Emulsions are everywhere. Learn some of the science behind emulsions and impress friends and family with your sauce game.

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Though we don’t often think about it, emulsions are everywhere in cooking. Mayonnaise, custard, hollandaise sauce, butter, sausages and even milk are all emulsions.

From a scientific perspective emulsions are also really interesting. Emulsions are the answer to the question of what happens when we mix two things that really don’t want to be mixed.

There’s a lot to the science of emulsions so in this post I want to look at the most humble emulsion: the vinaigrette.

Usually just something I put on lettuce to make it taste like something other than lettuce, the vinaigrette, in its simplest form, is just a mixture of oil and water. But even this gives us plenty of science to get on with.

What is a Food Emulsion?

An emulsion is simply a mixture of two substances that really don’t want to mix. In the kitchen this is almost always, with the exception of foams, an oil and something acidic; vinegar, cider or orange juice for example.

The Polarity of Water

What drives this reluctance to get together is the concept of hydrophobicity. I touched on this in the post on pasta and simply put molecules that don’t have an electrical charge on their surface don’t like interacting with water. They prefer to hang out with other uncharged molecules.

In water, this means hydrophobic molecules will clump together to minimise their contact with water.

They do this because water does have charges on its surface. Water is polar, it has a slightly positive end and a slightly negative end, kind of like a battery.

It has these partial charges because the oxygen and hydrogen atoms in water share electrons, but they don’t share them equally. The oxygen atom hugs these electrons a little closer to itself than the hydrogen atoms are able to. So, the oxygen is a little bit negative and the hydrogens are a little bit positive.

Because of its polarity, water is really good at interacting with things that are also polar, including itself.

Pure water forms a matrix in which opposite charges on different water molecules are interacting, these interactions between the partially charged hydrogen atoms and partially charged oxygen atoms are called hydrogen bonds.

Solutions vs. Emulsions: Why Salt Dissolves

Molecules that have whole or partial charges can dissolve in water.

Salt, for example, is just a sodium cation (ions are charged molecules or atoms and a cation is a positive ion) and a chlorine anion (anion is a negative ion).

When you mix salt and water the polar nature of water means it will spontaneously form electrostatic bonds with the sodium and chlorine. Once this has happened the salt is dissolved in the water and you can’t easily get your salt back out (without crystallising it or something).

In technical terms the salt is now a solute and the water is a solvent and both of them together are a solution. Sugar, carbon dioxide and a whole bunch of other things dissolve in water like this, but some things don’t.

Solutions vs. Emulsions: Why Oil Doesn’t Dissolve

Things that aren’t polar do not spontaneously dissolve into water because the non-polar molecules don’t have charges that can interact with the charges on water.

When you add oil, a non-polar molecule, to water they don’t mix. If you add oil to water and leave it alone the mixture will arrange itself to minimise the surface area between the oil and the water.

In practical terms this means that, in the presence of gravity, the oil will form a layer on top of the water.

So to make our emulsion we need to add energy to break the bonds between water molecules and ‘force’ our non-polar molecules in amongst the water.

For our vinaigrette we add this energy by shaking or vigorously whisking the oil and the water together.

Two Types of Emulsions in the Kitchen

Once mixed, in the absence of a stabilising force, the oil and water will begin to spontaneously separate. Eventually we’ll end up right back where we started with the oil floating in a single layer on top of the water.

What is happening here is that the oil and water emulsion is moving back to it’s lower energy state.

We put energy into the system by mixing and, without any stabilising force, that energy will gradually dissipate and the oil and water will separate again.

In thermodynamics terms, our mixture is following the second law of thermodynamics by tending towards the state with the lowest free energy (I’ve got an extra credit section below that describes this in more formal terms).

Oil-in-Water (Mayonnaise, Hollandaise, Custard)

Now you may say, what about mayonnaise? Mayonnaise is an emulsion and it seems pretty stable. Whats going on there?

In most of the emulsions we make in the kitchen what we end up with is a mixture of tiny droplets of oil dispersed in water. We call this kind of emulsion an oil-in-water emulsion and we say that water is the continuous phase and oil is the dispersed phase.

The more droplets of oil you get dispersed throughout the watery continuous phase the thicker and stiffer your emulsion will be, just like mayonnaise for example.

To make these kinds of emulsions, though, you need the help of molecules called emulsifiers.

I’ll talk much more about emulsifiers in later posts, but for now it is enough to know that they make it easier to form an emulsion and will then act as a stabiliser once it has formed.

Using emulsifiers, you can defy Newton and his second law and get a stable emulsion of three parts oil dispersed into one part water.

