The Periodic Table of Elements III: Electron configuration
Did you know that in the 1980s, there were more than half a million cases of AIDS, a deadly new disease without a cure? Scientists set to work building models of molecules that could stop the lifecycle of HIV, the virus that causes AIDS. After testing over 200 computer-simulated molecules, they found one that could keep HIV from replicating. Thanks to scientists’ understanding of electron configuration and valency, an HIV diagnosis is no longer the death sentence it used to be.
In June 1981, five previously healthy young men were treated for a strange pneumonia in Los Angeles, California. At the same time, health officials in New York noticed 26 cluster cases of aggressive, lesion-causing skin cancer. These cases would become the first reports of a terrible new disease – acquired immunodeficiency syndrome (AIDS). In 1981, there were 337 AIDS diagnoses made in the US, and the first known deaths occurred. By 1987, that number would be over 500,000 cases worldwide.
After diagnosis, death was pretty much a certainty within 18 months. People didn’t yet understand how this terrible disease was transmitted. All they knew was that most of the early AIDS victims were gay men, and that gave rise to misconceptions and panic. Panic ignited violence against the gay community. Even hospitals turned away dying people due to fear that the disease might be transmitted through casual contact.
Scientists identified the virus responsible for AIDS and called it human immunodeficiency virus, or HIV. But it would be many years before effective treatment arrived. Finally, a new chemistry technique – called structure-based drug design – brought about a breakthrough.
Scientists used X-ray crystallography to make a 3-D model of a crucial HIV protein called HIV protease. Think of HIV protease as the machinery the virus uses to replicate. The scientists’ goal? Throw a wrench in that machine!
If scientists could disable HIV protease, perhaps they could stop the virus’s spread. The most crucial part of HIV protease is the binding site (or active site). Like the name suggests, the binding site works by binding to other molecules. Scientists wanted to disable the binding site by blocking access to the molecules HIV protease wanted to bind. That should stop HIV from replicating in patients and making them sick.
So, scientists used computer-models to design hypothetical molecules that might bind HIV protease and knock out the virus.
And this is where we get to electron configuration. To model hypothetical molecules, scientists need to know how many bonds each atom in a molecule can form. This property is called valency. Different elements have different bonding capacities, or valence numbers. The reason for these differences comes down to only one of the three subatomic particles – electrons. Basically, the way electrons organize themselves around the nucleus determines an atom’s valence number. That valency-electron connection is the focus of this module.
Valency is a simple, yet powerful concept that scientists can use to do many things, including to create computer models of molecules – molecules that might not yet exist. Scientists can go on to predict shapes of hypothetical molecules and, just as important, anticipate what these molecules might be capable of binding.
Rhonda Levin was a chemist on one of these drug development teams focused on finding treatments for HIV. Levin's team synthesized and tested over two hundred of those computer-simulated molecules. "We were racing the clock to keep people alive,” she says.
One of those molecules worked. Compound number L-735,524 – later dubbed Crixivan – became one of the first protease inhibitors approved by the FDA to treat HIV. Figure 1 shows the 3-D shape of the drug and how it binds HIV protease.
Figure 1: This computer-generated image shows the antiretroviral drug Crixivan. The green structure in the center is the drug molecule, fitting snugly inside the HIV protease protein, which is shown in purple and orange. By binding to the protease molecule, the drug prevents the virus from reproducing. The drug Crixivan was first identified by computer modeling that relied on the concept of valency to predict binding and deactivation of HIV protease.
image © CC BY-SA 3.0 DEED A2-33
The valency of carbon
Recall from our Periodic Table II module that when the periodic table was first created, electrons had not yet been identified. So, scientists couldn’t have predicted that electrons played any role in the chemical properties of an element. 1800s chemists did, however, understand valency.
Clever experiments with the element hydrogen allowed chemists to work out how many bonds different elements could form. They knew that one hydrogen atom was capable of only one bond. And they worked out the valency of other elements by determining the number of hydrogen atoms with which an atom of another element could form bonds. For example, the element fluorine’s (F) valence number was 1, like hydrogen. Oxygen’s (O) was 2. Nitrogen (N) and phosphorus (P) showed higher bonding capacities, with valence numbers of 3. German chemist August Kekulé discovered the element carbon (C) had an even higher valence number of 4.
