivity_of_Amines/Amines_as_Bases

Amide formation

Acyl chlorides and acid anhydrides react with primary and secondary amines without the presence of heat to form amides. Tertiary amines cannot be acylated due to the absence of a replaceable hydrogen atom. With the much less active benzoyl chloride, acylation can still be performed by the use of excess aqueous base to facilitate the reaction.

Salt formation

Because amines are basic, they neutralize carboxylic acids to form the corresponding ammonium carboxylate salts. Upon heating to 200°C, the primary and secondary amine salts dehydrate to form the corresponding amides.

Neutralization

Amines R3N react with strong acids such as hydroiodic acid (HI), hydrobromic acid (HBr) and hydrochloric acid (HCl) to give ammonium salts R3NH+.

Reaction with nitrous acid

Nitrous acid with the chemical formula HNO2 is unstable. Usually it is produced indirectly in a mixture of NaNO2 and a strong acid such as HCl or H2SO4 in dilute concentration, so that the H+ ions will associate with the NO2 ions in solution.

Primary aliphatic amines with nitrous acid give very unstable diazonium salts which spontaneously decompose by losing N2 to form a carbenium ion. The carbenium ion goes on to produce a mixture of alkenes, alkanols or alkyl halides, with alkanols as the major product. This reaction is of little synthetic importance because the diazonium salt formed is too unstable, even under quite cold conditions.

NaNO2 + HCl → HNO2 + NaCl

 

  • Primary aromatic amines, such as aniline (phenylamine) forms a more stable diazonium ion at 0–5°C. Above 5°C, it will decompose to give phenol and N2. Diazonium salts can be isolated in the crystalline form but are usually used in solution and immediately after preparation, due to rapid decomposition on standing even with little ambient heat. Solid diazonium salts can be explosive on shock or on mild warming.

Reactions with ketones and aldehydes

  • Primary amines react with carbonyl compounds to form imines (see section 21.4.). Specifically, aldehydes become aldimines, and ketones become ketimines. In the case of formaldehyde (R’ = H), the imine products are typically cyclic trimers.
RNH2 + R’2C=O → R’2C=NR + H2O
  • Secondary amines react with ketones and aldehydes to form enamines. An enamine contains a C=C double bond, where the second C is singly bonded to N as part of an amine ligand.
R2NH + R'(R”CH2)C=O → R”CH=C(NR2)R’ + H2O
Sours: https://courses.lumenlearning.com/suny-potsdam-organicchemistry2/chapter/23-3-reactions-of-amines/
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Organic Nitrogen Compounds V: Amine Salts

Fortunately, like many other functional groups, amine salts have multiple IR features, and these come to the rescue here. In addition to NH+ stretching vibrations, amine salts also have NH+ bending vibrations as well. The NH3+ grouping of primary amine salts features two peaks from the asymmetric and symmetric bending vibrations, labeled B and C in Figure 2. In general, the asymmetric bend falls from 1625 to 1560, and the symmetric bend from 1550 to 1500. Strangely enough, these peaks are small in sharp contrast to the intense NH+ stretching envelope. This is all due to the difference in dµ/dx between the stretching and bending vibrations of the amine salt functional group.

Secondary Amine Salt Spectra

Secondary amine salts contain the NH2+ group. The IR spectrum of a secondary amine salt, diisopropylamine hydrochloride, is seen in Figure 5.


Figure 5: The infrared absorbance spectrum of a secondary amine salt, disopropylamine hydrochloride.

The NH+ stretching envelope is labeled A. Note that it is broad and strong, like the ones seen in Figures 3 and 4. As in Figure 3, the C-H stretches here also fall on top of the NH+ stretching envelope. It also has the expected complement of overtone and combination bands on its lower wavenumber side. For secondary amine salts in general, this envelope is found from 3000 to 2700. Note that there is some overlap between the envelopes of primary and secondary amines. However, secondary amine salts only have one NH+ bending band compared to primary amine salts. This feature typically falls from 1620 to 1560, and is labeled B in Figure 5. Thus the position and number of NH+ bending bands is what determines whether a sample contains a primary or secondary amine salt.

