Atomic Hook-Ups – Types of Chemical Bonds: Crash Course Chemistry #22
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Atomic Hook-Ups – Types of Chemical Bonds: Crash Course Chemistry #22

Humans, like chemicals, are really all about
bonds. Think about all the relationships in your
life. You’re a casual acquaintance to some people,
a colleague or friend to others, and maybe more to that someone special. Maybe you’re dating someone casually, or you’re
in a committed relationship, or you’re married. There are all kinds of different combinations
of people out there. And sometimes, you know, people fall for a
vampire or a werewolf. Who am I to judge? Fact is, each type of relationship requires
different things from you and the other person, but if you play your cards right, these relationships
allow you to relax and escape the stresses that come with the
constant search for affection. Distance is important in relationships too,
of course; too much distance makes it hard to stay focused on each other and requires a lot of effort to keep things together, and I may not have to tell you, too little
distance can be a problem as well. Everyone needs their space, and when you don’t
have any, you just end up pushing away from whatever’s
crowding you. In this way, atoms are a lot like us. We call their relationships bonds, just like
we do with our own relationships. And there are many different types. Each kind of atomic relationship requires
a different type of energy, but they all do best when they settle into
the lowest-stress situation possible. The nature of the bond between atoms is related
to the distance between them, and, like people (and vampires and werewolves, I suppose),
it also depends on how positive or negative they are. The difference is that, unlike human relationships, we can analyze exactly what makes different
kinds of chemical relationships work. And that’s what this episode is all about. But, people, please remember that we here at Crash Course do not dispense relationship advice. [Theme Music] First things first, why do atoms do this at
all? Well, like everything else in the universe, atoms do whatever they can to reduce their overall energy, and they reach their lowest energy by achieving a balance between attractive and repulsive forces, being neither too clingy nor too aloof. So when two atoms approach each other, the electrons of each are attracted to the protons of the other. This is the electrostatic force. Like charges repel, opposites attract, like
in real life, or at least Paula Abdul songs. I know, I’m old. So when one atom is attracted to another, just like Edward Cullen and Bella in chemistry
class, to use a slightly more timely reference, it gets stressed out by the attractive force and tries to relieve the stress by getting closer. We’ve all been there, right? That hot, nerdy
vampire girl in your chemistry class? It’s just, it’s intense. The pull is so strong that the stress level or energy rises when the two are separated, so they stay close. But sometimes, they can get a little too close. When that happens, the nuclei repel each other
because of their like charges, and the energy between them rapidly increases
and they both back off, just enough to find that perfect little distance
between them, and everyone relaxes. This ideal, wonderful distance is the bond
length. It’s the distance between two nuclei at the
point of minimum energy. In other words, where the attractive and repulsive
forces cancel each other out. The distance at which these two atoms of chlorine
reach their minimum energy, caught between the attraction of the electrons,
the nucleus, and the protons repelling the nuclei, is the
bond length. That energy minimum, which we know absolutely
is –239 kilojoules per mole (kJ/mol), occurs when the distance between the atoms
is 0.00199 nanometers (nm). That distance is the bond length of Cl2, chlorine
gas. Now because the electrons are attracted to
both nuclei in the molecule, they actually spend the majority of their
time in the space between them. This is often described as sharing electrons,
and we call this kind of bond a covalent bond. But not all sharing is equal.
I should know: I have an older brother. The strength with which an atom holds shared
electrons is called its electronegativity. The electronegativities of various elements are all super well known and waiting for you in tables on the Internet. If two atoms in a bond have very different
electronegativities, like, say,
hydrogen at 2.1 and oxygen at 3.5, the electrons are more attracted to the atom
with the higher electronegativity. The difference is so great that the electrons spend most of their time around the stronger atom and much less time around the other one. Like how all the neighborhood kids wanted
to hang out with John, my older brother, because he was more charismatic. When the electrons hang out closer to one
side of the bond, it creates a slight negative charge in that area and a slight positive charge around the other atom. This separation of charges is called polarity, and it’s the polarity of the molecule that these atoms form, H2O, that makes water the most important molecule
on Earth. Covalent bonds like this, where electrons
are attracted to one atom more than the other, causing a separation of charges, are called
polar covalent bonds. But when a covalent bond forms between two
identical atoms, like the two chlorine atoms in our graph earlier,
the electrons are distributed evenly. We call this a non-polar covalent bond. But you’ve also gotta consider the middle
option, where atoms aren’t identical, but have very similar electronegativities, like hydrogen, with an electronegativity of 2.1, and sulfur, at 2.5. The difference here is so tiny that the electrons
