Specific Heat of Unknown Metal
🔒 LAB RESOURCES:
Student Procedure [PDF] [DOC]
Teacher Annotated Procedure [PDF]
Complete Lab Guide [HERE]
GoogleSheet Data & Analysis [HERE]
SUMMARY/ OVERVIEW:
Students determine the specific heat capacity (c) of an unknown metal. A known mass of metal is heated to a measurable temperature (e.g., boiling water) and quickly transferred to a calorimeter containing water of known mass and temperature. By measuring the final equilibrium temperature, the heat transfer (q) is calculated, allowing for the determination of the metal's specific heat using $q=mc\Delta T$.
ESTIMATE TIME ⏰: 40-60 minutes
Initial Setup and Weighing (15-20 minutes)
Heating and Equilibration (30-45 minutes)
Transfer and Measurement (15-20 minutes)
SAFETY PRECAUTIONS:
Eye Protection is Mandatory: Always wear approved safety goggles throughout the entire experiment.
Burn Hazards: The metal sample will be heated in boiling water (near $100^\circ\text{C}$). The water, metal, and beaker will be extremely hot.
Prevent Splashing: Pour the hot water carefully. When quickly transferring the hot metal into the cooler water in the calorimeter, ensure the calorimeter is stable and do not splash the hot water onto yourself or others.
Thermometer Care: Handle thermometers gently. Do not use them as stirring rods. If a mercury thermometer is used (less common now), be immediately report any breakage to your instructor.
Water Disposal: Once the lab is complete, allow the water to cool significantly before disposing of it down the sink.
Background
The method used to determine the unknown metal's specific heat is calorimetry, the science of measuring heat transfer. This lab applies the Law of Conservation of Energy, which states that energy cannot be created or destroyed.
The Principle
When the hot metal is placed into the cooler water inside the calorimeter, the heat lost by the metal is equal to the heat gained by the water (assuming the calorimeter itself absorbs a negligible amount of heat):
$$q_{\text{metal}} = -q_{\text{water}}$$
The Working Equation
The heat transferred ($q$) is calculated using the following relationship:
$$q = m \cdot c \cdot \Delta T$$
Since the heat gained by the water equals the heat lost by the metal, we can set up the calculation as follows:
$$(m \cdot c \cdot \Delta T)_{\text{metal}} = - (m \cdot c \cdot \Delta T)_{\text{water}}$$
By measuring the masses (m), the initial and final temperatures ($\Delta T$), and using the known specific heat of water ($c_{\text{water}} \approx 4.184\ \text{J/g}\cdot^\circ\text{C}$), the only unknown left is the specific heat capacity of the metal ($c_{\text{metal}}$), which can then be solved algebraically.
IExperiment Reactions
As mentioned, chemical kinetics measures how fast a reaction is occurring. For most chemical reactions, the rate is so fast that special equipment is needed to measure it. For the iodine clock reaction, on the other hand, the rate can be easily measured by monitoring the color change of the reaction. To perform the iodine clock reaction, you will mix potassium iodide, sulfuric acid, starch, and thiosulfate. The time it takes for the reaction mix to turn blue will be measured with a timer.
The reactions that form the basis for the iodine clock reaction are shown below.
Equation 1
IO3– (aq) + HSO3– (aq) → I– (aq) + H+ (aq) + SO42– (aq)
Equation 2
H+(aq) + I–(aq) + IO3–(aq) → I2(aq) + H2O(l)
Equation 3
I2 (aq) + HSO3– (aq) + H2O (l) → I– (aq) + SO42–(aq) + H+(aq)
Equation 4
I2(aq) + Starch → Dark-blue colored complex
This experiment involves a reaction that is sometimes called an iodine clock reaction. There are a number of different combinations of chemicals that give a reaction of this type. What happens, essentially, is that there are two different reactions: one in which iodine is produced (a slow reaction) and one in which the iodine produced in the first reaction is used up (a fast reaction). By carefully controlling the quantities of reactants, you can obtain a situation in which the reactant in the second reaction is used up first, allowing iodine to form at that point. At very low concentrations the iodine then combines with starch to suddenly give a deep blue-black color, at a time determined by the conditions used. Hence the term “iodine clock”. The time elapsed from when the solutions were first mixed together until the point when the blue-black color appears is measured, and from this time measurement the rate of the reaction can be determined. You will alter the conditions of concentrations of reactants in Part I and temperature in Part II in order to determine their effect on reaction rate.