What is compost, and is it possible for a compost pile to spontaneously combust?

Before answering this initial question, define the following terms below:

The process of decomposing was mentioned in the last video.

We've looked at variables that could lead to spontaneous combustion and discussed decomposition, but what is occurring on the microscopic level?

You will observe and record the temperature change when dissolving an ionic compound in water. Recall that an ionic compound is composed of cations and anions in specific proportions to create a neutral ionic compound.

You will need the following resources to complete this assignment.

• HW Modeling Breaking Bonds

1. Record which ionic compound you are using in the table below.
2. Use a weighing boat to measure 1 gram of the ionic compound.
3. Use a graduated cylinder to measure 5 mL of water.
4. Record the initial temperature of the water.
5. In a small beaker, combine the water and ionic compound.
6. Place the temperature probe into the beaker and observe. Record the greatest difference in temperature from the starting temperature of the water/ionic compound mixture.
7. Dispose of water/ionic compound mixture in a sink and rinse out the beaker.
8. Repeat steps 1-7 with a total of five different ionic compounds.

In order to rank the five ionic compounds according to the temperature change in 5 mL of water, we will use an online data analysis platform called CODAP. Below, you will see a CODAP workbench. CODAP will allow us to visualize our data immediately upon entering it into the platform.

Let's begin with familiarizing ourselves with the CODAP environment:

1. The platform is below these instructions. Notice an example ionic compound (NaCl) has been entered for you. Enter your data for the five ionic compounds you tested below this example.
2. To better visualize the data, we can make a graph. In order to create a graph, drag the Ionic Compound variable to the x-axis of the plot as shown in the GIF on the right.
3. Drag the ∆T/gram variable to the y-axis of the plot as shown in the GIF on the right.
4. Observe the resulting plot. Then answer the questions below the CODAP workbench.

In order to convert our data from ∆T/gram to ∆T/mole, we need to find the molar mass of each compound. Then we can convert each ionic compound's mass from grams to moles. Finally, we would need to take our calculated temperature change and divide that by our moles, but that is so much work. What if we could make CODAP do some of that for us?

1. As before, the platform is below these instructions. The sodium chloride example has been expanded to include molar mass, moles, and ∆T/mole.
2. Calculate the molar mass for each ionic compound and enter it in the platform below.
3. CODAP can now calculate moles for you if you provide it with the correct formula. Hold the cursor over moles as seen in the gif at right and click on Edit Formula. Use the asterisk (*) for multiplication and the slash (/) for division.
4. Click Apply to close the formula editor. You should see values appear in the moles column.
5. Now follow the same procedure to calculate values for the ∆T/mole column. Again to better visualize the data, we can make a graph.
6. Drag the Ionic Compound variable to the x-axis of the plot.
7. Drag the ∆T/mole variable to the y-axis of the plot.
8. Observe the resulting plot. Then answer the questions below the CODAP workbench.

Last unit we explored how the amount of reactant will affect the amount of product produced. We want to expand upon this knowledge and explore how the amount of reactant affects other factors in chemical reaction. Read through the protocol below and determine which ionic compound you will use and the amounts of that compound you will use for each trial before starting your investigation.

Protocol Investigation 2.A – Vary the amount of ionic compound

• You will need to choose one ionic compound to test, from the ionic compounds tested in Investigation Part 1.
• You will need to vary the amount of ionic compound, staying within the 1-3 gram range.
• You will need to keep the amount of water constant, at 5 mL.
• You will need to record the initial temperature of each individually as well as the greatest difference in temperature from the starting temperature of the water/ionic compound mixture.
• Complete three trials.

Use the same ionic compound from Investigation 2.A for Investigation 2.B. Read through the protocol below and determine the amounts of water you will use for each trial before starting your investigation.

Protocol Investigation 2.B – Vary the amount of water

• You will need to use the same ionic compound used in the first part of Investigation 2.
• You will need to vary the amount of water, staying within the 2-10 mL range.
• You will need to keep the amount of ionic compound constant, at 1 gram.
• You will need to record the initial temperature of each individually as well as the greatest difference in temperature from the starting temperature of the water/ionic compound mixture.
• Complete three trials.

