A calorimeter is a device used to measure the heat flow of a chemical reaction or physical change. The process of measuring this heat is called calorimetry. A basic calorimeter consists of a metal container of water above a combustion chamber, in which a thermometer is used to measure the change in water temperature. However, there are many types of more complex calorimeters. The basic principle is that heat released by the combustion chamber increases the temperature of the water in a measurable way. The temperature change may then be used to calculate the enthalpy change per mole of substance A when substances A and B are reacted. The equation used is: q = Cv(Tf - Ti ) where:
The first ice calorimeters were built based on Joseph Black's concept of latent heat, introduced in 1761. Antoine Lavoisier coined the term calorimeter in 1780 to describe the apparatus he used to measure heat from guinea pig respiration used to melt snow. In 1782, Lavoisier and Pierre-Simon Laplace experimented with ice calorimeters, in which heat needed to melt ice could be used to measure heat from chemical reactions. Calorimeters have expanded beyond the original ice calorimeters.
Updated April 11, 2018 By Claire Gillespie
A calorimeter might sound like a fancy piece of science equipment, but it's actually a very simple heat-measuring device that you can make at home using two coffee cups. Often used in science project experiments, it measures the amount of heat involved in a chemical or physical process, such as heat transfer or specific heat of a substance.
The temperature of liquid changes when it gains or loses energy. A calorimeter measures the mass of liquid and the temperature change of the liquid to determine the quantity of energy gained or lost by the liquid.
A calorimeter has two vessels: an outer vessel and an inner vessel. The air between the two vessels acts as a heat insulator, meaning there is no (or minimal) heat exchange between what is inside the inner vessel and the outside environment. Calorimeters used in science labs have a fiber ring made of insulating material to hold the inner vessel in the center of the outer vessel. They include a thermometer to measure the temperature of the liquid in the inner vessel and a stirrer to stir the liquid and distribute heat throughout the vessel. It's easy to make a calorimeter at home with polystyrene cups, a cover, a thermometer and a stirrer. However, a "coffee cup" calorimeter allows more heat exchange with its surroundings and produces less accurate results.
If an exothermic reaction (a chemical reaction that releases energy by light or heat) happens in a solution in a calorimeter, the solution uses heat, which raises its temperature. If an endothermic reaction (a reaction that absorbs energy from its surroundings) happens, the solution loses heat, which lowers its temperature. The difference in temperature, together with the specific heat and mass of the solution, allows you to work out how much heat the reaction uses. For example, if you place a hot piece of copper in a cool amount of water within a calorimeter, heat will flow from the copper to the water. The temperature of the copper will decrease, and the temperature of the water will increase until they have the same temperature (thermal equilibrium). You don't gain or lose heat during the process because the calorimeter allows all heat transfer to occur between the two substances.
Specific heat is the amount of energy necessary to produce a temperature change of 1 degree Celsius per gram of substance, and it varies between substances. For example, the specific heat of water is 1.00 calorie/gram degrees Celsius. To determine the specific heat of an unknown metal, place a heated piece of the metal in water in the inner vessel of the calorimeter. Once you have measured the final temperature of both the metal and the water, such as the highest temperature reached by the water, you can work out the specific heat of the metal. First, multiply the mass of water by the specific heat of water by the temperature change of water, then multiply the mass of metal by the temperature change of metal. Divide your first answer by your second answer to establish the specific heat of the metal. Calorimetry is the process of measuring the amount of heat released or absorbed during a chemical reaction. By knowing the change in heat, it can be determined whether or not a reaction is exothermic (releases heat) or endothermic (absorbs heat). Calorimetry also plays a large part of everyday life, controlling the metabolic rates in humans and consequently maintaining such functions like body temperature.
