Thermochemical Reports and Hess Laws

I.       Title: THERMOCHEMICAL AND HESS LAWS

II.    Purpose:      

      1.      Measuring caloric reaction with simple tool.

      2.      Collecting and analyzing thermochemical data.

      3.       Apply the Hess law.

III. Practice Questions:

    1.      Provide an understanding of: (a) enthalpy; (B) isolated systems; (C) open system; (D) closed system; (E) the environment; (F) calorimeter; (G) exothermic.

Answer:

A.    Enthalpy is the sum of the energy of all forms of energy possessed by the substance whose quantities can be measured.

B.     An isolated system is a system that can not drain energy and material to its surroundings.

C.     Open system is a system that allows the flow or exchange of energy and material with its surroundings.

D.    A closed system is a system whose limits can be passed by energy but not to matter.

E.     The environment is outside the system.

F.      A calorimeter is a device used to measure the heat changes of a reaction.

G.    Exothermic is a chemical reaction that releases heat or gives heat to the environment.

      2.      What is the difference between enthalpy and inner energy (ΔE)?

Answer:

Enthalpy is the quantity of heat that is absorbed at a constant pressure, while the inner energy (ΔE) is the sum of both kinetic energy and potential energy possessed by a substance or system.

IV.            Theoretical Basis

Thermodynamics is a branch of science that studies changes in energy chemically or physically. In this experiment, we will investigate the energy changes in the form of heat, which accompany the chemical reaction (thermochemistry). According to the law of thermodynamics, the energy changes that accompany the change of being are expressed in the formula:

ΔE = Q – W

With Q = the heat absorbed by the system

        W = work done by the system

Most chemical reactions take place at a fixed working pressure formulated by the equation:

W = P. ΔV

      With P = gas pressure, ΔV = volume change for the gas system due to fixed pressure.

ΔE = Q - P. ΔV

      When ΔV = 0, then ΔE = Q

The quantity of heat absorbed at a fixed pressure is called enthalpy (ΔH) (Epinur et al., 2011: 36).

The heating studies produced or required by chemical reactions are called thermochemicals. Thermochemistry is a branch of thermodynamics because the test tube and its contents form the system. Thus, we can measure (indirectly, by measuring the work or rise in temperature) of energy generated by the reaction as heat and depending on whether it is with changes in energy or enthalpy changes. Conversely, if we know ΔV or ΔH a reaction, we can predict the amount of energy it produces as heat.

The law of Hess I "enthalpy as a whole is the number of enthalpy reactions and individual reactions that are part of a reaction". In the thermodynamics of this law the value of ΔH is independent of its course, in the sense that we can react certain reactants through various reactions that produce certain products and overall obtain the same enthalpy changes (P.W.Atkins, 1996: 47-54).

The application of the laws of thermodynamics I to chemical events is called thermochemistry, which deals with the heat accompanying chemical reactions. Chemical reactions including isothermal processes and when carried out in the open air then the heat of reaction "QP = ΔH", consequently the heat can be calculated from the change of enthalpy of reaction

"Q = ΔH reaction = H reaction product - H reeaksi",

 so there is uniformity must be set the standard state of temperature 25˚C and tea 1 atm. Thus, thermodynamic calculations are based on standard pressure: AB + CD AC + BD ΔH = x kj / mol  ΔH is the symbol (notation) of the reaction enthalpy changes in the food. Judging from the reaction type, there are 4 types of heat as follows:

1. Calor formation

2. Heat decomposition

3. Heat neutralization

4. Heat reaction (syukri, 1999: 84 - 85)

                                                                                          ( syukri, 1999 : 84 – 85 )

The reaction calorimeter experiments were determined with lavoiser and lapiance calorimeters in1780. It was realized that the heat absorbed in the compounding reaction already compounds is as large as the heat released in the formation reaction under the same conditions. That is, if a chemical reaction of its direction is written reverse then the quantity of ΔH reactions becomes the opposite too. Germain Henri Hess in 1048 suggests that the heat associated with a certain type of reaction is constant and independent of the path of reaction toward the reaction it enters, provided that it isother and at constant P, the condition is known as Hess's law. By Hess's law this is the way for us to calculate the value of ΔH from the kind we can not get experimentally (M. Suwandi, 1995: 77-78).

Energy is usually defined as the capacity to perform work. In an isolated system, the sum of the total energy is constant α through the process of the chemical energy being converted into other forms, such as mechanical energy (motion) of electrical energy, light energy, etc. This understanding raises the notion of energy conservation law. This law states that no energy can be created or destroyed but energy can be from one form to another, eg potential energy into electrical energy (sri sudiono, et al, 2005: 42).

