USING ENGLISH TO REPORT

this’s the example of report, the report can be made if someone has been doing a practicum and make it in the format below

 Experiment 9 ꜜ

Thermochemistry : Application of Hess’s Law and the Enthalpy of Formation of Magnesium Oxide

 Purpose:

Using Hess’s Law and simple calorimetry, the enthalpy of formation for magnesium oxide (∆HfMgO) can be determined to a first approximation.

 Background:

An understanding of the heat changes that accompany a chemical reaction is of fundamental importance in chemistry. Energy changes that occur during ordinary chemical reactions can be complex. Chemical reactions involve the breaking of chemical bonds in a given set of reactants, and the formation of other different chemical bonds in a given set of products. A thorough examination of the many roles that energy changes have in chemical processes can lead to insights into other chemical phenomena, such as:

♦   Chemical Kinetics → the physical mechanisms whereby reactants are converted to products.

♦   Reactions Rates → the “speed” of a given process that converts reactants into products.

♦   Chemical Thermodynamics → the changes in energy that occur during chemical processes as

    a function of absolute temperature.

A common measurement used when discussing the relationship between energy and chemical changes is the enthalpy change for a chemical process, H. The enthalpy change of a chemical reaction is defined as the amount of heat exchanged by that chemical reaction conducted at a constant pressure. It is a measure of the difference between the “heat contents” of the products and the reactants.

ΔH _rxn = ΔH _final - ΔH _initial = ΔH _products - ΔH _reactants

A chemical reaction occurs within a focused finite part of the universe, an environment called a “system.” Any energy released in the reaction or absorbed in order to initiate the reaction comes from an external environment called the “surroundings.” Thus, the reactants and products constitute the system and everything else, such as the reaction container, room, etc. make up the surroundings.

The sign of the quantity ∆H indicates the general direction of the energy flow into or out of a reaction system. If ∆H has a negative sign, heat has been transferred from the system to the surroundings in an “exothermic” reaction. If ∆H has a positive sign, heat has been transferred from the surroundings to the system in an “endothermic” reaction

Enthalpy, along with “internal energy (E)” and “entropy (S)” are referred to as being “state” functions. In thermodynamics, a state function, function of state, state quantity, or state variable is a property of a system that describes quantitatively the equilibrium of the system irrespective of how the system arrived in that state. Thus, the temperature, pressure, and volume of a system would also be considered state functions.

Enthalpy is also an “extensive” property, which means the amount of energy exchanged in the reaction is a function of the amount of substance being investigated.

Enthalpy is itself a general term for heat transfer. In the context of chemical reactions, it is common to be able to measure the following types of heat transfer.

♦  Enthalpy of Formation (∆Hf): The quantity of heat involved when forming 1 mole of a given

    substance in its standard state, directly from the elements that comprise the substance in their

    standard states (STP).

♦ Enthalpy of Combustion (∆H): The quantity of heat transferred per mole of a combustible

   substance, upon its reaction with excess oxygen, i.e., burning.

♦ Enthalpies of Solution (∆Hsol), Vaporization (∆Hvap), Fusion (∆Hfus), or Sublimation

   (∆Hsub): The quantity of heat involved when a substance changes its physical state (phase) as

    it melts, vaporizes, dissolves in water, or changes directly from a solid to a gas.

♦ Enthalpy of Neutralization (∆Hn): The quantity of heat involved when 1 mole of water is

   produced by the reaction of an aqueous acid and an aqueous base, in water.

Enthalpy, Internal Energy, and Heat

Enthalpy is a measure of the total energy of a thermodynamic system; the energy transferred between the system and its surroundings. It includes the internal energy, E, which is the energy required to create a system, and the amount of energy required to make room for it by displacing its environment and establishing its volume and pressure, i.e. the work (w) involved.

ΔH = ΔE + w = ΔE + PV

The internal energy of a system, E, is precisely defined as the heat at constant pressure (qp) plus any work (w = PV) done by the system.

ΔE = q _p + w = q _p + (- PV)

Internal Energy used to expand volume by increasing pressure is lost to the surroundings, thus the negative sign in the work term.

Thus,

q _p = ΔE + PV

and

p ΔH = q _p at constant pressure

Hess’s Law

This experiment is designed to test and verify Hess’s Law; which stated in words says:

The total energy change for any process that consists of a series of steps is equal to the sum of the enthalpies of the individual steps.

