Lesson Video: Metallic Bonding Chemistry

In this video, we will learn how to describe metallic bonding and the effect it has on the physical and chemical properties of metals.

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Video Transcript

Metallic Bonding

The buildings we live in, the vehicles we ride in, the device we’re watching this video on, they all harnessed the characteristics of metals in particular ways. But why do metals behave the way they do? Well, it has to do with the arrangement of their atoms and the electrons from those atoms; namely, that their valence electrons float freely in a sea of electrons. In this video, we will learn how to describe metallic bonding and the effect it has on the physical and chemical properties of metals.

Which elements are metals? Let’s take a look at the periodic table. We can divide the table into three sections: metals, nonmetals, and metalloids. Metals are the elements that tend to form positively charged ions. We find them in the left-hand portion of the periodic table. Nonmetals, which typically form negatively charged ions, are found on the right-hand side of the periodic table, while metalloids are in between the two other groups. In their elemental form, metals are shiny. They conduct heat and electricity. They are malleable, which means they can be formed into a specific shape, and they are ductile, meaning they can be pulled into a wire. Nonmetals tend to be the opposite. They are dull, poor conductors of heat and electricity, brittle, and nonductile.

Metalloids have some metallic properties and some nonmetallic properties. Metals include some familiar names, such as copper and iron, and they also include substances we may not think of as everyday metals, such as sodium and calcium. You may think to yourself, “Sodium is a metal! But table salt, sodium chloride, has sodium in it, and it’s a white crystal, not a shiny metal.” Well, to answer this question, we need to look at the different types of substances and compounds that can be formed by bonds between metal and nonmetal elements.

Before we look at metallic bonding, let’s recall a couple different types of commonly seen bonds and compare metallic bonding to these types. Ionic bonds occur between metal and nonmetal elements. An ionic bond typically occurs when a metal atom donates one or more electrons and a nonmetal atom receives one or more electrons. The result is two ions of opposite charge with a strong electrostatic attraction. While ionic compounds like sodium chloride may include metal elements, they’re not themselves metallic substances. Covalent bonds occur between atoms of nonmetal elements. In covalent bonds, the electrons are shared between the atoms rather than being donated completely to one side or the other. Simple examples of substances that contain covalent bonds include water and carbon dioxide.

Our substances of interest, metals, are made up of atoms of one or more metallic elements. Single-element examples include the elemental forms of iron, gold, and sodium. Metals that combine different types of metallic atoms include brass, a combination of copper and zinc, and bronze, a combination of copper and tin. We call the metallic substances made of two or more metal elements alloys.

The electrons and metallic substances flow freely in what we call a sea of delocalized electrons. Instead of the electrons being bound to a specific atom or shared between two atoms, each atom’s valence electrons move around freely in the space surrounding the atom. Let’s take a more detailed look at this arrangement of electrons. The particles in metallic bonds are arranged like this. The positive ions form a lattice, while the electrons flow freely in the remaining space. While the valence electrons move around freely, the other parts of the metal atoms, the positive ions, arrange themselves in repeated layers. The layers of different metals will have different geometric arrangements, like the square patterns of gallium or the hexagonal patterns of sodium that we see here.

The specific arrangement isn’t always important. But we should know that the positive metal ions made up of the positively charged nuclei and their inner electron shells arrange themselves in predictable repeated lattices. Within these lattices, the valence electrons or the electrons in the outermost electron shells move about freely. We typically describe this arrangement as a sea of delocalized electrons. Delocalized, meaning that the electrons don’t tend to belong to any one particular atom but are instead shared across the network.

In Ionic bonds, the electrons are gripped by the negative ions. In covalent bonds, each electron is shared between two particular atoms not shared by the whole network like in metallic bonding. These free-flowing electrons have interesting effects on the metal as a whole. As a reminder, in ionic bonds, the electrons are gripped by particular ions, while in covalent bonds, the electrons are held in place between particular pairs of atoms. So, the delocalized sea of electrons in metals gives the network of atoms flexibility and its constituent particles a lot of freedom of motion. Interestingly, the microscale freedom of motion of the particles translates into macroscale physical properties. Let’s take a look at how the arrangement of particles in metallic bonding leads to some of the physical properties of metallic substances.

First, metals are malleable. They can easily be formed into shape, like a bendy sheet of aluminum foil, rather than snap, like a brittle ionic compound such as a caped-together salt crystal. When a force is applied to a metal, the shape of the lattice can change, while the particles are still held together by an electrostatic attraction between the positive ions and the negative electrons. Thanks to the delocalized sea of electrons, the positive ions in the metal have more freedom of motion than in a brittle ionic compound. Metals are also ductile, meaning they can be pulled into a wire for similar reasons. When a tensile force or a pulling force is applied, the lattice can rearrange laterally to form the shape of a wire. Thanks again to the freedom of motion provided by the sea of delocalized electrons.

Metals conduct electricity as well. Electricity is the flow of charged particles. So when a metal is connected to a circuit, the free electrons will be drawn to the positive terminal. The flow of electrons through the metal, made easier by their loose arrangement, is what we call electricity. When there are no free charged particles, like in solid sodium chloride, electricity will not flow through the substance. But we can create free charged particles in the form of ions by dissolving the salt to create salt water, which will indeed conduct electricity.

