Episode 15 starts with a quick recap of electronegativity and how it determines the type of bond formed (1:10).
Episode 15 starts with a quick recap of electronegativity and how it determines the type of bond formed (1:10). Then it takes a closer look at covalent bonds and electronegativity, introducing polar and nonpolar bonds (2:20). Diving even deeper into covalent bonds, we focus on the relationship between distance between nuclei and potential energy and define bond length (5:00). The episode also recaps how the size of bonded atoms and the bond order influence the bond length and bond energy (6:28). Taking a closer look at ionic bonds, we apply Coulomb’s law to determine the strength of interaction between cations and anions (7:27).
Question: What type of bond is formed when atoms with low electronegativity combine? (8:49)
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Hi and welcome to the APsolute Recap: Chemistry Edition. Today’s episode will recap chemical bonds.
Lets Zoom Out:
Unit 2 - Molecular and Ionic Compound Structure and Properties
Topic - 2.1, 2.2
Big idea - Structure and Properties
Introduction:
Electronegativity and Coulomb’s Law are two central concepts in Chemistry. Coulomb’s Law, frankly, can explain almost anything in Units 1 and 2 - and beyond. And electronegativity - using our “tug of war” analogy - is a key component of chemical bonding. We will therefore now leave Unit 1 where we focused on the individual atom and take a closer look at the combination of atoms - compounds. Today’s episode will recap chemical bonding and dive a bit deeper into the strength of chemical bonds.
Let’s zoom in:
We are starting by recapping a recap: Our 4th episode in which we already did an Introduction to Chemical Bonding. If you haven’t had a chance to listen to that episode, stop here and go back! …
Okay, okay, since you are still here, we will briefly summarizing it:
In episode 4 we recap electronegativity, which is, as defined in episodes 4 and 15, the ability of an atom to attract a shared pair of electrons. We also know that electronegativity is a periodic property: It increases across a period due to the increasing number of protons, which results in a stronger attractive force between electrons and the nucleus. Since we are increasing the distance between the nucleus and the valence electrons when going down a group, electronegativity decreases within a family. This means that non-metals have a rather high electronegativity and metals a comparable low electronegativity.
We’ve also heard that valence electrons can be gained, lost and shared. When electrons, due to a large difference in electronegativity, are lost by a metal and gained by a non-metal we are forming ions of opposite charge - and therefore a compound with ionic bonding.
Two non-metals are usually pretty similar in electronegativity and are sharing valence electrons. They are forming a covalent bond.
When the two atoms that are forming a covalent bond are the same element (or if they only have a very small difference in electronegativity - like carbon and hydrogen) - the covalent bond is nonpolar. Our participants in the tug of war are of equal or have similar strength, since the number of protons and the distance from the nucleus is similar. Therefore, the electrons are shared equally.
But what if one of the two nonmetals is stronger - not strong enough to win the tug of war, but pulling stronger than the other team? Then both of our electrons that make up the bond are no longer shared equally, but closer to the atom of one element than the other. Let’s look at an example: Carbon Tetrafluoride. The central carbon has four valence electrons, which each are shared with one fluorine. Since fluorine has more protons than carbon, it has a stronger coulombic attraction of the shared valence electrons. The bond between carbon and fluorine is therefore a polar covalent bond, with the fluorine being partially negative, since it attracts not only its own valence electron, but also the carbon’s valence electron. We therefore have a higher electron density closer to the fluorine. The greater the difference in electronegativity is, the more polar is the bond. Within our “bond spectrum” from non-polar covalent to ionic bond, polar bonds are in between and have covalent as well as ionic characteristics.
Why do we form bonds, like the covalent bond? The idea is that the potential energy is lower when two atoms are bonded - and lower energy is better. So let’s take a closer look at the formation of the covalent bond in relationship to potential energy. This relationship is often visualized in a graph showing the distance between the nuclei vs. the potential energy. The graph is a bit unusual, because when describing it, we will actually start on the right and go to the left, because we want to start where the atoms are too far from each other to experience any attraction.
Moving closer, the graph has three areas: First, the area in which the attractive forces dominated. The atoms are getting closer and the potential energy decreases - energy is being released. Second, a point at which there is a minimum potential energy. Here the attractive forces are in equilibrium with the repulsive forces between the two nuclei. This point is the equilibrium bond length - the distance between two nuclei at lowest potential energy. The energy at this distance is called bond energy. It is the energy required to separate the atoms again. If you keep decreasing the distance between the two atoms, the potential energy rises again. In this third area, the repulsive forces are stronger than the attractive forces.
There are two factors influencing the bond length: the size of the atom and the bond order. The larger the atoms that are sharing the electron, the greater is, of course, the distance between the center of the bonded atoms, and therefore, the greater is the bond length. Bond order is the number of bonds between two atoms - we can have single bonds, where the atoms each contribute 1 valence electron, but also double bonds with 2 valence electrons per atom and triple bonds. When increasing the bond order, the bond length decreases. That means, a single bond is shorter than a double bond, which is shorter than a triple bond. The bond order also influences the bond energy: it takes more energy to break a triple bond than to break a double bond than to break a single bond because we have stronger attractive forces between the two atoms. It’s like when you are holding something with two hands it takes your sibling more energy to steal it from you than if you would only hold it with one hand.
Ionic bonds also vary in strength and here we can use Coulomb’s Law again: The greater the charge on each ion, the larger are the attractive forces. Example: The attractive forces between Na+ and Cl- are weaker than between Mg2+ and S2- or even between Mg2+ and 2 Chlorine anions. We also know that distance - or ionic size matters: The larger the ions are the smaller are the forces of attractions. Therefore, the strength of attractive forces between the sodium cation and the halogen anion decreases from sodium fluoride, to sodium chloride, sodium bromide and sodium iodide.
To recap……
In today’s episode we did a recap of how electronegativity determines the chemical bond and took a closer look at ionic and covalent bonds. Depending on the difference in electronegativity, we can distinguish between polar and nonpolar covalent bonds. A graph is used to show the relationship between atomic distance and potential energy. The smaller the atom and the higher the bond order, the shorter is the bond length. A higher bond order leads to higher bond energy. The higher the charge on the ions and the smaller the ions, the stronger is the attraction between ions.
Coming up next on the Apsolute RecAP Chemistry Edition: VSEPR Theory
Today’s Question of the day is about chemical bonds
Question: What type of bond is formed when atoms with low electronegativity combine?