Thus for a given reaction that can be expressed as the sum of four intermediate reactions, Hess’s Law would look something like this:

reaction #1 ----------- > ∆H1

reaction #2 ----------- > ∆H2

reaction #3 ----------- > ∆H3

reaction #4 ----------- > ∆H4

∆Htot = sum of ∆Hs of individual reactions = ∆H1 + ∆H2 + ∆H3 + ∆H4

Heat of Reaction and Heat of Formation

The total enthalpy change for a reaction involving several steps or multiple reactions requires the determination of the individual ∆H values. This can be done in two ways:

1) Measuring the heat of reaction using a calorimeter.

2) Tabulating the heat of reaction for a compound formed from its uncombined elements.

Heat of Reaction (experimental)

In order to measure the heat of reaction of a substance utilizing non-standard amounts of reagents, it is necessary to determine the actual heat generated in the reaction. The total heat change (qp) for a reaction carried in a calorimeter at constant pressure is defined as the product of the mass of material present (m, in grams), the specific heat capacity of that material (Cp, in J/g oC), and the change in temperature (∆t) during the reaction.

q _rxn = m x Cp x ΔT

The heat value, qrxn, is based on the actual amount of material present. The enthalpy change of the reaction (∆Hrxn) is based on the amount of heat generated per mole of substance. Thus, under conditions of constant pressure, as in this experiment:

Example:

How much heat is generated when 0.500 g calcium oxide (CaO) is placed in 50.0 mL of water (H2O) in a Styrofoam cup at 23.5oC and the temperature increased to 27.0oC?

balanced chemical equation:

CaO_(s) + H₂O_(l) → Ca(OH)_{2 (aq)}

mass of calcium = 0.500 g

mass of water is calculated from its volume and density (0.998 g/ml)

mass of water = 50.0 mL * 0.998 g/mL = 49.9 g

mass of the system is the mass of calcium plus the mass of water

mass of system = 0.500 + 49.9 = 50.4 g

temperature change (∆T) = 27.0oC - 23.5oC = 3.5oC

specific heat of water = 1.00 cal/g  •deg = 4.184 J/g  •deg

heat of reaction: qrxn = m * Cp * ∆T

rxn 4.184 J q = 50.4 g 3.5 deg = 738 J = 0.738 kJ g • deg

The The “Heat of Reaction” (∆Hrxn in kJ/mol), is the “Heat Evolved” adjusted to reflect the amount of heat that would have been generated by 1 mole of the reactant.

Since heat was liberated to the surroundings, an exothermic reaction, the products must contain less energy than the reactants. By convention then, Hrxn must be written as 82.8 kJ/mol.

Heat of Formation (tabulated)

If one or more of the intermediate reactions involves the formation of a compound from its uncombined elements, the heat of formation of the compound can be computed from Standard Heats of formation, symbolized by ∆H that have been derived and tabulated in reference lists (see the back of your general chemistry text book). Recall that the heat of formation for an uncombined element is by definition, zero (0).

The standard enthalpy of formation or standard heat of formation of a compound is defined as the change of enthalpy from the formation of 1 mole of the compound from its constituent elements, with all substances in their standard states at 1 atm (101.3 kPa) and 25oC (298 K).

For example, the heat of reaction to produce calcium oxide from its elements calcium and oxygen is just the heat of formation of calcium oxide.

The negative sign for the heat of formation value indicates the reaction is exothermic; energy is released to the surroundings.

The other compound needed in our original reaction is water. Its heat or formation from hydrogen and oxygen is -285.8 kJ/mol.

Thus, the heat of formation for Ca(OH)2 (aq) can now be computed.

The difference between the tabulated and experimental heats of reaction is due to experimental error. A percent error can be calculated.

The Experiment:

The goal of this experiment is to determine the heat of formation of magnesium oxide (MgO), using calorimetry and Hess’s Law. The experimental result will be compared to the known heat of formation (∆H ) and the percent error calculated.

The Hess’s Law components will consist of three reactions. The first two will involve the calorimetric determination of the heats of reaction for the reaction between magnesium (Mg)  metal and hydrochloric acid and the reaction between magnesium oxide (MgO) and hydrochloric acid.

Mg(s) + 2HCl(aq) →  MgCl2 (aq) + H2(g) ∆Hrxn (A) (1)

MgO(s) + 2HCl(aq) → MgCl2(aq) + H2O(l) ∆Hrxn (B) (2)

The 3rd reaction requires the standard heat of formation, H (kJ/mol) for the formation of water from hydrogen and oxygen.