Learning Outcomes
One technique we can use to measure the amount of heat involved in a chemical or physical process is known as calorimetry. Calorimetry is used to measure amounts of heat transferred to or from a substance. To do so, the heat is exchanged with a calibrated object (calorimeter). The temperature change measured by the calorimeter is used to derive the amount of heat transferred by the process under study. The measurement of heat transfer using this approach requires the definition of a system (the substance or substances undergoing the chemical or physical change) and its surroundings (all other matter, including components of the measurement apparatus, that serve to either provide heat to the system or absorb heat from the system). A calorimeter is a device used to measure the amount of heat involved in a chemical or physical process. For example, when an exothermic reaction occurs in solution in a calorimeter, the heat produced by the reaction is absorbed by the solution, which increases its temperature. When an endothermic reaction occurs, the heat required is absorbed from the thermal energy of the solution, which decreases its temperature (Figure 1). The temperature change, along with the specific heat and mass of the solution, can then be used to calculate the amount of heat involved in either case. Scientists use well-insulated calorimeters that all but prevent the transfer of heat between the calorimeter and its environment, which effectively limits the “surroundings” to the nonsystem components of the calorimeter (and the calorimeter itself.) This enables the accurate determination of the heat involved in chemical processes, the energy content of foods, and so on. General chemistry students often use simple calorimeters constructed from polystyrene cups (Figure 2). These easy-to-use “coffee cup” calorimeters allow more heat exchange with their surroundings, and therefore produce less accurate energy values. Commercial solution calorimeters are also available. Relatively inexpensive calorimeters often consist of two thin-walled cups that are nested in a way that minimizes thermal contact during use, along with an insulated cover, handheld stirrer, and simple thermometer. More expensive calorimeters used for industry and research typically have a well-insulated, fully enclosed reaction vessel, motorized stirring mechanism, and a more accurate temperature sensor (Figure 3). Before we practice calorimetry problems involving chemical reactions, consider a simpler example that illustrates the core idea behind calorimetry. Suppose we initially have a high-temperature substance, such as a hot piece of metal (M), and a low-temperature substance, such as cool water (W). If we place the metal in the water, heat will flow from M to W. The temperature of M will decrease, and the temperature of W will increase, until the two substances have the same temperature—that is, when they reach thermal equilibrium (Figure 4). If this occurs in a calorimeter, ideally all of this heat transfer occurs between the two substances, with no heat gained or lost by either the calorimeter or the calorimeter’s surroundings. Under these ideal circumstances, the net heat change is zero: [latex]{q}_{\text{ substance M}}+{q}_{\text{ substance W}}=0[/latex] This relationship can be rearranged to show that the heat gained by substance M is equal to the heat lost by substance W: [latex]{q}_{\text{ substance M}}=-{q}_{\text{ substance W}}[/latex] The magnitude of the heat (change) is therefore the same for both substances, and the negative sign merely shows that [latex]{q}_{\text{ substance M}}[/latex] and [latex]{q}_{\text{ substance W}}[/latex] are opposite in direction of heat flow (gain or loss) but does not indicate the arithmetic sign of either q value (that is determined by whether the matter in question gains or loses heat, per definition). In the specific situation described, [latex]{q}_{\text{ substance M}}[/latex] is a negative value and [latex]{q}_{\text{ substance W}}[/latex] is positive, since heat is transferred from M to W.
A 360-g piece of rebar (a steel rod used for reinforcing concrete) is dropped into 425 mL of water at 24.0 °C. The final temperature of the water was measured as 42.7 °C. Calculate the initial temperature of the piece of rebar. Assume the specific heat of steel is approximately the same as that for iron, and that all heat transfer occurs between the rebar and the water (there is no heat exchange with the surroundings). Check Your LearningA 248-g piece of copper is dropped into 390 mL of water at 22.6 °C. The final temperature of the water was measured as 39.9 °C. Calculate the initial temperature of the piece of copper. Assume that all heat transfer occurs between the copper and the water. Check Your LearningA 248-g piece of copper initially at 314 °C is dropped into 390 mL of water initially at 22.6 °C. Assuming that all heat transfer occurs between the copper and the water, calculate the final temperature. This method can also be used to determine other quantities, such as the specific heat of an unknown metal.
A 59.7 g piece of metal that had been submerged in boiling water was quickly transferred into 60.0 mL of water initially at 22.0 °C. The final temperature is 28.5 °C. Use these data to determine the specific heat of the metal. Use this result to identify the metal. Check Your LearningA 92.9-g piece of a silver/gray metal is heated to 178.0 °C, and then quickly transferred into 75.0 mL of water initially at 24.0 °C. After 5 minutes, both the metal and the water have reached the same temperature: 29.7 °C. Determine the specific heat and the identity of the metal. (Note: You should find that the specific heat is close to that of two different metals. Explain how you can confidently determine the identity of the metal). When we use calorimetry to determine the heat involved in a chemical reaction, the same principles we have been discussing apply. The amount of heat absorbed by the calorimeter is often small enough that we can neglect it (though not for highly accurate measurements, as discussed later), and the calorimeter minimizes energy exchange with the surroundings. Because energy is neither created nor destroyed during a chemical reaction, there is no overall energy change during the reaction. The heat produced or consumed in the reaction (the “system”), [latex]{q}_{\text{ reaction}}[/latex], plus the heat absorbed or lost by the solution (the “surroundings”),[latex]{q}_{\text{ solution}}[/latex], must add up to zero: [latex]{q}_{\text{ reaction}}+{q}_{\text{ solution}}=0[/latex] This means that the amount of heat produced or consumed in the reaction equals the amount of heat absorbed or lost by the solution: [latex]{q}_{\text{ reaction}}=\text{-}{q}_{\text{ solution}}[/latex] This concept lies at the heart of all calorimetry problems and calculations.
