Chemical bonds hold molecules together and create temporary connections that are essential to life. Types of chemical bonds including covalent, ionic, and hydrogen bonds and London dispersion forces.
Living things are made up of atoms, but in most cases, those atoms aren’t just floating around individually. Instead, they’re usually interacting with other atoms (or groups of atoms).
For instance, atoms might be connected by strong bonds and organized into molecules or crystals. Or they might form temporary, weak bonds with other atoms that they bump into or brush up against. Both the strong bonds that hold molecules together and the weaker bonds that create temporary connections are essential to the chemistry of our bodies, and to the existence of life itself.
Why form chemical bonds? The basic answer is that atoms are trying to reach the most stable (lowest-energy) state that they can. Many atoms become stable when their valence shell is filled with electrons or when they satisfy the octet rule (by having eight valence electrons). If atoms don’t have this arrangement, they’ll “want” to reach it by gaining, losing, or sharing electrons via bonds.
Ions and ionic bonds
Some atoms become more stable by gaining or losing an entire electron (or several electrons). When they do so, atoms form ions, or charged particles. Electron gain or loss can give an atom a filled outermost electron shell and make it energetically more stable.
Ions come in two types. Cations are positive ions formed by losing electrons. For instance, a sodium atom loses an electron to become a sodium cation, . Negative ions are formed by electron gain and are called anions. Anions are named using the ending “-ide”: for example, the anion of chlorine () is called chloride.
When one atom loses an electron and another atom gains that electron, the process is called electron transfer. Sodium and chlorine atoms provide a good example of electron transfer.
Sodium (Na) only has one electron in its outer electron shell, so it is easier (more energetically favorable) for sodium to donate that one electron than to find seven more electrons to fill the outer shell. Because of this, sodium tends to lose its one electron, forming Na.
Chlorine (Cl), on the other hand, has seven electrons in its outer shell. In this case, it is easier for chlorine to gain one electron than to lose seven, so it tends to take on an electron and become Cl.
When sodium and chlorine are combined, sodium will donate its one electron to empty its shell, and chlorine will accept that electron to fill its shell. Both ions now satisfy the octet rule and have complete outermost shells. Because the number of electrons is no longer equal to the number of protons, each atom is now an ion and has a +1 (Na) or –1 (Cl) charge.
In general, the loss of an electron by one atom and gain of an electron by another atom must happen at the same time: in order for a sodium atom to lose an electron, it needs to have a suitable recipient like a chlorine atom.
Making an ionic bond
Ionic bonds are bonds formed between ions with opposite charges. For instance, positively charged sodium ions and negatively charged chloride ions attract each other to make sodium chloride, or table salt. Table salt, like many ionic compounds, doesn't consist of just one sodium and one chloride ion; instead, it contains many ions arranged in a repeating, predictable 3D pattern (a crystal).
Certain ions are referred to in physiology as electrolytes (including sodium, potassium, and calcium). These ions are necessary for nerve impulse conduction, muscle contractions and water balance. Many sports drinks and dietary supplements provide these ions to replace those lost from the body via sweating during exercise.
Another way atoms can become more stable is by sharing electrons (rather than fully gaining or losing them), thus forming covalent bonds. Covalent bonds are more common than ionic bonds in the molecules of living organisms.
For instance, covalent bonds are key to the structure of carbon-based organic molecules like our DNA and proteins. Covalent bonds are also found in smaller inorganic molecules, such as , , and . One, two, or three pairs of electrons may be shared between atoms, resulting in single, double, or triple bonds, respectively. The more electrons that are shared between two atoms, the stronger their bond will be.
As an example of covalent bonding, let’s look at water. A single water molecule, , consists of two hydrogen atoms bonded to one oxygen atom. Each hydrogen shares an electron with oxygen, and oxygen shares one of its electrons with each hydrogen:
The shared electrons split their time between the valence shells of the hydrogen and oxygen atoms, giving each atom something resembling a complete valence shell (two electrons for H, eight for O). This makes a water molecule much more stable than its component atoms would have been on their own.
Polar covalent bonds
There are two basic types of covalent bonds: polar and nonpolar. In a polar covalent bond, the electrons are unequally shared by the atoms and spend more time close to one atom than the other. Because of the unequal distribution of electrons between the atoms of different elements, slightly positive (δ+) and slightly negative (δ–) charges develop in different parts of the molecule.
