Carbon has 6 protons, not 12. The atomic mass of carbon-12 is 12 because it has 6 protons and 6 neutrons. The common isotope of oxygen, oxygen-16, has 8 protons and 8 neutrons. The number of protons and the number of electrons match for a neutral atom: that’s why oxygen has 8 electrons.
Chava
Rats! You beat me to it. Here’s the link:
http://www.straightdope.com/mailbag/moxygen2.html
Another correction: Adding a proton to a C nucleus will, indeed, give you nitrogen, but with an atomic number of 7, not 13.
Rick
SDStaff Hawk avoids the question why we can’t breathe 03. Here’s my shot at an answer:
The confusion here stems from the use of the word “oxygen.” When we talk about O3 we call it ozone, but the word “oxygen” means both the oxygen atom, O, and the oxygen molecule we breathe, O2. Most everyone knows that animals need oxygen, but it’s not the oxygen atom that is required, it’s the oxygen molecule, O2. There are specific chemical processes that take place when animals convert stored energy (sugars) into other forms of energy (heat, motion, etc.) and those chemical reactions rely upon the O2 molecule. O3, with its different structure, bonding energies, etc, won’t work in animals adapted to use O2. Similarly, ethanol (C2H5OH, ie, booze) reacts very differently in the human body than glucose (C6H12O6) even though they’re made of the same sorts of atoms.
So, why can we breathe O2 instead of O3? Because that’s what we’ve evolved to breathe. And why did we evolve that way? The answer is in SDStaff Hawk’s column: Because most free oxygen is O2. The pressures of natural selection favored organisms that used the plentiful O2 molecule. There may at one time have been some form of life that randomly mutated to use O3, but it probably immediately suffocated due to the scarcity of O3.
Just to be thorough, I ought to point out that it isn’t just animals that use O2. Plants create O2 when they photosynthesize, turning light and CO2 into sugars and O2. When they metabolize those sugars, they use O2, basically reversing the process. But that’s getting off the subject of the original question.
-SengkelatOK, there was a minor problem with atomic weights versus atomic numbers, and a few things were glossed over, but other than that it was a very good explanation of covalent bonding. I say this from the viewpoint of someone who has taught this subject. I may use this explanation myself sometime.
A very good explanation of covalent bonding, yes, but as far as I can tell, I don’t think I saw any answers to the questions asked by “Terry in Vegas”. It always frustrates me when the SDSAB answers related question without answering any of the actual questions (and this ain’t the first time).
The only one that comes close is the question:
“Why does it have to be in a bond of two oxygen molecules [sic] for us to be able to use it?”
But I think Hawk answered this question instead:
“Why is most of the oxygen we use in the form of two bonded oxygen atoms, and how does that bond work?”
And there’s no coverage at all about “whatever happened to O1” – I’m not a chemist, and I always wanted to know what exactly is the difference between singlet oxygen and free oxygen atoms, and what would happen if we tried to breathe them. And in this field, it can get pretty hard for a layman to distinguish between science:
http://www.meto.umd.edu/~owen/METO123/OZHOLE/lecture.html
and pseudo-science:
http://www.asc-alchemy.com/oxygen.html
So, c’mon, SDSAB, give us some information that we didn’t learn in freshman chemistry. Pretty please??
I have to say I agree with onigame. While the article was a pretty good description of covalent bonding and how oxygen molecules form, it didn’t really address the questions about why we breathe O[sub]2[/sub] not O[sub]3[/sub], or anything about O[sub]1[/sub].
The first supplied link by onigame gives some discussion of what is ozone, how it differs from regular diatomic oxygen, and how it is beneficial and detrimental. However, it does not discuss monatomic oxygen - O[sub]1<[/sub].
Monatomic oxygen does not form well in the lower atmosphere because it is unstable. It “wants” to fill the outer electron shells, and has plenty of other oxygen molecules around to do so. Also, the UV light is blocked, which is the wavelength/energy required to break the oxygen molecules into individual atoms. However, in the very upper atmosphere (above the stratosphere, at the level that the space shuttle and station orbit), there is monatomic oxygen, and UV. The UV is a light more intense, providing energy to break the molecules apart, and the oxygen atom density is much lower. The equilibrium of the two gives more energy to break apart than chances for the atoms to join together. However, under the ozone layer where more of the UV is blocked and oxygen atoms are greatly more abundant (packed much tighter together), the frequency of bonding is much higher and fewer molecules are given the proper energy to break apart, maintaining an equilibrium of mostly O[sub]2[/sub].
