FAQFAQ   SearchSearch   MemberlistMemberlist   UsergroupsUsergroups 
 ProfileProfile   PreferencesPreferences   Log in to check your private messagesLog in to check your private messages   Log inLog in 
Forum index » Science and Technology » Chem
Nitrogen as an oxidizer
Post new topic   Reply to topic Page 1 of 1 [15 Posts] View previous topic :: View next topic
Author Message
The_Man
science forum addict


Joined: 21 May 2006
Posts: 52

PostPosted: Thu Jul 20, 2006 7:19 pm    Post subject: Re: Nitrogen as an oxidizer Reply with quote

lucasea@sbcglobal.net wrote:
Quote:
joepad1991@yahoo.com> wrote in message
news:1153092116.402565.295190@s13g2000cwa.googlegroups.com...
I am a young science student and I have a question about diatomic
nitrogen gas. I always read that N2 is inert. Nitrogen is
electronegative, just as Oxygen is. It is only a single atomic number
removed from Oxygen. With O2, when the temperature is high enough, the
double bound holding the two Oxygen atoms together breaks, leaving 2
atoms of oxygen, each with two highly reactive unpaired electrons
(which I suppose is another way of saying that Oxygen is
electronegative). When the temperature reaches this point (different
for different fuels), we say that combustion sets in.

Well, why doesn't the same thing happen with Nitrogen? Why is Nitrogen
inert? In fact, one might think it more reactive than Oxygen. If the
temperature rises to such a point to break the triple bond, you would
have atomic Nitrogen with THREE unpaired electrons. This might make
Nitrogen more reactive than Oxygen (though granted it is less
electronegative than oxygen).

So why isn't Nitrogen gas an oxidizer?

Excellent question, but it is based on several misconceptions that would be
a natural result of the simplifications taught in early chemistry classes.
First, the bonding picture in O2 is not the simple O=O that you would expect
based on things like the VSEPR model, and quantum mechanics is the culprit.

The VSEPR model is extremely simple, and it is a surprise that it is as
effective as it is. It leads to numerous mistakes, though, especially
with regard to "resonance". M.O. theory, while more complex, always
gives the right answers.

Quote:
Ground-state O2 is a singly-bonded diradical (best represented as .O-O.) It
just happens to be how the molecular orbital energy levels work out. In N2,
the doubly-degenerate pi bonding orbitals are both filled, giving it a
closed-shell fully bonded configuration with three full N-N bonds. O2 has
two more electrons in essentially the same set of orbitals, and these extra
electrons go one each into the two doubly-degenerate pi antibonding
orbitals, cancelling some of the double-bond character from the pi bonding
orbitals. Second, your description of how combustion occurs is not correct.
The unpaired electrons in O2 are capable of abstracting a H atom from a
hydrocarbon, leaving a C radical which is capable of further chemistry that
ultimately ends up at CO2 and H2O.

Everyone has overlooked a very important principle here, that ordinary
dioxygen is a TRIPLET. Reactions between triplet states and singlets
(like most hydrocarbons) are spin forbidden, and spin-forbidden
processes, though not "impossible", have low probablility. While the
ground state of dioxygen is a triplet, there is an accessible excited
state of oxygen which is a singlet. Singlet oxygen is a very powerful
oxidizing agent. It is a good thing that the ground state of oxygen is
a triple; otherwise combustion would proceed far more rapidly than it
does.

Quote:

The reactions actually involved form an extremely long and complex radical
chain reaction, but as I understand it, some of the more important ones are
as follows (R is anything bonded to carbon, including H)

R3C-H + .O-O. --> R3C. + HO-O.
R3C. + .O-O. --> R3C-O-O.
R3CO-O. + R3C-H --> R3CO-OH + R3C.
R3CO-OH --> R3CO. + .OH
.OH + R3C-H --> HOH + R3C.
R3CO. --> R2C=O + R.
R2C=O --> etc. --> CO, CO2, etc.