In a water molecule (above), the bond connecting the oxygen to each hydrogen is a polar bond. Oxygen is a much more electronegative atom than hydrogen, meaning that it attracts shared electrons more strongly, so the oxygen of water bears a partial negative charge (has high electron density), while the hydrogens bear partial positive charges (have low electron density).
In general, the relative electronegativities of the two atoms in a bond – that is, their tendencies to "hog" shared electrons – will determine whether a covalent bond is polar or nonpolar. Whenever one element is significantly more electronegative than the other, the bond between them will be polar, meaning that one end of it will have a slight positive charge and the other a slight negative charge.
Nonpolar covalent bonds
Nonpolar covalent bonds form between two atoms of the same element, or between atoms of different elements that share electrons more or less equally. For example, molecular oxygen () is nonpolar because the electrons are equally shared between the two oxygen atoms.
Another example of a nonpolar covalent bond is found in methane (). Carbon has four electrons in its outermost shell and needs four more to achieve a stable octet. It gets these by sharing electrons with four hydrogen atoms, each of which provides a single electron. Reciprocally, the hydrogen atoms each need one additional electron to fill their outermost shell, which they receive in the form of shared electrons from carbon. Although carbon and hydrogen do not have exactly the same electronegativity, they are quite similar, so carbon-hydrogen bonds are considered nonpolar.
Hydrogen bonds and London dispersion forces
Covalent and ionic bonds are both typically considered strong bonds. However, other kinds of more temporary bonds can also form between atoms or molecules. Two types of weak bonds often seen in biology are hydrogen bonds and London dispersion forces.
Not to be overly dramatic, but without these two types of bonds, life as we know it would not exist! For instance, hydrogen bonds provide many of the life-sustaining properties of water and stabilize the structures of proteins and DNA, both key ingredients of cells.
In a polar covalent bond containing hydrogen (e.g., an O-H bond in a water molecule), the hydrogen will have a slight positive charge because the bond electrons are pulled more strongly toward the other element. Because of this slight positive charge, the hydrogen will be attracted to any neighboring negative charges. This interaction is called a hydrogen bond.
Hydrogen bonds are common, and water molecules in particular form lots of them. Individual hydrogen bonds are weak and easily broken, but many hydrogen bonds together can be very strong.
London dispersion forces
Like hydrogen bonds, London dispersion forces are weak attractions between molecules. However, unlike hydrogen bonds, they can occur between atoms or molecules of any kind, and they depend on temporary imbalances in electron distribution.
How does that work? Because electrons are in constant motion, there will be some moments when the electrons of an atom or molecule are clustered together, creating a partial negative charge in one part of the molecule (and a partial positive charge in another). If a molecule with this kind of charge imbalance is very close to another molecule, it can cause a similar charge redistribution in the second molecule, and the temporary positive and negative charges of the two molecules will attract each other.
Hydrogen bonds and London dispersion forces are both examples of van der Waals forces, a general term for intermolecular interactions that do not involve covalent bonds or ions. Some textbooks use the term "van der Waals forces" to refer only to London dispersion forces, so make sure you know what definition your textbook or teacher is using.
How does that work in a cell?
Both strong and weak bonds play key roles in the chemistry of our cells and bodies. For instance, strong covalent bonds hold together the chemical building blocks that make up a strand of DNA. However, weaker hydrogen bonds hold together the two strands of the DNA double helix. These weak bonds keep the DNA stable, but also allow it to be opened up for copying and use by the cell.
More generally, bonds between ions, water molecules, and polar molecules are constantly forming and breaking in the watery environment of a cell. In this setting, molecules of different types can and will interact with each other via weak, charge-based attractions. For instance, a Na ion might interact with a water molecule in one moment, and with the negatively charged part of a protein in the next moment.
What's really amazing is to think that billions of these chemical bond interactions—strong and weak, stable and temporary—are going on in our bodies right now, holding us together and keeping us ticking!
Want to join the conversation?
- In the third paragraph under "Ionic Bonds", it says that there is no such thing as a single NaCl molecule. Why can't you have a single molecule of NaCl?(79 votes)
- This is because sodium chloride ionic compounds form a gigantic lattice structure due to the electrostatic attractions between the individual ions. In this case, each sodium ion is surrounded by 4 chloride ions and each chloride ion is surrounded by 4 sodium ions and so on and so on, so that the result is a massive crystal. This particular ratio of Na ions to Cl ions is due to the ratio of electrons interchanged between the 2 atoms.(118 votes)
- Regarding London dispersion forces, shouldn't a "dispersion" force be causing molecules to disperse, not attract?(35 votes)
- What is the typical period of time a London dispersion force will last between two molecules?
