Water gushes out of the faucet. Honey oozes out of a squeeze bottle. Gasoline flows out of the pump. These are just three examples of a highly diverse state of matter: liquids. One of the key defining properties of liquids is their ability to flow. Beyond this feature, though, the behaviors of different liquids span a broad range. Some liquids flow relatively easily, like water or oil, while others, like honey or molasses, flow quite slowly. Some are slippery, and some are sticky. Where do these different behaviors come from?
When it comes to interactions between different liquids, some mix well: Think of a Shirley Temple, made of ginger ale and grenadine. Others, though, don’t seem to mix at all. Consider oil spills, where the oil floats in a sticky, iridescent layer on top of the water. You may also notice a similar phenomenon in some salad dressings that separate into an oil layer that rests atop a layer of vinegar, which is primarily water. Why don’t these liquids mix well?
These varied behaviors arise primarily from the different types of intermolecular forces that are present in liquids. In this module we’ll first discuss liquids in the context of the other two main states of matter, solids and gases. Then we will go through a brief overview of intermolecular forces, and finally we’ll explore how intermolecular forces govern the way that liquids behave.
Liquids, solids, and gases
Liquids flow because the intermolecular forces between molecules are weak enough to allow the molecules to move around relative to one another. Intermolecular forces are the forces between neighboring molecules. (These are not to be confused with intramolecular forces, such as covalent and ionic bonds, which are the forces exerted within individual molecules to keep the atoms together.) The forces are attractive when a negative charge interacts with a nearby positive charge and repulsive when the neighboring charges are the same, either both positive or both negative. In liquids, the intermolecular forces can shift between molecules and allow them to move past one another and flow. (See Figure 1 for an illustration of the various intermolecular forces and interactions.)
Contrast that with a solid, in which the intermolecular forces are so strong that they allow very little movement. While molecules may vibrate in a solid, they are essentially locked into a rigid structure, as described in the Properties of Solids module. At the other end of the spectrum are gases, in which the molecules are so far apart that the intermolecular forces are effectively nonexistent and the molecules are completely free to move and flow independently.
At a molecular level, liquids have some properties of gases and some of solids. First, liquids share the ability to flow with gases. Both liquid and gas phases are fluid, meaning that the intermolecular forces allow the molecules to move around. In both of these phases, the materials don’t have fixed shapes and instead are shaped by the containers holding them.
Solids are not fluid, but liquids share a different important property with them. Liquids and solids are both held together by strong intermolecular forces and are much more dense than gases, leading to their description as “condensed matter” phases because they are both relatively incompressible. (Figure 2 shows the differences of gases, liquids, and solids at the atomic level.)
Most substances can move between the solid, liquid, and gas phases when the temperature is changed. Consider the molecule H20: It takes the form of ice, a crystalline solid, below 0° C; water, a liquid, between 0° and 100° C; and water vapor, or steam, a gas, above 100° C. These transitions occur because temperature affects the intermolecular attraction between molecules. When H20 is converted from a liquid to gas, for instance, the rising temperature makes the molecules’ kinetic energy increase such that it eventually overcomes the intermolecular forces and the molecules are able to move freely about in the gas phase. However, the intramolecular forces that hold the H20 molecule together are unchanged; H20 is still H20, regardless of its state of matter. You can read more about phase transitions in the States of Matter module.
Now that we’ve discussed how liquids are similar to and different from solids and gases, we can focus on the wide world of liquids. First, though, we need to briefly introduce the different types of intermolecular forces that dictate how liquids, and other states of matter, behave.
As we described earlier, intermolecular forces are attractive or repulsive forces between molecules, distinct from the intramolecular forces that hold molecules together. Intramolecular forces do, however, play a role in determining the types of intermolecular forces that can form. Intermolecular forces come in a range of varieties, but the overall idea is the same for all of them: A charge within one molecule interacts with a charge in another molecule. Depending on which intramolecular forces, such as polar covalent bonds or nonpolar covalent bonds, are present, the charges can have varying permanence and strengths, allowing for different types of intermolecular forces.
