Imagine a piece of metal. Now imagine all of the atoms in that piece of metal. It’s easy to think of a bunch of spheres packed neatly next to each other, but how do they all stay together? What about molecules of water? Don’t molecules want friends? You may think they all bond the same way, but these chemicals do so in different ways. They stick together using different chemical bonds. We’ve talked about what everything is made of, but let’s talk about how it stays together.
Atoms with these bonds will be friends till the end. These are the strongest chemical bonds and unsurprisingly, they are responsible for holding most solids together. Some examples include salt (ionic), diamond (covalent), and aluminum (metallic).
1. Ionic Bonds
As the name suggests, these bonds form between ions. An ion is simply an atom with an electric charge. This happens when an atom loses or gains an electron. When an atom loses electrons, it becomes positively charged and is now a cation. When an atom gains electrons, it becomes negatively charged and is now an anion.
Atoms, like humans, like stability. We can imagine an atom having an outer shell of electrons, called a valence shell. Depending on how many electrons are in this shell, the atom will want it to be filled or emptied; whichever’s easier. For example, Fluorine on the right of the picture has seven valence electrons. It only needs one more electron to fill its valence shell and achieve stability. It’s easier to take one than to give seven, so it takes from whatever element is willing to give an electron. In this case, Lithium on the left gives Fluorine its one valence electron. Noble gases already have full valence electrons, so they don’t care about interacting with other lowly atoms.
How does this make them ionic? The Fluorine gains an electron and becomes an anion, and the Lithium atom loses one and becomes a cation. The bond itself is caused by the electrostatic force from the oppositely charged ions. Opposites attract, right? Note: they are not sharing the electron. Ionic is a give-and-take kind of bond. However, sharing is caring and that’s how many atoms like to work. These atoms prefer covalent bonds.
2. Covalent Bonds
For covalent bonds, the same principle about stability and valence electrons applies: atoms want stability, and to do that, their valence shells need to be filled or emptied. Covalent bonds involve atoms sharing electrons. “Co-” meaning “with” and “valent” having to do with the outer shell. In a way, the atoms share a valence shell.
Covalent bonds occur when two atoms have enough electrons to connect and fill each other’s valence shells. It also happens when one atom has a couple electrons to spare and it shares with an atom which needs some. It’s a win-win situation for the atoms! Covalent bonds do involve sharing, but another way to see how they bond is thinking of the atoms “clicking” together. The pictures below emphasize the sharing and the “clicking” idea, respectively.
3. Metallic Bonds
The name is not here to deceive you, these really do happen between metals. Just metals. This type of bonding includes a “sea of electrons.” The metal atoms come together and their electrons are not lost, given, nor taken away. They delocalize. This means that those atoms do not need to have their original electrons, but just need to have the same number of valence electrons as before. The electrons in metals “swim” in the block of metal, passing by metal atoms. As we can see in the picture, it is possible to think of the metal atoms as becoming ions and needing electrons in the “sea” to neutralize their charges. Since conductivity is understood in terms of electron movement, it makes sense that metallically bonded materials have high conductivities.
Primary Bond Properties:
1. Ionic Bonds:
- Crystalline solids
- Hard, brittle as a solid
- High melting/boiling points
- Poor electrical/heat conductors in solid state
- Good electric conductor in liquid state or in a solution
- Soluble in water (not in nonpolar solvents)
2. Covalent Bonds:
- Gases, liquids, solids (molecular, softer)
- Low melting/boiling points
- Electrical/heat insulators
- Insoluble in water (most are soluble in nonpolar solvents)
3. Metallic Bonds:
- Malleable solids
- High melting/boiling points
- Electrical/heat conductors
- Insoluble in water and nonpolar solvents
Atoms with these bonds are like friends you meet at a party. After the party, you pretty much forget they existed. Though these bonds are very weak, they are very important: especially for geckos.
1. Van der Waals
This is sort of an umbrella term for all of the intermolecular forces. There are two types: dispersion forces and dipole-dipole interactions.
These are also called London Dispersion Forces, named after Fritz London, who postulated their existence. Electrons move around in materials constantly. The charge of the electrons causes the charges of the molecules to fluctuate. “Dipole” just means that something has two poles. In this case, a molecule has two areas of opposite charge. So, momentarily, these molecules will have oppositely charged ends, and thus be attracted to each other. Since these dipoles are caused by the motions of electrons, they are called induced dipole moments. Note: all molecules and materials experience these forces. The strength of these forces depends on the size and shape of the molecules, but we won’t get into that here.
Dipole-Dipole Interactions (DDIs)
What’s the difference between these and dispersion forces? To start, these forces are between molecules which are already polar. They already have dipoles, they are not induced. However, these DDIs occur simultaneously with dispersion forces, not separately. Since these molecules are electrically polar, they interact with other polar molecules. Oppositely charged ends attract; like charges repel.
To summarize: dispersion forces involve fluctuating/induced dipoles, DDIs involve permanent dipoles. Both involve interactions between molecules.
2. Hydrogen Bonds
Hydrogen bonds occurs in any chemical with hydrogen atoms. Typically, hydrogen bonding is not seen as its own type of bond, but for completion’s sake, I will include it in this post. Hydrogen bonds are stronger than van der Waals forces, but are still too weak to be called primary bonds. Essentially, Hydrogen bonds are glorified DDIs. If the ends of molecules include hydrogen atoms, those ends are more positively charged. These ends bond with the negative ends of other molecules. Water, shown above, is a wonderful example of hydrogen bonding. The positively charged hydrogen ends bond with the negatively charged oxygen ends of the water molecules.
- Hydrogen bonds increase viscosity, surface tension, and boiling point (significantly).
- Van der Waals interactions are linked to slightly higher boiling points.
Too Long; Didn’t Read
- Ionic- give and take
- Covalent- sharing is caring
- Metallic- sea of electrons
- Van der Waals
- Dispersion forces- fluctuating/induced dipoles
- Dipole-dipole interactions- permanent dipoles
- Hydrogen Bonds- dipoles between hydrogen and other dipoles
- Ionic- brittle, solid insulator, liquid conductor, high melting/boiling points, water soluble
- Covalent- softer, low melting/boiling points, polarly soluble, insulators
- Metallic- malleable, high melting/boiling points, insoluble, conductors
- Van der Waals- slightly raises boiling points, bonds molecules of gases and liquids (more significantly than in solids)
- Hydrogen- dramatically higher melting/boiling points, higher viscosity and surface tension in liquids
Ionic Bond Diagram: http://www.daviddarling.info/encyclopedia/I/ionic_bond.html
Metallic Bond Diagram: http://www.chemguide.co.uk/atoms/bonding/metallic.html
Polar Water Molecule: https://en.wikipedia.org/wiki/Chemical_polarity
Hydrogen Bonds of Water: http://www.physicsofmatter.com/NotTheBook/Talks/Ice/Image7.gif