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4.8: Polarity & Hydrogen Bonding

  • Page ID
    49887
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    Skills to Develop

    • Explain how polar compounds differ from nonpolar compounds.
    • Determine if a molecule is polar or nonpolar.
    • Identify whether or not a molecule can exhibit hydrogen bonding.
    • List important phenomena which are a result of hydrogen bonding.
    • Given a pair of compounds, predict which would have a higher melting or boiling point.

    Introduction

    The ability of an atom in a molecule to attract shared electrons is called electronegativity. When two atoms combine, the difference between their electronegativities is an indication of the type of bond that will form. If the difference between the electronegativities of the two atoms is small, neither atom can take the shared electrons completely away from the other atom and the bond will be covalent. If the difference between the electronegativities is large, the more electronegative atom will take the bonding electrons completely away from the other atom (electron transfer will occur) and the bond will be ionic. This is why metals (low electronegativities) bonded with nonmetals (high electronegativities) typically produce ionic compounds.

    Polar Covalent Bonds

    So far, we have discussed two extreme types of bonds. One case is when two identical atoms bond. They have exactly the same electronegativities, thus the two bonded atoms pull exactly equally on the shared electrons. The shared electrons will be shared CK12 Screenshot 4-8-1.pngexactly equally by the two atoms.

    The other case is when the bonded atoms have a very large difference in their electronegativities. In this case, the more electronegative atom will take the electrons completely away from the other atom and an ionic bond forms.

    What about the molecules whose electronegativities are not the same but the difference is not big enough to form an ionic bond? For these molecules, the electrons remain shared by the two atoms but they are not shared equally. The shared electrons are pulled closer to the more electronegative atom. This results in an uneven distribution of electrons over the molecule and causes slight charges on opposite ends of the molecule. The negative electrons are around the more electronegative atom more of the time creating a partial negative side. The other side has a resulting partial positive charge. These charges are not full \(+1\) and \(-1\) charges, they are fractions of charges. For small fractions of charges, we use the symbols \(\delta+\) and \(\delta-\). These molecules have slightly opposite charges on opposite ends of the molecule and are said to have a dipole or are called polar molecules.CK12 Screenshot 4-8-2.png

    When atoms combine, there are three possible types of bonds that they can form. In the figure, molecule A represents a covalent bond that would be formed between identical atoms. The electrons would be evenly shared with no partial charges forming. This molecule is nonpolar. Molecule B is a polar covalent bond formed between atoms whose electronegativities are not the same but whose electronegativity difference is less than 1.7, making this molecule polar. Molecule C is an ionic bond formed between atoms whose electronegativity difference is greater than 1.7.

    CK12 Screenshot 4-8-3.png

    Polar molecules can be attracted to each other due to attraction between opposite charges. Polarity underlies a number of physical properties including surface tension, solubility, and melting and boiling points. The more attracted molecules are to other molecules, the higher the melting point, boiling point, and surface tension. We will discuss in more detail later how polarity can affect how compounds dissolve and their solubility.

    In order to determine if a molecule is polar or nonpolar, it is frequently useful to look at Lewis structures. Nonpolar compounds will be symmetric, meaning all of the sides around the central atom are identical - bonded to the same element with no unshared pairs of electrons. Polar molecules are asymmetric, either containing lone pairs of electrons on a central atom or having atoms with different electronegativities bonded.

    Example: Label each of the following as polar or nonpolar.

    1) Water, \(\ce{H_2O}\): CK12 Screenshot 4-8-4.png

    2) Methanol, \(\ce{CH_3OH}\): CK12 Screenshot 4-8-5.png

    3) Hydrogen Cyanide, \(\ce{HCN}\): CK12 Screenshot 4-8-6.png

    4) Oxygen, \(\ce{O_2}\): CK12 Screenshot 4-8-7.png

    5) Propane, \(\ce{C_3H_8}\): CK12 Screenshot 4-8-8.png

    Solution:

    1) Water is polar. Any molecule with lone pairs of electrons around the central atom is polar.

    2) Methanol is polar. This is not a symmetric molecule. The \(\ce{-OH}\) side is different from the other 3 \(\ce{-H}\) sides.

    3) Hydrogen cyanide is polar. The molecule is not symmetric. The nitrogen and hydrogen have different electronegativities, creating an uneven pull on the electrons.

    4) Oxygen is nonpolar. The molecule is symmetric. The two oxygen atoms pull on the electrons by exactly the same amount.

    5) Propane is nonpolar, because it is symmetric, with \(\ce{H}\) atoms bonded to every side around the central atoms and no unshared pairs of electrons.

    While molecules can be described as "polar covalent" or "ionic", it must be noted that this is often a relative term, with one molecule simply being more polar or less polar than another. However, the following properties are typical of such molecules. Polar molecules tend to:

    • have higher melting points than nonpolar molecules
    • have higher boiling points than nonpolar molecules
    • be more soluble in water (dissolve better) than nonpolar molecules
    • have lower vapor pressures than nonpolar molecules

    Hydrogen Bonding

    When a hydrogen atom is bonded to a very electronegative atom, including fluorine, oxygen, or nitrogen, a very polar bond is formed. The electronegative atom obtains a negative partial charge and the hydrogen obtains a positive partial charge. These partial charges are similar to what happens in every polar molecule. However, because of the big difference in electronegativities between these two atoms and the amount of positive charge exposed by the hydrogen, the dipole is much more dramatic. These molecules will be attracted to other molecules which also have partial charges. This attraction for other molecules which also have a hydrogen bonded to a fluorine, nitrogen, or oxygen atom is called a hydrogen bond.

