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Alton Brown’s Jet Cream Ice Cream

Because you are currently reading a blog about science and food, there is a high probability that you have seen or at least heard of Alton Brown: host of Good Eats and about five other Food Network television shows. There is also a significant probability that you’re a mega-fan of Alton Brown, and if so, that’s something you and I have in common. I have been watching the bespectacled nerd-chef (I say that admiringly) since I was thirteen, and he has largely inspired my food science endeavors. On March 19th I had the absolute pleasure of attending Alton Brown Live! The Incredible Inevitable Tour in Napa, California.

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Alton describes the show content as all the things he can’t do on TV, including stand-up comedy, live music, and most excitingly, showing off his insane kitchen inventions. Because even the thought of burning myself on MegaBake terrifies me, we’re going to work through the science behind his colder contraption: Jet Cream.

Making ice cream is usually a simple process. Once you have your cream mixture, it simply needs to be repeatedly cooled and agitated. If we simply froze ice cream base, we’d get huge ice crystals, which aren’t necessarily bad. Dessert shops like Blockheads and Chilly Ribbons sell “Snow Cream,” that results from shaving fine sheets from a block of frozen milk or cream. But if we want ice cream, as Alton clearly does, we must continually add air to the cream and disrupt the crystallization process to make tiny crystals that are barely perceivable on the tongue. That’s why ice cream is smooth and unctuous, while frozen milk is crisp and icy. Whether you’re shaking a container of cream surrounded by ice by hand or using an industrial ice cream machine, the goal is to keep ice crystals small.

Alton’s goal is no different. To make ice cream, all he needs to do is simultaneously freeze and agitate his chocolate cream. His Jet Cream machine is an extravagant way to do a huge batch all at once, and in less than ten seconds. Rather than use ice and salt in a bason like pioneers did, or use liquid nitrogen like the modern gastronome, he uses compressed carbon dioxide via fire extinguisher.

When the fast-flying molecules of carbon dioxide gas are compressed into the extinguisher, they are stored at a very high pressure, typically 825 pounds per square inch. [1] A fire erupts on the stove, or you have a sudden urge for ice cream, so you pull the lever. The pressure is released; the gas flies out, and the nozzle and surrounding air become extremely cold, as tends to happen when a  gas suddenly expands from a high pressure to a low pressure. The change in temperature divided by the change in pressure makes a ratio (∆T/∆P) known as the Joule-Thomson coefficient.[2] The nozzle and surrounding air are chilled because the gas’ pressure change occurs too quickly for significant heat transfer to occur. For many gases at room temperature, as the CO2 in the extinguisher is, the ∆T/∆P ratio is positive, so a pressure drop is accompanied by a temperature drop. The molecules that were once speeding around inside the canister are now so low-energy that they form solid CO2, or dry ice. Dry ice is much, much colder than regular H2O ice because carbon dioxide freezes at -109 degrees Fahrenheit, while water freezes at 32 degrees. [3] Colder temperature = faster crystallization = quicker ice cream.

Photo Credit: David Allen, The Eater

Photo Credit: David Allen, The Eater

Now for the agitation: At the other end of Alton’s Jet Cream contraption is a typical water fire-extinguisher filled with chocolate cream. When this lever is pulled, a high-pressure spray of chocolate ensues. Between the two extinguishers are office water cooler jugs that act as the reaction chamber for the CO2 and cream. If the two levers are pulled exactly at the same time (synchronicity is very important in avoiding a catastrophic mess, stresses Alton), the blasts of cold and cream will collide in the coolers, providing the continual disturbance of the freezing process, as well as the incorporation of air, necessary to make tiny tasty ice crystals.

After plunking a scoop into a sugar cone and applying a generous coat of rainbow sprinkles, Alton hands off his creation to his volunteer assistant and asks if it is not the best ice cream he has ever had. Volunteer assistant replies that it is “So good.”

So there you have it. If you want ice cream that is “so good,” and you want a gallon of it fast, Jet Cream is the contraption for you.

