Primary Bond Types

Summary

In this article, Erik Dunmire elaborates on various types of bonding found in solid materials. He categorizes them into three main types: metallic, covalent, and ionic, while also introducing the concept of mixed bonding scenarios, such as in metalloids. Metallic bonding involves metals with loosely-held electrons contributing to their properties. Covalent bonding is between nonmetals that share electrons, often resulting in poor electrical conductivity. Ionic bonding involves electron transfer from metals to nonmetals, creating charged ions and often leading to poor conductivity. Dunmire also explains the spectrum between ionic and covalent, providing examples such as cesium fluoride and magnesium oxide. He introduces the primary bonding triangle, illustrating the relationship between electronegativity and bond type.

Highlights

• Dive into primary bonds: metallic, covalent, and ionic! 🎢
• Metals have delocalized electrons, lending them unique properties. ⚡
• Covalent bonds are all about sharing electrons and strong local bonds. 🤲
• Ionic bonds create charged interactions but impact material brittleness. 🌌
• Mixed bonding showcases the spectrum between ionic and covalent types. 🌐
• Use the primary bonding triangle to chart where bonds lie on the spectrum. 📈

Key Takeaways

• Explore the world of primary bonds: metallic, covalent, and ionic in solid materials! 🔍
• Metallic bonds feature a 'sea' of electrons, perfect for ductility and conductivity. 🌊
• Covalent bonds are all about sharing, which means poor conductivity but strong bonds. 🤝
• Ionic bonds involve electron transfers leading to charged atoms, impacting conductivity. ⚡
• Most materials exist on a spectrum, blending ionic and covalent characteristics. 🌈
• The primary bonding triangle helps identify a bond's position between ionic, covalent, and metallic. 🔺

Overview

In this insightful exploration, Erik Dunmire takes us through the fascinating realm of primary bonds in solid materials. He explains the three major types—metallic, covalent, and ionic—each with unique characteristics that dictate material properties. Metallic bonds, common in metals, involve a 'sea' of electrons leading to high conductivity and malleability. Covalent bonds between nonmetals involve electron sharing, resulting in strong localized bonds but usually poor conductivity. Ionic bonds feature electron transfer, creating charged ions affecting conductivity.

Dunmire delves deeper by illustrating how real-world materials often do not fit neatly into one of these categories. Instead, they exist on a spectrum of bonding types. Mixed bonding examples include metalloids that show characteristics of both metals and nonmetals. This blended bond nature, particularly in ceramics, often results in unique material properties not purely ionic or covalent.

Finally, he introduces the intriguing concept of the primary bonding triangle. This scientific model allows chemists to visualize how bonds span from fully ionic to completely covalent, depending on the atom's electronegativity and its difference. For instance, cesium fluoride stands at the ionic apex, while fluorine represents a covalent bond. This triangle helps predict how material characteristics may adjust through the shifting balance between these fundamental bond types.

Chapters

• 00:00 - 00:30: Introduction to Bonding in Solids In this chapter, the focus is on different types of bonding in solid materials. It delves into both primary and secondary bonding. Primary bonding includes metallic, covalent, and ionic bonds, particularly highlighting metallic bonds that occur among metal elements.
• 00:30 - 01:00: Primary Bonding Types In this chapter, the focus is on primary bonding types, specifically covalent, ionic, and mixed bonding. Covalent bonding occurs between nonmetal elements, while ionic bonding occurs when a metal pairs with a nonmetal. There are also scenarios that do not clearly fit into these types and may be classified as mixed bonding, such as the bonding in metalloids, which are elements that are neither purely metals nor nonmetals but have characteristics of both.
• 01:00 - 02:30: Metallic Bonding This chapter explores the concept of metallic bonding, explaining how metals have weak holds on their electrons due to low effective nuclear charges. Consequently, valence electrons in metals are loosely held and highly mobile, often described as a 'sea of electrons'.
• 02:30 - 04:30: Ionic Bonding This chapter discusses the concept of ionic bonding, focusing on the role of valence electrons that are donated in the process. It highlights the unique properties of materials resulting from this type of bonding, specifically noting that the bonding is nondirectional. Consequently, the crystal structures of metals are characterized by densely packed atoms, similar to packing oranges in a box, without specific angles between the atoms, due to the delocalized nature of the electrons.
• 04:30 - 05:30: Covalent Bonding The chapter on 'Covalent Bonding' explains the properties of metals related to the presence of free electrons among the atoms. This contributes to metals' excellent electrical and thermal conductivity. Furthermore, the ductility and malleability are attributed to the ability of atoms to slide past one another. Contrary to other materials, metals are not composed of molecules but are made up of enormous numbers of atoms.
• 05:30 - 09:30: Spectrum of Ionic to Covalent Bonding The chapter titled 'Spectrum of Ionic to Covalent Bonding' discusses the nature of different types of chemical bonds, focusing on ionic and covalent bonds. It begins by describing how atoms form structures in large numbers, leading to crystal formations. Ionic bonding is explained as a process involving a metal and a nonmetal, where the metal with a weak grip on its valence electron due to low effective nuclear charge interacts with a nonmetal which has a high effective nuclear charge, thereby attracting electrons away. This forms the basis for understanding how differing attractions between elements lead to various bonding types, ranging from fully ionic to covalent in nature.
• 09:30 - 13:00: Primary Bonding Triangle This chapter covers the concept of electron transfer between a metal and a nonmetal atom. When a metal atom loses an electron, it acquires a net positive charge, while the nonmetal atom gains an electron and obtains a net negative charge. This electron transfer results in the formation of charged spheres - one positively charged and the other negatively charged, illustrating the concept of ionic bonding.