As for the description by Sengkelat of why the body requires O[sub]2[/sub], I’m not sure I agree. I don’t know the specifics of body chemistry, but it is my understanding that how oxygen enters the body is through absorbtion through the lung lining and dissolving into the bloodstream, being picked up by the hemoglobin. Perhaps someone could describe how hemoglobin works, and how the oxygen gets through the lung tissue into the blood. That would provide insight into how O[sub]2[/sub] is used and processed. I suspect O[sub]3[/sub] is unusable by the body because it takes more energy (or perhaps a different reaction, per Sengkelat?) to break the molecules apart, and the first break forms the diatomic molecule plus one free atom, vs. two free atoms. Thus the toxic effect is the replacement of abundant free oxygen atoms with a lot more diatomic O[sub]2[/sub], reducing lung efficiency. That would match with the effects listed on that site mentioned above:
Unfortunately, that site does not provide any more information on the health effects. Even the link about “bad ozone” really just talks about the effects of increased UV radiation, not the effects of breathing ozone.
I hope someone with more information on this topic will pick up and fill in the gaps.
This was the first time I actually felt like my intelligence was insulted by an answer on the Straight Dope. Wan? No, actually, the stuff you explained wasn’t that hard. Why don’t you finish?
I was wondering about this quote
It seems to me that how an atom reacts to other atoms is determined by the number of electrons that it has. That is in turn affected by how many protons it has, but the electrons are the direct cause. In high school, we were asked what determines how an atom reacts, and I gave this answer. I was told that I was wrong. I argued for a while, but eventually said “Well, okay, whatever.” Then I got to college, was asked this question, and gave the answer I learned in high school. I told that I was wrong. Arggh! One of the many reasons that I hate chemistry with a passion. So which one is it? If you have an atom with six protons and seven electrons, would it act more like carbon or like nitrogen?
The Ryan,
The correct answer is protons. The number of protons determines the element. In a neutral atom the number of electrons equals the number of protons, but in an ionized atom the number of electrons is different than the number of protons. However, the loss or gain of an electron in an ionized atom does not change what element it is.
For example, Sodium has 11 protons. Neutral sodium has 11 electrons, but when you ionize sodium by stripping the electron in the 3s orbital you get an atom with 11 protons and 10 electrons. This atom is still sodium, it does not become Neon.
Protons determine the element. The number of protons determines the number of electrons in a neutral atom of the element, which determines its chemical properties. However, there is more to atoms than chemical reactions. There are nuclear reactions, which are determined by the protons (and neutrons).
The Ryan,
I think of it this way. (It may not be correct but
nothing’s contradicted me yet!)
When we speak of an element, such as “carbon”, we are
actually talking about many different things. We
could be talking about a single atom; we could be
talking about a highly organized diamond, we could
be talking about charcoal. We could be talking
about radioactive carbon (the one used in carbon-dating),
we could even be talking about a carbon atom in
a complex DNA molecule. It’s pretty obvious that
even though we call all of these “carbon”, they act
in different ways.
So why do we call them all “carbon”? Simple – because
the involved atoms all have a nucleus with six protons.
That is the definition of “carbon”. Calling something
“carbon” has very little to do with how it acts on
a macroscopic level.
So, when you ask “If you have an atom with six protons
and seven electrons, would it act more like carbon or
like nitrogen?” the question is not very meaningful –
because “acting like carbon” and “acting like nitrogen”
do not represent simple actions.
To put it another way, an atom with six protons and
seven electrons acts exactly like a carbon atom with
an extra electron. It doesn’t act like a carbon
atom with six electrons, and it doesn’t act like
a nitrogen atom with seven electron. It probably
acts more like a silicon atom with fifteen electrons
than any other (different type of) atom you can
think of.
Or, here’s another example. Say you have a happy
little atom, with 18 protons and 18 electrons.
It doesn’t want to interact with anything. It’s
really stable. Say we have another atom with
18 protons and 17 electrons. This one’s a vicious
little thing, because it really really wants another
electron, and it will do almost anything to get it.
Do these two atoms act differently because they have
a different number of electrons? Sure they do.
Does this mean they’re of different elements?
Certainly not. And does the second one act like
an atom with 17 protons and 17 electrons? Well,
not really.
