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## Lewis Acid and Base Molecules

Lewis bases are electron-pair donors, whereas Lewis acids are electron-pair acceptors.

Learning Objectives

Recognize Lewis acids and bases in chemical reactions.

Key Takeaways
Key Points
A Lewis acid is an electron -pair acceptor; a Lewis base is an electron-pair donor.
Some molecules can act as either Lewis acids or Lewis bases; the difference is context-specific and varies based on the reaction.
Lewis acids and bases result in the formation of an adduct rather than a simple displacement reaction, as with classical acids and bases. An example is HCl vs H+: HCl is a classical acid, but not a Lewis acid; H+ is a Lewis acid when it forms an adduct with a Lewis base.
Key Terms
covalent bond: a chemical bond in which two atoms are connected to each other by sharing two or more electrons
nucleophile: literally “lover of nuclei,” Lewis bases are often referred to as this because they seek to donate their electron pairs to electron-poor species, such as H+

A Lewis acid is defined as an electron-pair acceptor, whereas a Lewis base is an electron-pair donor. Under this definition, we need not define an acid as a compound that is capable of donating a proton, because under the Lewis definition, H+ itself is the Lewis acid; this is because, with no electrons, H+ can accept an electron pair.

A Lewis base, therefore, is any species that donates a pair of electrons to a Lewis acid. The “neutralization” reaction is one in which a covalent bond forms between an electron-rich species (the Lewis base) and an electron-poor species (the Lewis acid). For this reason, Lewis bases are often referred to as nucleophiles (literally, “lovers of nuclei”), and Lewis acids are sometimes called electrophiles (“lovers of electrons”). This definition is useful because it not only covers all the acid-base chemistry with which we are already familiar, but it describes reactions that cannot be modeled by Arrhenius or Bronsted-Lowry acid-base chemistry. For now however, we will consider how the Lewis definition applies to classic acid-base neutralization.

Applying the Lewis Definition to Classical Acid-Base Chemistry

Consider the familiar reaction of NaOH and HCl:

$\text{NaOH}(\text{aq})+\text{HCl}(\text{aq})\rightarrow \text{NaCl}(\text{aq})+\text{H}_2\text{O}(\text{l})$

We have previously described this as an acid-base neutralization reaction in which water and a salt are formed. This is still completely correct, but the Lewis definition describes the chemistry from a slightly different perspective. When considering Lewis acids and bases, the only real reaction of interest is the net ionic reaction:

$\text{OH}^-(\text{aq})+\text{H}^+(\text{aq})\rightarrow \text{H}_2\text{O}(\text{l})$

Under the Lewis definition, hydroxide acts as the Lewis base, donating its electron pair to H+. Thus, in this version of the neutralization reaction, what interests us is not the salt that forms, but the covalent bond that forms between OH– and H+ to form water. A significant hallmark for Lewis acid-base reactions is the formation of such a covalent bond between the two reacting species. The reaction’s final product is known as an adduct, because it forms from the addition of the Lewis base to the Lewis acid.

Lewis acids and bases: Lewis acids (BF3, top, and H+, bottom) react with Lewis bases (F–, top, NH3, bottom) to form products known as adducts. Note that the first reaction cannot be described by Arrhenius or Bronsted-Lowry acid-base chemistry.

Beyond Classical Acid-Base Chemistry

By treating acid-base reactions in terms of electron pairs instead of specific substances, the Lewis definition can apply to reactions that do not fall under other definitions of acid-base reactions. For example, a silver cation behaves as a Lewis acid with respect to ammonia, which behaves as a Lewis base, in the following reaction:

$\text{Ag}^+(\text{aq}) + 2\;\text{NH}_3 \rightarrow [\text{Ag}(\text{NH}_3)_2]^+$

This reaction results in the formation of diamminesilver(I), a complex ion; it is perfectly described by Lewis acid-base chemistry, but is unclassifiable according to more traditional Arrhenius and Bronsted-Lowry definitions.

Application to Organic Chemistry

In organic chemistry, it is useful to understand that nucleophiles are Lewis bases and electrophiles are Lewis acids. Nearly all reactions in organic chemistry can be considered Lewis acid-base processes.

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What are acids and bases?: This lesson continues to describe acids and bases according to their definition. We first look at the Bronsted-Lowry theory, and then describe Lewis acids and bases according to the Lewis Theory.

