What is the strongest covalent bond?

Strengths of Ionic and Covalent Bonds

A bond’s strength describes how strongly each atom is joined to another atom, and therefore how much energy is required to break the bond between the two atoms. In this section, you will learn about the bond strength of covalent bonds, and then compare that to the strength of ionic bonds, which is related to the lattice energy of a compound.

Bond Strength: Covalent Bonds

Stable molecules exist because covalent bonds hold the atoms together. We measure the strength of a covalent bond by the energy required to break it, that is, the energy necessary to separate the bonded atoms. Separating any pair of bonded atoms requires energy (see [link]). The stronger a bond, the greater the energy required to break it.

The energy required to break a specific covalent bond in one mole of gaseous molecules is called the bond energy or the bond dissociation energy. The bond energy for a diatomic molecule, D_X–Y, is defined as the standard enthalpy change for the endothermic reaction: $$XY(g)⟶X;(g)+Y;(g)qquad D_=ΔH°$$

For example, the bond energy of the pure covalent H–H bond, D_H–H, is 436 kJ per mole of H–H bonds broken: $$H_2(g)⟶2H(g)qquad D_=ΔH°=436;kJ/mol$$

Molecules with three or more atoms have two or more bonds. The sum of all bond energies in such a molecule is equal to the standard enthalpy change for the endothermic reaction that breaks all the bonds in the molecule. For example, the sum of the four C–H bond energies in CH₄, 1660 kJ, is equal to the standard enthalpy change of the reaction:

The average C–H bond energy, D_C–H, is 1660/4 = 415 kJ/mol because there are four moles of C–H bonds broken per mole of the reaction. Although the four C–H bonds are equivalent in the original molecule, they do not each require the same energy to break; once the first bond is broken (which requires 439 kJ/mol), the remaining bonds are easier to break. The 415 kJ/mol value is the average, not the exact value required to break any one bond.

The strength of a bond between two atoms increases as the number of electron pairs in the bond increases. Generally, as the bond strength increases, the bond length decreases. Thus, we find that triple bonds are stronger and shorter than double bonds between the same two atoms; likewise, double bonds are stronger and shorter than single bonds between the same two atoms. Average bond energies for some common bonds appear in [link], and a comparison of bond lengths and bond strengths for some common bonds appears in [link]. When one atom bonds to various atoms in a group, the bond strength typically decreases as we move down the group. For example, C–F is 439 kJ/mol, C–Cl is 330 kJ/mol, and C–Br is 275 kJ/mol.

Bond Energies kJ/mol

Average Bond Lengths and Bond Energies for Some Common Bonds

We can use bond energies to calculate approximate enthalpy changes for reactions where enthalpies of formation are not available. Calculations of this type will also tell us whether a reaction is exothermic or endothermic. An exothermic reaction (ΔH negative, heat produced) results when the bonds in the products are stronger than the bonds in the reactants. An endothermic reaction (ΔH positive, heat absorbed) results when the bonds in the products are weaker than those in the reactants.

Ionic Bond Strength and Lattice Energy

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a compound is a measure of the strength of this attraction. The lattice energy (ΔH_lattice) of an ionic compound is defined as the energy required to separate one mole of the solid into its component gaseous ions. For the ionic solid MX, the lattice energy is the enthalpy change of the process:

The strength of a covalent bond is measured by its bond dissociation energy, that is, the amount of energy required to break that particular bond in a mole of molecules. Multiple bonds are stronger than single bonds between the same atoms.

Chemistry End of Chapter Exercises

Which bond in each of the following pairs of bonds is the strongest?

7 Chemical Bonds

Atoms can form several types of chemical bonds. These bonds are interactions between two atoms that hold the atoms together. It is important to understand the various types of bonds because they help determine how different molecules function within an organism. There are four types of bonds or interactions: covalent, ionic, hydrogen bonds, and van der Waals interactions.

Covalent Bonds

Another type of strong chemical bond between two or more atoms is a covalent bond. These bonds form when an electron is shared between two elements. Covalent bonds are the strongest (*see note below) and most common form of chemical bond in living organisms.