So, where do these charges come from? In some cases, molecules are held together by polar covalent bonds – which means that the electrons are not evenly distributed between the bonded atoms. (This type of bonding is described in more detail in the Chemical Bonding module.) This uneven distribution results in a partial charge: The atom with more electron affinity, that is, the more electronegative atom, has a partial negative charge, and the atom with less electron affinity, the less electronegative atom, has a partial positive charge. This uneven electron sharing is called a dipole. When two molecules with polar covalent bonds are near each other, they can form favorable interactions if the partial charges align appropriately, as shown in Figure 3, forming a dipole-dipole interaction.
Hydrogen bonds are a particularly strong type of dipole-dipole interaction. (Note that although they are called “bonds,” they are not covalent or ionic bonds; they are a strong intermolecular force.) Hydrogen bonds occur when a hydrogen atom is covalently bonded to one of a few non-metals with high electronegativity, including oxygen, nitrogen, and fluorine, creating a strong dipole. The hydrogen bond is the interaction of the hydrogen from one of these molecules and the more electronegative atom in another molecule. Hydrogen bonds are present, and very important, in water, and are described in more detail in our Water: Properties and Behavior module.
Hydrogen bonds and dipole-dipole interactions require polar bonds, but another type of intermolecular force, called London dispersion forces, can form between any molecules, polar or not. The basic idea is that the electrons in any molecule are constantly moving around and sometimes, just by chance, the electrons can end up distributed unequally, creating a temporary partial negative charge on the part of the molecule with more electrons. This partial negative charge is balanced by a partial positive charge of equal magnitude on the part of the molecule with fewer electrons, with the positive charge coming from the protons in the nucleus (Figure 4). These temporary partial charges in neighboring molecules can interact in much the same way that permanent dipoles interact. The overall strength of London dispersion forces depends on the size of the molecules: larger molecules can have larger temporary dipoles, leading to stronger London dispersion forces.
Now, you might ask, if molecules can develop temporary partial charges that interact with each other, these temporary charges should also be able to interact with permanent dipoles, right? And you would be correct. These interactions are called, very creatively, dipole-induced dipole interactions. The partial charge of the polar molecule interacts with the electrons in the nonpolar molecule and “induces” them to move so they’re not evenly distributed anymore, creating an induced dipole that can interact favorably with the polar molecule’s permanent dipole (Figure 5).
As you might have guessed, London dispersion forces and dipole-induced dipole interactions are generally weaker than dipole-dipole interactions. These forces, as well as hydrogen bonds, are all van der Waals forces, which is a general term for attractive forces between uncharged molecules.
There’s a lot more to intermolecular forces than what we’ve covered here, but with this brief introduction, we’re ready to get back to the main event: liquids, and how intermolecular forces determine their properties and behavior.
Properties of liquids
If you’ve ever used oil for cooking or working on a car, you know that it’s nice and slippery. That’s probably why you used it: it keeps stir-fry pieces from sticking to each other or the pan, and it helps engine pistons and other moving parts slide easily.
One of the reasons oils are good for these applications is because they have low cohesion: the liquid molecules don’t interact particularly strongly with each other because the intermolecular forces are weak. The primary intermolecular forces present in most oils and many other organic liquids – liquids made predominantly of carbon and hydrogen atoms, also referred to as non-polar liquids – are London dispersion forces, which for small molecules are the weakest types of intermolecular forces. These weak forces lead to low cohesion. The molecules don’t interact strongly with each other, so they can slide right past one another.
On the other end of the cohesion spectrum, consider a dewdrop on a leaf in the early morning (Figure 6). How can such a thing exist if, as explained earlier, liquids flow and take the shape of the container holding them? As described above and in the Water module, water molecules are held together by strong hydrogen bonds. These strong forces lead to high cohesion: The water molecules interact with each other more strongly than they interact with the air or the leaf itself. (The interaction of the water with the leaf is an example of adhesion, or the interaction of a liquid with something other than itself; we’ll discuss adhesion in the next section.) Because of water’s high cohesion, the molecules form a spherical shape to maximize their interactions with each other.
