It is very rare for both sides to be evenly matched in a tug of war. Inevitably, one side will be stronger. The ribbon tied around the middle of the rope will be pulled closer to one side, rather than the other.
This ribbon represents the shared pair of electrons in a polar bond. Instead of being found exactly halfway between the two bonded atoms, the electrons are pulled over to one side. Let's explore why.
- This article is about polar and non-polar covalent bonds.
- We will look at the difference between polar and non-polar bonds.
- We'll explore what causes bond polarity and the characteristics of polar and non-polar covalent bonds.
- We'll then look at bond polarity as a whole, with consideration of ionic character.
- Finally, we'll provide you with a list of examples of polar and non-polar covalent bonds.
What are Polar and Non-Polar Covalent Bonds?
A covalent bond is nothing but a shared pair of electrons. A covalent bond is formed when atomic orbitals from two atoms, usually non-metals, overlap, and the electrons within them form a pair that is shared by both atoms. The bond is held together by strong electrostatic attraction between the negative electrons and the atoms' positive nuclei.
If the two atoms involved in the covalent bond are the same, they share the electron pair evenly between them. This forms a non-polar bond.
A non-polar covalent bond is a bond in which the electron pair is shared equally between the two bonded atoms.
One example is hydrogen gas, H2. The two hydrogen atoms are identical, so the bond between them is non-polar.
Fig. 1. A non-polar H-H bond.
But if the two atoms involved in the covalent bond are different, the electron pair might not be shared evenly between them. One atom could attract the shared pair of electrons more strongly than the other atom, pulling the electrons over towards itself. The electron pair is shared unequally between the two atoms. We call this a polar bond.
A polar covalent bond is a bond in which the electron pair is shared unequally between the two bonded atoms.
Now we know that a polar bond is formed when an electron pair is shared unequally between two atoms. But what causes this uneven distribution?
What Causes Polar Bonds?
We've learned that polar covalent bonds are formed when one atom in a covalent bond attracts the shared pair of electrons towards itself more strongly than the other. This is all to do with the atom's electronegativity.
We measure electronegativity on the Pauling scale. It runs from 0.79 to 3.98, with fluorine being the most electronegative element, and francium the least electronegative. (The Pauling scale is a relative scale, so don't worry about how we get these numbers for now).
Fig. 2. The Pauling scale.
You can read more about this topic over at Electronegativity.
When it comes to covalent bonds, the more electronegative atom attracts the shared pair of electrons more strongly than the less electronegative atom. The more electronegative atom becomes partially negatively charged, and the less electronegative atom becomes partially positively charged. For example, you can see in the table above that oxygen is a lot more electronegative than hydrogen. This is why the oxygen atom in an O-H bond becomes partially negatively charged, and the hydrogen atom becomes partially positively charged.
In general, we can say the following:
- When two atoms with the same electronegativity share a pair of valence electrons, they form a non-polar bond.
- When two atoms with different electronegativities share a pair of valence electrons, they form a polar bond.
Characteristics of Polar and Non-Polar Covalent Bonds
Now that we know what polar and non-polar covalent bonds are, let's look at their characteristics. In the section above, you learned that polar covalent bonds are formed between two elements with differing electronegativities. This gives polar covalent bonds the following characteristics:
One example of a polar bond is the O-H bond, such as in water, or H2O. Oxygen attracts the shared pair of electrons much more strongly than hydrogen, resulting in a polar bond. Let's use this example to explore the characteristics of polar covalent bonds a little further.
Partial Charges
Look at our example, the O-H bond. Oxygen is more electronegative than hydrogen and so attracts the shared pair of electrons towards itself more strongly. Because the negative pair of electrons is found much closer to oxygen than hydrogen, the oxygen becomes partially negatively charged. The hydrogen, which is now electron-deficient, becomes partially positively charged. We represent this using the delta symbol, δ.
Fig. 3. The polar O-H bond.
Dipole Moments
You can see in the example above that the uneven distribution of electrons in a polar bond causes an uneven distribution of charge. One atom involved in the bond becomes partially negatively charged, while the other is partially positively charged. This creates a dipole moment. Asymmetrical molecules with dipole moments form dipole molecules. (You can explore this in more detail in Dipoles, and Dipole Moment.)
In contrast to polar bonds, the atoms in a non-polar covalent bond have no partial charges and form completely neutral molecules without any dipole moments.
The Difference Between Polar and Non-Polar Covalent Bonds
The basic difference between a polar and non-polar covalent bond is that a polar covalent bond has an unequal distribution of charges, whilst in a non-polar bond all atoms have the same charge distribution. This is because in polar bonds some of the atoms have higher electronegativity than others, whilst in the non-polar bonds all atoms have the same electronegativity value.
However, in real-life examples, when it comes to bonding, it is hard to draw a line between polar, non-polar, and indeed even ionic bonding. To understand why, let's look more closely at one particular bond: the C-H bond.
Carbon has an electronegativity of 2.55; hydrogen has an electronegativity of 2.20. This means that they have an electronegativity difference of 0.35. We might guess that this forms a polar bond, but in actual fact, we consider the C-H bond to be non-polar. This is because the electronegativity difference between the two atoms is so small that it is essentially insignificant. We can assume that the electron pair is shared equally between the two atoms.