## Metal Cations that Act as Lewis Acids

Transition metals can act as Lewis acids by accepting electron pairs from donor Lewis bases to form complex ions.

Learning Objectives

Recognize metals that function as Lewis acids.

Key Takeaways
Key Points
A Lewis acid is an electron pair acceptor; because metal ions have one or more empty orbitals, they act as Lewis acids when coordinating ligands.
Examples of metals that can act as Lewis acids include Na+, Mg2+, and Ce3+.
Metal ions rarely exist uncoordinated; they often have to dissociate from weaker ligands, like water, before complexing with other Lewis bases.
Key Terms
coordinate bond: a type of covalent bond in which two shared electrons originate from the same atom; a dative bond
ligand: the species that coordinates with a metal cation to form a complex ion
Complex ion: a compound consisting of a metal ion coordinated to various ligands in solution

The modern-day definition of a Lewis acid, as given by IUPAC, is a molecular entity—and corresponding chemical species—that is an electron-pair acceptor and therefore able to react with a Lewis base to form a Lewis adduct; this is accomplished by sharing the electron pair furnished by the Lewis base. Classically, the term “Lewis acid” was restricted to trigonal planar species with an empty p orbital, such as BR3 where R can be an organic substituent or a halide. However, metal ions such as Na+, Mg2+, and Ce3+ often form Lewis adducts upon reacting with a Lewis base.

Complex Ion Formation

Ligands create a complex when forming coordinate bonds with transition metals ions; the transition metal ion acts as a Lewis acid, and the ligand acts as a Lewis base. The number of coordinate bonds is known as the complex’s coordination number. Common ligands include H2O and NH3 ; examples of complexes include the tetrachlorocobaltate(II) ion, [CoCl4]2- and the hexaqua-iron(III) ion, [Fe(H2O)6]3+.

Usually, metal complexes can only serve as Lewis acids after dissociating from a more weakly bound Lewis base, often water. For instance, Mg2+ can coordinate with ammonia in solutions, as shown below:

$[\text{Mg}(\text{H}_2\text{O})_6]^{2+} + 6\text{NH}_3 \rightarrow [\text{Mg}(\text{NH}_3)_6]^{2+} + 6\text{H}_2\text{O}$

Nearly all compounds formed by the transition metals can be viewed as collections of the Lewis bases—or ligands—bound to the metal, which functions as the Lewis acid. The product is known as a complex ion, and the study of these ions is known as coordination chemistry. One coordination chemistry’s applications is using Lewis bases to modify the activity and selectivity of metal catalysts in order to create useful metal-ligand complexes in biochemistry and medicine.

Examples of metal-ligand coordination complexes: Examples of several metals (V, Mn, Re, Fe, Ir) in coordination complexes with various ligands. All these metals act as Lewis acids, accepting electron pairs from their ligands.

Lewis acids and bases — Wikipedia Republished // WIKI 2 – Some of the most studied examples of such Lewis acids are the boron trihalides and organoboranes, but other compounds exhibit this behavior Lewis acids and bases are commonly classified according to their hardness or softness.The theory of acids and bases, like many other chemical theories, has Such acids and bases are called 'secondary' by Lewis as distinct from his 'primary' acids and bases The Lewis Theory covers more completely substances that show the qualitative attributes normally associated with acids.Acids and Bases – CliffsNotes Study Guides Acids and Bases. (n) A regional or social variety of a language distinguished by pronunciation, grammar, or vocabulary, especially a variety of speech differing from the standard literary language or speech pattern of the culture in which it exists…

PDF A solution then is not necessarily neutral at – Lewis acid and Lewis base are most closely associated with ELECTRON TRANSFER.A Lewis acid is any substance which can accept a pair of non bonding electron. Thus, in Lewis theory of acid and base reaction, bases donate pairs of electrons and acids accept the pairs of electrons.What are Lewis Acids and Bases? A Lewis acid is an electron pair acceptor and a Lewis base is an electron pair donor. They can react with each another such that a The most common Lewis bases are anions. The strength of Lewis basicity correlates with the pKa of the parent acid: acids with high…An acid is a molecule or ion capable of donating a proton (hydrogen ion H+) (a Brønsted-Lowry acid), or, alternatively, capable of forming a covalent bond with an electron pair (a Lewis acid). The first category of acids are the proton donors, or Brønsted-Lowry acids.