Tertiary Amine Salt Spectra

Tertiary amine salts contain the NH+ group, as seen in Figure 1. The IR spectrum of a tertiary amine salt, 2,2',2''-trichloroethylamine hydrochloride, is seen in Figure 6.


Figure 6: The infrared absorbance spectrum of a tertiary amine salt, 2,2',2''-trichloroethylamine hydrochloride.

The NH+ stretching envelope is Figure 6 is labeled A. Note that it is lower in wavenumber than for primary and secondary amine salts, and that the C-H stretches fall as shoulders to left of the envelope peak. Given that the NH+ stretching envelope for tertiary amines falls squarely in the overtone-combination range from 2800 to 2000, these peaks show up on top of and as shoulders to the right of the NH+ stretching envelope. In general, for tertiary amine salts, this envelope falls from 2700 to 2300. The size, width, and position of this peak is practically unique in IR spectroscopy-in my decades of experience, I have never seen a peak like it (10). Thus, this peak by itself is strongly indicative of their being a tertiary amine salt in a sample. Tertiary amine salts do not have any NH+ bending peaks, so the lack of peaks from 1625 to 1500 can also be used to distinguish tertiary amine salts from primary and secondary salts.

We discussed previously that tertiary amines have no strong, unique peaks, and thus are difficult to detect using IR spectroscopy (12). This contrasts with tertiary amine salts, whose NH+ stretching envelope sticks out like a sore thumb. A way of detecting a tertiary amine in a sample then is to treat 1 mL of liquid tertiary amine, or tertiary amine dissolved in an organic solvent, with 1 mL of 50:50 HCl in ethanol. If there is a tertiary amine present, the amine salt will form and precipitate as a solid from solution (12). Collect the precipitate via filtration, dry, and measure its IR spectrum. If you see a big, whopping NH+ stretching envelope like the one seen in Figure 6, your original sample contained a tertiary amine.

The group wavenumber peaks for amine salts are listed in Table I.

Conclusions

Amine salts are made by reacting amines with strong acids. Primary amine salts contain the NH3+ group, secondary amine salts the NH2+ group, and tertiary amine salts the NH+ group. Amine salts are important, because they are used to make drug substances water soluble, and hence more bioavailable.

All amine salts contain an intense, broad NH+ stretching envelope that is a rather unique infrared feature. The envelope position overlaps for primary and secondary amine salts, but is unique for tertiary salts. Primary and secondary amine salts can be distinguished by the number and position of NH+ bending peaks.

References

(1) B.C. Smith, Spectroscopy34(7), 18–21, 44 (2019).

(2) B.C. Smith, Spectroscopy34(5), 22–26 (2019).

(3) B.C. Smith, Spectroscopy34(3), 22–25 (2019).

(4) B.C. Smith, Spectroscopy34(1), 10–15 (2019).

(5) A. Streitweiser and C. Heathcock, Introduction to Organic Chemistry (MacMillan, New York, New York, 1st ed., 1976).

(6) B.C. Smith, Spectroscopy30(1), 16–23 (2015).

(7) https://en.wikipedia.org/wiki/Cocaine

(8) B.C. Smith, Spectroscopy31(7), 30–34 (2016).

(9) B.C. Smith, Spectroscopy30(4), 18–23 (2015).

(10) B.C. Smith, Infrared Spectral Interpretation: A Systematic Approach (CRC Press, Boca Raton, Florida, 1999).

(11) B.C. Smith, Spectroscopy33(1), 14–20 (2018).

(12) B.C. Smith, Spectroscopy33(3), 16–20 (2018).

(13) B.C. Smith, Spectroscopy31(11), 28–34 (2016).

(14) B.C. Smith, Spectroscopy31(5), 36–39 (2016).