Metals also conduct heat. Heat is the transfer of energy due to the movement and vibration of atoms. When a substance is heated, its particles begin to vibrate. For metals, their positive ions are more free to move back and forth, more readily transferring energy to nearby particles, eventually transferring the energy to more and more particles all the way across the substance. Think of a pan or a kettle on the stove as a simple example of a metal quickly conducting lots of heat.

Metals also have high melting points. The substance as a whole is held together by the electrostatic attraction between the positive ions and the negative valence electrons. Each ion is attracted to multiple surrounding electrons, and each electron is attracted to multiple surrounding ions. So, it requires a lot of energy to separate a positive ion and its valence electrons from the many neighboring particles of opposite charge that attract them. The more energy it takes to separate the particles, the higher the melting point will be. We can also look at how different characteristics of different metal atoms would affect the attractions in this held-together network, thereby raising or lowering the boiling point.

One thing that can affect the melting point of a metal is the number of valence electrons that each atom contributes to the sea of delocalized electrons. For example, sodium atoms contribute one electron to become Na+ ions, whereas magnesium atoms contribute two electrons to become Mg2+ ions. So, how does this difference in the number of valence electrons affect the melting point? Well, more valence electrons means there are more particles for the positive ions to be attracted to. This increases the strength of the metallic bond which leads to a higher melting point.

There is another pattern we can look at as we move down a group in the periodic table. If we compare sodium to the element directly below it on the periodic table, potassium, we can ask the question, “How might the presence of more electron shells affect the melting point of the metal?” It is worth noting that there are many exceptions to this pattern. But atoms with more electron shells have more separation between their nucleus and their valence electrons. In general, this means that atoms further down in a group on the periodic table have weaker metallic bonds and lower melting points. But again, there are many exceptions to this pattern. So, don’t be surprised if you see an element further down on the periodic table that ends up having a higher melting point. As you can see, the micro and macro properties of metals are intricately linked by the structure we call metallic bonding.

Now that we’ve reviewed metallic bonding and the physical properties of metals, let’s take a look at some practice problems.

Which of the following diagrams best represents metallic bonding?

The diagram that accurately represents metallic bonding will be the one that has a lattice of positive ions with a sea of delocalized electrons. Answer (B) is the correct answer, as it represents the red positive ions in the form of a lattice as well as the sea of delocalized electrons, the black dots floating in the space between.

Choice (C) has a lattice of positive ions as well, but there are negative ions mixed into this lattice. Choice (C) depicts ionic bonding. Choice (A) depicts hydrogen bonding that occurs between the hydrogen atoms and the oxygen atoms of water molecules. The diagram in choice (D) depicts network covalent bonding. In this specific case, it depicts the repeated arrangement of carbon atoms to form diamond. So, the diagram that best represents metallic bonding is choice (B).

Which of the following metals has the strongest metallic bonding? (A) Sodium, (B) lithium, (C) beryllium, (D) magnesium, or (E) aluminum.

This question is asking us to find the strongest metallic bond. The strongest bond is the one that has the strongest attractions between the particles. There are two key characteristics that can affect the strength of the attractions from metal to metal. First, the more valence electrons that a metal has, the stronger the bond. The extra electrons will add to the overall attractive forces acting on the positive ions. The other relationship that could come into play is that, in general, more electron shells means a weaker bond, although there are many exceptions to this pattern.

The separation between the nucleus and the valence electrons that the extra electron shells provide weakens the attractive forces between the positively charged and negatively charged particles. However, this second relationship is not as strong as the relationship between valence electrons and bond strength. So, let’s investigate that relationship first by finding the number of valence electrons of each of these metals. How do we find the number of valence electrons? Well, we can take a look at the periodic table.

As members of group one, sodium and lithium have one electron in their outermost electron shell or one valence electron. One column to the right, we find beryllium and magnesium in group two with two valence electrons. Aluminum, the third element in its row in group 13, has three valence electrons. Since it has the most valence electrons, aluminum, choice (E), is the correct answer. In the case of aluminum, each positive ion has more negative electrons around it, leading to more attractions and a stronger bond.

In this problem, we did not end up using the second relationship between the number of electron shells and the bond strength. If there were multiple metals that were tied for the highest number of valence electrons among the choices, we could use this second relationship as a sort of tiebreaker. However, since there are many exceptions to this pattern, it would be useful to confirm any assumptions about bond strength or melting points gleaned from this pattern by looking them up. Of the choices, the metal with the strongest metallic bonding is choice (E) aluminum.

Now that we’ve done some practice, let’s review the key points of metallic bonding. Metal elements are the elements that tend to form positive ions. They’re found on the left-hand side of the periodic table. Substances are considered metallic if they contain only atoms of metal elements. If a substance is made up of two or more metallic elements, we call that substance an alloy. The structure of metallic bonding involves a specific arrangement of the constituent parts. The positive ions form a lattice, while the electrons flow free in a sea of delocalized electrons. This structure gives rise to many physical properties of metals such as malleability, ductility, electrical conductivity, heat conductivity, and high melting points.

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