PDF Acids And Bases Study Guide – Lewis acids and bases are defined in terms of being able to accept or donate electron pairs. This means that acids can accept a lone pair of electrons from a Lewis base because the acid has vacant valence orbitals.Lewis acid is a chemical species that reacts with a Lewis base to form a Lewis adduct. A Lewis base, then, is any species that donates a pair of electrons to a Lewis acid to form a Lewis adduct. For example, OH− and NH3 are Lewis bases, because they can donate a lone pair of electrons.A Lewis acid is defined as as a chemical specie that can accept an electron pair from a Lewis base to form a Lewis adduct while a Lewis base is defined as a specie which can donate an electron pair to a Lewis acid or an acceptor compound. Thus, Lewis base and acid has to do with electrons transfer.

Acidity of Aromatic Compounds Vid 8 Orgo Acids and Bases – Leah here from leah4sci.com and in this video
we'll take a more advance look at Acids and Bases in Organic Chemistry by looking at the
acidity of aromatic compounds.
You can find my entire series on orgo Acids
and Bases along with my practice quiz and cheat sheet by visiting my website leah4sci.com/AcidBase. Say you're asked to compare two different
sets of molecules to determine which is stronger in each pair. You'll recognize that they have
very similar structures but the right molecule on each pair has a benzene or aromatic ring.
The first thing that comes to mind is Resonance. And that's not wrong but you want to make
sure that the resonance is a possibility. For the first set we have Cyclohexanol and
Phenol. If we deprotonate the acidic hydrogen, we get a conjugate base with a negative Oxygen
where the one on the left is not aromatic,it can't resonate making the Oxygen the sole
carrier of the negative charge. But the right one has the negative Oxygen on the benzene
ring which means it can resonate into the ring distributing that charge and making it
more stable. So for the first set, the molecule on the
right is the stronger acid and the molecule on the left is the weaker acid. If you're
not comfortable with this, go back to the resonance video on the series. For the second
set, we have an acidic Hydrogen once again sitting on an Oxygen but the functional group
here is the Carboxylic acid. So even though there is a resonance within the carboxylic
acid itself you can't resonate the conjugate base into the benzene ring. If you don't believe
me, try it! But make sure you recognize you cannot resonate into the benzene ring. So
what's the difference between them? For a problem like this you have to look at the
inductive effect of the Benzene ring rather than the resonance. Benzene, having those extra electrons resonating
in the ring, makes it more of an electronegative group and therefore a stronger influencer
or the negative charge and that means the molecule on the right is the stronger acid.
The molecule on the left just has a standard carbon hydrogen chain, it has some induction
but not as much as the benzene ring making the molecule on the left the weaker acid.
So we've established that Phenol is dis and acid because the negative Oxygen can resonate
into the ring. But now what happens if I put Nitrogen containing group onto that Benzene
ring. For the first one, we'll put in NH2 in the
para position, for the second one we'll put an NO2 in the para position, and for the third
one we'll put in NO2 in the meta position. If you'll look at the CARIO mnemonic and just
compare the negative conjugate base, we have charge negative one, atom, oxygen holds the
charge, resonance – they can all resonate into the benzene ring, Inductive effect – we
have something to look out with the Nitrogen atoms and the orbitals are the same. This
is where a simple mnemonic is not enough, you actually have to practice the resonance
to see what happens. The NO2 is a little bit of a suspicious group
because Nitrogen is an electronegative atom so you want to think that it has those electron
donating group and therefore puts negativity into the molecule but not in NO2. NO2 are
Nitro group has a Nitrogen double bound to an Oxygen and single bound to a second Oxygen.
A quick formal charge will show you that the single bound Oxygen has the charge of negative
one and the Nitrogen has a charge of positive one. That's your first clue. Having a positive
group on the benzene ring, near an Oxygen that wants to resonate into the ring is a
good thing because a negative charge resonating towards a positive charge is much more stable
than say a negative charge resonating towards an electronegative Nitrogen. So think of a
lone pair on the NH2 as depleting or trying to push away that negative Oxygen resonance
making this one the worst or the weakest acid. If resonance makes it acidic and you're discourage
resonance then you're discouraging the acidity and that's the weakest acid. So we have the
same NO2 on the meta position which means we have the same concept of positive Nitrogen,
negative Oxygen.They're both decent or stronger acids but which one is stronger? Iis it a
question of proximity or something else? And in this case it's not proximity, it's resonance
again. So watch what happens. If I take the Oxygen and I resonate it, notice we have these
two electrons which move over this way and then we'll show these two electrons, but where
do they go? Standard resonance is to continue around the ring so that we have a total of
4 resonance structures and we can show the same thing here. These electrons go down,
these electrons move over, these electrons move to the next position and once again for
resonance. But if we look back to the meta structure, those black electrons did not have
to continue in the ring, there is a bonus resonance when you have the Nitro group in
the para position and those electrons can actually go down to form a pi bond between
carbon and Nitrogen. Nitrogen would now have 5 bonds violating
the octet rule so we'll show the electrons between the nitrogen double bound to oxygen
collapsing onto the oxygen and that right there gives you a bonus resonance structure.
Remember, resonance gives the molecule the ability to distribute the charge over more
atoms. So if you have four resonance that means four atoms are carrying the charge and
here you have five resonance that means the molecule with more resonance has the charge
distributed even more and that makes it even more stable. If the conjugate base is more
stable then it comes from the most willing or simply the strongest acid. And when you
have slightly less resonance, it's slightly less stable and that means it's a weaker acid.
It's still a decent acid where compare to phenol but it's slightly stronger. So if we
throw phenol into the mix and then rank it from weakest to strongest, the weakest acid
would be the one with the NH2 because the NH2 is slowing down the negative resonating
into the ring. Then we have the phenol itself because we don't have any group to make it
stronger or weaker than phenol itself. Then for number three we have the one with
the Nitro group in the meta position because the nitro is positive and that means it stabilizes
the positive coming into the ring but it doesn't give us a bonus resonance structure. And finally
a Nitro group in the para position with a bonus resonance structure is the strongest
acid in this sequence. Be sure to join me in the next video where
we look at the basicity of aromatic compounds. Are you struggling with Organic Chemistry?
Are you looking for resources and information to guide you through the course and help you
succeed? If so, then I have a deal for you, a FREE copy of my ebook “10 Secrets to Acing
Organic Chemistry”. Use the link below or visit orgosecrets.com to grab your free copy.
updates with Cheat Sheets, reaction guides, study tips and so much more. You’ll also
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many videos over the course of the semester so if you haven’t subscribed to my channel
yet, do so right now to be sure that you don’t miss out. .