Brian C. Smith, PhD, is founder and CEO of Big Sur Scientific, a maker of portable mid-infrared cannabis analyzers. He has over 30 years experience as an industrial infrared spectroscopist, has published numerous peer reviewed papers, and has written three books on spectroscopy. As a trainer, he has helped thousands of people around the world improve their infrared analyses. In addition to writing for Spectroscopy, Dr. Smith writes a regular column for its sister publication Cannabis Science and Technology and sits on its editorial board. He earned his PhD in physical chemistry from Dartmouth College. He can be reached at: [email protected]

Sours: https://www.spectroscopyonline.com/view/organic-nitrogen-compounds-v-amine-salts
(L-20) Amine reaction with HNO2 -- Diazonium Salt Formation -- with Mechanism by Arvind Arora

The easiest way of looking at the basic properties of amines is to think of an amine as a modified ammonia molecule. In an amine, one or more of the hydrogen atoms in ammonia has been replaced by a hydrocarbon group.

Replacing the hydrogens still leaves the lone pair on the nitrogen unchanged - and it is the lone pair on the nitrogen that gives ammonia its basic properties. Amines will therefore behave much the same as ammonia in all cases where the lone pair is involved.

The reactions of amines with acids

These are most easily considered using the Bronsted-Lowry theory of acids and bases - the base is a hydrogen ion acceptor. We'll do a straight comparison between amines and the familiar ammonia reactions.

A reminder about the ammonia reactions

Ammonia reacts with acids to produce ammonium ions. The ammonia molecule picks up a hydrogen ion from the acid and attaches it to the lone pair on the nitrogen.

If the reaction is in solution in water (using a dilute acid), the ammonia takes a hydrogen ion (a proton) from a hydroxonium ion. (Remember that hydrogen ions present in solutions of acids in water are carried on water molecules as hydroxonium ions, H3O+.)

If the acid was hydrochloric acid, for example, you would end up with a solution containing ammonium chloride - the chloride ions, of course, coming from the hydrochloric acid.

You could also write this last equation as:

. . . but if you do it this way, you must include the state symbols. If you write H+ on its own, it implies an unattached hydrogen ion - a proton. Such things don't exist on their own in solution in water.

If the reaction is happening in the gas state, the ammonia accepts a proton directly from the hydrogen chloride:

This time you produce clouds of white solid ammonium chloride.

The corresponding reactions with amines

The nitrogen lone pair behaves exactly the same. The fact that one (or more) of the hydrogens in the ammonia has been replaced by a hydrocarbon group makes no difference.

For example, with ethylamine:

If the reaction is done in solution, the amine takes a hydrogen ion from a hydroxonium ion and forms an ethylammonium ion.

Or:

The solution would contain ethylammonium chloride or sulphate or whatever.

Alternatively, the amine will react with hydrogen chloride in the gas state to produce the same sort of white smoke as ammonia did - but this time of ethylammonium chloride.

These examples have involved a primary amine. It would make no real difference if you used a secondary or tertiary one. The equations would just look more complicated.

The product ions from diethylamine and triethylamine would be diethylammonium ions and triethylammonium ions respectively.

The reactions of amines with water

Again, it is easiest to use the Bronsted-Lowry theory and, again, it is useful to do a straight comparison with ammonia.

A reminder about the ammonia reaction with water

Ammonia is a weak base and takes a hydrogen ion from a water molecule to produce ammonium ions and hydroxide ions.

However, the ammonia is only a weak base, and doesn't hang on to the hydrogen ion very successfully. The reaction is reversible, with the great majority of the ammonia at any one time present as free ammonia rather than ammonium ions.

The presence of the hydroxide ions from this reaction makes the solution alkaline.

The corresponding reaction with amines

The amine still contains the nitrogen lone pair, and does exactly the same thing.

For example, with ethylamine, you get ethylammonium ions and hydroxide ions produced.

There is, however, a difference in the position of equilibrium. Amines are usually stronger bases than ammonia. (There are exceptions to this, though - particularly if the amine group is attached directly to a benzene ring.)

Sours: https://www.chemguide.co.uk/organicprops/amines/base.html

Hcl with amine reaction

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Tests for Amines - MeitY OLabs

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