The hydrogen and oxygen atoms that combine to form water molecules are bound together by strong covalent bonds. The electron from the hydrogen atom shares its time between the hydrogen atom and the oxygen atom. In order for the oxygen atom to be stable, two electrons from two hydrogen atoms are needed, hence the subscript “2” in H₂O. H₂O means that there are 2 hydrogen atoms bonded to 1 oxygen atom (the 1 is implied below the O in the chemical formula). This sharing makes both the hydrogen and oxygen atoms more chemically stable.

There are two types of covalent bonds: polar and nonpolar (Figure 3). Nonpolar covalent bonds form between two atoms that share the electrons equally so there is no overall charge on the molecule. For example, an oxygen atom can bond with another oxygen atom. This association is nonpolar because the electrons will be equally shared between each oxygen atom. Another example of a nonpolar covalent bond is found in the methane (CH₄) molecule. The carbon atom shares electrons with four hydrogen atoms. The carbon and hydrogen atoms all share the electrons equally, creating four nonpolar covalent bonds (Figure 3).

In a polar covalent bond, the electrons shared by the atoms spend more time closer to one atom than to the other. Because of the unequal distribution of electrons between the atoms, a slightly positive (δ+) or slightly negative (δ–) charge develops. The covalent bonds between hydrogen and oxygen atoms in water are polar covalent bonds. The shared electrons spend more time near the oxygen than they spend near the hydrogen. This means that the oxygen has a small negative charge while the hydrogens have a small positive charge.

*BUT WAIT! In chemistry, I learned that Ionic bonds are stronger than covalent bonds! What’s up with that?

Turns out that chemists and biologists measure bond strength in different ways. Chemists measure the absolute strength of a bond (kind of like the theoretical strength). Ionic bonds are very strong when measured this way. Biologists are more interested in how the bond behaves in a biological system, which is usually aqueous (water-based). In water, ionic bonds come apart much more readily than covalent bonds, so biologists would say that they are weaker.

So what’s the right answer? Depends on whether you’re in a chemistry or a biology class! If you look in a biology textbook, it will almost always tell you that covalent bonds are stronger. If you look in a chemistry textbook, you’ll see something different. This is a great example of how the same information can lead to different answers depending on the perspective that you’re viewing it from.

So what answer should you give for this class? Because this is a biology class, you should say that covalent bonds are stronger than ionic bonds because they act stronger in aqueous solutions.

Ionic Bonds

Atoms normally have an equal number of protons (positive charge) and electrons (negative charge). This means that atoms are normally uncharged because the number of positively charged particles equals the number of negatively charged particles. When an atom does not contain equal numbers of protons and electrons, it will have a net charge. An atom with a net charge is called an ion. Positive ions are formed by losing electrons. Negative ions are formed by gaining electrons. Atoms can lose and donate electrons in order to become more stable.

When an element donates an electron from its outer shell, as in the sodium atom example above, a positive ion is formed (Figure 2). The element accepting the electron is now negatively charged. Because positive and negative charges attract, these ions stay together and form an ionic bond, or a bond between ions. The elements bond together with the electron from one element staying predominantly with the other element. When Na and Cl combine to produce NaCl, an electron from a sodium atom goes to stay with the other seven electrons in the chlorine atom, forming a positively charged sodium ion and a negatively charged chlorine ion. The sodium and chloride ions attract each other.

Hydrogen Bonds

Ionic and covalent bonds are strong bonds that require considerable energy to break. However, not all bonds between elements are ionic or covalent bonds. Weaker bonds can also form. These are attractions that occur between positive and negative charges that do not require much energy to break. Two weak bonds that occur frequently are hydrogen bonds and van der Waals interactions. These bonds give rise to the unique properties of water and the unique structures of DNA and proteins.