Organic Chem Basicity: How to compare Basicity of Organic Compounds – Hi everyone this is Maverick Puah the Chemistry Guru.
Now in this video we want to go through the basicity of organic compounds. Now for organic chem usually the concept for us to explain basicity is usually using the Lewis acid-base concept, where our lewis base is a lone pair or an electron pair donor. So therefore basicity in organic chemistry is actually fairly straightforward. We just focus on the lone pair availability. Now in organic chem all the bases that we have are nitrogen compounds so it becomes a very straightforward issue of just focusing on the lone pair on my N. If the lone pair on my N is more available then it makes it a better base, a better lewis base, now if the lone pair is less available then it makes it a weaker lewis base so let's take a look at the comparison between the basicity of different N compounds. Now the trend that we have is basically here so what you notice is our tertiary amine is the most basic functional group as compared to a secondary amine, a primary amine and then followed by ammonia now you notice the number of R groups actually affect the basicity so this one is explained using our electronic effect which is related to the number of donating or withdrawing groups attached to my functional group so therefore it affects the reactivity of that functional group. Now what we notice in this case our phenylamine is less basic than ammonia so later we will talk about this using the resonance effect and our amide is less basic than phenylamine. In fact our amide is neutral so it actually doesn't function as a base and the N doesn't donate the lone pair so later we'll also be able to explain this using the resonance effect. Now the electronic effect is fairly straightforward. Because R group is a donating group, so what R group will do is it will push electron to N so it'll increase the lone pair availability on my N so the more R groups we have and therefore the more available the lone pair will be on N and it makes it more basic. So the more electron donating groups we have, so N becomes more electron rich, so the lone pair on my N becomes more available for donation, it makes it a stronger base. So if I compare a tertiary amine you have more electron donating groups, 3 R groups that are pushing electrons to N so the lone pair is very available for donation whereas if you look at ammonia, ammonia because it doesn't have any donating group so the basicity in terms of ammonia as compared to my amine will be weaker next if I look at phenylamine what we notice in this case is my lone pair on my N is delocalised into my benzene now benzene has this delocalised pi system, this orange ring here, it's the delocalised pi system involving benzene, so what N would do is it'll use the lone pair to interact with this delocalised pi system and because it has to spend some time to interact with this delocalised pi system the lone pair becomes less available for donation. So we can say that the lone pair on my N, becaused it's delocalised into the pi electron system of benzene, so it's less available for donation so it makes it a weaker base as compared to ammonia now for amide functional group it also has this delocalisation but it's delocalised into the acyl group or the acid group so what this means is effectively the lone pair on my N actually can move away from N and towards and into my O so therefore the lone pair is actually delocalised very well between my N, C and O. One way we can do is we can show using this structure here where this curve will represent the delocalised lone pair and this lone pair is delocalised very well between my N, C and O because O is very electronegative and N is also pretty electronegative so the lone pair is actually spread out very well between the 2 electronegative atoms. So because the lone pair on my N is delocalised extensively between the electronegative N and O we don't use it for donation, amides actually don't use the lone pair on my N at all for donation so therefore amides are neutral I hope you find this video lesson useful. If you like it please give me the thumbs up, like this video and share it with your friends. Now for more video lessons please subscribe to my YouTube Channel. If you have any questions or concepts that you want me to clarify, please put them at the comment section below so that I can address them in future. Thank you for watching I'll see you next time. .