H2 + ½ O2 →  H2O ∆H)f(H2O) (3)

The net reaction for the formation of magnesium oxide (MgO) from magnesium (Mg) and oxygen (O) is obtained from the summation of reactions 1 ,2, and 3. Before the summation can be done, the equations must balanced and, if necessary, reversed in order to cancel out all intermediate components that do not participate in the net equation. The only change in this regard is to reverse equation 2. As a result of this reversal, the sign of the ∆Hrxn (B) value must also be reversed.

The Heat of Reaction for the formation of MgO, H (MgO), from oxygen and magnesium is computed from the individual heats of reaction from reactions 1, 2, 3.

(∆H)f (MgO) = ∆H1 + ∆H2 + ∆H3

Note: The reactions of both Mg and MgO with HCl result in the release of hydrogen and heat; thus, they are exothermic reactions and the ∆Hrxn values are negative. However, since reaction (2) was reversed, the original ∆Hrxn (2) value must also be reversed, i.e., it is now positive.

The standard heat of formation for water, ∆H (H2O) (3), is obtained from standard reference tables. Its value is -285.8 kJ/mol

Since ∆H usually does not change significantly with temperature and the data will be obtained at close to standard conditions (1atm, 25oC), H’s and Ho’s can be used interchangeably.

Sample Calculation:

Assume the ∆H(1) and ∆H(2) values from the calorimeter measurements were -496 kJ/mol and -195 kJ/mol, respectively. Taking into account the reversal of the 2nd reaction, the overall Hess’s Law expression for the ∆Hf of magnesium oxide from its elements would look like the following:

Materials and Equipment: Materials	Equipment
Styrofoam cups (2), with plastic cover	thermometer
weighing tray	metal spatula or glass stirring rod
magnesium metal	electronic balance
magnesium oxide	calculator
hydrochloric acid, 1 M

Procedure:

1. Obtain 2 Styrofoam cups and plastic cover with hole

2. Form a calorimeter by placing one cup into the other cup

3. Add about 100 mL (precisely measured to nearest 0.1 mL) of 1.00 M HCl to the calorimeter

Note: Hydrochloric acid is in excess and 100 ml of the acid should be sufficient for all 4 samples (2 samples of magnesium metal and two samples of magnesium oxide).

4. Cover the calorimeter and place a thermometer through the hole in the cover

5. Record and continue to monitor the temperature of the solution

6. Weigh out about 0.2 g of Magnesium (Mg) metal precisely measured to the nearest 0.001 g 

7. Add the metal to the calorimeter all at once and quickly cover the calorimeter

8. Stir the mixture with gentle swirling

9. Monitor the temperature of the solution until a temperature maximum has been reached

10. Record the final temperature

11. Repeat this process for a 2nd weighed sample of Magnesium

12. Weigh out a sample of about 0.5 g of Magnesium Oxide (MgO) precisely measured to the nearest 0.001 g

13. Add the MgO to the HCl solution

14. Stir the mixture with gentle swirling

15. Monitor the temperature of the solution until a temperature maximum has been reached

16. Record the final temperature

17. Repeat the process for a 2nd sample of MgO

18. Clean out the reaction vessel, flushing the solution down the drain with water

Calculations:

Compute the heat of reaction (qrxn) for each trial.

q _rxn = m x Cp x  ΔT

Where: m = mass of the system (metal + HCl soln)

Cp = Specific Heat of HCL soln (water) = 4.184 J/g-deg

∆T = Change in temperature

Note: The aqueous hydrochloric acid solution is the absorbing mass in the calorimeter, but it can be assumed that the specific heat is the same as for water, i.e., both water and HCl absorb the same amount of heat per gram of their mass for each degree of temperature change.

Data Processing:

Use the printed Pre-lab as the laboratory notebook to record the experimental results of the experiment in the appropriate procedure results block. Follow the instructions below to populate the spreadsheet file and setup the algorithms for the Hess’s Law computations.  Summarize the measured and computed laboratory results in the printed copy of the “Hess’s Law Results Summary Table.” If required by the instructor, transfer the laboratory results to the electronic files and finalize the laboratory report.

Spreadsheet Processing:

Use a lab computer computers and the web-based data entry form as shown in Figure 9.1 below to enter the Hess’s Law laboratory results into an Excel spreadsheet.

source :

http://mason.gmu.edu/~jschorni/chem211lab/Chem%20211-212%20Hess%27s%20Law.pdf