In this lesson, we will focus on the transfer of energy that occurs when bonds are broken. In the previous lesson, you added several different solid ionic compounds to water. When ionic compounds are added to water, ionic bonds are broken due to the attraction of the cations and anions that make up ionic compounds to the partially charged sides of water molecules.

You will need the following resources to complete this assignment.

• HW Conservation of Energy

Let's review what occurred in lesson 2 by looking at Temperature vs Time graphs for three of the ionic compounds we tested.

         KCl: 

    NH₄Cl:  

  NaNO₃:  

Observe the simulation by hitting "setup" and then “go/stop”. After you feel you have learned all the features of the simulation move on to answer questions 2.1 and 2.2 below.

The simulation you were working with models dissolving KI in water. You have control over the exact number of water molecules represented in the “beaker”. You will begin by looking at the overall bulk system.

Directions:

• Adjust the slider bar to the maximum number of water molecules possible. Then press "setup". Press "go" and observe the water molecules moving around the beaker in liquid form.
• Now press “add KI” and observe what happens to the ionic compound and the temperature of the beaker.
• Allow the simulation to run until all KI is split into K⁺ cations and I⁻ anions.
• You may repeat the procedure as needed to answer the questions below.

Now you will use the same simulation, but this time you will be observing individual molecules.

Directions:

• Adjust the slider bar to 10 water molecules. Then press "setup" and "go".
• Now press “add KI” and allow the simulation to run until all KI is split into K⁺ cations and I⁻ anions.
• You may repeat the procedure as needed to answer the questions below.

A GIF of the simulation you have been working with is shown below. Observe that a water molecule slows down and seems to lose kinetic energy when it collides with the ionic compound KI. This collision causes the ionic compound KI to split into a K⁺ cation and an I⁻ anion.

Where does the kinetic energy of the water molecule go? We will explore this question further in your homework and again in the next lesson.

In this lesson, we will observe a reaction where bonds are formed called a synthesis reaction. We will again focus on the transfer of energy that occurs by modeling the process of bond formation.

Before we observe a synthesis reaction, let's review the transfer of energy that occurred in lesson 3 when an ionic compound split apart in water.

In that simulation we saw that a water molecule slowed down and seemed to lose kinetic energy when it collided with the ionic compound KI. This collision caused the ionic compound KI to split into a K⁺ cation and an I⁻ anion.

If we put together two concepts from the homework and lesson 3, we can come to a conclusion about where the kinetic energy from the water went. 

• Potential energy can be converted to kinetic energy, and kinetic energy can be converted back to potential energy.
• Energy can be transferred from the system to the surroundings, or from the surroundings into the system.

Recall that the water temperature at the beginning of our simulation experiment contained a certain amount of kinetic energy. That kinetic energy decreased the moment the water collided with the ionic compound and it subsequently broke apart. That kinetic energy must have transferred into the system.

In Lesson 2: Breaking Ionic Bonds Lab Activity, we recorded the temperature of water as different ionic solids were added to the beaker. We observed a temperature decrease for each of the five ionic solids.

The ionic bonds of ionic solids (the system) are broken when placed in water (the surroundings) causing a decrease in the overall temperature of the water. The decrease in temperature is is due to the transfer of kinetic energy from the surroundings to the system in order to break the bonds of the ionic compound. Temperature decreasing when bonds are broken is an endothermic process.

Definitions:

Endothermic - An endothermic process absorbs energy from its surroundings, usually in the form of heat.

Exothermic - An exothermic process releases energy from the system to its surroundings, usually in the form of heat.

Your instructor will conduct an experiment using solid zinc and solid iodine to yield solid zinc iodide as shown in the chemical equation below. This type of reaction is known as a synthesis reaction because bonds are only being formed in the reaction, no bonds are broken.