Primary Bond Types Transcription

• 00:00 - 00:30 We can now take a closer<br>look at the different types of bonding that exist in solid materials, including both the three types of primary bonding that we've already mentioned, as well as secondary bonding that exists in molecular compounds. Again, with primary bonding, there are really three types; metallic, covalent, and ionic. Metallic whenever we have metal elements bonded together.
• 00:30 - 01:00 Covalent when we have nonmetal elements bonded together. And ionic whenever we have a metal paired up with a nonmetal. We can also imagine situations in which we don't neatly fall into one of those three categories but lie somewhere in between, something we might classify as mixed bonding. And one example scenario you might think of is those of the metalloids where the elements themselves are neither metals nor nonmetals but rather lie somewhere in between.
• 01:00 - 01:30 Let's take a look at metallic bonding. Again, remember here we have all atoms that have relatively weak grips on their electrons due to low effective nuclear charges. As a result, the valence electrons are held very loosely and are relatively mobile. We tend to think of metals as having a sea
• 01:30 - 02:00 of donated valence electrons. This gives a rise to the nature of bonding and properties in those materials. The bonding is nondirectional. And a a result, crystal structures of metals tend to involve simply packing atoms as densely as possible, like packing oranges in a box, without much regard to the angles between those different spheres. Because of the delocalized nature of the electrons
• 02:00 - 02:30 that are free to roam around among those atoms, helps to explain the good electrical and thermal conductivity that we see in metals, as well as the ductility and malleability of those solid materials since the atoms are relatively free to slide pass each other. It's important to emphasize that in metals there are no molecules. Rather, there are enormous numbers of atoms,
• 02:30 - 03:00 uncountable numbers of atoms packed together in regular arrangements, crystal structures. Ionic bonding on the other hand, again involves a metal and a nonmetal; the metal having a weak grip on its valence electron or low effective nuclear charge, and the nonmetal having a high effective nuclear charge or a strong attraction for electrons. As a result, the nonmetal attracts an electron away
• 03:00 - 03:30 from the metal atom, causing an electron transfer and net charges on each of the atoms. The metal, because it's lost an electron, now has a net positive charge. The nonmetal, because it's gained an electron, has a net negative charge. As a result, we can now think of these two atoms as simply being charged spheres with positive and negative charges.
• 03:30 - 04:00 This type of simplified model for bonding allows us to use mathematical models from physics for positive/negative attractions to describe this type of bonding. Again, the bonding is fairly nondirectional. We're just packing charged spheres together. But in this case, the bonding is very localized since the electrons in both atoms are held strongly. Remember that the electrons
• 04:00 - 04:30 in the metal atom are now only the core electrons since the valence electrons have been stripped away, and while those in the nonmetal atom are held strongly because of its effective nuclear charge. As a result, because the electrons are not free to move about, these materials tend to have poor electrical conductivity and poor thermal conductivity. They also tend to be quite brittle since the atoms cannot move pass each other easily.
• 04:30 - 05:00 As with metals, there are no molecules in the case of ionic materials, but rather enormous arrays of uncountable numbers of these ions packed together in some regular crystal arrangement. The third type of primary bonding, covalent bonding, which we see within the chains of polymers and also
• 05:00 - 05:30 within a number of semiconductors in ceramic materials, involves nonmetal atoms that are sharing electrons. Again in this case, both types of atoms have strong effective nuclear charges, and therefore strong grip on their valence electrons. As a result, neither donates the electron to the other, but rather they share those valence electrons in order to fill their outer shells.
• 05:30 - 06:00 If you remember, and which we'll review shortly, the bonding in these nonmetal systems, covalent bonding, is highly directional in its nature. Whenever we have three or more atoms, there are predictable angles between those bonds that are governed by hybridization theory. Angles of 180, 120, 109.5 degrees, and so on.
