REPORT : TERMOKIMIA AND LEGAL HESS

PRACTICAL TITLE
TERMOKIMIA AND LEGAL HESS
OBJECTIVES OF EXPERIMENT
1. Measuring the heat of reaction with a simple tool.
2. Collecting and analyzing thermochemical data.
3. Apply the Hess law.
THEORETICAL BASIS
Thermodynamics explains the relationship between heat and other forms of energy. Its development, which was an important scientific achievement in the 19th century, was due to the efforts of physicists and engineers who wanted to achieve high efficiency in a heat engine. The interest in improving the engine once again becomes important because of the need to use fossil fuels effectively. However, in the last 75 years, the important application of thermodynamics is in the field of chemistry. The law of thermodynamics is an important tool for studying chemical reactions. Thermochemistry is the influence of the heat that accompanies the chemical reaction. The second law of thermodynamics is primarily the basis for deriving the equilibrium constant of the properties of thermodynamic properties, in the third law of thermodynamics will be unveiled the starting point for looking at the properties of experimental thermodynamic properties (Petrucci.1987: 225).

Each system has energy because the material particles (solid, liquid, or gas) always move randomly and diverse. There is translational motion, rotation, and ubrasi (vibrate). In addition, there may be a shift in the energy levels of electrons in atoms or molecules. Each movement, influenced by many factors and can change shape when colliding with each other. As a result, the energy of a particle's ganga will be different from the others. The total energy of all the particles will be different from the others. The total energy of all the particles in the system is called the inner energy (U). Therefore, the absolute value of U can not be calculated.
The first law of thermodynamics discusses the energy changes that accompany the event, and is useful for calculating the incoming or outgoing heat of the system. By equation: q = aU - w. The second law, to be discussed about spontaneous and non-spontaneous changes. The second law of thermodynamics sounds the natural process of adding natural entropy or entropy as it increases, and the third thermodynamic law sounds a pure element or compound in the form of perfect crystals having zero entropy at 0 ° C (Syukri.1999: 74).

The application of the first law of thermodynamics to chemical events is called thermodynamics, which deals with the heat that accompanies chemical reactions. Chemical reactions include isothermal processes, and when done in the open air then the reaction calories qp = ΔH. Consequently, the heat can be calculated from the enthalpy change of the reaction q = ΔH = Result H-reaction. So that there should be a uniformity of standard conditions, ie temperature 25 ° C and pressure 1 Atm. Thus, thermodynamic calculations are based on standard conditions (Shukri.1999: 84).

Thermochemistry is a part of thermodynamics that studies heat changes that follow chemical reactions. The amount of heat that arises or is required in a chemical reaction is called reaction heat. The heat of the reaction at P remains the same as the change in the ental strength, and the heat of reaction in U remains the same as the change in power.
The magnitude of the reaction pans depends on the type of reaction, the phase state of the substances in the reaction, the amount of the reacting agent, and the reaction temperature. In thermodynamic equations, the amount of substances in the reaction is expressed in moles while the heat is expressed in Kilocalories (Sukardjo 1990: 192).
· Change of enthalpy (ΔH) and internal energy changes (ΔU) in chemical reactions. The reaction at constant pressure ΔH and heat of reaction at constant volume, ΔU is connected through the equation: ΔU = ΔH-PΔU If the heat of reaction is carved under constant pressure conditions at a constant temperature of 298 K, it is -566.0 KJ, indicating that the 566.0 KJ energy has left the system as heat ΔH = -566.0 KJ. To evaluate the pressure-volume work PΔU = P (Vt -Vi) Then we can use the ideal gas equation. The kinetic equation PΔV = RT (nf-ni) Nf is the mole of gas in this product (2 moles CO2) is the number of moles of reactant solid gas (2 moles CO + 1 mole O2), so PΔV = 0,0083145 KJ / mol K-1. 298 K × [2 (2 + 1)] mol         = -2.5 KJ Internal energy change is ΔU = ΔH-PΔV      = -566,0KJ - (- 2,5KJ)      = -563.5KJ · Change of enthalpy (ΔH) accompanying material changes. When the liquid comes into contact with the atmosphere, energized molecules on the liquid surface can overcome the attraction with each other and enter the form of gas or steam. · Indirect determination ΔH: Hess's law of equilibrium of enthalpy changes ΔH is an extensive nature Hess's law concerning the sum of the constant heat (Petrucci.1992: 239-244). G.H.Hess a Swiss chemist in 1840 investigated whether a chemical reaction taken through several streets would affect the heat of his reaction. To answer the problem Hess did some experiments. The following presented data from experiments investigating the heat of sulfur formation reaction. S (P) + 11 / 2O2 → SO3 (Q) This SO3 formation reaction can be done through several ways. From each experiment will be investigated the amount of reaction heat generated. 1. Direct way S (S) + 3 / 2O2 (g) → SO3 (g) ΔH = -395.73 2. The indirect way S (S) + O2 (g) → SO2 (g) ΔH = -296.83 SO2 (g) + 1/2 O2 (g) → SO3 ΔH = -98.9 S (S) + 3 / 2O2 (g) → SO3 (g) ΔH = -395.73 In general hhhuuukum Hess can be expressed by the following reaction equation A D                            ΔH1  ΔH2 ΔH3        B ΔH4 C     ΔH = ΔH2 + ΔH3 + ΔH4 For the reaction of A to D, the change in the ental strength is equal to the amount of change in enthalpy of each reaction step (Agus.1987: 123). Standard enthalpy change (ΔHo) Some types of standard enthalpy changes, namely:
A. A. Change of enthalpy of standard formation (ΔHfo) It is an enthalpy change that occurs in the formation of 1 mole of a compound of the most stable elements in the satandar state. B. Change of standard decomposition enthalpy (ΔHdo) It is an enthalpy change that occurs on the decomposition of 1 mole of a compound into its most stable elements in the standard state. C. Change of standard burning enthalpy (ΔHoc) Is the enthalpy change that occurs in the burning of 1 mole of a substance perfectly. Burning is the reaction of a substance with oxygen, thus when a substance is completely burned and it contains: - C à CO2 - H à H2O - S à SO2 (Susanto.2003: 46).