The O-O bond probably doesn't break until after the alkyl radical and an O2
molecule combine to form an alkyl peroxyl radical R-O-O. or an alkyl
hydroperoxide, RO-OH, at which point you have a very weak and unstable O-O
single bond. N2 doesn't have any of this radical character that O2 uses to
start the whole thing off, and it has an extremely strong, stable triple

In fact, Lavoisier, who discovered nitrogen, called it "azote" - "not
alive" to refelct its relative inertness.

Quote:
bond. If you look at compounds that have N=N double bonds (azines) and
especially N-N single bonds (hydrazines), the N=N bonds and especially N-N
bonds are much weaker than the triple bond in N2, and are generally more
reactive as oxidants. Also, keep in mind that N is close to O in the
periodic table, but it's also right next to C, which is usually not
considered an oxidant. Its reactivity is somewhere in the middle. Finally,
also keep in mind that, depending on reaction conditions and co-reagents,
hydrazine itself (N2H4) can behave as either an oxidant (by breaking the N-N
bond) or it can be oxidized by a stronger oxidant, by losing H atoms to form
N2. This is the basis of its use as a rocket fuel.

H2N-NH2 + O2 --> N2 + 2 H2O

When we say that Oxygen is an oxidixer, do we really mean that diatomic
oxygen (which I would think is relatively unreactive) is an oxidixer,
or do we mean that atomic oxygen (with two unpaired electrons) is an
oxidizer?

Generally, yes, diatomic oxygen is the reactive species.

Eric Lucas
Back to top
Craig
science forum beginner


Joined: 23 Jun 2005
Posts: 23

PostPosted: Tue Jul 18, 2006 12:53 am    Post subject: Re: Nitrogen as an oxidizer Reply with quote

lucasea@sbcglobal.net wrote:
Quote:
You are asking good questions, but I'd like to introduce a different way of
thinking about the different explanations you see in chemistry. Learning
any complex discipline like chemistry has to be like peeling an onion.
...
But that's the really cool thing
about chemistry--you get to decide how deeply you want to understand
something, and you never need to stop peeling the onion!

Eric Lucas

Eric, thank you for that metaphor! That is a better way of describing
some of what I skirted around so much more clumsily. Smile

- Craig
Back to top
Craig
science forum beginner


Joined: 23 Jun 2005
Posts: 23

PostPosted: Tue Jul 18, 2006 12:44 am    Post subject: Re: Nitrogen as an oxidizer Reply with quote

joepad1991@yahoo.com wrote:
Quote:
Hi everyone,

This got me thinking again. I know Fluorine is a strong oxidizer.
According to VSEPR, Fluorine exists as a diatomic gas with a single
bond, each atom with 3 lone pairs (which are presumably unreactive,
except possible as a lweis base).

Based on all of the explanation above, the VSEPR model is probably
wrong.

What is the actual electron structure like in Fluorine that makes it an
oxidizer?

Is the study of actual electron structure covered in courses such as
Physical Chemistry?

Thanks again everyone,
Joe


VSEPR may not be the best possible description of non-bonding lone
pairs. Nevertheless, it does its job well enough. Di-oxygen has a
bond order of 2, di-nitrogen has a bond order of 3, and di-fluorine has
a bond order of 1, just as the elementary Lewis model predicts. Don't
toss out the baby with the bath water! Smile VSEPR does a pretty good job
of describing where the *atoms* are, and it shouldn't be expected to do
more.

The simplest explanation of why elemental fluorine is such a strong
oxidizer is that fluorine has a very high electronegativity. In other
words, it "wants" electrons "really badly." Electronegativity is a
property purely of the fluorine atom (e.g. 7 electrons in a shell that
can hold Cool, not what the atom is bound to. Fluorine is just as
electronegative in F2 as it is in HF, NaF or any other compound. In
simple terms, fluorine (and oxygen for that matter) is highly
electronegative, and that helps make F2 such a strong oxidizer. It
"wants" electrons so badly that it can oxidize some very stubborn
things in order to get them (or release large amounts of energy as it
does so). Of course in compounds like NaF, fluorine already has what
it "wants" (it has already taken the electron from sodium) so it is
relatively unreactive, unlike in F2 (fluorine has to share electrons
with another fluorine; it can't outright swipe one, because the other
fluorine atom pulls it back too strongly). Briefly, a strong oxidizer
"wants" electrons really badly but doesn't yet get to keep them.