When 50.0 mL of 0.10 M HCl(aq) and 50.0 mL of 0.10 M NaOH(aq), both at 22.0 °C, are added to a coffee cup calorimeter, the temperature of the mixture reaches a maximum of 28.9 °C. What is the approximate amount of heat produced by this reaction [latex]\text{HCl}\left(aq\right)+\text{NaOH}\left(aq\right)\rightarrow\text{NaCl}\left(aq\right)+{\text{H}}_{\text{2}}\text{O}\left(l\right)[/latex] Check Your LearningWhen 100 mL of 0.200 M NaCl(aq) and 100 mL of 0.200 M AgNO3(aq), both at 21.9 °C, are mixed in a coffee cup calorimeter, the temperature increases to 23.5 °C as solid AgCl forms. How much heat is produced by this precipitation reaction? What assumptions did you make to determine your value?
When working or playing outdoors on a cold day, you might use a hand warmer to warm your hands (Figure 5). A common reusable hand warmer contains a supersaturated solution of NaC2H3O2 (sodium acetate) and a metal disc. Bending the disk creates nucleation sites around which the metastable NaC2H3O2 quickly crystallizes (eventually we will investigate saturation and supersaturation in more detail). The process [latex]{\text{NaC}}_{\text{2}}{\text{H}}_{\text{3}}{\text{O}}_{\text{2}}\left(aq\right)\rightarrow{\text{NaC}}_{\text{2}}{\text{H}}_{\text{3}}{\text{O}}_{\text{2}}\left(s\right)[/latex] is exothermic, and the heat produced by this process is absorbed by your hands, thereby warming them (at least for a while). If the hand warmer is reheated, the NaC2H3O2 redissolves and can be reused. Another common hand warmer produces heat when it is ripped open, exposing iron and water in the hand warmer to oxygen in the air. One simplified version of this exothermic reaction is [latex]2\text{Fe}\left(s\right)+\dfrac{3}{2}{\text{O}}_{\text{2}}\left(g\right)\rightarrow{\text{Fe}}_{\text{2}}{\text{O}}_{\text{3}}\left(s\right)[/latex]. Salt in the hand warmer catalyzes the reaction, so it produces heat more rapidly; cellulose, vermiculite, and activated carbon help distribute the heat evenly. Other types of hand warmers use lighter fluid (a platinum catalyst helps lighter fluid oxidize exothermically), charcoal (charcoal oxidizes in a special case), or electrical units that produce heat by passing an electrical current from a battery through resistive wires. This Wikimedia video shows the precipitation reaction that occurs when the disk in a chemical hand warmer is flexed.
When solid ammonium nitrate dissolves in water, the solution becomes cold. This is the basis for an “instant ice pack” (Figure 6). When 3.21 g of solid NH4NO3 dissolves in 50.0 g of water at 24.9 °C in a calorimeter, the temperature decreases to 20.3 °C. Calculate the value of q for this reaction and explain the meaning of its arithmetic sign. State any assumptions that you made. Check Your LearningWhen a 3.00-g sample of KCl was added to 3.00 × 102 g of water in a coffee cup calorimeter, the temperature decreased by 1.05 °C. How much heat is involved in the dissolution of the KCl? What assumptions did you make? If the amount of heat absorbed by a calorimeter is too large to neglect or if we require more accurate results, then we must take into account the heat absorbed both by the solution and by the calorimeter. The calorimeters described are designed to operate at constant (atmospheric) pressure and are convenient to measure heat flow accompanying processes that occur in solution. A different type of calorimeter that operates at constant volume, colloquially known as a bomb calorimeter, is used to measure the energy produced by reactions that yield large amounts of heat and gaseous products, such as combustion reactions. (The term “bomb” comes from the observation that these reactions can be vigorous enough to resemble explosions that would damage other calorimeters.) This type of calorimeter consists of a robust steel container (the “bomb”) that contains the reactants and is itself submerged in water (Figure 7). The sample is placed in the bomb, which is then filled with oxygen at high pressure. A small electrical spark is used to ignite the sample. The energy produced by the reaction is trapped in the steel bomb and the surrounding water. The temperature increase is measured and, along with the known heat capacity of the calorimeter, is used to calculate the energy produced by the reaction. Bomb calorimeters require calibration to determine the heat capacity of the calorimeter and ensure accurate results. The calibration is accomplished using a reaction with a known q, such as a measured quantity of benzoic acid ignited by a spark from a nickel fuse wire that is weighed before and after the reaction. The temperature change produced by the known reaction is used to determine the heat capacity of the calorimeter. The calibration is generally performed each time before the calorimeter is used to gather research data.
