CHM 431 PHYSICAL CHEMISTRY
TITLE
: CALORIMETRY: HESS’S LAW
NAME
: NURFADHILAH BINTI JAAFAR
ID NUMBER
: 2016675256
GROUP
: AS 246 3B
SUBMITION DATE
: 12TH AUGUST 2016
LECTURER’S NAME : MADAM ZARILA MOHD SHARIFF
TITLE Calorimetry : Hess’s Law OBJECTIVE 1. To compare the heat capacities of a coffee cup calorimeter and a copper calorimeter. 2. To determine the standard enthalpy of formation magnesium oxide ΔHof. INTRODUCTION Calorimetry is an accurate process for determining an enthalpy change in a reaction. Calorimetry, as defined by the Nelson Chemist text book, is “the technological process of measuring energy changes in a chemical system.” By performing an experiment using calorimetric based, one can measured the change in heat (kinetic) energy for exothermic reaction (released heat) or endothermic reaction (absorbed heat) in a thermodynamic system. The heat gain and released cannot be directly measured during the reaction, hence energy heat represents as q can be measured in Joules in the equation q=mcΔT. The change of heat can be calculated based on mass, temperature change and specific heat of the substance that absorbing or losing heat. This experiment is conducted to determine the standard molar enthalpy of formation of magnesium oxide Mg(s) + 2HCl(aq) MgO(s) + 2HCl(aq)
MgCl2 (aq) + H2(g) MgCl2(aq) + H2O(l)
The 3rd reaction requires the standard heat of formation,
Hrxn (A) (1) Hrxn (B) (2)
H (kJ/mol) for the formation of water
from hydrogen and oxygen. H2 + ½ O2
H2O
H)f(H2O) (3)
Hess’s law also states that it does not matter how many steps it takes for a reactant to become a product, the enthalpy change will remain the same. The change in enthalpy, H, is the heat (qp) associated with a reaction at constant pressure, in which no work is involved other than "expansion work" or "compression work", associated with volume changes of the system.
PROCEDURE The experiment is performed at standard conditions (atmospheric pressure and 25oC). All the experiments are done in a done in a double Styrofoam cup (coffee cup) calorimeter and copper calorimeter. A) Heat capacity of calorimeter Make sure the double nested Styrofoam cups are clean and dry 1. A burette was used to deliver the exactly 50 cm3 of tap water into the calorimeter. The cover and the thermometer was replaced. The water temperature was recorded for 4 minutes at one minute intervals. 2. 50 cm3 of hot water (40- 50oC above the room temperature) was measured by graduated cylinder and was poured into a beaker. By the used of another thermometer, the temperature of the hot water was recorded and then was poured completely into the calorimeter containing cold liquid at the 5th minute. The lid was replaced and carefully the water was stirred with the thermometer. The temperature was recorded every 15 seconds for the next 3 minutes. 3. Steps (1) and (2) was repeated by using the copper calorimeter.
B) Reaction 2: Magnesium with hydrochloric acid 1. 1.0-1.1 g of magnesium powder was weighed. The exact weight used was recorded. 2. 50 cm3 of 2 M HCl was drained from a burette into the calorimeter. The cover and thermometer were replaced. The temperature of the HCl was recorded every minute for four minutes. At the fifth minute, the magnesium powder was quickly poured into the HCl. The lid was replaced and the contents were stirred carefully in the calorimeter with the thermometer. The temperature for the next 3 minutes at 15 seconds intervals.
C) Reaction3: Magnesium oxide with hydrochloric acid 1. The magnesium oxide was weighed between 1.6-1.8 g. The exact weigh used was recorded. 2. A step (2) in B was repeated.
