Heat+of+solution

=Determination of Heat of solution of different solutes, comparing values and coming up with appropriate explanations= Project outline [|Heat of Solution.pdf] Final Report: Heat of Solution   (Draft deleted)

Project Objective (from Heat of Solution.pdf)  W      hen certain compounds are dissolved in water there is an enthalpy change (heat change) associated with this process. In this experiment you will determine the heat of solution of compounds like potassium nitrate and ammonium chloride. Compare their results and come up with logical arguments as to the reason for the observation. This experiment can be performed using improvised calorimeters (Styrofoam cups) >  discuss if the materials required are available and the lab activity can be performed.   //For the sake of organization, please follow the "Entry #, Topic Title, and Posted By" format to avoid any mix-up!// The process of dissolving is a process which involves the //breaking and making of bonds//, and that //involves energy//. - The //breaking of bonds requires or absorbs energy//. Using energy like that is called //endothermic//. - The //formation of bonds releases energy//. This is called //exothermic//. - Dissolution overall can be either endothermic or exothermic, depending on whether more energy was used to break the bonds, or more energy was released when new bonds or more energy was released when new bonds were formed. If more energy is released in making bonds than is used in breaking bonds, the process is exothermic. If more energy is used than is released, the process is endothermic. ^ That's just basic information for reference and I also typed up the project objective so it won't be a hassle to keep referring to the pdf :) I didn't manage to find too much on the topic, but the links should provide a little bit of background info. //30/09/08 Posted by Yiing Hu//
 *  Research and come up with a possible lab procedure for setting up your equipments and performing the experiment.
 *  Include safety concerns that should be addressed
 *  Make a list of equipments that may be used to perform the experiment.
 *  Submit a draft of the procedure adopted for the lab to your teacher; and
 *  Choose a convenient time for your experiment, record your data and present your findings as instructed.
 * __ENTRY 1__: Basic Information and Links**
 * link:** http://dl.clackamas.edu/ch105-03/heatof.htm
 * enthalpy change of a solution:** http://en.wikipedia.org/wiki/Enthalpy_change_of_solution
 * potassium nitrate:** <span style="color: rgb(129, 223, 194);">[|http://web.njit.edu/~grow/Heatsolution/HeatofSolution.html]
 * ammonium chloride:** <span style="color: rgb(129, 223, 194);">http://www.bjservices.com/website/ps.nsf/0/135A6D69D98D0B5F86256A5400349D73/$file/ST-AmmoniumChloride.pdf
 * will refer to textbook later to find additional info, 加油!

**__ENTRY 2:__ Basic Info Continued** <span style="color: rgb(251, 203, 209);"><span style="font-size: 110%; color: rgb(112, 235, 196); font-family: 'Times New Roman',Times,serif;"><span style="color: rgb(0, 0, 0);">- Expressed in **kJ/mol** at constant pressure <span style="color: rgb(251, 203, 209);"><span style="font-size: 110%; color: rgb(112, 235, 196); font-family: 'Times New Roman',Times,serif;"><span style="color: rgb(0, 0, 0);"> - the heat of solution of a substance is defined as the __sum of the energy absorbed (positive kJ/mol) and the energy released (negative kJ/mol)__ - According to http://dl.clackamas.edu/ch105-03/heatof.htm, when using the term "**heat of solution**", the word "**solution**" in this case doesn't refer to the actual physical mixture that is formed, but rather the process of dissolving

Dissolution takes place in three steps: 1. Breaking solute-solute attraction (endothermic) 2. Breaking solvent-solvent attractions (endothermic) 3. Forming solvent-solute attractions (exothermic) through solvation (read more [|here])

//04/10/08 Posted by Nicole Pun//

__**ENTRY 3:** Measuring Heat of Solution__ - a <span style="color: rgb(230, 121, 121);">**calorimeter** is used to measure the heat of chemical reactions, physical changes, and heat capacity - to find heat change of a substance made of two liquids, add the liquids to the calorimeter and take note of the initial temperature, as well as the final temperature when the reaction is finished; multiply temperature change by mass and specific heat capacity of the two reactants to find the energy given off during the reaction (if the reaction is exothermic); divide energy change by how many moles of the substance was present [this method doesn't take into account heat loss through container or the heat capacity of the thermometer and container itself; we can use this as SOURCES OF ERROR]

Basically, we need to create a simple calorimeter to roughly measure the heat of solution of potassium nitrate or ammonium chloride (or similar compounds). I suppose we'll need a relatively good thermometer, Styrofoam cup(s), the compounds (in their separate form) and specific heat capacities, depending on what we end up using.