Watch this video on how a bomb calorimeter is prepared for action. You can view the transcript for “Physical Chemistry iBook – Bomb Calorimetry” here (opens in new window). The Oxygen Bomb Calorimeter website shows calorimetric calculations using sample data.
When 3.12 g of glucose, C6H12O6, is burned in a bomb calorimeter, the temperature of the calorimeter increases from 23.8 °C to 35.6 °C. The calorimeter contains 775 g of water, and the bomb itself has a heat capacity of 893 J/°C. How much heat was produced by the combustion of the glucose sample? Check Your LearningWhen 0.963 g of benzene, C6H6, is burned in a bomb calorimeter, the temperature of the calorimeter increases by 8.39 °C. The bomb has a heat capacity of 784 J/°C and is submerged in 925 mL of water. How much heat was produced by the combustion of the glucose sample? Since the first one was constructed in 1899, 35 calorimeters have been built to measure the heat produced by a living person. These whole-body calorimeters of various designs are large enough to hold an individual human being. More recently, whole-room calorimeters allow for relatively normal activities to be performed, and these calorimeters generate data that more closely reflect the real world. These calorimeters are used to measure the metabolism of individuals under different environmental conditions, different dietary regimes, and with different health conditions, such as diabetes. In humans, metabolism is typically measured in Calories per day. A nutritional calorie (Calorie) is the energy unit used to quantify the amount of energy derived from the metabolism of foods; one Calorie is equal to 1000 calories (1 kcal), the amount of energy needed to heat 1 kg of water by 1 °C.
In your day-to-day life, you may be more familiar with energy being given in Calories, or nutritional calories, which are used to quantify the amount of energy in foods. One calorie (cal) = exactly 4.184 joules, and one Calorie (note the capitalization) = 1000 cal, or 1 kcal. (This is approximately the amount of energy needed to heat 1 kg of water by 1 °C.) The macronutrients in food are proteins, carbohydrates, and fats or oils. Proteins provide about 4 Calories per gram, carbohydrates also provide about 4 Calories per gram, and fats and oils provide about 9 Calories/g. Nutritional labels on food packages show the caloric content of one serving of the food, as well as the breakdown into Calories from each of the three macronutrients (Figure 8). For the example shown in (b), the total energy per 228-g portion is calculated by: [latex]\left(5\text{ g protein}\times 4\text{ Calories/g}\right)+\left(31\text{ g carb}\times 4\text{ Calories/g}\right)+\left(12\text{ g fat}\times 9\text{ Calories/g}\right)=252\text{ Calories}[/latex] So, you can use food labels to count your Calories. But where do the values come from? And how accurate are they? The caloric content of foods can be determined by using bomb calorimetry; that is, by burning the food and measuring the energy it contains. A sample of food is weighed, mixed in a blender, freeze-dried, ground into powder, and formed into a pellet. The pellet is burned inside a bomb calorimeter, and the measured temperature change is converted into energy per gram of food. Today, the caloric content on food labels is derived using a method called the Atwater system that uses the average caloric content of the different chemical constituents of food, protein, carbohydrate, and fats. The average amounts are those given in the equation and are derived from the various results given by bomb calorimetry of whole foods. The carbohydrate amount is discounted a certain amount for the fiber content, which is indigestible carbohydrate. To determine the energy content of a food, the quantities of carbohydrate, protein, and fat are each multiplied by the average Calories per gram for each and the products summed to obtain the total energy.
Calorimetry is used to measure the amount of thermal energy transferred in a chemical or physical process. This requires careful measurement of the temperature change that occurs during the process and the masses of the system and surroundings. These measured quantities are then used to compute the amount of heat produced or consumed in the process using known mathematical relations. Calorimeters are designed to minimize energy exchange between the system being studied and its surroundings. They range from simple coffee cup calorimeters used by introductory chemistry students to sophisticated bomb calorimeters used to determine the energy content of food.
Glossarybomb calorimeter: device designed to measure the energy change for processes occurring under conditions of constant volume; commonly used for reactions involving solid and gaseous reactants or products calorimeter: device used to measure the amount of heat absorbed or released in a chemical or physical process calorimetry: process of measuring the amount of heat involved in a chemical or physical process nutritional calorie (Calorie): unit used for quantifying energy provided by digestion of foods, defined as 1000 cal or 1 kcal surroundings: all matter other than the system being studied system: portion of matter undergoing a chemical or physical change being studied |