When polar covalent bonds containing a hydrogen atom form, the hydrogen atom in that bond has a slightly positive charge. This is because the shared electron is pulled more strongly toward the other element and away from the hydrogen nucleus. Because the hydrogen atom is slightly positive (δ+), it will be attracted to neighboring negative partial charges (δ–). When this happens, a weak interaction occurs between the δ+ charge of the hydrogen atom of one molecule and the δ– charge of the other molecule. This interaction is called a hydrogen bond. This type of bond is common; for example, the liquid nature of water is caused by the hydrogen bonds between water molecules (Figure 4). Hydrogen bonds give water the unique properties that sustain life. If it were not for hydrogen bonding, water would be a gas rather than a liquid at room temperature.

Hydrogen bonds can form between different molecules and they do not always have to include a water molecule. Hydrogen atoms in polar bonds within any molecule can form bonds with other adjacent molecules. For example, hydrogen bonds hold together two long strands of DNA to give the DNA molecule its characteristic double-stranded structure. Hydrogen bonds are also responsible for some of the three-dimensional structure of proteins.

van der Waals Interactions

Like hydrogen bonds, van der Waals interactions are weak attractions or interactions between molecules. They occur between polar, covalently bound, atoms in different molecules. Some of these weak attractions are caused by temporary partial charges formed when electrons move around a nucleus. These weak interactions between molecules are important in biological systems.

References

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The Science of PFAS: Finding Strength in the Single Bond

Article-The Science of PFAS: Finding Strength in the Single Bond

EREF Staff | May 13, 2021

To date, the large variation in the properties of PFAS compounds make it difficult to have a concrete understanding of exactly how they will behave coupled with the difficulty of being able to track and understand the over 4,600 PFAS compounds currently registered. If there is anything that can be learned from a quick search of PFAS it is that the carbon-fluorine (C-F) bond is one of the main properties that we can thank for the conveniences that these compounds offer, but also for the headaches that we experience when trying to figure out how to manage them. What is unique about the C-F bond and why is it different from any other chemical bond?

One of the main attributes of the C-F bond and also one of the main reasons for the love-hate relationship we have with PFAS, is the bond strength. In general, the bond strength dictates how easy it may or may not be for that bond to be broken. As the saying goes, there is strength in numbers and a similar statement can be made about the number of bonds in chemistry. As you increase the number of bonds from single, to double, to triple the strength of the bond increases. One way to think about this concept is attaching one rubber band to two balls and seeing how easy it is to pull apart those balls. What if you added another rubber band or even two to simulate the triple bond? What do you think happens? There is more resistance making it more difficult to pull apart relative to the balls attached with one rubber band.

Without increasing the number of bonds, chemists are able to use the various properties of elements to manipulate the strength of single bond. Organic chemists consider the C-F bond to the strongest single bond possible in organic chemistry. Therefore, the bond strength is the primary reason PFAS is known to be persistent in the environment, bioaccumulates, does not degrade in conventional treatment processes for both wastewater and drinking water, and is thermally stable. These properties are the reason PFAS is considered the “forever chemical.” Subsequent pieces will breakdown the science on how the number of carbons and functional groups also contribute to the complex nature of PFAS and how this impacts the management of these compounds.

What is it about the addition of fluorine that makes the bond so strong?

The reason for the strength of this bond is the electronegativity of fluorine and its relative attraction to carbon. Chemists use electronegativity as the measure of the ability for an atom (fluorine) to pull electrons away from another atom (carbon). This behavior is similar to a magnetic force pulling the electrons away from the carbon and towards fluorine causing a stronger bond and fluorine not wanting to give up those electrons. More specifically fluorine has a partial negative charge and carbon has a partial positive charge as shown below and as expected opposites will attract.

Looking further at a typical C-H bond relative to a C-F bond, we see that there is a 17% increase in the energy required to break the single bond when you replace H with F. Comparing the H-F and C-F bond to other common single bonds you see that the bond strength for C-F is still at the upper end relative to the other examples shown below.

The interaction between C-F has created both unique properties as well as challenges in the management of PFAS compounds. Unlocking ways to break the bond barrier will be key advancements needed towards finding affective ways to manage these compounds in the environment.