Water-in-Oil (Vinaigrettes, Butter)

If we get back to our vinaigrette, we can now see that it is the opposite of mayonnaise, it is actually a water-in-oil emulsion.

The usual ratio of oil to water in a vinaigrette is still two or three parts oil to one part water. But without emulsifiers there is no way you are going to get that much oil into the dispersed phase.

In this scenario, the water is the dispersed phase, and it forms droplets in the oil, the continuous phase.

With much less dispersed phase the emulsion that forms is pretty runny when compared to something like mayonnaise. But this is ideal for the usual application of a vinaigrette as the dressing for a salad. The runny vinaigrette is better able to coat the relatively large surface area of the salad than a thicker, creamier oil-in-water emulsion.

Emulsion Type	Continuous Phase	Dispersed Phase	Texture / Viscosity	Common Examples
Oil-in-Water (O/W)	Water	Oil	Thick, creamy, rich	Mayonnaise, Hollandaise, Crème Anglaise, Milk
Water-in-Oil (W/O)	Oil	Water	Runny, pourable	Vinaigrettes, Butter, Chimichurri

Why Vinaigrettes Split (and How to Fix It)

As a water-in-oil emulsion with little in the way of emulsifiers, your vinaigrette is going to eventually split and separate into oil and vinegar again.

Really this isn’t much of problem, most of the time you make a vinaigrette and put it straight on the food you are saucing. But you can do a few things to stabilise your vinaigrette a bit.

Firstly, and this applies to all emulsions, you can stabilise an emulsion by generating smaller droplets in the dispersed phase. Larger droplets will make the emulsion split faster as when they collide and coalesce each collision represents a larger proportion of the dispersed phase.

Larger droplets will also rise to the surface (or drop to the bottom if water is the dispersed phase) of the emulsion more quickly, also hastening coalescence.

There are a few things you can do to get smaller droplets in your vinaigrette but using a blender to mix the emulsion is probably the easiest.

As I mentioned above emulsifiers can stabilise an emulsion. And there are plenty of aromatics that contain emulsifiers that can be added to a vinaigrette.

Both mustard and garlic, for example, contain emulsifiers that will stabilise your vinaigrette and stop it from splitting as quickly as it would without them.

My favourite vinaigrette, because it goes on steak, is chimichurri. Chimichurri is a heavily herbed vinaigrette of olive oil and vinegar with minced garlic added as an aromatic but which also contains emulsifiers.

Of course, if your vinaigrette splits all you need to do is shake it again to reform the emulsion. So I wouldn’t get too concerned with stabilising a vinaigrette, particularly if it’s going to change the flavour profile you’re aiming for.

Also, if you add enough emulsifiers you’ll end up with an oil-in-water emulsion, like mayonnaise, which may not be a texture you are looking for.

Quick Comparison: Kitchen Emulsions

What I haven’t talked about much here is flavour. A vinaigrette is really all about getting some fatty luxuriousness contrasted with some type of acid and flavoured with some aromatics.

You can make vinaigrettes with any type of oil or fat; bacon fat, olive oil and neutral oils will all taste different.

There are a multitude of acids that you can use: red vinegar, malt vinegar, lemon and lime juice and verjuice are all good options. There are also plenty of aromatics you can add: all the herbs of the world, garlic, shallots, chillies and spices.

The only real problem with vinaigrettes is option paralysis. You just need to experiment and see what you like.

Conclusion

That’s it for vinaigrettes. For something that is just a mixture of oil and water there is an awful lot of science going on (especially so if you read the discussion below). But it’s also science that we will revisit when we come to things like mayonnaise, hollandaise and custards.

We have completely ignored the effects of temperature on our emulsions and have barely touched on important topics like emulsifiers. I will get onto these topics, starting here in the second post on emulsions, but understanding the rich science behind a thing as simple as a vinaigrette will hold us in good stead when get there.

I’ve also got a section below that goes through some of the physics of emulsions if you are a sucker for punishment.

Deep Dive: Gibbs Free Energy & Emulsion Thermodynamics (Advanced)

I’ve tried to keep the discussion above reasonably free of physics, but emulsions are pretty good examples of free energy and the second law of thermodynamics.

If you aren’t interested in this skip it, the explanation above will still work for future posts, but if you want to go a bit deeper read on.

This is the science behind the statements I make like “charged molecules like hanging out together”. Molecules don’t like or dislike anything, why they group together comes down to Gibbs free energy.

Gibbs Free Energy

The solubility, or not, of a molecule in water can be described in terms of the changes in Gibbs free energy between its undissolved and dissolved states.