Kekulé suggested carbon’s high valence number explains why carbon makes up some of the largest molecules in existence. Today, we know that there are more types of carbon-containing molecules than all non-carbon molecules put together! That’s why the chemistry of carbon-containing molecules gets its own name: organic chemistry.
Carbon has the unique ability to link together and form chains, and even rings. Kekulé proposed the first ring-shaped carbon structure in 1872: a 6-carbon molecule called benzene (Figure 2). Over a century after Kekulé’s benzene work, Levin’s point of focus for the anti-HIV drug project would become one such ring structure on the drug Crixivan.
Figure 2: Kekulé’s proposal for the structure of benzene, 1872.
image ©Public DomainKekulé proposed that “carbon skeletons” form the backbones of the large molecules that make life possible. Years after Kekulé’s time, chemists would work out the reason for carbon’s high bond capacity. It all came down to the way electrons organize around an atom’s nucleus. This is called electron configuration. Each element has a unique electron configuration at the atomic level that determines how many bonds an atom can form.
What do atoms look like?
Carbon atoms form four bonds, thanks to the way carbon’s electrons organize around the nucleus. Of the three subatomic particles – protons, electrons, and neutrons – electrons are the ones most involved in chemical bonds. As discussed in our Atomic Theory II module, atoms contain a small positively charged nucleus surrounded by a large electron field consisting of mostly empty space. Within this space, imagine electrons as residing in shell-like layers. Each layer fits up to a certain number of electrons. Overflow electrons must go to outer shells. Thus, more electrons mean the atom must have more electron shells.
Electrons fill the electron shells closest to the nucleus before filling the next one out. We know how many electrons reside in each shell, so if we know how many electrons an atom has, we know how many shells the atom needs to hold them all.
Each electron shell maxes out at a different electron capacity. See Table 1 for the electron capacity of the first three electron shells.
| Electron shell | Electron capacity |
|---|---|
| 1st electron shell | 2 |
| 2nd electron shell | 8 |
| 3rd electron shell | 18 |
Table 1: Table shows the electron capacity of the three electron shells nearest the nucleus.
Shells further from the nucleus, such as the 3rd electron shell, fit more electrons than those nearer. This makes sense if you imagine the atom as a 3D sphere, like a balloon. As a balloon's surface expands, its surface area increases. In atoms, that means room for more electrons.
Niels Bohr first proposed a model of the atom that contained electron shells in 1913 (to learn more about this discovery, refer to our Atomic Theory II module. And his models can help us envision this electron-shell-filling concept. We see in Figure 3 that only 2 electrons fit into the 1st shell. And the second fills up at 8, leaving 7 for the 3rd shell (Figure 3).
Figure 3: This Bohr model of a chlorine atom shows the electrons distributed in the first three electron shells. The letters s and p represent subshells, which we will discuss later in the module.
image ©VisionlearningAlthough the Bohr model can help us understand electron shells, it oversimplifies the orbit of electrons. For example, the Bohr model depicts all electron shells as circular. Today, we know the electron shells consist of electrons taking differently shaped pathways. But despite the Bohr model’s limitations, it remains a powerful tool for conceptualizing electron capacity for the different electron shells. That’s important because electron shells explain why atoms react and form bonds in the first place, especially the outermost shell.
The outermost electron shell is so important that it gets a special name: the valence shell. And all the electrons that reside within it are valence electrons. Like a collector dislikes incomplete sets, we often say that atoms “dislike” incomplete electron shells. So atoms accept, donate, or share electrons with one another to either (a) fill up their outermost shell or (b) empty it completely. That’s what’s happening when chemicals react.
Valence electrons explain why carbon’s valence number is 4. A carbon atom contains 6 electrons. Two of those electrons fill the 1st electron shell. That leaves 4 electrons for the 2nd electron shell, carbon’s valence shell (Figure 4). Table 1 tells us the 2nd electron shell can hold 8 electrons. Not only is carbon capable of fitting four more electrons, but carbon atoms are more stable with four more. Carbon atoms form bonds to complete their valence electron shell. For example, a carbon atom may share electrons with other atoms, essentially borrowing other atoms electrons to achieve a full set. Since carbon has 4 out of 8 electrons in its outer shell, it can form up to four bonds with other atoms. And that’s why carbon can form up to four bonds.