Joe, it seems that you may have a misconception that this relates
somehow to the electronic structure of F2 and how it gets oxidized.
This is not so. The strength of an oxidizer does not depend on *how*
the material gets oxidized, in a microscopic sense, but only on the
difference between the beginning and ending chemical state. In other
words, don't read too much into the fact that F2 has one bond and N2
has 3 bonds or precisely where the non-bonding electrons are. Having
to break these bonds doesn't affect the element's strength as an
oxidizer. What matters is that when each element reacts with, say,
carbon, fluorine will do a much better job of winning the "tug-of-war"
for electrons. Nitrogen "wins" a little bit over carbon, but fluorine
"wins" a lot! The triple bond in nitrogen slows its rate of reaction
down tremendously, but that isn't why it is considered a weaker
oxidizer than oxygen. Going back to your original question (why is
oxygen such a strong oxidizer?), the fact that di-oxygen has unpaired
electrons is just a red herring. Look at electronegativity and how
strongly reduced the starting material (e.g F2, O2 = not at all; F in
NaF = very much) is, not at the detailed structure of the non-bonding
electrons (e.g. lone pairs, or some unpaired in O2).

Yes, this whole topic is addressed in much greater depth in Physical
Chemistry, and perhaps in the second semester of General (college)
Chemistry to some extent.

In the words of Einstein, I have tried to "make everything as simple as
possible, but not simpler." I have taken an element's
electronegativity as a given, rather than trying to explain why that
number might be high - that's a whole other discussion! I hope this
helps and doesn't muddy the waters.

- Craig
Back to top
Fred Kasner
science forum beginner


Joined: 27 Nov 2005
Posts: 25

PostPosted: Tue Jul 18, 2006 12:36 am    Post subject: Re: Nitrogen as an oxidizer Reply with quote

lucasea@sbcglobal.net wrote:
Quote:
"donald haarmann" <donald-haarmann@worldnet.att.net> wrote in message
news:kEAug.402968$Fs1.160585@bgtnsc05-news.ops.worldnet.att.net...
joepad1991@yahoo.com

|I am a young science student and I have a question about diatomic
| nitrogen gas. I always read that N2 is inert.

[snip]


----------
Inert?! Forsooth. Mg will burn in N2.

It certainly takes quite a bit of heat to get Mg to burn in N2, but Li will
do so spontaneously. It's quite a feeling to have flames inside a
supposedly inert glovebox!

Mg will also burn in CO2 (for example, when lit in air and placed between
two slabs of dry ice). Quite impressive!

Eric Lucas


And note that if the common first year lab experiment of burning Mg

ribbon in air the product is not pure MgO. The ash of the experiment is
heated slightly and a drop or two of water dropped onto it. There is a
distinct odor of ammonia [ NH3 ] produced. Some Mg3N2 was produced as
well and decomposes when heated with added water. In other words N2 was
capable of oxidizing the hot Mg.
FK
Back to top
<lucasea@sbcglobal.net
science forum addict


Joined: 25 Mar 2006
Posts: 94

PostPosted: Tue Jul 18, 2006 12:32 am    Post subject: Re: Nitrogen as an oxidizer Reply with quote

<joepad1991@yahoo.com> wrote in message
news:1153156438.956635.186600@i42g2000cwa.googlegroups.com...
Quote:
Hi everyone,

This got me thinking again. I know Fluorine is a strong oxidizer.
According to VSEPR, Fluorine exists as a diatomic gas with a single
bond, each atom with 3 lone pairs (which are presumably unreactive,
except possible as a lweis base).

Based on all of the explanation above, the VSEPR model is probably
wrong.

What is the actual electron structure like in Fluorine that makes it an
oxidizer?

