What is the real cause of the boiling (forming of bubbles) of water? I've got a question about the boiling of water. I'm a first year physics student and from the Netherlands. 
I've searched alot about the boiling of water and this confused me. Everyone said something else about the cause of the boiling. Let me explain it further. 
Let's say you want to cook some eggs. You put on the gas. 
1. As the temperature of the water increases, the evaporation increases. 
2. When the evaporation increases, the vapor pressure will increase too. When the vapor pressure is equal to the external pressure, there will form a bubble.
So, my question is: is this a chain/link of causes? So the first link causes the next one? So the temperature increase causes the evaporation to increase which causes the vapor pressure to increase which causes the forming of a bubble (the actual boiling)?
I doubt if it is a link of causes (the one thing causes the other) because they happen at the same time. And in my opinion a cause happens BEFORE the consequence. 
https://www.youtube.com/watch?v=CfagHzOtIDM
In this video they say boiling has more causes. What are these causes? 
When the boiling point is reached, there happen two things:
- 100 degree celcius
- vapor pressure that equals the external pressure 
Are they both causes of the boiling (coming up of the bubbles)? Or is the vapor pressure that equals the external pressure the cause?
 A: 
When the vapor pressure is equal to the external pressure, there will form a bubble.

Not true.   Instead, when the vapor pressure is equal to the external pressure, then any existing bubbles will begin growing continuously.
And, if no bubbles are already present, then the water will superheat far above the boiling temperature, yet no bubbles will appear.  For vapor pressure to exist within water, first gas pockets must exist within water.  We need gas-fluid interfaces.  Without these, all boiling takes place silently, at the surface where water touches air.
Beware of common misconceptions.  Boiling-bubbles are typically not seeded by dirt or contamination. ("Mythbusters" show got it wrong![1]) Instead, boiling is seeded by existing micro-bubbles trapped in small crevices.  So, yes, your coffee heated in a microwave oven can superheat and explode, even though coffee is very impure water. If the surface of your ceramic mug lacks air-filled micro-scratches, the coffee will not boil until its temperature is raised far above 100C. But liquids may refuse to boil even when up against a very rough surface, if that surface has been previously wetted with water over 100C. The hot water fills the microscopic roughness with steam, which then condenses, removing any air pockets that let the rough surface act like a "seed" for roiling boil.  A fully-wetted rough surface won't prevent superheating.
On a typical stove with a metal pot, the metal bottom will be heated far above 100C, even though the water has not yet approached 100C.  The metal surface will be covered with spontaneous steam pockets, but these bubbles cannot grow, since they're right against cooler, under-100C water. When the main volume of water reaches 100C, seed-bubbles are already present on the hot metal, so the pot will immediately begin a visible boil.  But without this large temperature excursion at the metal surface, visible boiling may not commence.
Often with new glass cookware, (with no scratches,) and with water heated on a gas stove (with no tiny hot-spots,) the water won't boil.  Instead it superheats far above 100C, then unexpectedly produces a few spontaneous micro-bubbles, and exhibits the boiling-explosions called "bumping."  The explosions may splash boiling water out of the container.  Sometimes they're violent enough to shatter glass.  To prevent this, use 'boiling stones' sold in laboratory supply catalogs.  Or with microwaved coffee, provide a dry, air-filled wooden stir-stick.  These cause a roiling boil at 100C, which cools the bulk liquid and halts superheating.
The 'boiling stones' commonly used to prevent bumping-explosions in laboratory glassware won't work if cooled and then quickly re-used.  This happens because all their internal crevices become filled with water.  No small bubbles, no "boiling seeds." Restore your 'boiling stones' before re-use, by allowing them to dry thoroughly so the air again fills all of their internal pockets.  Or, use a 150C drying oven to rapidly boil off their trapped water.   Similarly, a sodden, well-boiled wooden stick in your microwave coffee may stop working.  So turn it over and use the dry end!
[1]Big caveat: if micro-bubbles aren't present anywhere, then superheated water will continue to rise in temperature, finally boiling spontaneously with explosive violence. The "seed" or "trigger" for this boiling can be:  intrusion of a dry object, vibration and sound waves, thermal fluctuations, ionization from background radiation, and yes, suspended particles, surface roughness, and contamination.  Mythbusters(tm) did have it right, but only regarding extremely superheated water.  When we first eliminate all surface micro-bubbles and then raise the temperature far past boiling, eventually something will trigger a steam-explosion. The center of this explosive vapor-production might be a tiny grain of dirt, or a pencil tapping against the container, or even a cosmic-ray strike.