Figure 4: This Bohr model shows the electrons distributed in a carbon atom’s electron shells.
image ©VisionlearningWinning the war on HIV
Carbon’s high bond capacity makes possible the molecules that Levin and her team modeled and tested. By continuing to model, create, and test new molecules with a backbone of carbon rings, the team was able to create a molecule that could bind to the protease and help save the lives of AIDS patients.
Protease inhibitors changed the tide in the fight against AIDS, especially when combined with other drugs that attacked different parts of the HIV viral life cycle. When the team finally demonstrated that replication of the HIV virus could be inhibited by almost 95 percent, it was “a magic moment," according to Levin: "People were dying, and then we were keeping them alive."
It is estimated that between 1995 and 2015, over 9.5 million lives were saved from Crixivan and other such retroviral drugs. Levin says:
I really credit the advances that we made to being open to a variety of thought and … voices at all different levels and from all different groups. Everybody's voice was listened to.… My input as a really entry-level bench scientist was listened to with the same open-mindedness as people with far more experience – all the medicinal chemists, the biologists, the virologists, and the molecular modelers.
The molecular model for Levin’s team’s HIV protease inhibitor is shown in Figure 5. Carbon, represented by gray spheres, forms a maximum of four bonds with different elements. We see other elements in the molecule as well: Hydrogen atoms (white spheres) form only one bond; oxygen atoms (red spheres) form up to two bonds; and nitrogen (blue spheres) form up to three bonds. That means that the valency for hydrogen, oxygen, nitrogen, and carbon is 1, 2, 3, and 4, respectively. Valency is how many electrons an atom of a certain element needs to gain, share, or lose – whatever is easier – for that atom to be stable.
Figure 5: This ball-and-stick model shows the structure of the antiretroviral drug Crixivan. Gray spheres = carbon atoms. White spheres = hydrogen atoms. Red spheres = oxygen atoms. Blue spheres = nitrogen atoms. As shown in the model, these four elements have different valencies, and thus form a different number of bonds.
image ©Public DomainElectron shells, subshells, and orbitals
Outer shell electrons explain an element’s valence number. But they are not the whole picture when it comes to electron configuration. Each electron shell is composed of a number of subshells. And understanding how each electron shell uses subshells allows us to connect the invisible, nanoscopic world of atoms and electrons with the familiar macroscopic world around us.
Up to now, we’ve imagined electron shells as onion-like layers composed of sphere-shaped electron orbitals. But despite the word “orbital” suggesting a roundish shape, orbitals come in four primary shapes. Each shape makes up part of a subshell. A “subshell” is all the electron orbitals in the same electron shell with the same shape.
Think of the relationship between electron shells, subshells, and electron orbitals like this: Electron shells are made of subshells. Subshells are made of electron orbitals. And electron orbitals contain the individual electrons.
There are several types of electron subshells: s, p, d, f and g. In this module, we’ll discuss only two: s and p. They are the most common subshells. In fact, all elements up to calcium (Ca) use only these two types of subshells. S-type subshells and p-type subshells differ in two major ways:
- their electron capacity
- the “shape” of the paths taken by electrons
Only the 1st of those differences – electron capacity – will be used in this module. S-type subshells hold a maximum of 2 electrons. P-type subshells hold a maximum of 6 electrons. That explains the electron capacities of the electron shells.
Recall from Table 1 that the 1st electron shell maxes out at 2 electrons total. That’s because the 1st electron shell only contains an s-type subshell.
Recall also that the 2nd electron shell can hold 8 electrons total. That’s because the 2nd electron shell contains an s-type subshell and a p-type subshell. Take a look at Table 2 to see it all together. Though we won’t discuss the 3rd electron shell much, notice that it can hold 18 electrons total. And that’s because it contains an s-type, p-type, and d-type subshell.
| Electron shell | Presence of an s-subshell? | Presence of a p-subshell? | Presence of a d-subshell? | Maximum electron capacity for shell |
|---|---|---|---|---|
| 1st electron shell | Yes (up to 2 electrons) | No (0 electrons) | No (0 electrons) | 2 electrons total (2+0) |
| 2nd electron shell | Yes (up to 2 electrons) | Yes (up to 6 electrons) | No (0 electrons) | 8 electrons total (2+6) |
| 3rd electron shell | Yes (up to 2 electrons) | Yes (up to 6 electrons) | Yes (up to 10 electrons) | 18 electrons total (2+6+10) |
Table 2: This chart shows the subshells and the electrons each subshell can hold for the 1st, 2nd, and 3rd electron shells.