• 06:00 - 06:30 Because there are only nonmetal atoms present, the bonding is quite localized between those atoms. And so we see very poor electrical conductivity and poor thermal conductivity in most covalently-bonded materials. Unlike the other two types of primary bonding that we saw, it is possible in the case of covalent bonding to have molecules. That is, discreet packages of a small number of atoms
• 06:30 - 07:00 that are covalently-bonded together. But it's also possible in some cases to have large networks of enormous numbers of atoms that are all bonded together through these strong covalent bonds. Examples being things like silicon dioxide, diamond, silicon carbide, and so forth.
• 07:00 - 07:30 In reality, most substances fall somewhere on the spectrum between ionic and covalent bonding. At one extreme, we could have something like cesium fluoride that is almost purely ionic bonding, having an element with a very low electronegativity paired up with one with very high electronegativity. At the far end of the other spectrum, we might have something like a fluorine-fluorine bond
• 07:30 - 08:00 where both atoms in the bond pull equally on the electrons. But of course, in between we could have a wide range such as magnesium oxide, which we would generally think of as a ionic bond but not as ionic as cesium fluoride. We could have something like zinc sulfide that falls somewhere in the middle between ionic and covalent bonding.
• 08:00 - 08:30 And we could have a bond like carbon to chlorine that we certainly think of as a covalent bond but where the two atoms do not pull equally; the chlorine pulling slightly more strongly. And so we think of this as a polar-covalent bond. So in reality, there is not really ionic or covalent bonding, but rather a spectrum from ionic to covalent bonding. And most ceramic materials fall somewhere along this spectrum.
• 08:30 - 09:00 Chemists often try to quantify where in the spectrum we are by calculating a percent or a fraction ionic character for a bond using the electronegativity difference between the two atoms in the bond. For example, if we wanted to calculate the percent ionic character of a magnesium oxide bond, we could look
• 09:00 - 09:30 up the electronegativity for each of these elements. 1.2 for magnesium. 3.5 for oxygen. And we could calculate their electronegativity difference, often represented as delta x. And in this case, the electronegativity difference would be 2.3. As indicated in the equation above, we can then calculate the percent ionic character
• 09:30 - 10:00 in the bond by taking 1 minus the exponent of negative 0.25 times the electronegativity difference of 2.3 squared. And then of course, converting this to a percentage. So we then get a value of 0.73.
• 10:00 - 10:30 Or in other words, 73 percent ionic, 27 percent covalent. So primarily ionic bonding but with a slight covalent character. For an even more general look at all of the different types
• 10:30 - 11:00 of bonding and the mixed bonding characteristics they can have, material scientists sometimes talk in terms of a primary bonding triangle. Here, we can actually make a graph with average electronegativity on one axis, and difference in electronegativity on the other axis. And create a graph, if you will, using this triangle of ionic, covalent, and metallic bonding.
• 11:00 - 11:30 At one extreme, you could have cesium where in cesium metal there is of course no difference in electronegativity. All of the atoms are the same, so the difference in electronegativity is zero. The average electronegativity is the lowest of all possible elements at 0.7. And so this corner of the triangle corresponds to cesium.
• 11:30 - 12:00 At the covalent corner of the triangle, we have fluorine, F2, where again the difference in electronegativity is zero since the atoms are the same. And the average electronegativity is the maximum possible, 4.0, for fluorine. And then finally on the third corner of the triangle,
• 12:00 - 12:30 we have the compound cesium fluoride where the difference in electronegativity is the maximum possible, 3.3. And the average electronegativity falls exactly in between the two values shown. So 2.35. All other possible bonds between elements on the periodic table will fall somewhere within this triangle, and so we can get some sense of how close any particular bond is
• 12:30 - 13:00 to ionic, covalent, or metallic character, or where in between those extreme categories a particular bond lies.