From Wikipedia: <span style="font-size: 90%; font-family: 'Times New Roman',Times,serif;">A **constant-pressure calorimeter** measures the change in [|enthalpy] of a reaction occurring in [|solution] during which the [|atmospheric pressure] remains constant. An example is a coffee-cup calorimeter, which is constructed from two nested [|Styrofoam] cups and holes through which a [|thermometer] and a stirring rod can be inserted. The inner cup holds the solution in which of the reaction occurs, and the outer cup provides [|insulation]. Then //C////p// = (//W// * //D////H// / (//M// * //D////T//)) where
 * <span class="mw-headline" style="color: rgb(194, 97, 162); font-family: 'Times New Roman',Times,serif;">Constant-pressure calorimeter **

DH = Enthalpy of solution DT = change of temperature W = weight of solute M = molecular weight of solute

<span style="font-size: 90%; font-family: 'Times New Roman',Times,serif;">**A. Finding the Heat of Solution from Calorimetry** (taken from [|http://www.wwnorton.com/college/chemistry/gilbert/concepts/chapter13/ch13_1.htm)****] <span style="font-size: 90%; font-family: 'Times New Roman',Times,serif;">Review calorimetry from [|Chapter 11]. This chapter tends to use specific heat instead of heat capacity. Using specific heat, Equation 11.8 becomes

<span style="font-size: 90%; font-family: 'Times New Roman',Times,serif;">heat energy (//q//) = (mass)(specific heat)(change of temperature)

<span style="font-size: 90%; font-family: 'Times New Roman',Times,serif;">Because there is normally much more solvent than solute, it is usually assumed that the specific heat is that of the solvent. For water (the most common solvent),the specific heat is 4.184 J/g•°C. Remember that the mass will refer to the mass of the entire solution, since it is the entire solution undergoing the temperature change. <span style="font-size: 90%; font-family: 'Times New Roman',Times,serif;">The heat of solution refers to the dissolution of a solute. For an ionic compound, he reaction is the salt forming its composite ions. Although water is required for the reaction, it is not written as part of the reaction. For example,

<span style="font-size: 90%; font-family: 'Times New Roman',Times,serif;">

<span style="font-size: 90%; font-family: 'Times New Roman',Times,serif;">So heat of solution is calculated from



<span style="font-size: 90%; font-family: 'Times New Roman',Times,serif;">Since the heat (//q//) comes from the chemical reaction of the solute, the //H// value is per moles solute, not per mol solution. <span style="font-size: 90%; font-family: 'Times New Roman',Times,serif;">The value of //H// is positive for endothermic dissolutions(temperature decreases) and negative for exothermic dissolutions (temperature increase).