The individual electrons in each subshell are organized in electron orbitals. Electron orbitals refer to regions in a subshell where electrons are likely to be found, and up to two individual electrons fit in one electron orbital. Take note that “electron orbitals” is not a word for individual electrons.
In Table 3, see that the electron capacity for each type of subshell is two times the number of electron orbitals it can contain.
| Subshell types | # of electron orbitals | Maximum electron capacity |
|---|---|---|
| s-type | 1 | 2 |
| p-type | 3 | 6 |
| d-type | 5 | 10 |
Table 3: s-type, p-type, and d-type subshells and their different electron capacities
The second difference between subshells – the shape of an electron’s path around the nucleus – is beyond the scope of this module. That’s not to say these shape differences are unimportant, though. On the contrary. In fact, these differences explain why these subshells have different electron capacities in the first place, and they come into play when exploring organic chemistry, though we won’t discuss that in this module.
Let’s pause here and organize the terms we’ve learned. First, electrons can reside in different electron shells, and these shells are numbered 1, 2, 3, etc. Within electron shells, electrons reside within subshells, and these subshells are labeled by letters s, p, and others. Finally, subshells are made up of electron orbitals. An electron orbital can hold up to two individual electrons.
Punto de Comprensión
Energy: time to level up
Understanding how electrons fill up electron shells and subshells explains why elements have the properties they do. (We will connect electron configuration to chemical properties in detail in our module Periodic Table IV. For any atom, the first electrons will fill the 1st s-subshell. The notation for referring to the 1st s-subshell is shown below.
1s
“1” indicates the electron shell.
“s” indicates the subshell.
One can indicate the number of electrons in a particular subshell by adding a superscript. For example, the notation for a one-electron atom is 1s1. See Figure 6 for details.
Figure 6: The figure identifies the symbolism in the electron configuration notation.
image ©VisionlearningThe notation above shows the complete electron configuration for a hydrogen (H) atom. See the illustration for this hydrogen atom in Figure 7.
Figure 7: The image depicts a Bohr model of a hydrogen atom, also called a “planetary model.”
image ©VisionlearningFor a two-electron atom, we need only adjust the superscript.
1s2
This is the complete electron configuration for helium (He). See the illustration for helium in Figure 8.
Figure 8: The image depicts a Bohr model of a helium atom.
image ©VisionlearningFor a three-electron atom – lithium (Li) – the last electron must go to the 2s subshell. Therefore, we would write lithium’s electron configuration like this:
Li 1s2 2s1
See the illustration for lithium in Figure 9.
Figure 9: The image depicts a Bohr model of a lithium atom.
image ©VisionlearningElectrons organize into electron shells and subshells based on different energy levels. In fact, another way chemists refer to electron shells is by calling them energy levels. Electron shells closer to the nucleus are less energetic than those farther away. And the next step up energy-wise from 1s orbital is 2s.
Check out the increasing energy level of the subshells, which corresponds to the order in which they fill, in Figure 10.
Figure 10: The arrow indicates increasing energy levels of the subshells, which indicates the order in which these subshells will fill with electrons.
image ©VisionlearningAfter 2s fills, next up is 2p. Then comes 3s, and so forth.
Punto de Comprensión
Most elements have more complex electron configurations than hydrogen, helium, and lithium since higher atomic numbers indicate more electrons. In the animation below, explore the atomic structure of the first 12 elements.
Animación Interactiva: Atomic and ionic structure of the first 12 elements
One could continue to predict electron configurations by using Figure 10. Just count up through the electrons. But for large elements with many electrons, that method can become cumbersome. For example, sodium contains 11 electrons (Figure 11). Imagine trying to write the electron configuration for radium with its 88 electrons!
Figure 11: The image depicts a Bohr model for a sodium atom.
image ©VisionlearningThanks to “the father of the periodic table,” Dimitri Mendeleev (and others after him), there’s an easier way. You can use your cheat sheet – the periodic table.
Mendeleev gets credit for arranging the elements according to similar chemical properties. His 1871 work evolved into what we recognize today as the modern periodic table. Despite almost no knowledge of atomic structure, Mendeleev still managed to group elements with similar valence shell configurations. And that’s because valence shell configuration drives chemical properties, and thus elements with similar valences behave similarly.