    Hydrogen bonds in water

    The most important, most common, and perhaps simplest example of a hydrogen bond is found between water molecules. This interaction between neighboring water molecules is responsible for many of the important properties of water.

    Hydrogen bonding strongly affects the crystal structure of ice, helping to create an open hexagonal lattice. The density of ice is less than water at the same temperature; thus, the solid phase of water floats on the liquid, unlike most other substances in which the solid form would sink in the liquid form.

    Water also has a high boiling point (\(100^\text{o} \text{C}\)) compared to the other compounds of similar size without hydrogen bonds. Because of the difficulty of breaking these bonds, water has a very high boiling point, melting point, and viscosity compared to otherwise similar liquids not conjoined by hydrogen bonds.

    Water is unique because its oxygen atom has two lone pairs and two hydrogen atoms, meaning that the total number of bonds a water molecule is up to four. For example, hydrogen fluoride - which has three lone pairs on the \(\ce{F}\) atom but only one \(\ce{H}\) atom - can form only two bonds: \(\ce{H-F}\)--\(\ce{H-F}\)--\(\ce{H-F}\) (ammonia has the opposite problem: three hydrogen atoms but only one lone pair).

    Have you ever experience a belly flop? This is also due to the hydrogen bonding between water molecules, causing surface tension. On the surface of water, water molecules are even more attracted to their neighbors than in the rest of the water. This attraction makes it difficult to break through, causing belly flops. It also explains why water striders are able to stay on top of water and why water droplets form on leaves or as they drip out of your faucet.

    CK12 Screenshot 4-8-9.png

    Hydrogen bonds in DNA and proteins

    Hydrogen bonding also plays an important role in determining the three-dimensional structures adopted by proteins and nucleic bases, as found in your DNA. In these large molecules, bonding between parts of the same macromolecule cause it to fold into a specific shape, which helps determine the molecule's physiological or biochemical role. The double helical structure of DNA, for example, is due largely to hydrogen bonding between the base pairs, which link one complementary strand to the other and enable replication. It also plays an important role in the structure of polymers, both synthetic and natural, such as nylon and many plastics.

    As a result of the strong attraction between molecules that occurs in a hydrogen bond, the following properties can be summarized. Molecules with hydrogen bonding tend to:

    • have higher melting points than polar molecules
    • have higher boiling points than polar molecules
    • be more soluble in water (dissolve better) than polar molecules

    Example: Label each of the following as polar or nonpolar and indicate which have hydrogen bonding.

    a) \(\ce{H_2O}\), CK12 Screenshot 4-8-10.png

    b) Ammonia, CK12 Screenshot 4-8-11.png

    c) \(\ce{CH_4}\), CK12 Screenshot 4-8-12.png

    d) Acetone, \(\ce{CH_3COCH_3}\), CK12 Screenshot 4-8-13.png

    Solution:

    a) This molecule is polar (the unshared pairs of electrons make a polar asymmetric shape), and hydrogen bonding (hydrogen is bonded to \(\ce{O}\)).

    b) This molecule is polar (the unshared pairs of electrons make a polar asymmetric shape), and hydrogen bonding (hydrogen is bonded to \(\ce{N}\)).

    c) This molecule is nonpolar (the molecule is symmetric with \(\ce{H}\)'s bonded to all four sides of the central atom), and does not have hydrogen bonding (hydrogen is not bonded to \(\ce{N}\), \(\ce{O}\), or \(\ce{F}\)).

    d) This molecule is polar (the \(\ce{O}\) is not the same as the \(\ce{CH_3}\) bonded to the central atom) and does not have hydrogen bonding (hydrogen is not bonded DIRECTLY to \(\ce{N}\), \(\ce{O}\), or \(\ce{F}\)).

    Example: For each pair of molecules, indicate which you would expect to have a higher melting point. Explain why. Also, refer to the Lewis structures given to you in the previous example.

    a) \(\ce{H_2O}\) vs. acetone

    b) \(\ce{CH_4}\) vs. acetone

    Solution:

    a) \(\ce{H_2O}\) (polar, hydrogen bonding) vs. acetone (polar, no hydrogen bonding). \(\ce{H_2O}\) will have a higher melting point because compounds with hydrogen bonding tend to have higher melting points than polar compounds.

    b) \(\ce{CH_4}\) (nonpolar, no hydrogen bonding) vs. acetone (polar, no hydrogen bonding). Acetone will have a higher melting point because polar molecules tend to have higher melting points than nonpolar molecules.

    Lesson Summary

    • Covalent bonds between atoms that are not identical will produce polar bonds.
    • Molecules with polar bonds and non-symmetrical shapes will have a dipole.
    • Hydrogen bonding is a special interaction felt between molecules, which is a stronger interaction than polar-polar attraction.
    • Hydrogen bonding occurs between molecules in which a hydrogen atom is bonded to a very electronegative fluorine, oxygen, or nitrogen atom.
    • Compounds with hydrogen bonding tend to have higher melting points, higher boiling points, and greater surface tension.
    • The unique properties of water are a result of hydrogen bonding.
    • Hydrogen bonding plays roles in many compounds including DNA, proteins, and polymers.

    Vocabulary

    • Electronegativity: The tendency of an atom in a molecule to attract shared electrons to itself.
    • Polar covalent bond: A covalent bond in which the electrons are not shared equally because one atom attracts them more strongly than the other.

    Further Reading/Supplemental Links

    Contributors


    4.8: Polarity & Hydrogen Bonding is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts.

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