<span style="color: rgb(251, 203, 209);"><span style="font-size: 110%; color: rgb(112, 235, 196); font-family: 'Times New Roman',Times,serif;"><span style="color: rgb(0, 0, 0);"> //04/10/08 Posted by Nicole Pun// <span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;"> <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"> __**ENTRY 4:** Enthalpy Change of Solution__ Summarized from Wikipedia's **[|Enthalpy Change of Solution]** ||   ||  <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">**-** || <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">17.89 ||   ||  <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">**+** || <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">6.14 ||   ||  <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">**-** || <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">7.29 ||   ||  <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">**-** || <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">13.77 ||   ||  <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">**-** || <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">17.10 ||   ||  <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">**+** || <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">3.89 ||   ||  <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">**+** || <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">9.89 ||   ||  <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">**-** || <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">0.360 || <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"> <span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;"> -Charissa 2008/10/04
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">enthalpy of solution = enthalpy of dissolution = enthalpy change associated with the dissolution of a substance in a solvent at constant pressure
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">enthalpy: heat content (denoted by //H//, //h//, or rarely as //χ//)
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">dissolution: the process by which a solid, gas, or liquid is dispersed homogeneously in a gas, solid, or, esp., a liquid.
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">one of the three dimensions of solubility analysis
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">kJ/mol, at constant temperature
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">heat of solution of a substance is defined as the <span style="color: rgb(231, 24, 24);">**sum of the energy absorbed**, or <span style="color: rgb(231, 24, 24);">**endothermic energy** , and <span style="color: rgb(12, 16, 202);">**energy released** , or <span style="color: rgb(12, 16, 202);">**exothermic energy**
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;"><span style="color: rgb(231, 24, 24);">endothermic = //absorption of energy (i.e. heat)// = positive kJ/mol
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;"><span style="color: rgb(12, 16, 202);">exothermic = //liberation of energy (i.e. heat)// = negative kJ/mol
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;"><span style="color: rgb(12, 16, 202);">heating decreases solubility of gases, exothermic
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">gas dissolves in liquid solvent, temperature decreases, solution releases E
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">an effect of the increase in heat (heat = E needed to //attract// solute and solvent molecules; outweighs the E needed to separate solvent molecules)
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">__THEREFORE: "COMPLETE" DISSOLUTION = HEAT OF SOLUTION AT MAX__
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">see Nicole's post for three steps of dissolution
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">value of overall enthalpy change = sum of individual enthalpy changes of each step
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">solutions with <span style="color: rgb(0, 24, 255);">negative heats of solution form <span style="color: rgb(0, 24, 255);">stronger bonds and have <span style="color: rgb(15, 15, 179);"><span style="color: rgb(15, 15, 179);"><span style="color: rgb(34, 19, 195);">lower vapor pressure
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;"><span style="color: rgb(0, 0, 0);">EXAMPLES:
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">**Heat of solution for some selected compounds**
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">[|hydrochloric acid]
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">[|ammonium nitrate]
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">[|ammonia]
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">[|potassium hydroxide]
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">[|caesium hydroxide]
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">[|sodium chloride]
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">[|potassium chlorate]
 * <span style="font-size: 77%; font-family: 'Times New Roman',Times,serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;">[|acetic acid]

<span style="color: rgb(251, 203, 209);"><span style="font-size: 110%; color: rgb(112, 235, 196); font-family: 'Times New Roman',Times,serif;"><span style="color: rgb(0, 0, 0);">//   <span style="color: rgb(0, 0, 0);">__**ENTRY 5:** Compounds__ Background info on the compounds we'll be experimenting with //   http://en.wikipedia.org/wiki/Potassium_nitrate <span style="font-size: 9pt; color: rgb(112, 235, 196); font-family: 'Trebuchet MS';">
 * Potassium nitrate ** is a <span style="font-size: 9pt; color: rgb(0, 43, 184); font-family: 'Trebuchet MS';">  [|chemical compound] with the <span style="font-size: 9pt; color: rgb(0, 43, 184); font-family: 'Trebuchet MS';"> [|chemical formula]  <span style="font-size: 9pt; color: rgb(0, 43, 184); font-family: 'Trebuchet MS';">K [|N] [|O]3. A naturally occurring mineral source of <span style="font-size: 9pt; color: rgb(0, 43, 184); font-family: 'Trebuchet MS';"> [|Nitrogen], KNO3constitutes a critical <span style="font-size: 9pt; color: rgb(0, 43, 184); font-family: 'Trebuchet MS';"> [|oxidizing] component of <span style="font-size: 9pt; color: rgb(0, 43, 184); font-family: 'Trebuchet MS';"> [|black powder] / <span style="font-size: 9pt; color: rgb(0, 43, 184); font-family: 'Trebuchet MS';"> [|gunpowder] . In the past it was also used for several kinds of burning fuses, including <span style="font-size: 9pt; color: rgb(0, 43, 184); font-family: 'Trebuchet MS';"> [|slow matches] . Since potassium nitrate readily <span style="font-size: 9pt; color: rgb(0, 43, 184); font-family: 'Trebuchet MS';"> [|precipitates] , <span style="font-size: 9pt; color: rgb(0, 43, 184); font-family: 'Trebuchet MS';"> [|urine] was a significant source, through various malodorous means, from the <span style="font-size: 9pt; color: rgb(0, 43, 184); font-family: 'Trebuchet MS';"> [|Late Middle Ages] and <span style="font-size: 9pt; color: rgb(0, 43, 184); font-family: 'Trebuchet MS';"> [|Early Modern era] through the 19th century.[<span style="font-size: 9pt; color: rgb(0, 43, 184); font-family: 'Trebuchet MS';">[|citation needed] ] <span style="font-size: 9pt; color: rgb(112, 235, 196); font-family: 'Trebuchet MS';">