Zn (s) + I₂ (s) → ZnI₂ (s)          

1. Zinc and iodine will be combined in a 1:1 molar ratio as indicated in the balanced equation above. Corresponding gram amounts are shown in the table below.
2. After weighing out appropriate amounts of zinc and iodine, the two solids will be ground together using a mortar and pestle (your instructor will demonstrate).
3. The contents will then be transferred to a test tube mounted to a ring stand inside the fume hood.
4. A surface temperature probe will then be adhered to the bottom of the test tube to record the initial temperature before any reaction occurs.
5. Exactly 1 mL of distilled water will be dispensed into the pulverized mixture of zinc (s) and iodine (s).
6. The greatest difference in temperature from the starting temperature (of the surroundings) must be recorded so the overall temperature change can be calculated.

Consider the following as you model the interaction occurring in trials 1 and 3 below:

• The drawing space is meant to represent a zoom-in of the test tube, so there is no need to illustrate the test tube itself in your model. Only illustrate the contents inside the test tube.
• Temperature that was measured was of the surroundings, not of the system itself.
• Temperature correlates directly with kinetic energy. If temperature increases, kinetic energy should also increase.
• Follow the law of conservation of energy: the total number of energy boxes (potential energy + kinetic energy) filled in the before portion of your model must match the total number of energy boxes filled in the after portion.

So far in this unit, we have learned that energy is required to break bonds and that energy is released when bonds are made, but how much energy? How can we actually measure the energy transferred in one or both of these processes?

Because the energy transfer involved in both breaking and making bonds involves heat, or thermal energy, we can employ a method called calorimetry to measure the heat transfer. Calorimetry is the science of measuring heat. In aqueous solutions, the energy is transferred to or taken away from the water, therefore, the thermal energy transferred may be calculated using the change in temperature of the water itself.

A calorimeter can be used to measure the amount of thermal energy transferred when bonds are broken or formed. A calorimeter will consist of a container with a cover, and a thermometer or temperature probe. The ideal calorimeter (for our purposes) must prevent the transfer of thermal energy between the calorimeter's container and its environment. This allows for an accurate determination of the thermal energy that has either been transferred to or from the water within the calorimeter.

Experiment with the simulation by observing two different chemical processes.

• The separation of the ionic compound potassium iodide (KI) in water:

KI (s) → K⁺ (aq) + I⁻ (aq)

• The synthesis of calcium hydroxide, Ca(OH)₂ , and subsequent production of hydrogen gas when solid calcium metal (Ca) is added to water:

Ca(s) + 2 H₂O (l)→ Ca(OH)₂ (aq) + H₂ (g)

Begin by hitting "setup” and then “run/pause”. In order to monitor the temperature of your experiment you will need a sensor in the container. Press “place sensor” in order to add a sensor inside your container with your cursor.

In order to observe the ionic compound separate in water, you will need to press the "Add KI" button; to observe the synthesis of calcium hydroxide and production of hydrogen gas, you will need to press the "Add Ca" button.

Within this unit, we have observed changes in thermal energy by observing changes in temperature within a closed system. We saw this in a temperature decrease with ionic solids dissolving in water in lessons 2 and 3. We also saw this with a temperature increase through the synthesis of an ionic compound in lesson 4.

The thermal energy transferred when bonds are broken or formed is referred to as the enthalpy change (∆H). For our purposes, energy transfer will always occur within a closed system.

Because any energy transfer discussed in this class will occur within a closed system, the following relationship will be true for these energy transfers:

Enthalpy change = heat energy

∆H = q

In other words, as long as the energy is transferred within a closed system the enthalpy change (∆H) will be equal to the heat energy (q) that is transferred:

q = m c ∆T

∆H = m c ∆T

where:

q = heat energy in joules (J), so ∆H = energy transferred in joules (J)

m = mass of sample in grams (g)

c = specific heat (J/g^OC)

ΔT = change in temperature (final temperature - initial temperature)

Specific heat, (c), is an intensive physical property that has a different, characteristic value for every substance. For example, the specific heat for water is 4.18 J/g^OC. This would mean that it will take 4.18 joules of energy to raise the temperature of 1 gram of water 1^OC.

