Thermochemistry and hess's law lab conclusion

Hess’ Law

Peter Jeschofnig, Ph.D.

Version 42-0158-00-01

Review the safety materials and wear goggles when

working with chemicals. Read the entire exercise

before you begin. Take time to organize the materials

you will need and set aside a safe work space in

which to complete the exercise.

Experiment Summary:

Students will have the opportunity to measure

temperature changes taking place in a calorimeter

during neutralization reactions and use the

measurements to calculate enthalpy of reaction.

They will illustrate the validity of Hazy’ Law by

comparing the values of enthalpy of two chemical

reactions.

Objectives

●● To measure temperature changes taking place in a calorimeter during neutralization reactions

and use the measurements to calculate enthalpy of reaction.

●● To compare the enthalpy of two chemical reactions and use these measured values to illustrate

the validity of Hess’ Law.

Materials

Materials From: Label or

Box/Bag: Qty Item Description:

Student Provides Distilled water

Watch

Coffee cups

Paper towels

From LabPaq 1 Thermometer - Digital

1 Goggles-Safety

4 Cup, Styrofoam, 8 oz

1 Cylinder-25-mL

From Experiment Bag

Hess' Law 2 Ammonia , NH3 (comes as aqueous

ammonia, NH4OH), - 2 M - 10 mL

2 Ammonium chloride, NH4Cl - 2M - 10mL

2 Hydrochloric acid, HCl - 2 M - 20 mL

2 Pipet, Long Thin Stem

2 Sodium hydroxide, NaOH - 2M - 20 mL

Note: The packaging and/or materials in this LabPaq may differ slightly from that which is listed

above. For an exact listing of materials, refer to the Contents List form included in the LabPaq.

Discussion and Review

Thermochemistry is the study of the heat energy involved in chemical reactions and changes of physical state. Nearly all chemical reactions involve the release or absorption of heat, a form of energy. The burning of any fuel such as gasoline, coal, or wood is an example of a heat-releasing reaction. Heat energy is called thermal energy, and it is always spontaneously transferred from hotter to colder matter.

The First Law of Thermodynamics is the Law of Energy Conservation. It states that the total energy of the universe must remain constant. Therefore, all energy transferred between a system and its surroundings must be accounted for as heat or work.

The standard S.I. unit for heat energy is the joule, J. It takes 4.184 joules, the equivalent of 1

calorie, to raise the temperature of one gram of water by 1° C. The kilojoule, kJ, is commonly used in many applications: 1000 joule = 1 kilojoule.

When a chemical reaction takes place in a stable environment where the temperature and

pressure remain constant, the system defined by the reactants and products either produces or

releases heat energy.

●● If the reacting system releases heat energy to its surroundings, a concurrent increase in

surroundings temperature is observed, and the reaction is exothermic

●● If the system absorbs heat energy from its surroundings, a decrease in the surroundings

temperature is observed, and the reaction is endothermic.

●● A measure of the amount of heat given off or absorbed in any chemical reaction is called the

enthalpy change or heat of reaction, and is given the symbol H.

When thermodynamic measurements are carried out at standard-state conditions where the

pressure is constant at 1 atm and the temperature is constant at 25oC, the reaction enthalpy is

designated as the standard enthalpy change or ΔH°. It is important to have standardized values because the enthalpy of a reaction can vary with different reaction conditions.

The following reaction for the formation of water from its constituents is exothermic:

H2(g) + ½ O2(g) à H2O(l); ΔH °f = -286 kJ

For every mole of H2O (l) formed at standard-state conditions, 286 kilojoules of heat energy are

released. When the standard enthalpy change of reaction describes the formation of 1 mol of

compound directly from its elements in their standard states as in this example, the value of ΔH of is called the standard heat of formation.

To determine the enthalpy change for a given reaction (ΔH°rxn), the summation of the heats of

formation (ΔH° f ) for the reactants are subtracted from the summation of the heats of formation ( ΔH ° f ) for the products.

ΔH° rxn = [n ΔH°f (products)] - [n ΔH°f (reactants)]

Tables containing the standard heats of formation for a number of compounds are available in the appendices of any general chemistry textbook.

Hess's Law states that if a reaction is the sum of two or more other reactions, the ΔH for the

overall process must be the sum of the ΔH values of the constituent reactions.

Enthalpy change (ΔH) is independent of the path that a reaction follows to move from reactants

to products. It only depends on the relative energy difference between the reactant and product

molecules at constant pressure. Enthalpy change is referred to as a state function due to its

independent of pathway. Since the enthalpy of a substance is not commonly determined, the

change in enthalpy when reactants are converted to products is often used to describe a chemical

or physical process.

The thermal energy absorbed or produced by a chemical process reflects a difference between

the enthalpy between the reactants and products (ΔH). For example, in the decomposition of

liquid water into its component elements, H2 (g) and O2 (g), there are two successive changes.

First, the liquid water is vaporized. Second, the water vapor decomposes into its constituent

elements shown below. The ΔH value for this overall process can be determined by adding the

ΔH values from the equations for each step as shown below.

(1) H2O (l) àH2O (g); ΔH 1 = +44 kJ

(2) H2O (g) àH2 (g) + ½ O2 (g); ΔH 2 = +242 kJ

_______________________________________________________________

(1) + (2) H2O (l) àH2 (g) + ½ O2 (g); ΔHnet = +286 kJ

In order to determine ΔH for the reaction NH3 + HCl àNH4Cl in this experiment, ΔH rxn for the

following two reactions will be measured:

1. NaOH (aq) + HCl (aq) àH2O (l) + NaCl (aq)

2. NaOH (aq) + NH4Cl (aq) àNH3 + NaCl + H2O (l)

Comparison of the calculated results for different parts of the experiment will verify the

generalization known as Hess's Law of Constant Heat Summation. In this case the target reaction NH3 + HCl àNH4Cl can also be performed directly and the results compared to reactions 1 and 2.

A Styrofoam coffee cup calorimeter will be used to measure the amount of heat energy evolved

or absorbed during the chemical reactions of this experiment. A digital thermometer is used to

measure the change in temperature between the final and initial temperatures of the solutions.

Unfortunately, it is impossible to have perfect insulation and some of the heat energy will be lost to the surroundings, including to the material from which the calorimeter is constructed.

Calibrating the calorimeter before using it to make measurements on an unknown system usually solves the problem of heat losses. A known amount of heat energy from a known process is released into the calorimeter system, and the temperature change is measured. A simple calculation is done to determine the amount of heat energy loss, called the heat capacity of the calorimeter or calorimeter constant. For this experiment it assumed that the heat capacity of the calorimeter is insignificant and it is ignored.

Another practical problem is that heat energy exchanges do not occur instantaneously; i.e., it takes time for energy to move from a hot object to a cold one. An acceptable solution to this problem is to obtain a cooling curve for the heat energy exchange in question and then extrapolate the data back to the exact time that the exchange began.

Below is a sample graph from hypothetical data. Notice that at the time of combining the

two solutions, their starting temperature is 20oC. Since the starting temperatures are at room

temperature no initial temperature adjustment is needed. From 0 to 40 seconds the temperature

rises rapidly to 34.2oC. The temperature then drops gradually 31.1oC and will continue to drop.

Usually recording the temperature in 20-20 second intervals for 5 minutes is enough to provide a

good cooling curve. Extrapolation of these data backward in time determines what the temperature

at the time of mixing would have been if the temperature of the reaction had been instantaneous

and the calorimeter had warmed instantaneously. In this example, the temperature at the time

of mixing determined by extrapolation is 34.3oC.