** Ammonium chloride ** ([|N][|H]4[|Cl]) (also ** Sal Ammoniac **, ** salmiac **, ** nushadir salt **, ** zalmiak **, ** sal armagnac **, ** sal armoniac **, ** salmiakki **, [|salmiak]and ** salt armoniack **) is, in its pure form, a clear white water-soluble crystalline [|salt] of [|ammonia]. The aqueous ammonium chloride solution is mildly acidic. [|Sal ammoniac] is a name of natural, mineralogical form of ammonium chloride. The mineral is especially common on burning coal dumps (formed by condensation of coal-derived gases), but also on some volcanoes. http://en.wikipedia.org/wiki/Ammonium_chloride Sorry for the huge font, the visual editor is very troublesome. <span style="color: rgb(251, 203, 209);"><span style="font-size: 110%; color: rgb(112, 235, 196); font-family: 'Times New Roman',Times,serif;"><span style="color: rgb(0, 0, 0);"> 04/10/08 Posted by Yiing Hu

<span style="color: rgb(0, 0, 0);">__**ENTRY 6:** Calorimeter__ ** Okay, this is some information on making an "improvised calorimeter", since we'll be using styrofoam cups instead of the real thing. ** <span class="Apple-style-span" style="word-spacing: 0px; font-family: verdana; font-style: normal; font-variant: normal; font-weight: normal; font-size: 13px; line-height: normal; text-transform: none; color: rgb(0, 0, 0); text-indent: 0px; white-space: normal; letter-spacing: normal; border-collapse: separate; font-size-adjust: none; font-stretch: normal; orphans: 2; widows: 2;">Calorimetry is used to determine the heat released or absorbed in a chemical reaction. The calorimeters shown here can determine the heat of a solution reaction at constant (atmospheric) pressure. The calorimeter is a double styrofoam cup fitted with a plastic top in which there is a hole for a thermometer. (It's crude, but very effective!) Key techniques for obtaining accurate results are starting with a dry calorimeter, measuring solution volumes precisely, and determining change in temperature accurately. [|http://www.dartmouth.edu/~chemlab/techniques/calorimeter.html]

<span class="Apple-style-span" style="word-spacing: 0px; font-family: 'times new roman'; font-style: normal; font-variant: normal; font-weight: normal; font-size: 16px; line-height: normal; text-transform: none; color: rgb(0, 0, 0); text-indent: 0px; white-space: normal; letter-spacing: normal; border-collapse: separate; font-size-adjust: none; font-stretch: normal; orphans: 2; widows: 2;">Calorimetric measurements led to the discovery that **every substance requires a characteristic amount of heat to change its temperature over a temperature interval.** The amount of heat required is approximately, but not exactly, uniform over any reasonable range of temperatures. The amount of heat required per unit of mass came to be known as the **specific heat** of the substance. The specific heat of a substance is an intensive property characteristic of the substance.