Gibbs free energy describes the thermodynamic potential of a system. The formula for the difference in free energy of two different states goes as follows:

dG = dH – TdS

Where delta G (dG) is the change in Gibbs free energy, delta H (dH) is the change in enthalpy, T is temperature and delta S (dS) is the change in entropy.

Enthalpy can be thought of as the internal energy of the system. Entropy can be a bit harder to grasp, but I just think of it as the number of potential microstates there are in a given system.

For example, if you increase the heat of a solution you increase the movement of molecules in the solution which increases the entropy because the increased ‘range’ of the molecules provide more potential states for the system to be in.

Likewise conformational constraints on molecules can also cause a decrease in entropy, like if you freeze water.

Spontaneous Reactions Because the Free Energy of the End State is Less Than Free Energy of the Initial State

Gibbs free energy is important because for any reaction to be spontaneous the new state of the system needs to have a Gibbs free energy that is smaller than the original state.

That is the dG going from the first state to the second needs to be negative.

The reaction will be spontaneous because the second law of thermodynamics tells us that things will tend to move towards a state that minimises free energy.

The second law also forbids spontaneous movement from a state with a lower free energy to one with a higher free energy.

So, anything that dissolves spontaneously in water must have less free energy in its dissolved state than its undissolved state.

An example

Lets consider something like calcium chloride (CaCl₂). Breaking a bond requires energy (it is endothermic) and forming a bond releases energy (it is exothermic).

When calcium chloride dissolves in water a bunch of bonds are broken and a new bunch of bonds are formed.

For calcium chloride the sum of the energy fluctuations during all this bond breaking and reforming is negative (the precise value is −80 kJ/mol at 27$\degree$C).

This means that when calcium chloride dissolves energy is released and the solution containing it becomes a little warmer.

This is the enthalpy part of the free energy equation, so in this case the value of the dH term will be negative.

The complex of water with the calcium and chlorine is more structured than that of the of undissolved calcium chloride so we actually get a negative entropy value for the reaction (-44.7JK^-1) but the heat produced by the enthalpy of the reaction also increases the entropy of the solution so the dS term ends up only slightly negative.

So, overall dG ends up negative and we get the spontaneous dissolution of calcium chloride in water.

Salt and Entropy Driven Reactions

Salt, a complex of Na⁺ and Cl^–, is interesting.

We all know that salt spontaneously dissolves in water but the enthalpy of this reaction (dH) is actually positive. The reason that salt will dissolve in water at all is that there is an increase in entropy that outweighs the positive enthalpy term.

The dissolution of the two ions of salt and the relatively unstructured complex they form with water increases the number of microstates in the solution so the entropy has increased enough to cancel out the positive dH.

This means that salt dissolving in water is an entropy driven process.

When you put salt in water it will spontaneously dissolve (overall dG is negative) but because the enthalpy is positive it will actually cause a drop in temperature of the water because it needs thermal energy from the water to power the reaction.

A Formalish Definition of Solubility

If salt and calcium chloride will spontaneously dissolve in water because the free energy change is negative then the reverse, salt and calcium chloride coming out of solution under the same conditions of temperature and pressure, will have a positive dG.

So, this will not happen unless conditions are somehow changed.

We can now define a solution as a mixture of water and a solvent that has a negative free energy of dissolution in water.

Looking at An Emulsion in Terms of Free Energy

Now let’s look at an oil or fat. To get an oil to mix with water we need to break bonds between water molecules but, because the oil is non-polar, no new bonds are formed. This means we’ll have a positive enthalpy (dH).

In plain speech, to break the bonds binding water molecules to other water molecules energy is required but we aren’t releasing any energy by forming new bonds to compensate so we just end up needing energy.

We will also have a decrease in entropy (dS) as water forms a ‘cage’ around non-polar molecules (called a clathrate¹) that restricts the movement of water molecules and the oil molecules are also constrained and can’t move around as freely as the would in the separated state.

A positive dH and a negative dS means we end up with a very positive dG for a non-polar molecule to dissolve in water.

So, it won’t happen unless we force it to happen by adding energy. Shaking or whisking the mixture does this nicely.

Conversely, once mixed the dG for moving to the separated state will be very negative meaning that, unless stabilised, our emulsion will spontaneously split as it moves towards a state with a lower free energy.

It’s the second law of thermodynamics in action.

Footnotes

1. OK – I know I haven’t included the hydrogen bonds between water molecules forming the clathrate when considering the enthalpy, but this post is already way too long. Also people are still publishing papers trying to work out the details of the hydrophobic effect so I’m leaving it like this for now. ↩︎