are pretty much still evenly distributed, and we call that a non-polar covalent bond
as well. There’s a huge world of important chemicals
that have these kinds of bonds. So many, in fact, that we will dedicate a
couple of separate episodes to them. Covalent bonds tend to form from non-metals
and sometimes metalloids, those elements that have both metallic and
non-metallic characteristics. That’s because most of them hold their electrons
so tightly that they’re more likely to share them with
another atom than to gain or lose them altogether. Metals, on the other hand, have loosely-held
outer electron shells, so they’re constantly dropping electrons and
becoming positive ions. And when positive ions come across negative ions, like those formed from halogens, for instance, you have to know what’s gonna happen. They are attracted to each other, which means
energy is required to keep them apart, which means that they’re gonna bond if they can, creating that oh-so-wonderful point of minimum energy. This type of bond is unsurprisingly called an ionic bond, a bond formed between a positive ion and a negative ion. Because the ions are formed when one atom
loses electrons and the other gains them, we often say that an ionic bond is formed by the transfer of electrons from one atom to another. And we can calculate the amount of energy
that exists in a bond between ions at a given distance using a formula called Coulomb’s
law. Note that this only works for ionic bonds because the calculation requires the charges of the ions, which covalent bonds don’t have. Coulomb’s law says that the energy been two
ions equals the product of the two charges, which are represented by capital Qs, because
why not, divided by the distance, or radius, between
the two nuclei, all multiplied by a constant, 2.31 × 10–19
joules per nanometer (J•nm). Of course, the radius also has to be expressed in nanometers — you gotta make the units match. Let’s see how it works with something simple:
sodium chloride, or table salt. We know that the normal charge on a sodium ion is +1 and the normal charge on chloride is -1. These are Q1 and Q2. The length of a stable NaCl bond is 0.276
nm, so we put that in for the radius, and finally a quick calculation tells us that
the bond contains –8.37 × 10–19 J of energy. Remember, that negative number represents a decrease in the energy of the system due to an attractive force, which certainly makes sense here. Sodium and chloride ions are strongly attracted
to each other due to their opposite charges. Of course, you may have noticed that –8.37
× 10–19 J looks like a tiny, tiny number, but keep in mind we’re talking about one single
pair of ions. The –239 kJ that we got for chlorine? That
was for a whole mole of molecules. When multiplied by the 10^17 or so ions in
a single grain of table salt and then by the thousands of grains of salt
in a mole, the energy becomes much more significant. The NaCl bond is, in fact, quite strong. And because they are formed by a positive
ion and a negative ion, two charges completely separated between two
different particles, ionic bonds are extremely polar, way more
polar than polar covalent bonds. And so those are our three types of bonds: non-polar covalent, formed by the equal or
nearly equal sharing of two electrons between non-metal or metalloid atoms; polar covalent, formed by the uneven sharing
of electrons between two non-metals or metalloids; and ionic, formed by the transfer of electrons
from a metal to a non-metal. It’s important to remember, though, that there aren’t only three designations for chemical bonds. Just like human relationships, bonds don’t
always have really well-defined boundaries. Everything is a continuum. Labels are useful,
but they can only take us so far. There are, however, certain properties that each kind of bond tends to have that you should know. For instance, ionic compounds are often crystalline
in their solid form because of the way the ions pack together, like salt is. They’re generally soluble in water because
the two ions interact separately with the positively and negatively charged areas on
a water molecule. And once they’re separated or dissolved, the
ions allow the solution to conduct electricity. Covalent compounds, on the other hand, tend to be softer solids, liquids, or gases like Cl2 is at room temperature. They’re often not soluble in water, and even when they are, the solutions don’t conduct electricity. The differences in these properties stem mostly
from the differences in their polarities. So yeah, polarity is crazy important. So important that we’ll be doing a whole episode on it soon. Until then, I want to thank you for the bond
that you have to Crash Course Chemistry, whether it’s casual observer, faithful viewer,
or committed subscriber. Today, if you were paying attention, you learned that chemical bonds form in order to minimize the energy between two atoms or ions. You’ve also learned that the chemical bonds
may be covalent if the atoms share electrons, and that covalent bonds can share those electrons
evenly or unevenly. Bonds can also be ionic if the electrons are
transferred, and you learned how to calculate the energy transferred in an ionic bond using Coulomb’s law. This episode of Crash Course Chemistry was written by Edi González and edited by Blake de Pastino and myself. Our chemistry consultant is Dr. Heiko Langner. It was filmed, edited, and directed by Nicholas
Jenkins. Our script supervisor is Michael Aranda. He is also our sound designer, and our graphics
team, as always, is Thought Café.