Generally, elements in the same column tend to form the same number of bonds. For example, group 8 elements on the far-right column – called “noble gases” – don’t easily react or bond with other elements. However, just one column away group 7 elements – including fluorine, chlorine, and bromine – all react violently with elements from the left side of the periodic table, such as sodium and lithium.
The modern organization of the periodic table lets us quickly predict an element’s electron configuration by its placement on the table. Each row of the periodic table, called a period, corresponds to the filling of an electron shell. In other words, an element’s period tells you which electron shell is its valence shell.
For example, note in Figure 12 that the top period (the first row) – mostly empty – consists of only hydrogen (H) and helium (He). This is because, as we discussed, the first electron shell contains only one orbital – a 1s orbital – which fits a maximum of two electrons. Thus hydrogen, with one electron, and helium, with two, are the only elements that have the first electron shell as their valence shell.
Figure 12: The image shows the periodic table with the different color-coded orbital blocks. Red = s-block. Blue = d-block. Yellow = p-block. Green = f-block
image ©VisionlearningNow let’s take a look at the second period. All elements in the second period (Li, Be, B, C, N, O, F, Ar) have the 2nd electron shell for their valence shell. The second period starts with lithium. We already established lithium’s electron configuration using a count-up method, reasoning that with 3 total electrons, 2 must go to the 1st electron shell, leaving 1 electron for the 2nd shell. Here is lithium’s electron configuration again as a reminder.
Li 1s2 2s1
We could have figured out lithium’s electron configuration another way. We could have used the periodic table.
Using the periodic table to predict electron configuration
The periodic table tells us about an element’s valence shell. Once we know that, we can just backfill everything that came before the valence shell. Let’s use lithium again as an example. Since lithium is in period two of the periodic table, we know that its second shell is its valence shell. And since lithium is the first element (furthest left) in this period, we know it contains only one electron. Thus, we know lithium’s valence configuration, 2s1, which in turn tells us that the 1s shell must have filled first. So we could “backfill” that subshell.
There are three rules for working out electron configurations using the periodic table.
Rule #1: To determine which electron shell is the valence shell, look to the period. Lithium lies in the second period. That indicates its valence shell is in the 2nd electron shell.
Rule #2: To determine subshells, look to the periodic table “block.” In Figure 13, take note of the color-coded sections on the periodic table. These are periodic table “blocks.”
Figure 13: The image shows the different subshell blocks on the periodic table.
image ©VisionlearningFor all elements in the s-block (Columns 1 & 2), their valence electrons occupy an s-orbital subshell. (That’s where lithium resides; thus, its valence orbital is 2s.)
For all elements in the p-block (Columns 13-18), their valence electrons occupy the p-orbital subshell, plus the s-orbital that filled before it.
For all elements in the d-block (Columns 3-12), their valence electrons wind up in a d-orbital subshell, plus the s-orbitals that filled before it.
Rule #3: To determine the number of electrons in each subshell, count how many spaces that the element lies from the start of its block.
Another example is the element beryllium (Be) – a rare metal that is used to make the hexagon mirrors of NASA’s $10 billion James Webb Space Telescope due to its resistance to temperature-related warp (Figure 14).
Figure 14: Dave Chaney inspects six of the James Webb Space Telescope’s gold-plated primary mirrors at NASA’s Marshall Space Flight Center in Huntsville, Ala. The mirror’s beryllium base resists warping in extreme temperatures.
image © CC BY 2.0 DEED NASA's James Webb Space TelescopeBeryllium lies to the right of lithium. It is the second element from the start of s-block. That placement indicates beryllium has two electrons in its 2s-subshell. Its configuration is shown below.
Be 1s2 2s2
If you're willing to work electron configurations out backward, starting from the valence shell, then the periodic table can guide you. Being such simple elements, lithium and beryllium don’t show the full value of this hack. Let’s use this technique to work out electron configurations for some of the elements of life – carbon, oxygen, and sulfur.