Not quite. Ozone is unstable in that it easily gives up the third oxygen atom. When you breath air with lots of ozone, the ozone reacts with the molecules in the lung tissue by giving them the excess oxygen atoms. So it’s basically burning the lungs and that leaves scar tissue.
Single oxygen atoms are extremely reactive, much more so than O[sub]2[/sub]. You won’t find them in a natural environment on Earth.
Sheesh.
OK, the discussion of covalent bonding was good. But an understanding of covalent bonding, while necessary for answering the question is not sufficient for answering the question.
Singlet oxygen (O1), molecular oxygen (O2), and ozone (O3) are all breathed in and out by all of us every day. They all exist in the atmosphere, and they all can cross membranes from alveoli into capillaries in the lungs. The difference is what happens after that.
Singlet oxygen is one of the most reactive forms of matter (singlet fluorine is the most reactive, at least in my 1970s era Chem classes). It it formed when sufficient energy is added to a molecule of O2 to break the covalent bonds. High in the atmosphere (the ozone layer), this energy comes from the sun. Closer to earth, this happens whan an O2 molecules collides with another molecule with enough energy transfer to break the bonds. In air, as soon as an atom of singlet oxygen collides with anything, though, it tries to covalently bond with that other molecule to fill its outer shell. The molecule it most often does this with is O2, and that is how ozone is formed.
If you’ll go back to Hawk’s explanation of covalent bonds, and try to arrange the 18 electrons among the 3 nuclei so that each nucleus keeps 8 electrons in its shells, you’ll get frustrated - it doesn’t work out neatly like it did for O2 - electrons move around in a constant effort to find a stable arrangement, but there is not one as stable as for O2. So ozone is a more reactive, less stable for of oxygen than O2, but it is more stable, and less reactive than singlet oxygen.
Inside the body, it turns out that singlet oxygen and ozone can’t substitute for O2. They are physically different - remember that atoms take up a 3 dimensional volume, and don’t fit into the same “pockets” that O2 does in enzymes and and other large molecules that use oxygen. Additionally they are chemically different - they behave in an entirely different manner.
They (ozone and singlet oxygen) are termed “reactive oxygen species” and can lead to to formation of other unstable things like free radicals - for example, an oxygen covalently bonded to a hydrogen with a total of 7 electrons in the outer shells, instead of the 8 that lead to stability. These free radicals quickly react with other molecules, and more or less permanently alter the structure of DNA, proteins, and other crucial molecules. Since alterations in structure lead to alterations in function, this is bad. The body has a repair mechanism to repair this so-called oxidative damage (with the help of Vitamin, C, vitamin E, flavinoids and other antioxidants) but if too much damage happens too fast, the repair mechanisms can’t keep up and cell death, or other irreparable harm is done to the body.
First I would like to commend Facts R Us on an insightful and informative post. He beat me to much of my reply, but here’s my $0.02.
First, the concentrations of monotomic oxygen and ozone are extremely low at breathable atmosphereic pressures. Although it is theoretically possible to inhale these species, they would almost certainly not dissolve in the bloodstream. The same is true for most atmospheric gases (including oxygen)- at normal pressures they will not dissole into blood to any appreciable degree. The method our body uses for transporting gas is the protein hemoglobin (in the red blood cell). As Facts R Us stated, proteins and enzymes work based on three-dimensional shapes – called lock-and-key functionality. Hemoglobin has practically no affinity for either ozone or O1- because their ‘shapes’ are wrong. Thus even if either of these species was available in appreciable amounts, our body simply wouldn’t be able to pick them up.
As further info, the affinity of hemoglobin for common gasses from strongest to weakest is cabon-monoxide, carbon-dioxide, oxygen. Hence the ease of suffocation with CO.
Warning, a lot of chemistry jargon follows.
This may seem like a minor detail, but actually pretty important for explaining oxygen’s reactivity. Your average O2 molecule in the air is not exactly “sharing” 4 electrons, and you don’t really have a double bond between the two oxygen atoms. Rather, you have a single bond with two unpaired electrons. So O2 is really a “diradical” (a molecule with an unpaired elecron is a “radical” or “free radical”).
Now, in high school chemistry, which is about the level presented in the original answer, we are taught that we always pair up the electrons as much as possible, so it would seem natural to assume a double bond between the two oxygen atoms. But forming a double bond between oxygen atoms would actually result in an “antibonding” molecular orbital-- forming the double bond would destabilize the molecule.