Is the study of actual electron structure covered in courses such as
Physical Chemistry?

The VSEPR model isn't wrong--like any other model, it just has its
limitations (more on this in a minute). In the case of F2, it turns out
that the VSEPR model works fine. It doesn't have the diradical character
that O2 has that makes it so reactive, but it does have a couple of other
things going for it. 1) The F-F bond is extremely weak, and in some cases,
formation of F. is involved in the reactions of F2. 2) Another reaction
mechanism that F2 has is to act as an electrophile, with a nucleophile
attacking one F, and expelling the other F as F-, in a sort of an Sn2 type
of reaction. O2 doesn't really have this option, because O= or O-. are not
especially good leaving groups. If you've had organic chemistry classes,
you might recall that leaving group strength is exactly opposite to
basicity--the weaker a base that a leaving group is, the better a leaving
group it is. So, since O= is a pretty strong base, it's not a good leaving
group. F- is a rather weak base (conjugate acid of HF, a moderately strong
acid), so it's a fairly good leaving group. Also, keep in mind that most of
the aspects of O2 that you mentioned, and most of the ones that people have
responded with, dealt with the reaction mechanisms that O2 uses to oxidize
things. When most people talk about oxidizing "strength", they're talking
about the thermodynamics--that is, the energy O2 + (the things being
oxidized) is much higher than the energy of H2O + (the oxidized products).
Keeping mechanism (reaction kinetics) and thermodynamics (energy
differences) straight is a big struggle in understanding why things are the
way they are in chemistry.

You will begin to get exposure to other models of molecular electronic
structure in p-chem. If you want to really understand it, you might want to
take some classes in quantum mechanics, advanced organic chemistry
(especially physical organic chemistry), advanced inorganic chemistry, and
maybe computational chemistry if they're available.

You are asking good questions, but I'd like to introduce a different way of
thinking about the different explanations you see in chemistry. Learning
any complex discipline like chemistry has to be like peeling an onion. You
learn simple "models" to understand and explain simple aspects of reality
first--like VSEPR, which does a fantastic job of explaining about 99% of
organic chemistry, and even a reasonable fraction of inorganic chemistry.
As you absorb and understand those simple models, you begin to see cases
where they are incomplete and that better/more accurate models are
needed--in this case, there are big gaps in inorganic chemistry that aren't
adequately described by VSEPR, and one of the first that most people run
across is O2. Ultimately, no model is "the truth", or "actual", they are
all just successively more and more accurate representations of reality.
All models are imperfect, it just depends on how well you want to understand
something how sophisticated a model you need to use. Just think if they
started out first quarter freshman year teaching the state-of-the-art models
of advanced theoretical chemistry--that would be a complete disaster.
Nobody would have any basis for understanding why things have to be so
complicated, and everyone would get frustrated and leave. However, if you
see the simple models, and see why they don't work in all situations, then
the more complicated models make sense. But that's the really cool thing
about chemistry--you get to decide how deeply you want to understand
something, and you never need to stop peeling the onion!

Eric Lucas
Back to top
joepad1991@yahoo.com
science forum beginner


Joined: 16 Jul 2006
Posts: 3

PostPosted: Mon Jul 17, 2006 5:14 pm    Post subject: Re: Nitrogen as an oxidizer Reply with quote

Hi everyone,

This got me thinking again. I know Fluorine is a strong oxidizer.
According to VSEPR, Fluorine exists as a diatomic gas with a single
bond, each atom with 3 lone pairs (which are presumably unreactive,
except possible as a lweis base).

Based on all of the explanation above, the VSEPR model is probably
wrong.

What is the actual electron structure like in Fluorine that makes it an
oxidizer?

Is the study of actual electron structure covered in courses such as
Physical Chemistry?

Thanks again everyone,
Joe


joepad1991@yahoo.com wrote:
Quote:
Thanks everyone. This really sets me straight.

Joe


Bob wrote:
On 16 Jul 2006 16:21:56 -0700, joepad1991@yahoo.com wrote:

Eric Lucas has posted an excellent reply. Just a couple of additional
comments here.