To me it seems it last far shorter than a second. If so, has it got time to do anything productive with the molecules? Do this London dispersion forces come in bursts? Electrons rapidly moving back and forth between two atoms in a molecule creating many London dispersion forces so fast it acts like a long-lasting bond between molecules?(16 votes)
- The London dispersion forces occur so often and for little of a time period so they do make somewhat of a difference. In my biology book they said an example of van der Waals interactions is the ability for a gecko to walk up a wall. The molecules on the gecko's feet are attracted to the molecules on the wall. You could think of it as a balloon that sticks to a wall after you rub if on your head due to the transfer of electrons. The bond is not long-lasting however since it is easy to break.(30 votes)
- If enough energy is applied to mollecular bonds, they break (as demonstrated in the video discussing heat changing liquids to gasses). Usually, do intermolecular or intramolecular bonds break first? Is there ever an instance where both the intermolecular bonds and intramolecular bonds break simultaneously?(14 votes)
- Intermolecular bonds break easier, but that does not mean first. Statistically, intermolecular bonds will break more often than covalent or ionic bonds. Yes, they can both break at the same time, it is just a matter of probability. Water, for example is always evaporating, even if not boiling. But, then, why no hydrogen or oxygen is observed as a product of pure water? Because water decomposes into H+ and OH- when the covalent bond breaks. The concentration of each of these ions in pure water, at 25°C, and pressure of 1atm, is 1.0×10e−7mol/L... that is: covalent bonds are breaking all the time (self-ionization), just like intermolecular bonds (evaporation). By the way, that is what makes both pH and pOH of water equal 7.
Note that even though H+ and OH- are naturally produced in water, they also recombine back into H2O. The speed of bonds breaking and the speed of recombination "fight" one another, until they are in chemical equilibrium, that is when both speeds are the same. See https://en.wikipedia.org/wiki/Chemical_equilibrium
For instance, this kind of "fight" also happens with evaporation inside a closed recipient. The equilibrium is reached when the number of molecules escaping from the liquid phase is the same as the number of molecules entering it.(17 votes)
- Are hydrogen bonds exclusive to hydrogen?(9 votes)
- A hydrogen-bond is a specific type of strong intermolecular dipole-dipole interaction between a partially positively-charged hydrogen atom and a partially negatively-charged atom that is highly electronegative, namely N, O, and F, the 3 most electronegative elements in the periodic table.(14 votes)
- what determines a molecule's electronegativity?(11 votes)
- The electronegativity of an atom depends on its effective nuclear charge (the positive charge experienced by an electron in the atom's outermost shell) and the distance between the outermost shell and the nucleus. Elements with a high effective nuclear charge and a small atomic radius tend to have high electronegativities. Fluorine, which has the highest electronegativity value of 4.0 on the Pauling scale, has a small atomic radius and a high effective nuclear charge.(12 votes)
- what are metalic bonding(3 votes)
- Metallic bonding occurs between metal atoms. In this type of bond, the metal atoms each contribute their valence electrons to a big, shared, cloud of electrons. Because the electrons can move freely in the collective cloud, metals are able to have their well-known metallic properties, such as malleability, conductivity, and shininess.(5 votes)
- In the section about nonpolar bonding, the article says carbon-hydrogen bonds are relatively nonpolar, even though the same element is not being bonded to another atom of the same element.
How is one supposed to know which bonds are nonpolar if the same atom isn't being bonded to the same atom (e.g. O2)(4 votes)
- Oils are nonpolar as a molecule due to their symmetry, however, there can be polar bonds (bond dipoles) that make it up. In this situation can there be times when the partial positive end of one of the oil molecules attracts the partial negative end of the other oil molecules? Can this give a dipole-dipole force, or can there only be dispersion forces? Thanks.(4 votes)
- What is the sense of 'cell' in the last paragraph?(4 votes)
- Because it is the compartment "biology" and all the chemistry here is about something that happens in biological world. So it's basically the introduction to cell structures.(1 vote)