Carbon’s 6 electrons would partially fill its 2p orbitals, which we know based on carbon’s placement in the p-block of the 2nd period. Furthermore, since carbon is the second element from the beginning of p-block (start counting at boron, B), we predict that carbon has two electrons in its 2p orbital. That means carbon’s valence subshell looks like this:
2p2
The other subshells are easy. We know subshells 1s and 2s had to fill up before the 2p, so, we can just backfill those. That leaves us with carbon’s electron configuration:
C 1s22s22p2
Another example is oxygen. With its 8 electrons, oxygen would also fill through the 2p orbitals. But oxygen is four elements from the beginning of p-block, and so has four electrons in the 2p subshell. With a capacity of six electrons in the 2p subshell, oxygen is two electrons short of filling its orbital and tends to form two bonds:
O 1s22s22p4
Yet another example is sulfur, which has 16 electrons. According to Figure 6, the last electrons spill into the 3p orbitals:
S 1s22s22p63s23p4
That means that sulfur is also two electrons short of filling its valence shell, and so like oxygen, will tend to form two bonds when reacting.
Punto de Comprensión
Let’s return to the big picture. We’ve discussed details regarding individual electrons and how they fill electron shells. But don’t miss the developing pattern: each electron shell holds more electrons than those that came before because it contains an extra type of subshell. For example, the 1st electron shell contains only an s-type subshell. The 2nd electron shell contains an s-type and a p-type subshell. Based on this pattern, it would be logical to assume that the 3rd electron shell contains an s-type, a p-type, plus one more. And you would be right. It’s called d-type.
We won’t delve into these elements, however. Tricky things start happening as atoms get more complex. Electron configurations for elements with d-type subshells are beyond the scope of this module. However, we will use one d-block element (iron, Fe) to learn a shortcut.
A shortcut to long configurations
Take a look at the complete electron configuration for the element iron, or Fe.
Fe 1s22s22p63s23p64s23d6
With 26 electrons, iron’s electron configuration seems long. (Imagine writing an electron configuration for an element such as Seaborgium with 106 electrons!) Luckily, there’s a shortcut. An abbreviated notation for electron configuration replaces all core electrons (non-valence electrons) with the last noble gas in brackets.
For iron, that noble gas is argon (Ar). We know that noble gases’ valence shells are full, so their configurations are automatic. Argon's electron configuration is as follows:
Ar 1s22s22p63s23p6
Instead of writing out the full configuration for iron, we can replace the overlap with the noble gas. Note the overlap in iron and argon full notations:
Ar 1s22s22p63s23p6
Fe 1s22s22p63s23p64s23d6
We can replace that overlap with argon, or [Ar]. Doing so gives us Iron’s abbreviated notation, which looks like this:
Fe [Ar] 4s2 3d6
Abbreviated notations become almost crucial as we move down the periodic table, working with increasingly more complex elements.
Keep in mind also that exceptions to predicted electron configurations exist. Copper is one example. Its electron configuration would be expected to be as follows:
[Ar] 4s2 3d9
But instead, copper’s configuration is the following:
[Ar] 4s1 3d10
The reason for these exceptions is beyond the scope of this module. (Hint: It’s because some subshells are very close to one another in energy level.) What’s important however, is that you know such exceptions exist.
Conclusion
Chemical valency explains the everyday chemistry we depend on, from the common table salt that makes food enjoyable to medications that can save lives. Structure-based drug development’s first big success story was the development of anti-HIV drugs such as Crixivan. Thanks to the ability to create new molecules based on knowledge of chemical valency, HIV is no longer the death sentence it once was in many parts of the world.
In 2005, pharmaceutical companies released HIV drugs around the world. Today, 90 percent of diagnosed HIV patients (estimated to be 73% of all infected persons) are in viral repression – meaning the disease is essentially dormant in their bodies. After Crixivan, even healthier drugs were developed to treat HIV. Today about a dozen protease inhibitors are available to people. Levin says, "There was a progression made towards living your full life. That's really where we are now in HIV therapy."
This technology is made possible in part by the knowledge of chemical valency, which stems from electron configuration. Electron configuration bridges the world we cannot see, the sub-microscopic one, with the world can observe.
Tabla de Contenido
Active el resaltado de términos del glosario para identificar fácilmente los términos clave dentro del módulo. Una vez resaltados, puede hacer clic en estos términos para ver sus definiciones.
Active las anotaciones NGSS para identificar fácilmente los estándares NGSS dentro del módulo. Una vez resaltados, puede hacer clic en ellos para ver estos estándares.