I am a young science student and I have a question about diatomic
nitrogen gas. I always read that N2 is inert.

That is a generalization, not an absolute. At room temperature, N2 is
usually almost as inert as a noble gas. I suppose that of the common
gases (other than noble gases) or the common diatomic element, it is
the most inert. But it can react. And remember that the noble gases,
at least the heavier ones, can also react.

The biological process of nitrogen fixation involves reduction of N2
to ammonia, at organism temperature -- soil temperature in this case,
to a large extent.


Nitrogen is
electronegative, just as Oxygen is.

Remember, electronegativity (EN) needs a number, not yes/no. Saying it
is EN is vague. The EN of N is actually about half way between C and
O.

bob
Back to top
Lloyd Parker
science forum Guru


Joined: 08 May 2005
Posts: 657

PostPosted: Mon Jul 17, 2006 1:39 pm    Post subject: Re: Nitrogen as an oxidizer Reply with quote

In article <3VAug.122059$H71.95866@newssvr13.news.prodigy.com>,
<lucasea@sbcglobal.net> wrote:
Quote:

joepad1991@yahoo.com> wrote in message
news:1153092116.402565.295190@s13g2000cwa.googlegroups.com...
I am a young science student and I have a question about diatomic
nitrogen gas. I always read that N2 is inert. Nitrogen is
electronegative, just as Oxygen is. It is only a single atomic number
removed from Oxygen. With O2, when the temperature is high enough, the
double bound holding the two Oxygen atoms together breaks, leaving 2
atoms of oxygen, each with two highly reactive unpaired electrons
(which I suppose is another way of saying that Oxygen is
electronegative). When the temperature reaches this point (different
for different fuels), we say that combustion sets in.

Well, why doesn't the same thing happen with Nitrogen? Why is Nitrogen
inert? In fact, one might think it more reactive than Oxygen. If the
temperature rises to such a point to break the triple bond, you would
have atomic Nitrogen with THREE unpaired electrons. This might make
Nitrogen more reactive than Oxygen (though granted it is less
electronegative than oxygen).

So why isn't Nitrogen gas an oxidizer?

Excellent question, but it is based on several misconceptions that would be
a natural result of the simplifications taught in early chemistry classes.
First, the bonding picture in O2 is not the simple O=O that you would expect
based on things like the VSEPR model, and quantum mechanics is the culprit.
Ground-state O2 is a singly-bonded diradical (best represented as .O-O.)

Huh? O2 has a bond order of 2. It's just that the 2 highest-energy electrons
are not paired that makes it paramagnetic.
Back to top
joepad1991@yahoo.com
science forum beginner


Joined: 16 Jul 2006
Posts: 3

PostPosted: Mon Jul 17, 2006 7:09 am    Post subject: Re: Nitrogen as an oxidizer Reply with quote

Thanks everyone. This really sets me straight.

Joe


Bob wrote:
Quote:
On 16 Jul 2006 16:21:56 -0700, joepad1991@yahoo.com wrote:

Eric Lucas has posted an excellent reply. Just a couple of additional
comments here.


I am a young science student and I have a question about diatomic
nitrogen gas. I always read that N2 is inert.

That is a generalization, not an absolute. At room temperature, N2 is
usually almost as inert as a noble gas. I suppose that of the common
gases (other than noble gases) or the common diatomic element, it is
the most inert. But it can react. And remember that the noble gases,
at least the heavier ones, can also react.

The biological process of nitrogen fixation involves reduction of N2
to ammonia, at organism temperature -- soil temperature in this case,
to a large extent.


Nitrogen is
electronegative, just as Oxygen is.

Remember, electronegativity (EN) needs a number, not yes/no. Saying it
is EN is vague. The EN of N is actually about half way between C and
O.

bob
Back to top
Bob111
science forum Guru Wannabe


Joined: 13 Jan 2006
Posts: 115

PostPosted: Mon Jul 17, 2006 4:17 am    Post subject: Re: Nitrogen as an oxidizer Reply with quote

On 16 Jul 2006 16:21:56 -0700, joepad1991@yahoo.com wrote:

Eric Lucas has posted an excellent reply. Just a couple of additional
comments here.