V. Tools and Materials

A. Tool

Ø Measuring cups

Ø Calorimeter

Ø Mixer

Ø Cup of trophies

B. material

Ø 40 mL of distilled water

Ø 40 mL HCl 1 M

Ø 40 mL NaOH 1M

Ø 1M acetic acid

Ø Sodium Hydroxide 1M

Ø Sodium Acetate 1M

Ø 1M Nitrate Acid

Ø Ammonia 1M]

VI.Works Procedures

       A.     Determination of calorimeter constants

40 ml of distilled water

- Measured with measuring cups

- Poured chalorimeter

The completed calorimeter & stirrer tool is closed

- Recorded its contents

40 ml of distilled water

- Measured again with measuring cups

- Poured into dry cup

- Heated 60oc-70oc

- Measured hot water temperature appropriately

- Rapidly displaced kekalorimeter

- Recorded every 15 seconds while stirring

The temperature of the maximum solution and slowly decreases, recorded every 1 minute until constant

      B.     Determination of ΔH of neutralization for acid-base

The calorimeter is dried

40 ml of 1M NaOH solution

- Measured and inserted calorimeter

40 ml 1M HCl

- Measured in 150ml cup glass

- Close near the calorimeter

- Temperature of acid and base solutions is measured

- The temperature of the two solutions should not be at odds of 0.5oc, when different temperatures are equalized

- When the temperature is the same, put it quickly into the calorimeter

Observation result

          VII. Observation Data

      A. Determination of calorimeter constants

	Ulangan 1	Ulangan 2	Rata - rata
Hot water temperature,°C Cold water temperature,°C Mixed temperature,°C	60°C 28°C 39°C	57°C 26°C 38°C	58,5°C 27°C 38,5°C

       B. Determination of ΔH of neutralization for acid-base

1.      NaOH + HCl 1M

	Ulangan 1
Temperature of acid solution The temperature of the base solution Mixed temperature	29,5°C 30°C 39°C

VIII. DISCUSSION

A. The calorimeter constant

In this experiment the calorimeter is cleaned and dried. Enter 20 mL of distilled water into the calorimeter, record the weight and measure the temperature, then take 40 mL of distilled water with a measuring cup, heat and record the hot water weight. Combine hot water and cold water into the calorimeter, note the temperature. Result of experiment of determination of calorimeter constant, by formula:

C.Mp (Tp-Tm)      =   C.Md(Tm-Td) + W(Tm-Td)          

   4,184 J/g°C.20(60°C-39°C)     =   4,184 J/g°C.20(39°C-28°C) + W

                                                         (39°C-28°C)

                4,184 J/g°C (21°C)      =   4,148 J/g°C (11°C) + W (11°C)

                                 87,864 J      =   46,024 J + 11 W

 11 W                                           =   87,864 J – 46,024 J

 11 W                                           =   41,84

                                           W      =   3,803 J°C

B. Determination of ΔH of neutralization for acid-base

In this experiment we made observations of different mixtures of acid-base solutions. After conducting an experiment in accordance with work procedures, the data obtained are:

The temperature of the acid solution (HCl) 1M = 29.5 ° C

The temperature of the basic solution (NaOH) = 30 ° C

Mixed temperature = 39 ° C

And get it:

Qreaksi = C. M. (Tf-Ti) + W (Tf-Ti)

            = 4.184.80 (30 -29.75) + (39 -29.75) .3,80

              = 334.72. (9,25) + 35,17

              = 3131.33 J

Qreaksi = -Range

-Round about = 3131.33 J

IX. Conclusion

1. The heat of the reaction can be measured using a device called a calorimeter

2. By using the calorimeter obtained analysis and data calorimeter constant is 52.997 J / ˚C

3. To determine the calorimeter constant then we need the weight data of a substance, heat of that type of substance. The equation temperature of the equation determines the calorimeter constant based on the black principle ie Qepas = Qterima.

4. The calorific equation of the reaction can be done by using:

A. Hess's Law

B. Standard enthalpy entries

C. Average bond energy

XIII. Bibliography

Afkins, P.W.1999. Chemical Physics Issue 4 Volume 1. Oxford: University Press

Epinur et al. Guide to Basic Chemical Practicum. Jambi: University of Jambi

Suwandi, W.1985. Chemistry Physics. Jakarta: Pt. Bina Aksara

Sudiono, Sri.1990. Physical Chemistry And Problems. Jakarta: Universirtas Indonesia

Shukri.1999. Basic chemistry. Bandung: Bandung Institute of Technology