 

Photo credits: instagram.com/altonbrown/

Photo credit: instagram.com/altonbrown/

References cited:

  1. “CO2 Fire Extinguishers.” Fire Extinguisher Guide. N.p., n.d. Web. 06 Apr. 2015.
  2. Joule Thomson Effect.” Wright State University – Department of Chemistry.  N.p., n.d. Web. 06 Apr. 2015
  3. “UCSB Science Line.” UCSB Science Line. N.p., n.d. Web. 06 Apr. 2015.

Elsbeth SitesAbout the author: Elsbeth Sites is pursuing her B.S. in Biology at UCLA. Her addiction to the Food Network has developed into a love of learning about the science behind food. Read more by Elsbeth Sites


Food Science Careers & Bad Apples

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Kirsten Schimoler at Ben & Jerry’s talks about food science as a career and scientists figure out how to keep produce healthier.
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Homemade Ice Cream

Phase transitions—transformations from one state of matter to another—are ubiquitous in food and cooking. Butter’s phase transition from a solid to a liquid results in flaky pie crusts, while water’s phase transition from a liquid to a gas can be used to steam vegetables. There are various ways to manipulate these phase transitions, such as by altering temperature, pressure, or salt content. In this classic home experiment, we will make ice cream by using salt to alter the phase behavior of water.


Objectives

  • Understand how solutes (salt) affect the phase behavior of a solvent (water).
  • Use freezing point depression to make a batch of amazing ice cream.


Materials

  • 1 cup cream
  • 1/2 cup sugar
  • 200 grams ice
  • Kosher salt
  • 1 quart Ziploc bag
  • 1 gallon Ziploc bag
  • Thermometer
  • Scale


Part 1: Use salt to lower the melting point of ice

To successfully freeze ice cream without the help of a freezer, we need a way to efficiently transfer heat out of the ice cream. Liquid water is much better than solid ice at transferring heat, so an ice-water bath will absorb heat from our ice cream better than solid ice. To effectively freeze ice cream, however, we need stable temperatures well below 0˚C.

How is it possible to have a mixture of water and ice at a temperature below 0˚C, water’s freezing point?

When you take ice straight out of the freezer, the ice will be roughly the same temperature as the freezer itself. The temperature in a home freezer is typically between 0˚C and -20˚C. As the ice sits out, it will absorb heat from its surroundings and slowly get warmer until it reaches 0˚C and begins to melt. Adding impurities like salt to ice will lower its melting point.  This means that salted ice will start melting at temperatures below 0˚C. As a result, a salty ice-water bath can stay liquid at temperatures well below 0˚C and efficiently freeze our ice cream. We refer to this phenomenon as “freezing point depression.”

We can use the freezing point depression equation to calculate how much a solvent’s freezing point will drop as a solute is added:

∆Tf = b · Kf  · i

∆T    Freezing point depression, defined as Tf of  pure solvent – Tf of solution.
K f        Cryoscopic constant of the solvent. This is an intrinsic property of the solvent.
b          Molar concentration of the solute: the number of moles of solute per kilogram of solvent.
i           Number of ion particles per molecule of solute, also known as the “Van’t Hoff factor”.
Salt is made up of one sodium ion and one chloride ion, so its Van’t Hoff factor is 2.

  1. Use the freezing point depression equation to calculate how much salt (our solute) is needed to decrease the freezing point of water (our solvent) from (a) 0˚C to -5˚C, (b) 0˚C to -10˚C, (c) 0˚C to -15˚C, and (d) 0˚C to -20˚C.
  2. Plot the magnitude of freezing point depression (ΔTf) versus salt concentration (Results from 1a, b, c, and d). Remember to use units!
  3. Based on your answer from 1d, calculate how many grams of salt are required to create a -20˚C freezing point depression for 200g of ice. This is the amount of salt you will use in Part 2.

Some useful values:
Freezing point (Tf) for pure water: 0˚C.
Cryoscopic constant (Kf) for water: 1.853 ˚C*kg/mol.
Molecular weight of salt (NaCl): 58.44 g/mol.

Click here to check your answers.