<span class="Apple-style-span" style="word-spacing: 0px; font-family: 'times new roman'; font-style: normal; font-variant: normal; font-weight: normal; font-size: 16px; line-height: normal; text-transform: none; color: rgb(0, 0, 0); text-indent: 0px; white-space: normal; letter-spacing: normal; border-collapse: separate; font-size-adjust: none; font-stretch: normal; orphans: 2; widows: 2;">The ice calorimeter is simply a large insulated container of ice and water with a basket which can be used to remove the ice for weighing. The amount of heat evolved in whatever reaction takes place within the calorimeter is equal to the mass of ice melted multiplied by the heat of fusion of ice, 333.51 kJ/kg. The water calorimeter is often used in undergraduate student laboratories in the form shown as the Figure below, the "coffee cup" calorimeter. [|http://www.ualberta.ca/~jplambec/che/p101/p01073.htm]

On Page 133 - Read more on the website, for more information on making a coffee cup calorimeter http://books.google.ca/books?id=6PazDs_r0LMC&pg=PA133&lpg=PA133&dq=styrofoam+calorimeter&source=web&ots=ZVQYoLdIWk&sig=yZW1rWiP3Qg2o1k_Ac7lXpuA4qA&hl=en&sa=X&oi=book_result&resnum=6&ct=result

<span class="Apple-style-span" style="word-spacing: 0px; font-family: 'times new roman'; font-style: normal; font-variant: normal; font-weight: normal; font-size: 16px; line-height: normal; text-transform: none; color: rgb(0, 0, 0); text-indent: 0px; white-space: normal; letter-spacing: normal; border-collapse: separate; font-size-adjust: none; font-stretch: normal; orphans: 2; widows: 2;"> Most reactions we perform in chemistry 30 are done in solution. The cheapest and most effective calorimeter to use to study solution chemists is a simple Styrofoam coffee cup. Water and the Styrofoam are the surroundings and the Styrofoam provides adequate insulation since most reactions in solution occur relatively rapidly. A thermometer is used to measure the temperature change in the water solution. Water is typically used because of it's high specific heat capacity, which means it can store large amounts of heat without large temperature changes occurring. (It can also release large amounts of heat for the same reason). Other substances, like oil can be used in place of the water. The coffee cups may be nested (placed one inside another) to provide addition insulation if needed. http://www.scs.sk.ca/cyber/elem/learningcommunity/sciences/chemistry30/curr_content/chem30/modules/module3/lesson4/typesandlabs.htm

<span style="color: rgb(251, 203, 209);"><span style="font-size: 110%; color: rgb(112, 235, 196); font-family: 'Times New Roman',Times,serif;"><span style="color: rgb(0, 0, 0);">// 04/10/08 Posted by Yiing Hu // **__ Entry 7: <span style="color: rgb(0, 255, 130);">Tentative Procedure and Materials __** 1. Trace the outline of the thermometer’s bulb onto the lid. Carefully cut it out with an Exacto knife. Ensure that the thermometer slides through the lid easily. 2. Measure 100mL of water with a graduated cylinder and pour it into the Styrofoam cup. 3. Measure 10mL of the chemical and pour it into the water gently. Stir. 4. Put the lid onto the cup. 5. Place the thermometer through the lid and into the cup. Ensure that the bulb is below the water surface. 6. For the next few minutes (est. 3?), observe and take note of the changes in temperature. 7. Repeat. 8. Calculate heat flows. //Note: Take note of the experimental errors that take place during the experiment, and any outside factors that might have influenced the results.//
 * Procedure**

100mL x (# of compounds used) of water 10mL of each compound 2 Styrofoam cups (to put one in the other to reduce heat loss) 1 lid that fits the cup 1 thermometer 1 Exacto knife
 * Materials:**

-Charissa 2008/10/04

=<span style="color: rgb(0, 4, 255);"><span style="color: rgb(0, 4, 255);"><span style="color: rgb(0, 4, 255);"><span style="color: rgb(0, 4, 255);"><span style="color: rgb(0, 4, 255);"><span style="color: rgb(0, 4, 255);"><span style="color: rgb(0, 4, 255);"><span style="color: rgb(103, 175, 24);">Entry 8       =

<span style="color: rgb(24, 50, 205);">**Determining Heat of Solution - Rough Outline** To determine (molar) heat of solution for Ammonium nitrate:

In this part of the experiment, the calorimeter is filled with 60.0 g of water. A small sealed glass bulb containing 5.00 g of pure NH4NO3 is placed in the calorimeter. The reaction is initiated by breaking the glass bulb, allowing the NH4NO3 to dissolve in the water. The heat capacity of the calorimeter (//Ccal//) is 153. J oC-1. The formula weight of NH4NO3 is 80.04. Repeat the procedure in Part 1 for determining the heat of solution of ammonium nitrate using different values for the mass of ammonium nitrate and/or the mass of water. You should obtain the same heat capacity as you did in Part 1.
 * Part 1**
 * Part 2**

We can adopt the method; it's here to give you an idea For more info go to http://www.chm.davidson.edu/chemistryApplets/calorimetry/HeatOfSolutionOfAmmoniumNitrate.html <span style="color: rgb(37, 62, 233);"> = = Go to: http://www.emsb.qc.ca/laurenhill/science/calorimetry.html
 * The purpose of Calorimetry and Molar Enthalpy**
 * The purpose of calorimetry = determine the enthalpy of a substance undergoing chemical change
 * bomb calorimeter = enthalpy of combustion measured
 * This is how the caloric content of foods is determined
 * Since the heat absorbed/released is proportional to the amount of reactant used, //molar enthalpy// = D H///n// is a more meaningful and characteristic quantity.

Calorimeter ** The following example will demonstrate how this can be done. The temperature of a calorimeter increases 0.10 K when 7.52 J of electric energy is used to heat it. What is the heat capacity of the calorimeter? //Solution://** Dividing the amount of energy by the temperature increase yields the heat capacity, //C//, //C// = 7.52 / 0.10 = 75.2 J/K.
 * <span style="color: rgb(58, 77, 233);">Calorimetry - Measuring Heats of Reactions **
 * Any process that results in heat being generated and exchanged with the environment is a candidate for a calorimetric study
 * has a very broad range of applicability, with examples ranging from drug design in the pharmaceutical industry to the study of metabolic rates in biological (people included) systems
 * two types of calorimetry: measurements based on constant pressure and measurement based on constant volume.
 * <span style="color: rgb(41, 46, 224);">
 * device used to measure heat of reaction
 * A simple styrofoam cup is sufficient
 * because it is a container with good insulated walls to prevent heat exchange with the environment
 * In order to measure heats of reactions, we often enclose reactants in a calorimeter, initiate the reaction, and measure the temperature difference before and after the reaction
 * The temperature difference enables us to evaluate the heat released in the reaction.
 * Calorimeter may be operated under constant (atmosphere) pressure, or constant volume
 * Whichever kind to use, we must know its heat capacity
 * **heat capacity** = amount of heat required to raise the temperature of the entire calorimeter by 1 K
 * determined experimentally before or after the actual measurements of heat of reaction
 * heat capacity of the calorimeter is determined by transferring a known amount of heat into it and measuring its temperature increase
 * B/C temperature differences = VERY small, sensitive thermometers are needed
 * Example 1

For more information, go to: [|http://www.science.uwaterloo.ca/~cchieh/cact/c120/calorimetry.html]

-Thamy G. 05/10/08

__**Entry 9**__ Just a note: I added two more compounds from Wikipedia onto the draft because Mr. Vincent said we need at least three compounds in order to compare them.

-Charissa 2008/10/05

<span style="color: rgb(234, 42, 42);">__**Entry 10**__
We need to update our materials list. Here are the things that we used in our lab:

- safety goggles - scoopula - 2 Styrofoam "bowls" - 1 thermometer - graduated cylinder - water - 10g each of NaOh, NH4Cl, NaCl, K2NO3

Please add more if you guys remember something I missed.

And just a quick reminder about what our objective is:

//**Objective: When certain compounds are dissolved in water there is an enthalpy change (heat change) associated with this process. In this experiment you will determine the heat of solution of compounds like potassium nitrate and ammonium chloride. Compare their results and come up with logical arguments as to the reason for the observation. This experiment can be performed using improvised calorimeters (Styrofoam cups)**//

//Posted by Nicole Pun [2008/11/23]//