On the other hand, consider the Na-Cl bond. Sodium has an electronegativity of 0.93; chlorine has an electronegativity of 3.16. This means that they have an electronegativity difference of 2.23. This bond is polar. However, the electronegativity difference between the two atoms is so great that the electron pair is essentially completely transferred from sodium to chlorine. This transfer of electrons forms an ionic bond.
Visit Ionic Bonding for more on this subject.
Bonding falls on a spectrum. At one end, you have completely non-polar covalent bonds, formed between two identical atoms with the same electronegativity. At the other end, you have ionic bonds, formed between two atoms with an extremely large difference in electronegativity. Somewhere in the middle, you find polar covalent bonds, formed between two atoms with an intermediate difference in electronegativity. But where do we draw the limits?
- If two atoms have an electronegativity difference of 0.4 or less, they form a non-polar covalent bond.
- If two atoms have an electronegativity difference between 0.4 and 1.8, they form a polar covalent bond.
- If two atoms have an electronegativity difference of more than 1.8, they form an ionic bond.
We can say that the bond has an ionic character proportional to the difference in electronegativity between the two atoms. As you might be able to guess, atoms with a larger difference in electronegativity show more ionic character; atoms with a smaller difference in electronegativity show less ionic character.
Fig. 4. Non-polar, polar, and ionic bonds are shown with the electronegativities of the atoms.
Predicting Bonding from Elemental Properties
Although bonding falls on a spectrum, it is often easier to classify a bond as non-polar covalent, polar covalent, and ionic. Generally, a bond between two non-metals is a covalent bond, and a bond between a metal and a non-metal is an ionic bond. But this isn't always the case. For example, take SnCl4. Tin, Sn, is a metal, and chlorine, Cl, is a non-metal, so we'd expect them to bond ionically. However, they actually bond covalently. We can use their properties to predict this.
- Ionic compounds have high melting and boiling points, are brittle, and can conduct electricity when molten or aqueous.
- Covalent small molecules have low melting and boiling points and don't conduct electricity.
Let's look at our example above: SnCl4 melts at -33°C. This gives us a pretty good indication that it bonds covalently, not ionically.
You might wonder: Why don't we just look at the difference in electronegativity when determining the nature of a bond? While it is a useful guide most of the time, this system doesn't always work.
We learned that SnCl4 forms polar covalent bonds. Indeed, a look at the two elements' electronegativities confirms this: Tin has an electronegativity of 1.96, while chlorine has an electronegativity of 3.16. Their electronegativity difference is therefore 1.2, well within the range for polar covalent bonding. However, tin and chlorine don't always bond covalently. In SnCl2, the two elements actually form ionic bonds.
Once again, the compound's properties help us deduce this: SnCl2 melts at 246°C, a much higher boiling point than that of its cousin SnCl4. But like all rules of thumb, this doesn't work for all compounds. For example, some giant "covalent network solids" such as diamond consist entirely of non-polar covalent bonds but have very high melting and boiling points.
To summarise, ionic bonding is generally found between metals and non-metals, and covalent bonding is generally found between two non-metals. Electronegativity differences also give us an indication of the bonding present in a molecule or compound. However, some compounds break these trends; looking at properties is a more reliable way of determining the bond.
List of Polar and Non-Polar Covalent Bonds (Examples)
Let's end with some examples of polar and non-polar covalent bonds. Here's a handy table that should help you.
| Non-polar covalent bond | Example | Polar covalent bond | Application |
| Any bond between two atoms of the same element | Cl-Cl, used to disinfect water | O-H | Two essential liquids: H2O and CH3CH2OH |
| C-H | CH4, a troublesome greenhouse gas | C-F | Teflon, the non-stick coating that you find on pans |
| Al-H | AlH3, used to store hydrogen for fuel cells | C-Cl | PVC, the world's third-most widely produced plastic polymer |
| Br-Cl | BrCl, an extremely reactive golden gas | N-H | NH3, which serves as a precursor to 45% of the world's food |
| O-Cl | Cl2O, an explosive chlorinating agent | C=O | CO2, a product of respiration and the source of bubbles in fizzy drinks |
That's all! You should now be able to state the difference between polar and non-polar covalent bonding, explain how and why polar bonds are formed, and predict whether a bond is polar or non-polar based on the properties of the molecule.
Polar and Non-Polar Covalent Bonds - Key takeaways
- A covalent bond is a shared pair of electrons. A non-polar covalent bond is a bond in which the electron pair is shared equally between the two bonded atoms, while a polar covalent bond is a bond in which the electron pair is shared unequally between the two bonded atoms.
- Polar bonds are caused by differences in electronegativity. The more electronegative atom becomes partially negatively charged, and the less electronegative atom becomes partially positively charged.
- Bonding is a spectrum, with non-polar covalent bonding at one end and ionic bonding at the other. Most bonding falls somewhere in between, and we say that these bonds show ionic character.
- We can use differences in electronegativity to predict the dipole moment. However, this isn't always the case; looking at a molecular species' physical properties can be a more accurate way of determining its bonding.
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