TOOLS AND MATERIALS
        A. Tool
Ø Measuring cups
Ø Calorimeter
Ø Mixer
Ø Cup of trophies
B. material
Ø 40 mL 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

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:

1. 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

The equation of the reaction

HCl + NaOH → NaCl + H2O

Then ΔH = reaction -Range around

                                    The reacting mole

                               = 3131.33 J

                                      0.02 mol

                               = 156566,5J / mol

2. The temperature of 1M CH3COOH acid solution: 26 ° C The temperature of 1M NaOH base solution is: 30 ° C Mixed temperature: 27 ° C To get Qreaksi used formula: Qreaksi = C. M. (Tf-Ti) + W (Tf-Ti)               = 4.184J / g ° C.80gr (27 ° C-28 ° C) + (27 ° C-28 ° C)                    . 3,803 J / ° C               = -334,72 J / ° C. (-1 ° C) + (- 3,803 J)               = -338,523 J -Qseconds = Qreaksi                    = -338,523 J The reacting mole CH3COOH + NaOH → NaCH3COO + H2O Then ΔH = reaction -Range around                                     The reacting mole                                = 338,523 J                                       0.02 mol                                = 16926.15 J / mol 3. The temperature of the acid solution (NaCH3COO) 1M: 23 ° C The temperature of 1M: 23 ° C basic solution (HCl) Mixed temperature: 25 ° C To get Qreaksi used formula: Qreaksi = C. M. (Tf-Ti) + W (Tf-Ti)               = 4.184J / g ° C.80gr (25 ° C-23 ° C) + (25 ° C-23 ° C)                    . 3,803 J / ° C               = 334.72 J / ° C. (2 ° C) +3,803 J / ° C (2 ° C)               = 677,046 J The equation of the reaction NaCH3COO + HCl → NaCl + CH3COOH Then ΔH = reaction -Range around                                     The reacting mole                                = 677,046 J                                       0.02 mol                                = 33852,3 J / mol

4. Temperature of 1M HNO3 acid solution: 31 ° C The aqueous solution temperature of NaOH is 1M: 32 ° C Mixed temperature: 31 ° C To get Qreaksi used formula: Qreaksi = C. M. (Tf-Ti) + W (Tf-Ti)               = 4.184J / g ° C.80gr (31 ° C-31,5 ° C) + (31 ° C-31,5 ° C)                    . 3,803 J / ° C               = 334.72 J / ° C. (-0.5 ° C) + (- 1,9015 J)               = -169,2615 J -Qseconds = Qreaksi                    = -169,2615 J Mol that reacts → mol = MV                                                   = 1 mol / V. 0.021 V                                                   = 0.021 mol The equation of the reaction HNO3 + NaOH → NaNO3 + H2O Then ΔH = reaction -Range around                                     The reacting mole                                = -169,2615 J                                       0.02 mol                                = 8403,075 J / mol 5. Temperature of 1M HCl acid solution: 31 ° C The temperature of NH3 base solution is 1M: 27 ° C Mixed temperature: 29 ° C To get Qreaksi used formula: Qreaksi = C. M. (Tf-Ti) + W (Tf-Ti)               = 4.184J / g ° C.80gr (29 ° C-29 ° C) + (29 ° C-29 ° C)                    . 3,803 J / ° C               = 0 J The equation of the reaction HCl + NH3 → NH4Cl Then ΔH = reaction -Range around                                     The reacting mole                                                 = 0 J / mol

CONCLUSION

1. In measuring the reaction can use a simple calorimeter.

2. To determine the calorimeter constant, it takes data in the form of weight, heat type, and temperature of the substance. To determine the calorimeter constant based on the black principle, the equation is:

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

information:

C = Water heat type 4,148 J / g ° C

Mp = Weight of hot water

Md = Cold water weight

Tp = Hot water temperature

Td = Cold water temperature

Tm = mixed temperature

W = The calorimeter constant J / g

3. For ΔH the reaction is used equation

ΔH = reaction

Where, Q = = =

                                  = Qlarutan + Qkalorimeter

                                  = C. Mtot (Tf-Ti) + W (Tf-Ti)

4. Calculation of heat of reaction can be done by using:

A. Hess's Law

B. Standard enthalpy entries

C. Average reactant energy

DAFTAR PUSTAKA

Agus. 2009.Kimia Dasar Universitas.Jakarta:Erlangga.

Petrucci, Raip H. 1992. Kimia Dasar. Jakarta: Erlangga

Sukardjo. 1990. Kimia Dasar. Jakarta : Erlangga

Susanto. 2003. Panduan Belajar Sukses SPMBPTN. Yogyakarta: UGM

Syukri. 1999. Kimia Dasar I. Bandung: ITB