This high cohesion also creates surface tension. You may have noticed insects walking on water on an outdoor pond (Figure 7), or seen a small object such as a paperclip resting on water’s surface instead of sinking; these are two examples of water’s surface tension in action. Surface tension results from the strong cohesive forces of some liquids. These forces are strong enough to be maintained even when they experience external forces like the weight of an insect walking across its surface.
Adhesion is the tendency of a compound to interact with another compound. (Remember that, in contrast, cohesion is the tendency of a compound to interact with itself.) Adhesion helps explain how liquids interact with their containers and with other liquids.
One example of an interaction with high adhesion is that between water and glass. Both water and glass are held together by polar bonds. Therefore, the two materials can also form favorable polar interactions with each other, leading to high adhesion. You may have even seen these attractive adhesive forces in action in lab. When water is in a glass graduated cylinder, for example, the water creeps up the sides of the glass, creating a concave curve at the top called a meniscus, as shown in the figure below. Water in graduated cylinders made out of some types of non-polar plastic, on the other hand, forms a flat meniscus because there are neither attractive nor repellant cohesive forces between the water and the plastic. (See Figure 8 for a comparison of polar and non-polar graduated cylinders.)
At the beginning of the module, we said that one of the defining features of liquids is their ability to flow. But among liquids there is a huge range in how easily this happens. Consider the ease with which you can pour yourself a glass of water, as compared to the relative challenge of pouring thick, slow-moving motor oil into an engine. The difference is their viscosity, or resistance to flow. Motor oil is quite viscous; water, not so much. But why?
Before we dive into the differences between water and motor oil, let’s compare water with another liquid: pentane (C5H12). While we don’t think of water as viscous, it’s actually more viscous than pentane. Remember, water molecules form strong hydrogen bonds with each other. Pentane, on the other hand, made up of just hydrogen and carbon atoms, is nonpolar, so the only types of intermolecular forces it can form are the relatively weak London dispersion forces. The weaker intermolecular forces mean that the molecules can more easily move past each other, or flow – hence, lower viscosity.
But both water and pentane are relatively small molecules. When we’re looking at liquids made of up bigger molecules, size comes into play as well. For example, compare pentane to motor oil, which is a complex mixture of large hydrocarbons much larger than little pentane, and some with dozens or even hundreds of carbons in a chain. If you’ve ever poured motor oil into an engine, you know it’s pretty viscous. Both liquids are nonpolar, and so have relatively weak intermolecular forces; the difference is the size. The big, bendy motor oil hydrocarbons can literally get tangled with their neighbors, which slows the flow. It’s almost like a pot of spaghetti: If you don’t prepare it correctly, you can end up with a blob of tangled noodles that are very hard to serve because they’re all stuck together – in a sense, it’s a viscous pasta blob. Shorter noodles – or smaller molecules – don’t tangle as much, so they tend to be less viscous (Figure 9).
Returning to our original comparison of motor oil versus water, even though water has such strong intermolecular forces, the much larger size of the molecules in the motor oil makes the oil more viscous.
There’s one more piece to the story: temperature. Warming a liquid makes it less viscous, as you may have observed if you’ve ever experienced how much easier it is to pour maple syrup onto your pancakes when the syrup has been warmed than when it is cold. This is the case because temperature affects both of the factors that determine viscosity in the first place. First, increasing the temperature increases the molecules’ kinetic energy, which allows them to overcome the intermolecular forces more easily. It also makes the molecules move around more, so those big molecules that got tangled up when they were cold become more dynamic and are more able to slide past each other, allowing the liquid to flow more easily.
When you think of water, you might think of its chemical formula, H2O. This formula describes a pure liquid composed only of H2O molecules, with absolutely no other components. The reality, though, is that the vast majority of liquids we encounter are complex mixtures of many compounds.
Solutions are made of a liquid solvent in which one or more solutes are dissolved. Solutes can be solids, liquids, and gases. There are many, many common solutions that use water as the solvent, including salt water and pretty much any type of flavored drink. Carbon dioxide (CO2) gas is a common gaseous solute in carbonated drinks, and ethanol is a liquid solute in any alcoholic drink. Although solutions are mixtures of multiple compounds, the properties discussed in the previous section still apply.