Part 2: Use freezing point depression to make ice cream

  1. Combine cream and sugar in the quart-size bag and mix well. Place this bag inside the gallon bag.
  2. Record the initial temperatures of the ice and the cream mixture.
  3. In the gallon bag, pack the ice around the quart-size bag, and then sprinkle the calculated amount of salt over the ice. Be careful that the salt does not fall into the cream mixture.
  4. Gently shake the bag until the cream mixture solidifies into ice cream.
  5. Record the final temperatures of the ice-salt-water mixture and the ice cream.
  6. Enjoy your homemade ice cream!


Questions

  • What was the final temperature of the ice cream? Did it end up below 0˚C? How does its temperature compare to the temperature of the salt-ice-water mixture?
  • What was the final temperature of the ice-salt-water mixture? Is warmer or colder than the ice you started with? How does the temperature compare to the freezing point depression you calculated in Part 1?


Discussion

In this experiment, we used salt to lower the freezing point of water. By adding salt to ice, we were able to achieve a salt-ice-water mixture that was able to freeze our ice cream.

Why does ice cream need temperatures colder than the freezing point of water in order to freeze?

When water freezes, it forms a well-ordered crystalline structure (an ice cube). This unique crystalline structure is what gives solid water a slightly lighter density. Although ice cream is a combination of  cream, sugar, and flavorings, it is still approximately 60% water. The remaining 40% is a mixture of sugar molecules, fat globules, and milk proteins [1]. This liquid mixture is emulsified: the water molecules are dispersed among sugar molecules, milk protein complexes, and large clusters of fat globules.. When this mixture is brought to the freezing temperature of water, the fats, proteins, and sugars hamper the freezing process by interrupting the formation of ordered crystal water structures. The ice cream mixture thus remains a liquid, requiring even colder temperatures below 0˚C to successfully solidify [2].

Structure of ice cream. (A) an electron micrograph of ice cream showing air bubbles, ice crystals, and the sugar solution [3]. Fat globules and milk proteins are not visible at this resolution. (B) Diagram of ice cream structure adapted from University of Guelph.

How did the salt in our experiment create a salt-ice-water mixture below 0˚C?

At 0˚C, ice and water are “at equilibrium” with each other. The total amount of water and ice remains relatively constant, but individual water molecules are constantly switching states: as some water molecules melt and become liquid, other water molecules freeze and become solid. Adding a solute like salt shifts this equilibrium. Solutes essentially “trap” water molecules in the liquid state, preventing them from readily switching back to the solid state. On a macroscopic scale, salt causes solid ice to melt faster and at temperatures below 0˚C, resulting in a salt-ice-water mixture below 0˚C. To get a better feel for how this process works at the molecular level, check out this interactive demonstration of how temperature and solutes affect the water-ice equilibrium.

Contrary to popular belief, the addition of salt to ice does not actually make the ice any colder!

The temperature that you recorded for the salt-ice-water mixture was probably colder than the temperature of the pure ice you started with. How is this possible? When you take the temperature of solid ice, you are not really measuring the temperature of the ice itself—you are measuring the average temperature of the ice, the air around the ice, and any water that has formed from the ice melting. The true temperature of the ice depends on the temperature  freezer it came from (typically between 0˚C and -20˚C) and the length of time the ice has spent out of the freezer.


Online Resources

  1. Interactive explanation of how temperature and solutes affect water-ice equilibrium
  2. “Ice Cream Structure” from University of Guelph

More from On Food and Cooking

  • McGee, Harold. On Food and Cooking. Scribner, 2004. (39–44).

References Cited

  1. Goff HD (1997) Colloidal aspects of ice cream—A review. International Dairy Journal 7: 363–373. doi:10.1016/S0958-6946(97)00040-X.
  2. Hartel RW (1996) Ice crystallization during the manufacture of ice cream. Trends in Food Science & Technology 7: 315–321. doi:10.1016/0924-2244(96)10033-9.
  3. Clarke C (2003) The physics of ice cream. Physics Education 38: 248–253. doi:10.1088/0031-9120/38/3/308.

Liz Roth-JohnsonAbout the author: Liz Roth-Johnson is a Ph.D. candidate in Molecular Biology at UCLA. If she’s not in the lab, you can usually find her experimenting in the kitchen.

Read more by Liz Roth-Johnson