D) Reaction4: Hydrogen gas with oxygen gas For safety reasons, this value will not be experimentally determined in the laboratory. It has been professionally determined and verified and the value is listed in standard reference sources. The assignment is to find the value, in kilojoules per mole (kJ/mol)
CHEMICALS 1. 2.0 M hydrochloric acid 2. Magnesium oxide powder 3. Magnesium powder
RESULT THE EXTRAPOLATED TEMPERATURE
Temperature of Coffee Cup Calorimeter with Hot Water 45 40 35 30 25 20 15 10 5 0
y = -0.2795x + 42.989
Temperature/ C Series2 Linear (Series2)
0
2
4
The extrapolated temperature is 42.98 o
6
8
10
Heat capacity of calorimeters Coffee cup calorimeter Calculation of data Part 1, Using Mg Δ T = Tfinal - Tinitial ΔT = 47 °C -28 °C = 10 °C 90.8 cm = 1.0926 g 90.8 cm = 1.0926 g 10.0 cm = x 90.8x = 10.93 x = 0.120 g Mg 0.120/24.31 g/mol = 0.00495 mol Mg q = mcΔT q = 50g HCl x 4.18 J/g·°C x 12 °C q = 2510 J 2.510kJ ΔH = -q ΔH= -2.510 kJ -2.510/0.00495 = -507.15 kJ/mol
Part 2, using MgO: to find the enthalpy of MgO(s) + 2H+(aq) → H2O (l) + Mg2+(aq) Δ T = Tfinal - Tinitial ΔT = 50°C -31 °C = 19°C 1.621 g MgO/ 40.31 = 0.0402 mol MgO q = mcΔT q = 50 g HCl x 4.184 J/g·°C x 19 °C q = 3974.8 J or 3.9748 kJ ΔH = -q ΔH= -3.9748. kJ -3.9748/ 0.0402 mol = -98.88 kJ/mol to find the enthalpy of Mg + 1/2O2 → MgO using these equations and their enthalpies: Enthalpy: -507.15 kJ/mol Mg(s) + 2H+(aq) → Mg2+(aq) + H2 (g) Enthalpy: +98.88 kJ/mol H2O (l) + Mg2+(aq) → MgO(s) + 2H+(aq) Enthalpy: -258.8 kJ/mol H2 + 1/2O2 → H2O ____________________________________________________________________________ Enthalpy: - 667.07 kJ/mol Mg + 1/2O2 → MgO By combining the 3 reactions above, and therefore their enthalpies, we can determine the enthalpy of the equation of the combustion of Mg (Mg + 1/2O2 → MgO) is -667.07 kJ/mol -507.15 + 98.88 + -258.8 = -667.07 kJ/mol
DISCUSSIONS A. The theoretical enthalpy value for the combustion of magnesium in the reaction Mg + 1/2O2 → MgO is -607.10 kJ/mol. Our value is -667.07 kJ/mol. Differences in these values could have been from incorrect lab procedure. The standard enthalpy for the reaction Mg + 1/2O2 → MgO is -607.1 kJ/mol. Finding the percent error: Percent Error = (-667.07- -607.10)
x 100 = 9.88 %
-607.10 Possible sources of error could have occurred from improper lab technique. The scale might be fluctuating slightly when measuring the magnesium oxide powder which possibly giving us the moderately incorrect weight of powder. This could have inflated or deflated the enthalpy calculated. Besides that, when moving the magnesium oxide powder from the weighing boat to the Styrofoam cup, a small amount of magnesium powder slipped onto the lab table. The enthalpy change might be inflated due to the decreased mass of the magnesium oxide for the reaction. Lastly, when pouring the 50 mL of HCl into the coffee cup, a small amount spilled the cup. This decreases the amount of acid added, resulting in an inflation of the enthalpy value found. We calculated the change in enthalpy using the specific heat capacity formula. We had used it before but we learned a lot by using it in a real life setting. At first, when calculating enthalpy, we made the mistake of putting in the grams of Magnesium and Magnesium Oxide for m (in the equation q = mcΔT). We were getting enthalpies close to zero. We had to go back to the equation and use the amount of HCl for m (in the equation q = mcΔT) to get the correct enthalpy and making this mistake at first in this experiment permanently taught us that lesson. This experiment taught how to use Hess’s law to calculate the energy change in a reaction through experimentation. We learned how to use the sum of the individual reactions in our reaction to calculate its change in energy with experimental data. In this lab we also learned how to convert and alternate units. We learned how to convert kilojoules to joules and also that a temperature changes in Kelvin is equal temperature change in Celsius.
CONCLUSIONS In conclusion, the data collected supported our hypothesis. The reaction in this experiment was exothermic with the negative enthalpy value. Hess’s law could be used to calculate the heat of combustion of magnesium. The magnesium’s heat of combustion values of our group are quite close from the expected value. There was a small percent error of 9.88% and that was due to incorrect lab procedures. The incorrect lab procedures might be the reason for the percent error of our results.
REFERENCES 1. Fowler,R., Guggenheim, E.A. (1939/1965). Statistical Thermodynamics. A Version of Statistical Mechanics for Students of Physics and Chemistry, Cambridge University Press, Cambridge UK, page 57 2. Preston, T (1894/1904). The Theory of Heat, Second Edition, revised by J.R. Cotter, Acillan, London, pages 700-701. 3. Polik, W (1997). Bomb Calorimetry. Retrieved from http://www.chem.hope.edu/~polik/chem345-2000/bombcalorimetry.htm