Not all solutes dissolve in all solvents. You can dissolve huge amounts of some solutes in some liquids, and other solutes are only marginally soluble in any solvent. The underlying explanation for solubility is that “like dissolves like.” Nonpolar solutes generally dissolve better in nonpolar liquids, and polar solutes dissolve better in polar liquids. For example, oil-based (and therefore nonpolar) paints require a non-polar solvent such as turpentine for clean up; they will not dissolve in water, which is polar. Table salt or sugar, on the other hand, both polar solids, easily dissolve at high concentrations in water.
More complex solutions include emulsions, colloids, and suspensions. Briefly, an emulsion is a well-dispersed mixture of two or more liquids that don’t normally mix. Mayonnaise, for example, is an emulsion of oil, egg yolk, and vinegar or lemon juice, which is made by very vigorous mixing.
Colloids and suspensions both consist of insoluble particles in a liquid. In a colloid, the miniscule insoluble particles are distributed in a liquid and won’t separate. And a suspension, on the other hand, is a liquid that contains larger insoluble particles that will eventually separate. Milk is a useful example of the difference between these two. Fresh milk is a suspension. It’s a complex mixture of components that don’t normally mix – water, fats, proteins, carbohydrates, and more – and if left alone the fat globules separate from the water-based portion of the mixture. (Remember the separation of vinegar and oil in salad dressing? The milk separation process is similar, with the oily fat separating from the water.) The milk at most grocery stores, on the other hand, is a colloid. The components don’t separate thanks to a process called homogenization, which breaks the fat globules into small enough particles that they can remain suspended in the liquid.
Beyond simple liquids
We’ve discussed a lot of different liquids, with varying cohesion, adhesion, and viscosity, as well as other properties. But in addition to this already wide variety, there are some substances that blur the distinction between liquid and solid. For example, as a kid you may have played with oobleck, a mixture of water and starch that gets its name from a Dr. Seuss book. Oobleck is a slimy substance that can flow between your fingers if you hold it gently in your hands but becomes hard and firm, almost solid, if you squeeze it.
For a more technical example, consider the material used in LCD television displays and other electronic screens. LCD stands for Liquid-Crystal Display. That doesn’t mean that the displays use both liquids and crystals; it means that they use a material that is both liquid and crystal, at the same time. This might sound like a contradiction – crystals are solids, not liquids, you say – but such materials exist.
The first liquid crystal discovered was a modified version of cholesterol, called cholesteryl benzoate. It’s a solid at room temperature and melts at around 150°C, but then things get weird. At about 180°C, it changes phase again, but not from liquid to gas; it goes from cloudy liquid to clear liquid. Austrian botanist and chemist Friedrich Reinitzer observed this unusual behavior in 1888 and discussed it with his colleague, German physicist Otto Lehmann. Lehmann then took over the investigation, studying cholesteryl benzoate and other compounds with similar double-melting behavior. When he looked at the cloudy phase under his microscope, he found that the material appeared crystalline, a defining feature of solids. But the phase also flowed, like a liquid. In 1904 he coined the term “liquid crystal” to describe this phase, with properties between those of a conventional liquid and crystalline solid. Liquid crystals play an important role in biology, particularly in membranes, which need to be fluid but also must retain a regular structure.
There are also some liquids that are so viscous you wouldn’t be blamed for thinking that they’re solid, such as pitch, a substance derived from plants and petroleum. It appears almost solid, and shatters if hit with a hammer, but if left to gravity it will flow extremely, extremely slowly. A few labs around the world are running so-called pitch drop experiments, in which they leave some pitch in a funnel and wait for it to drip; about 10 years pass between each drop (Figure 10).
These examples of substances behaving in ways that seem to defy the traditional definitions for the phases of matter illustrate the inherent complexity of science and the natural world, even when it comes to something as seemingly simple as determining whether a substance is a liquid or a solid. In this module we have focused on defining and explaining the basic properties of liquids, which provides the foundation for you to think about states of matter in all their complexity. In other modules we discuss the solid and gas phases to help you contrast the different physical properties of these states.
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