The unpaired electrons give oxygen some interesting properties. For example, it is paramagnetic (it is attracted by a magnetic field) because the unpaired electrons can alter their spins to align with the magnetic field. And the unpaired electrons make O2 very eager to react with other compounds. Compare to N2, which has no unpaired electrons and true triple-bond character between the N atoms. N2 is quite inert.
And now about some other terminology. “O1” or simply “O” and “singlet oxygen” are not the same thing. “O” would be “monatomic oxygen”. It would also exist as a diradical, but as soon as it runs into another monatomic oxygen, it will react to form an O2 diradical. The O2 is not especially “happy”, but it’s much “happier” than a monatomic oxygen. So if monatomic oxygen exists at all, it’s at such low concentrations that you would never notice it.
“Singlet Oxygen” is a special form of O2, often written as “1O2”. The standard O2 is actually “Triplet” oxygen, or “3O2”, reffering to the three different magnetic states possible because of the unpaired electrons. I think that if you excite triplet oxygen with exactly the right amount of energy (say from a photon of light) you can “excite” triplet oxygen, and give the unpaired electrons exactly enough energy to form that “antibonding” orbital, forming an O-O double bond. Now, all of the electrons are paired, and there is only one magnetic state possible, hence “singlet” oxygen. The singlet oxygen is a high-energy species, and will dump its energy at the first opportunity by reacting with whatever it bumps into.
For more info, see:
http://courses.chem.psu.edu/chem38/mol-gallery/oxygen/oxygen.html
Freightliner
Damn. I’ll try harder next time.
Few elements can exist as a single element. Only the the atoms which have all of their electron shells already filled are stable as independent atoms. If you look atthe periodic table you will see these atoms all the way to the right. They are called the noble gasses. So Oxygen is not unusual in its preference to share electrons. It just happens to get along really well with itself. It also gets along well with almost anything else. (ie rust, most minerals, sugars…)
Monomeric oxygen (a single oxygen atom) is so reactive it will likely react with the first atom it hits. That is why it does not exist in any appreciable quantity.
Now for ozone. If you read the previous examples you should know that oxygen has two electrons available for bonding. Try conecting three oxygen atoms using up all of these bonding electrons without making a triangle.(A triangle is too strained to exist.) It doesn't take much trying to realixe that no matter how you do it one bonding electron is always left free. This is why ozone is so reactive. It is ready to give up that oxygen to anything that will take it including almost anything in your body. You can't breath it, because it reacts with manythings in your body.
Ozone will also give up its oxygen to a CFC. Ozone is produced in our upper atmosphere by high energy reactions. CFC's very efficiently deplete this supply. Unfortunately we need ozone in our upper atmosphere to absorb UV rays.Ozone is actually destroyed by chlorine ions and not CFC’s directly. The original CFC has three chlorine atoms in it. When struck by UV radiation, it decomposes into a CFC with two atoms and a free chlorine radical. That chlorine radical combines with ozone to form chlorine oxide, ClO, and oxygen gas. ClO then pairs up with a free oxygen atom to form another chlorine radical and more oxygen gas. In essence, the chlorine atom is acting as a catalyst for the decomposition of ozone into oxygen gas, and can remain in the atmosphere for approximately 150 years, able to destroy ozone molecules over and over again. Luckily, ozone can also be produced naturally by using energy from that same UV radiation to combine oxygen gas and a free oxygen atom. Unfortunately, the damage that has been done will not go away fast because the results of our pollution are still up there in the atmosphere causing damage.
I was happy to see that someone (freightliner) corrected Hawk’s misstatement that O2 has a double bond. At least Hawk didn’t write O=O…
Sorry, Jill, too late. Best you can hope for is second.
I’d completely forgotten about oxygen having two half-filled pi-orbitals… it’s been way too long since I finished physical chemistry and started working in molecular biology. 
I wanted to comment on a few things:
Christopher, many elements other than noble gases can exist in “free” form. Most elements that exist in such a form are the unreactive metals. The metal atoms (they’re neutral) form a crystal lattice, but they can’t even be said to be part of an ionic compound 'cause there aren’t any other atoms near them. Gold, silver, copper, platinum, and so forth are all found this way in nature, sometimes even in large nuggets (which will generally be mixed with oxides as well, most especially in copper’s case).
As to the structure of ozone: in Lewis form, it’s generally considered to be a resonance structure consisting of two each single and a half-bonds. It’s been so long since I did MO theory that I can’t figure the orbital structures. <sigh>
LL