Change in enthalpy is always relative to the energy within a system. If the system gains energy, change in enthalpy (∆H) will be positive. If the system loses energy, the change in enthalpy (∆H) will be negative.

In lessons 2 and 3, we observed the dissolving of an ionic compound in water. A decrease in temperature of the surroundings was observed, and this was labeled as endothermic. In an endothermic reaction, the change in enthalpy (∆H) equals the energy in joules absorbed by the system from the surroundings. Because energy is absorbed into the system, ∆H is positive.

In lesson 4, we observed the synthesis of an ionic compound. An increase in temperature of the surrounding was observed, and this reaction was labeled as exothermic. In an exothermic reaction, the change in enthalpy (∆H) equals the energy in joules released by the system to the surroundings. Because energy is released into the surroundings, ∆H is negative.

The simulation you were working with on page 2 models a calorimeter. The simulation allows you to choose the material to use for your calorimeter using the dropdown menu labeled "container-material".

Begin by hitting "setup” and then “run/pause”. Adjust the "sensor-radius" slider bar to 48, then press “place sensor” to add this to the inside of the container ensuring the sensor does not contact to walls of the container. Now adjust the "sensor-radius" slider bar to 4, then press “place sensor” to add this to the outside of the container, again ensuring the sensor does not contact to walls of the container.

You will choose one of the chemical processes available (either the separation of the ionic compound potassium iodide (KI) in water or the synthesis of calcium hydroxide, Ca(OH)₂ , and subsequent production of hydrogen gas when solid calcium metal (Ca) is added to water) to determine the ideal material to use for your calorimeter. In order to observe the ionic compound separate in water, you will need to press the "Add KI" button; to observe the synthesis of calcium hydroxide and production of hydrogen gas, you will need to press the "Add Ca" button.

Now that you have experimented using several different container materials, you will determine which material is best suited for a calorimeter.

On page 1 of this lesson, it was stated that the ideal calorimeter must prevent the transfer of thermal energy between the calorimeter's container and its environment in order to allow for an accurate determination of the thermal energy transferred.

The data you acquired for aluminum, glass, and Styrofoam is shown below.

In the last lesson you were introduced to calorimetry: an experimental method used to measure the transfer of heat within a closed system. You also learned that this transfer of heat is equal to the enthalpy change (∆H) which is the thermal energy transferred when bonds are broken or formed. You were able to explore how a calorimeter functions and determine the ideal material to use for your own calorimeter.

You will now use what you learned in the last lesson to design and execute your own calorimetry experiment.

You will need the following resources to complete this assignment.

• Calorimetry Lab Student Rubric
• HW Calorimetry Calculations

Recall that overall chemical reactions can be categorized as endothermic or exothermic based on energy transferred into or out of the system. The chemical reaction is the system in this case.

Positive ΔH = energy moves into system from surrounding = endothermic

Negative ΔH = energy moves out system to surrounding  = exothermic

When a chemical reaction occurs, there is a characteristic change in enthalpy. The enthalpy change for a reaction is based on the overall balanced equation and is expressed in kilojoules per mole. This can be calculated using the same equation used to calculate enthalpy (q = m c ∆T), then dividing by the number of moles of a substance utilized.

The enthalpy of an exothermic reaction is shown below:

2 H₂ (g) + O₂ (g) → 2 H₂O (g)   ΔH = -483.6 kJ/mol

This means that 483.6 kilojoules of energy are released for every two moles of hydrogen gas and one mole of oxygen gas that react to produce two moles of water vapor.

The purpose of the lab is to measure the energy absorbed or released in a chemical reaction. You will utilize a calorimeter to measure the energy absorbed or released in the following chemical reaction:

Ca (s) + 2 H₂O (l) → Ca(OH)₂ (aq) + H₂ (g)

In this reaction you have two reactants, calcium and water. The amount of calcium will be your independent variable. The amount of water will be a controlled variable, as you will keep it constant for all trials. Energy absorbed or released will be your dependent variable. Anything that is constant throughout your experiment is a controlled variable.

