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When iron is left outside, over time it reacts with oxygen to form iron oxide or rust. But how does the oxygen react with iron, isn't iron held together by strong metallic bonds? Where does this required energy to initially begin the reaction of iron and oxygen come from?

Do all reactions initially require a input of energy?

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Do all reactions initially require a input of energy?

Yes, it is called activation energy. When you say "initially required", it does not mean just when the overall reaction starts. Each time another atom or molecule reacts, you need that activation energy again.

But how does the oxygen react with Iron, isn't Iron held together by strong metallic bonds.

Yes, for every oxygen that reacts, there has to be some activation energy to start prying apart the bonds. For every iron atom that reacts, there has to be some activation energy to start prying it from the metallic bonds. In the case of oxygen reacting with iron, you get that energy back (and more) once the bonds between iron and oxygen form.

When Iron is left outside, overtime it reacts with oxygen to form Iron Oxide or rust.

If you leave iron inside, this also happens because there is as much oxygen inside as outside. However, when it is wet and salty, the reaction proceeds at a higher rate, and it might be more humid and saltier outside (especially in the ocean). You can avoid this by providing a barrier between iron and oxygen, for example by painting your boat.

Where does this required energy to initially begin the reaction of Iron and Oxygen come from?

It comes from the thermal energy. Thermal energy is not equally distributed, there are always some particles that have higher energy. They might react, or give that thermal energy to the next particle, which might react. We put our food in the refrigerator or freezer to slow down the reactions (and the metabolism of microbes) that would spoil our food.

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    $\begingroup$ Why is it that "wet" and "salty" speed up the reaction? I also read somewhere that pure Iron is very hard to rust. Why might that be? $\endgroup$ May 21 at 3:35
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    $\begingroup$ Check out the answer by @AChem and the linked video. $\endgroup$ May 21 at 10:41
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When Iron is left outside, overtime it reacts with oxygen to form Iron Oxide or rust. But how does the oxygen react with Iron, isn't Iron held together by strong metallic bonds.

Let's forget about iron for now. Allow me to mention a very general trend. Any metal in the periodic table that is generally not found in its native form in the Earth's crust will eventually corrode when exposed to the open atmosphere. Using chemical reactions, humans input a lot of energy in the form of heat and electricity to extract metals out of the ore, but the metals would like to return to their original states, as they existed "naturally" in the combined state. Isolated metals such as sodium or potassium are extreme cases. They instantly corrode to sodium and potassium oxides, the instant air and moisture touches them, because these metals never occur naturally. The same holds true for iron. In the Earth's crust it occurs in a combined state with oxygen or sulfur. Thus pure iron would eventually combine with oxygen (as a matter of time). Copper reacts with the atmosphere, as does aluminum. However, gold which occurs in its native state does not corrode in the air. A crude example is that of a living human body. We come from the dust and eventually become dust. It is a matter of time.

Corrosion is not a simple process. People do PhDs in this subject trying to understand the process or slow it down. Corrosion is an spontaneous electrochemical reaction, i.e., in the case of iron and oxygen there is a redox reaction in the presence of moisture. This is beautifully demonstrated with the help of a galvanic corrosion demonstration. Royal Society of Chemistry Website

As you state in the comments, impurities do help in forming local electrochemical cells on the metal surface. Sometimes pure metals can form a uniform and non-porous protective oxide layer. That aspect slows downs corrosion drastically even though the metal itself is very reactive. A common example is that of aluminum otherwise we would not be able to see aluminum pans in the kitchen.

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    $\begingroup$ Of course, free oxygen is not exactly a "natural" thing either. Plants and algae spend a lot of energy (coming from the sun) to split carbon dioxide, releasing waste oxygen in the process. The reason iron "naturally" rusts into iron oxides (etc.) is tied to the fact that right now, the Earth has a heavily oxidative atmosphere. That wasn't always the case. $\endgroup$
    – Luaan
    May 22 at 14:11
  • $\begingroup$ @Luaan To add to that, isn't there a theory that oxygen levels in the atmosphere were kept artificially low for an extended geological period after the origin or photosynthesis as all the iron on the planet rusted? $\endgroup$
    – Dannie
    May 23 at 8:45
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    $\begingroup$ @Dannie Yup; most of the available iron was dissolved in the oceans - iron oxides are no longer soluble in water and precipitate out, which is where the massive banded iron formations we heavily rely on come from. The "bands" tell us a lot about how this proceeded, much like growth rings in trees. In fact, it's thought that this was a very lucky thing for the algae - early on, they couldn't live in oxygenated environments themselves, and the iron-rich waters got rid of the waste oxygen periodically. When algae developed oxygen resistance, the oxygen production really took off. $\endgroup$
    – Luaan
    May 23 at 9:27
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@Howard Stark. Rust is a punctual phenomena. It happens on the surface of an iron piece, but only if there is some conducting impurity (carbon grain for example) on this surface. If a drop of water falls on this contact $\ce{Fe-C}$, the following electrochemical phenomena start.

The first $\ce{Fe}$ atom touching the grain yields two electrons according to $$\ce{Fe -> Fe^{2+} + 2 e-}$$ These electrons are moving through the carbon grain up to the dissolved $\ce{O2}$ in the water drop, according to : $$\ce{O2 + 2 H2O + 4 e- -> 4 OH-}$$ Then both ions $\ce{Fe^{2+}}$ and $\ce{OH-}$ are attracted to one another and produce a precipitate $\ce{Fe(OH)2}$. This precipitate is quickly oxidized by oxygen form the air according to : $$\ce{4 Fe(OH)2 + O2 + 2H2O -> 4 Fe(OH)3}$$ Rust is mainly made of the brown precipitate $\ce{Fe(OH)3}$, which is later on partially dehydrated and even partially carbonated.

The main point here is the fact that this corrosion starts and continues at the very contact $\ce{Fe-C}$ and not in the bulk of the metal.

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