Quote:
I am a young science student and I have a question about diatomic
nitrogen gas. I always read that N2 is inert.

That is a generalization, not an absolute. At room temperature, N2 is
usually almost as inert as a noble gas. I suppose that of the common
gases (other than noble gases) or the common diatomic element, it is
the most inert. But it can react. And remember that the noble gases,
at least the heavier ones, can also react.

The biological process of nitrogen fixation involves reduction of N2
to ammonia, at organism temperature -- soil temperature in this case,
to a large extent.


Quote:
Nitrogen is
electronegative, just as Oxygen is.

Remember, electronegativity (EN) needs a number, not yes/no. Saying it
is EN is vague. The EN of N is actually about half way between C and
O.

bob
Back to top
<lucasea@sbcglobal.net
science forum addict


Joined: 25 Mar 2006
Posts: 94

PostPosted: Mon Jul 17, 2006 12:49 am    Post subject: Re: Nitrogen as an oxidizer Reply with quote

"donald haarmann" <donald-haarmann@worldnet.att.net> wrote in message
news:kEAug.402968$Fs1.160585@bgtnsc05-news.ops.worldnet.att.net...
Quote:
joepad1991@yahoo.com

|I am a young science student and I have a question about diatomic
| nitrogen gas. I always read that N2 is inert.

[snip]


----------
Inert?! Forsooth. Mg will burn in N2.

It certainly takes quite a bit of heat to get Mg to burn in N2, but Li will
do so spontaneously. It's quite a feeling to have flames inside a
supposedly inert glovebox!

Mg will also burn in CO2 (for example, when lit in air and placed between
two slabs of dry ice). Quite impressive!

Eric Lucas
Back to top
<lucasea@sbcglobal.net
science forum addict


Joined: 25 Mar 2006
Posts: 94

PostPosted: Mon Jul 17, 2006 12:33 am    Post subject: Re: Nitrogen as an oxidizer Reply with quote

<joepad1991@yahoo.com> wrote in message
news:1153092116.402565.295190@s13g2000cwa.googlegroups.com...
Quote:
I am a young science student and I have a question about diatomic
nitrogen gas. I always read that N2 is inert. Nitrogen is
electronegative, just as Oxygen is. It is only a single atomic number
removed from Oxygen. With O2, when the temperature is high enough, the
double bound holding the two Oxygen atoms together breaks, leaving 2
atoms of oxygen, each with two highly reactive unpaired electrons
(which I suppose is another way of saying that Oxygen is
electronegative). When the temperature reaches this point (different
for different fuels), we say that combustion sets in.

Well, why doesn't the same thing happen with Nitrogen? Why is Nitrogen
inert? In fact, one might think it more reactive than Oxygen. If the
temperature rises to such a point to break the triple bond, you would
have atomic Nitrogen with THREE unpaired electrons. This might make
Nitrogen more reactive than Oxygen (though granted it is less
electronegative than oxygen).

So why isn't Nitrogen gas an oxidizer?

Excellent question, but it is based on several misconceptions that would be
a natural result of the simplifications taught in early chemistry classes.
First, the bonding picture in O2 is not the simple O=O that you would expect
based on things like the VSEPR model, and quantum mechanics is the culprit.
Ground-state O2 is a singly-bonded diradical (best represented as .O-O.) It
just happens to be how the molecular orbital energy levels work out. In N2,
the doubly-degenerate pi bonding orbitals are both filled, giving it a
closed-shell fully bonded configuration with three full N-N bonds. O2 has
two more electrons in essentially the same set of orbitals, and these extra
electrons go one each into the two doubly-degenerate pi antibonding
orbitals, cancelling some of the double-bond character from the pi bonding
orbitals. Second, your description of how combustion occurs is not correct.
The unpaired electrons in O2 are capable of abstracting a H atom from a
hydrocarbon, leaving a C radical which is capable of further chemistry that
ultimately ends up at CO2 and H2O.

