HS Chemistry - Chemical Bonding

Metallic Bonding

Overview of The Page

This page will cover:

• How does metallic bonding work?
• What are some properties of metallic compounds?
• How are these properties different in Transition Metals?
• What are alloys?

There are three types of bonding between atoms: ionic bonding, covalent bonding, and metallic bonding. This page looks at metallic bonding.

Metallic bonds are formed between metal atoms. Metallic bonding is different from ionic and covalent bonding in that all the valence electrons in a metallic bond are free-floating among the metal atoms $they are **delocalized**$. They are not connected to any specific atom. They freely travel across all the atoms connected by the metallic bonds. Think of it like the metal atoms being in a sea of electrons. The atoms are held together by their attraction to the valence electrons floating around.

The metal atoms are still atoms, not ions. They haven't lost their valence electrons. Their valence electrons are still there, just free-floating and not always connected to them. But why do the electrons get delocalized in the first place?

Metal atoms are very, very tightly packed together in metallic bonding, so the valence shell of each atom intersects the valence shells of other atoms. This allows the valence electrons to easily move from one atom's valence shell to another atom's valence shell. Additionally, since metals have very few valence electrons $less than 4$, they have space to accommodate quite a few more electrons in their valence shell. These two things, unique to metallic bonding, are what allow the valence electrons to be delocalized.

This pattern of bonding is responsible for some of metal's properties.

• Since the valence electrons are free-floating, metals can easily conduct electricity - the electrons can freely move across the entire structure for conduction.

• The metal atoms are also all attracted to the valence electrons, and this, combined with their close proximity to one another in the metallic structure, gives metals their strength. They don't break easily because their bonds are not in a fixed direction. Their shape can be changed, and the bonds will still hold.

• Their tight structure also gives them high melting and boiling points, as the atoms are very tightly packed together, and it will take a lot of energy to separate them enough to change the state of matter.

• This tight structure also makes metals good conductors of heat. Heat is produced by the movement of atoms. If an atom moves or vibrates at one end of the metal, the movement/vibration will quickly spread throughout the metal, as the atoms are very tightly packed together, and will therefore quickly pick up and pass on any nearby vibrations. Thus, since heat easily and quickly spreads within metals, metals are good conductors of heat.

• However, this same structure also makes metals very malleable. Since they are held together by the atoms' attraction to the free-floating electrons, one layer of atoms can easily slide over another layer of atoms without breaking bonds. They're still connected to the same thing - the free-floating electrons. But now that one layer of atoms has slid over another layer, the shape of the metal has changed.

Two other properties of metals that should be noted are that:

1. Metals are lustrous $shiny$.
2. Metals are insoluble in water.

Alloys

Everything that's been said on this subpage so far is in regard to pure metals, metals made entirely of one element $e.g. an Iron sheet, a Silver sheet, a Gold brick, etc.$. Alloys are different. An alloy is a mixture of chemical elements, in which at least one is a metal. Alloys retain most of the characteristics of a metal, but aren't malleable. Alloys are often made of atoms of different sizes. The combination of atoms of different sizes disrupts the metal's regular structure. Since some of the atoms are larger than the other atoms, the layers of atoms can't easily slide over them without breaking bonds. This prevents the atoms from sliding over, making the material less malleable.

In the pure metal, we can see that when given a push, one layer of metal atoms easily slid over another layer, shifting $in this drawing$ two atoms to the right. In the alloy, however, when the top layer of metal atoms is given a push, they only slide one atom to the right before stopping. The larger atom blocks them - they can't cross it without disconnecting from each other and breaking the bonds between them. Thus, they don't shift as much, and the metal is less malleable.

A common example of an alloy is steel, which is mostly iron but contains small amounts of carbon and other materials.

Transition Metals

Transition Metals can also form metallic bond structures. For example, Iron $Fe$ is a Transition Metal, and it's one of the best known metals on the planet. When Transition Metals form metallic bond structures, they are almost identical to ordinary metals, but with one difference: when Transition Metals form metallic bond structures, the d-subshell electrons also become delocalized, rather than just the s- and/or p-subshell electrons. For example, Gold $Ag$ would have both the electrons in its 5s subshell and in its 4d subshells delocalized. This changes some of their characteristics, such as giving Transition Metals a higher melting point than other metals.

Practice

Why is Iron (Fe) a better conductor of electricity than Silicon (Si)? Because an Iron (Fe) atom is larger in size than an Silicon (Si) atom, and larger atoms conduct electricity better Because Iron (Fe) is a Transition Metal, and therefore more electrons are delocalized Because Iron (Fe) has 4 electron shells while Silicon (Si) has 3 electron shells, so Iron (Fe) has more electrons

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