Hydrogen fluoride – Hydrogen fluoride is a chemical compound with
the chemical formula HF.
This colorless gas or liquid is the principal industrial source
of fluorine, often in the aqueous form as hydrofluoric acid, and thus is the precursor
to many important compounds including pharmaceuticals and polymers. HF is widely used in the petrochemical
industry and is a component of many superacids. Hydrogen fluoride boils near room temperature,
while other hydrogen halides evaporate at much lower temperatures. Unlike other hydrogen
halides, HF is lighter than air and diffuses relatively quickly through porous substances.
Hydrogen fluoride is a highly dangerous gas, forming corrosive and penetrating hydrofluoric
acid upon contact with tissue. The gas can also cause blindness by rapid destruction
of the corneas. French chemist Edmond Frémy is credited with
discovering anhydrous hydrogen fluoride while trying to isolate fluorine, although Carl
Wilhelm Scheele prepared hydrofluoric acid in large quantities in 1771, and this acid
was known in the glass industry before then. Structure Near or above room temperature, HF is a colorless
gas. Below its melting point), HF forms orthorhombic crystals, consisting of zig-zag chains of
HF molecules. The HF molecules, with a short H–F bond of 95 pm, are linked to neighboring
molecules by intermolecular H–F distances of 155 pm. Liquid HF also consists of chains
of HF molecules, but the chains are shorter, consisting on average of only five or six
molecules. Hydrogen bonding
HF molecules interact through hydrogen bonds, thus creating extra clustering associations
with other HF molecules. Because of this, hydrogen fluoride behaves more like water
than like other hydrogen halides, such as HCl. This hydrogen bonding between HF molecules
gives rise to high viscosity in the liquid phase and lower than expected pressure in
the gas phase. Hydrogen fluoride does not boil until 20 °C in contrast to the heavier
hydrogen halides which boil between −85 °C and −35 °C.
Hydrogen fluoride is fully miscible with water, while the other hydrogen halides have large
solubility gaps with water. Hydrogen fluoride and water also form several compounds in the
solid state, most notably a 1:1 compound that does not melt until −40 °C, which is 44
°C above the melting point of pure HF. Acidity
Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid
in dilute aqueous solution. This is in part a result of the strength of the hydrogen-fluorine
bond, but other factors such as the tendency of HF, H
2O, and F− anions to form clusters. At high concentrations, HF molecules undergo homoassociation
to form polyatomic ions (such as bifluoride, HF−
2) and protons, thus greatly increasing the acidity. This leads to protonation of very
strong acids like hydrochloric, sulfuric, or nitric when using concentrated hydrofluoric
acid solutions. Although hydrofluoric acid is regarded as a weak acid, it is very corrosive,
even attacking glass when hydrated. The acidity of hydrofluoric acid solutions
vary with concentration owing to hydrogen-bond interactions of the fluoride ion. Dilute solutions
are weakly acidic with an acid ionization constant Ka = 6.6×10−4, in contrast to
corresponding solutions of the other hydrogen halides which are strong acids. Concentrated
solutions of hydrogen fluoride are much more strongly acid than implied by this value,
as shown by measurements of the Hammett acidity function H0. The H0 for 100% HF is estimated
to be between −10.2 and −11, comparable to the value −12 for sulfuric acid.
In thermodynamic terms, HF solutions are highly non-ideal, with the activity of HF increasing
much more rapidly than its concentration. The weak acidity in dilute solution is sometimes