The reactions actually involved form an extremely long and complex radical
chain reaction, but as I understand it, some of the more important ones are
as follows (R is anything bonded to carbon, including H)

R3C-H + .O-O. --> R3C. + HO-O.
R3C. + .O-O. --> R3C-O-O.
R3CO-O. + R3C-H --> R3CO-OH + R3C.
R3CO-OH --> R3CO. + .OH
..OH + R3C-H --> HOH + R3C.
R3CO. --> R2C=O + R.
R2C=O --> etc. --> CO, CO2, etc.

The O-O bond probably doesn't break until after the alkyl radical and an O2
molecule combine to form an alkyl peroxyl radical R-O-O. or an alkyl
hydroperoxide, RO-OH, at which point you have a very weak and unstable O-O
single bond. N2 doesn't have any of this radical character that O2 uses to
start the whole thing off, and it has an extremely strong, stable triple
bond. If you look at compounds that have N=N double bonds (azines) and
especially N-N single bonds (hydrazines), the N=N bonds and especially N-N
bonds are much weaker than the triple bond in N2, and are generally more
reactive as oxidants. Also, keep in mind that N is close to O in the
periodic table, but it's also right next to C, which is usually not
considered an oxidant. Its reactivity is somewhere in the middle. Finally,
also keep in mind that, depending on reaction conditions and co-reagents,
hydrazine itself (N2H4) can behave as either an oxidant (by breaking the N-N
bond) or it can be oxidized by a stronger oxidant, by losing H atoms to form
N2. This is the basis of its use as a rocket fuel.

H2N-NH2 + O2 --> N2 + 2 H2O

Quote:
When we say that Oxygen is an oxidixer, do we really mean that diatomic
oxygen (which I would think is relatively unreactive) is an oxidixer,
or do we mean that atomic oxygen (with two unpaired electrons) is an
oxidizer?

Generally, yes, diatomic oxygen is the reactive species.

Eric Lucas
Back to top
Oscar Lanzi III
science forum Guru Wannabe


Joined: 30 Apr 2005
Posts: 176

PostPosted: Mon Jul 17, 2006 12:26 am    Post subject: Re: Nitrogen as an oxidizer Reply with quote

Not inert. Suppose you batch-anneal aluminum-killed steel at too high a
temperature, specifically a temperature where austenite starts to form.
Then nitrogen diffuses into the austenite and precipitates aluminum
nitride. Changes the properties of the product. Batch annealing cycles
are designed (among other things) to prevent this.

--OL
Back to top
Fred Kasner
science forum beginner


Joined: 27 Nov 2005
Posts: 25

PostPosted: Mon Jul 17, 2006 12:25 am    Post subject: Re: Nitrogen as an oxidizer Reply with quote

joepad1991@yahoo.com wrote:
Quote:
I am a young science student and I have a question about diatomic
nitrogen gas. I always read that N2 is inert. Nitrogen is
electronegative, just as Oxygen is. It is only a single atomic number
removed from Oxygen. With O2, when the temperature is high enough, the
double bound holding the two Oxygen atoms together breaks, leaving 2
atoms of oxygen, each with two highly reactive unpaired electrons
(which I suppose is another way of saying that Oxygen is
electronegative). When the temperature reaches this point (different
for different fuels), we say that combustion sets in.

Well, why doesn't the same thing happen with Nitrogen? Why is Nitrogen
inert? In fact, one might think it more reactive than Oxygen. If the
temperature rises to such a point to break the triple bond, you would
have atomic Nitrogen with THREE unpaired electrons. This might make
Nitrogen more reactive than Oxygen (though granted it is less
electronegative than oxygen).

So why isn't Nitrogen gas an oxidizer?

Second question:

When we say that Oxygen is an oxidixer, do we really mean that diatomic
oxygen (which I would think is relatively unreactive) is an oxidixer,
or do we mean that atomic oxygen (with two unpaired electrons) is an
oxidizer?