A: While the answer of wbeaty is very interesting in showing points relevant in practice, I think all the answers are still missing an important and simple theoretical point, which you should consider to understand the process. 
vapour pressure does mean two different things as used above. First, the pressure, the existing water vapour would have (if it were alone) - let's call it the partial pressure of the water. And secondly, the maximal possible vapour pressure (a function of temperatue) - let's call it the saturation pressure. 
The latter is easy to understand qualitatively:
If you have a surface between water and vapour, you have two random processes: particles hop out of the surface into the gas, and particles from the gas get caught in the water. After some time, there will be equillibrium of these processes. The first process depends mainly on temperature (how fast are the water particles), the second mainly on the pressure (how often do vapour particles hit the surface).
There has thus to be some function $p(T)$ or $T(p)$ which describes these two at equillibrium. This is the abovementioned saturation pressure. 
Over every water surface there is some vapour, and after a while it has the saturation pressure - if nothing moves to blow it away. If you blow on your soup to cool it, that's the point: to remove vapour, reduce its partial pressure, and thus enable the water to evaporate further.
Note: the net pressure of the air is always atmospheric, you just change the part of it produced by water and the other part produced by the rest gases.
Now, what happens if the saturation pressure gets greater than the atmospheric?
There is no need to blow anymore! The vapour can just push the atmosphere away, it doesn't have to wait, it just takes the place it need to evaporate into. That's boiling. 

PS: concerning the concept of partial pressure.
The necessary asumption for this concept is, that we deal with ideal gases. The particles do not interact, they just bump against the walls and thus produce the pressure. The net pressure is just the sum of the pressures each part of the gas produces. I.e., each part of the gas would have if it were alone at that temperature in that volume. 
For an ideal gas you know: $pV=nRT$
So since $p$ is proportional to $n$ it means, that $n_1$ of oxygen and $n_2$ of nitrogen will produce the sum of the pressures each part does by itself.
Mark: this is not self-evident, it's a concequence of the proportionality between $p$ and $n$. 
A: It's the vapor pressure equaling the external pressure that is the cause, and not the temperature (except to the extent that it determines the vapor pressure).  When the vapor pressure is equal to the external pressure, bubbles can begin forming under the liquid, and the atmosphere can be pushed back at the upper surface of the liquid to accommodate the additional volume of the bubbles within the liquid.  The heating and temperature changes that take place before the temperature reaches the required value (for the vapor pressure to equal the external pressure) does not matter.
A: It seems like you are complicating the process a bit too much.  As noted in the earlier response: boiling occurs when the vapor pressure of the liquid reaches the surrounding pressure (e.g. of atmosphere or atmosphere + pressure of overlying liquid (heating usually done from bottom)).  Evaporation, in the context of initiation of boiling, serves only to carry heat away and reduce the temperature rate of rise ("a watched pot never boils").  You can attain boiling in two ways: heat the liquid or reduce the overlying pressure.  Put water into a sealed container and draw a vacuum on it, and it will boil at room temperature...
Water at room temperature has a vapor pressure (a few mm Hg pressure) - by heating it you are only raising it's vapor pressure until it reaches the atmospheric pressure (at your current elevation - at higher elevation, water boils at a lower temperature).  Once boiling is achieved, the temperature will not increase further - because you are removing heat through evaporation as fast as you are putting heat into the liquid from the external heat source.
There is one other added complexity as well: in order for a vapor bubble to form, the vapor pressure must overcome the surface tension of the water in order to make a "hole" in the water.  When the water is very pure, and the container surface is very smooth, it can be difficult for vapor bubbles to form and the water can become superheated - this is why, in chemistry lab, boiling chips are placed in beakers - the ceramic chips provide nucleation points for vapor bubbles to form and avoid superheating the water.
A: Boiling, liquid starts to vapourize which means sufficient heat is provided break the inter-molecular bonds. the temperature at which the inter molecular bond breaks will be called as boiling point of that liquid.Breaking of the bonds results in the formation of bubbles i.e liquid  is broken into gases.