attributed to the high H—F bond strength, which combines with the high dissolution enthalpy
of HF to outweigh the more negative enthalpy of hydration of the fluoride ion. However,
Giguère and Turrell have shown by infrared spectroscopy that the predominant solute species
is the hydrogen-bonded ion-pair [H3O+•F−], which suggests that the ionization can be
described as a pair of successive equilibria: H2O + HF [H3O+•F−]
[H3O+•F−] H3O+ + F− The first equilibrium lies well to the right
and the second to the left, meaning that HF is extensively dissociated, but that the tight
ion pairs reduce the thermodynamic activity coefficient of H3O+, so that the solution
is effectively less acidic. In concentrated solution, the additional HF
causes the ion pair to dissociate with formation of the hydrogen-bonded hydrogen difluoride
ion. [H3O+•F−] + HF H3O+ + HF2−
The increase in free H3O+ due to this reaction accounts for the rapid increase in acidity,
while fluoride ions are stabilized by strong hydrogen bonding to HF to form HF2−. This
interaction between the acid and its own conjugate base is an example of homoassociation. At
the limit of 100% liquid HF, there is self-ionization 3 HF H2F+ + HF2−
which forms an extremely acidic solution. The acidity of anhydrous HF can be increased
even further by the addition of Lewis acids such as SbF5, which can reduce H0 to −21.
Solvent Dry hydrogen fluoride readily dissolves low-valent
metal fluorides as well as several molecular fluorides. Many proteins and carbohydrates
can be dissolved in dry HF and recovered from it. In contrast, most non-fluoride inorganic
chemicals react with HF rather than dissolving. Production and uses
Hydrogen fluoride is produced as by the action of sulfuric acid on pure grades of the mineral
fluorite and also as a side-product of the extraction of the fertilizer precursor phosphoric
The anhydrous compound hydrogen fluoride is more commonly used than its aqueous solution,
hydrofluoric acid. HF serves as a catalyst in alkylation processes in oil refineries.
A component of high-octane petrol called "alkylate" is generated in alkylation units that combine
C3 and C4 olefins and iso-butane to generate petrol.
HF is a reactive solvent in the electrochemical fluorination of organic compounds. In this
approach, HF is oxidized in the presence of a hydrocarbon and the fluorine replaces C–H
bonds with C–F bonds. Perfluorinated carboxylic acids and sulfonic acids are produced in this
way. Hydrogen fluoride is an important catalyst
used in the majority of the installed linear alkyl benzene production in the world. The
process involves dehydrogenation of n-paraffins to olefins, and subsequent reaction with benzene
using HF as catalyst. Elemental fluorine, F2, is prepared by electrolysis
of a solution of HF and potassium bifluoride. The potassium bifluoride is needed because
anhydrous hydrogen fluoride does not conduct electricity. Several million kilograms of
F2 are produced annually. Acyl chlorides or acid anhydrides react with
hydrogen fluoride to give acyl fluorides. HF is often used in palynology to remove silicate
minerals, for extraction of dinoflagellate cysts, acritarchs and chitinozoans.
1,1-Difluoroethane is produced by the mercury-catalyzed addition of hydrogen fluoride to acetylene:
HC≡CH + 2 HF → CH3CHF2 The intermediate in this process is vinyl
fluoride, the monomeric percursor to polyvinyl fluoride.
Health effects Upon contact with moisture, including tissue,
hydrogen fluoride immediately converts to hydrofluoric acid, which is highly corrosive
and toxic, and requires immediate medical attention upon exposure. Breathing in hydrogen
fluoride at high levels or in combination with skin contact can cause death from an
irregular heartbeat or from fluid buildup in the lungs.