Thanks everyone for considering my questions.

Joe


It will probably put you the track for a better understanding of O2 if
you notice that the usual Lewis structure for O2 is incorrect. One pair
of electrons is not really a pair but are two unpaired electrons. Not so
in the case of N2. Evidence? O2 is paramagnetic. The extent of the
paramagnetism? That associated with two unpaired electrons.
FK
Back to top
donald j haarmann
science forum addict


Joined: 24 Mar 2005
Posts: 58

PostPosted: Mon Jul 17, 2006 12:15 am    Post subject: Re: Nitrogen as an oxidizer Reply with quote

<joepad1991@yahoo.com>

|I am a young science student and I have a question about diatomic
| nitrogen gas. I always read that N2 is inert.

[snip]


----------
Inert?! Forsooth. Mg will burn in N2.

NB "Oxidation" and it brother "reduction" are all 'bout electrons no oxygen is required.

Sometbing else to think about CO can be either an oxidizing or a reducing agent.


--
donald j haarmann - independently dubious
Back to top
joepad1991@yahoo.com
science forum beginner


Joined: 16 Jul 2006
Posts: 3

PostPosted: Sun Jul 16, 2006 11:21 pm    Post subject: Nitrogen as an oxidizer Reply with quote

I am a young science student and I have a question about diatomic
nitrogen gas. I always read that N2 is inert. Nitrogen is
electronegative, just as Oxygen is. It is only a single atomic number
removed from Oxygen. With O2, when the temperature is high enough, the
double bound holding the two Oxygen atoms together breaks, leaving 2
atoms of oxygen, each with two highly reactive unpaired electrons
(which I suppose is another way of saying that Oxygen is
electronegative). When the temperature reaches this point (different
for different fuels), we say that combustion sets in.

Well, why doesn't the same thing happen with Nitrogen? Why is Nitrogen
inert? In fact, one might think it more reactive than Oxygen. If the
temperature rises to such a point to break the triple bond, you would
have atomic Nitrogen with THREE unpaired electrons. This might make
Nitrogen more reactive than Oxygen (though granted it is less
electronegative than oxygen).

So why isn't Nitrogen gas an oxidizer?

Second question:

When we say that Oxygen is an oxidixer, do we really mean that diatomic
oxygen (which I would think is relatively unreactive) is an oxidixer,
or do we mean that atomic oxygen (with two unpaired electrons) is an
oxidizer?

Thanks everyone for considering my questions.

Joe
Back to top
Google

Back to top
Display posts from previous:   
Post new topic   Reply to topic Page 1 of 1 [15 Posts] View previous topic :: View next topic
The time now is Sun Jun 25, 2017 10:25 pm | All times are GMT
Forum index » Science and Technology » Chem
Jump to:  

Similar Topics
Topic Author Forum Replies Last Post
No new posts Liquid Nitrogen in a Swimming Pool - - - Experiment (2006) fufko@sbcglobal.net Chem 11 Sun Jun 18, 2006 2:37 am
No new posts How is the nitrogen in air accelerated in ramjets/scramjets? Robert Clark Mechanics 8 Fri May 19, 2006 4:28 pm
No new posts is liquid nitrogen reasonably affordable anywhere anymore? Mike Amling Chem 3 Sat Jul 16, 2005 11:33 am
No new posts SP hybridization in nitrogen Maleki Physics 2 Fri May 06, 2005 12:52 pm
No new posts sp hybridization for nitrogen Maleki Physics 13 Wed May 04, 2005 12:50 pm

Copyright © 2004-2005 DeniX Solutions SRL
Other DeniX Solutions sites: Electronics forum |  Medicine forum |  Unix/Linux blog |  Unix/Linux documentation |  Unix/Linux forums  |  send newsletters
 


Powered by phpBB © 2001, 2005 phpBB Group
[ Time: 0.0393s ][ Queries: 20 (0.0050s) ][ GZIP on - Debug on ]