100 thoughts on “Atomic Hook-Ups – Types of Chemical Bonds: Crash Course Chemistry #22

  1. Seriously who was in charge of this lighting? Hank looks like he as black demon eyes half the time. They’re so distracting I can barely focus on what he’s saying 🙄

  2. I was using this video in my Chemistry class to help my students visualize the bonds. I had set up the link in my lecture slides a few month ago. BUT I ended up watching the piece that Sunday Morning did on John from 2014 on YouTube the night before actually giving the lecture. Cue in 'just like the neighborhood kids…. part' – I laughed out loud. Sorry Hank…. Not good when your class sees their professor (me) start to laugh in the middle of a teaching spot. LOL. Don't forget to be awesome.

  3. I have question: First they tell us that atoms cannot exist if they don't have their outer shells complete. Then they say that in metallic bonding many electrons are free? And in the model of graphite one electron is free from each carbon atom then how does it exists?

  4. If you don't get all the info because of the speed, then go back and listen that part again. I can understand everything and english is not my first language. His work is amazing

  5. I think it needs a bit of structure and more attention to important definitions and properties. Tons of sentences on off-topic and one quick sentence defining covalent bonds, for instance. It makes it a fun video but are the learning objectives really being achieved that way? Or does it really pursue any?

  6. can u explain double bonds, triple bonds, the thick lines (as in cocaine), those wierd hexagonal structures that are shown in 2D structures?

  7. I dare you to watch this in 0.25 speed and not laugh. If you laugh, you have to restart the video.

  8. Seriously. I can't even watch an educational video without some jackass trying to promote homosexuality. Idiot.

  9. Why cracks don't stick back?
    Why there are grains of salt and not the salt rock? Why don't they all stick together?
    Why atoms stick mostly at some specific temperatures?

  10. Thank you so much; this is a lifesaver for my molecular bio class; just wish you had addressed hydrogen bonds

  11. OOOOH! THAT's why we compress fuel to make it explode! If it's "too close together", it becomes much more energetic! NOW it makes sense! So like the electrons around the nuclei and their spinning motion around the nuclei are just like gears! And each different atom's atomic number is like the number of teeth for their gear, so a Carbon "gear" has 12 teeth and spins around, interconnecting easily and at same speed as other carbon atoms, and if they combine with other atoms, which have a different number of teeth in the "gear", then they need to have a compatible number of teeth or else their speed or something will grind together and "not work".

    And also, each gear not only spins in the direction of its disc, but the disc also turns (spins) in a different axis, making like a 3D spin? And maybe there are some atoms whose "gear" spin in a different pattern? Maybe like some are disc-like and other are more helical-like?

    Well thank you for giving me a more "what happens if I do this?" kind of explanation! That's what I need! YES!

  12. Visualizing…………………… omg particles movment isnt random at all! Its just hypersensitive to gravity cuz theres electric balls spinning around each other and theyre sharing those and all wobbling and fitting into their space any way they can. It can be simulated. It is not random. Does that matter at all? lol could help create new procedures for examining particles. Its weird that people would want to describe that with the precision of numbers.

  13. I did not understand what we where going to have on my Science test and askes my friend and he showen me this. Of course I didn’t react…

  14. I kept wondering where I knew this guy from… lol turns out to be a celebrity i'm watching to pass my nursing test Monday lmao and I was already nervous..

  15. i have a revision exam for my chem final coming up and hell yeah for videos that explain the stuff i barely understood!!!!

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