As you run your experiment, you will use three different amounts of calcium. You will run each trial three times, for a total of nine trials. The amount of water must be held constant throughout your experiment.

Design an experiment to determine the enthalpy of reaction, ∆H°_rxn, for the reaction of calcium metal with water by answering the questions below.

Use the procedure you designed on pages 2-3 to complete three trials for your first decided upon amount of calcium to test.

Use the procedure you designed on page 2 to complete three trials for your second decided upon amount of calcium to test.

Use the procedure you designed on page 2 to complete three trials for your third decided upon amount of calcium to test.

Most chemical reactions involve both breaking and making bonds. Whether or not energy is absorbed or released in a chemical reaction can be understood by looking at the difference between the number of bonds broken versus the number of bonds formed.

Most chemical reactions involve both breaking and making bonds. Whether or not energy is absorbed or released in a chemical reaction can be understood by looking at the difference between the number of bonds broken versus the number of bonds formed.

You will need the following resources to complete this assignment.

• Element Cutouts

Your instructor will demonstrate the following chemical reaction using 32 grams of solid barium hydroxide, Ba(OH)₂, and 10 grams of solid ammonium chloride, NH₄Cl:

Ba(OH)₂ (s) + 2 NH₄Cl (s) → 2 NH₃ (g) + BaCl₂ (s) + 2 H₂O (l)

In order to describe the energy transfer in the reaction, first we must define both the system and the surroundings in the reaction as demonstrated.

• Get a large whiteboard and whiteboard markers for your table group.
• Create the starting molecules with the paper elements given. You should use all the paper materials you have (no elements should be left over). Draw bonds between atoms on the whiteboard.
• Now, rearrange the molecules you have in order to create the products by breaking as few bonds as possible. Again, draw bonds between atoms on the whiteboard
• Once you have a solid grasp on which bonds are not broken during the process, you can tape the atoms that stay bonded throughout the reaction together.

Comparing the total number of bonds broken to those formed is the first step in helping us determine if a reaction is endothermic or exothermic. On the previous page, you were provided with atom cutouts to better enable you to determine where bonds are located in the reactants and which bonds must be broken to form the products.

Drawing atom models is another way to go about this same process. In the questions below, you will be asked to draw atom models for both the reactants and the products in each chemical reaction. Then you will count both the bonds present in the reactants that must be broken in order to form the products, and the number of new bonds that must be formed to create the products.

You will then use the difference in the number of bonds broken vs. formed to determine if the overall chemical reaction is likely endothermic or exothermic. More bonds broken than formed correlates with an endothermic reaction as more energy must be added to the system to break the bonds, while fewer bonds broken than formed correlates with an exothermic reaction. An exothermic reaction is one is which energy is being released from the system.

As it turns out, all bonds are not created equal. Some bonds have more energy associated with them than others. A stronger, more energetic bond requires more energy to be broken, and releases more energy when it is formed. The end of the video below shows an example chemical reaction where two bonds are broken and two bonds are formed.

The Law of Conservation of Energy states that energy is neither created nor destroyed in any chemical reaction. Thus, we can determine if a chemical reaction is endothermic or exothermic by quantifying both the amount of energy transferred from the surroundings into the system to break bonds, and the amount of energy released from the system back into the surroundings when bonds are formed.

You will need the following resources to complete this assignment.

• Modeling a Chemical Reaction

Your instructor will begin by modeling an endothermic chemical reaction on the board for you using a set amount of energy (potential energy + kinetic energy) that will be traced from the surroundings into the system and back out again. The amount of energy absorbed or released will be based on the number of bonds that are broken versus the number of new bonds formed in the chemical reaction.

After your instructor models the endothermic reaction, you will model an exothermic reaction on your handout using the same principles and ideas demonstrated in your instructor's model on the board. Then you will practice modeling two more chemical reactions; you will need to determine if these reactions are endothermic or exothermic during the modeling process to ensure the correct amount of energy absorbed or transferred throughout the chemical reaction