enthalpy changes in chemical reactions

7/27/2019 Enthalpy Changes in Chemical Reactions

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ENTHALPY CHANGES IN CHEMICAL REACTIONS

Hess’s Law

Purpose

Test Hess’s law (which says the heat of reaction is the same whether the

reaction is carried out directly or in several steps) by measuring the

enthalpy changes for the same net chemical reaction carried out by two

different paths.

Make an approximate estimate, by an indirect measurement, of the

enthalpy change for the dissociation of water into hydronium andhydroxide ions.

Pre-Lab PreparationWhat accounts for the energy change in a chemical reaction? The simplest

explanation is that in a chemical reaction, the positively charged nuclei and

the negatively charged electrons of the reacting atoms rearrange

themselves. It is the energy change associated with the rearrangement (or

reconfiguration) of nuclei and electrons that produces the energy change of

a chemical reaction.

Thermochemistry deals with the thermal energy changes that accompany

chemical reactions. These energy changes are usually called heats of reaction. When the reaction is carried out at constant pressure, the heat of

reaction is called the enthalpy changes.

Enthalpy changes may be classified into more specific categories: (1) the

heat of formation is the amount of heat involved in the formation of one

mole of the substance directly from its constituent elements, (2) the heat of combustion is the amount of heat produced when a mole of a combustible

substance reacts with excess oxygen; (3) the heat of solution of a substance

is the thermal energy change that accompanies the dissolving of a

substance in a solvent; (4) the heats of vaporization, fusion and sublimationare related to the thermal energy changes that accompany changes in state;

(5) the heat of neutralization is the enthalpy change associated with the

reaction of an acid and base.

In this experiment, you will measure the enthalpy changes for the following

three reactions:

(Reactions)

In these reactions, the numbers in parentheses specify the molar concentrations of the reactants and products. Water is both the solvent and



a reaction product. In order to distinguish these two roles, we write H2O

when we mean a reaction product and write solvent when water is used to

dissolve the reactants and serve as the reaction medium. You will also see

how the measurement of the enthalpy change of reaction (3) leads to an

approximate estimate of the enthalpy of dissociation of water intohydronium ions and hydroxide ions.

Reactions (1) and (3) can be regarded as neutralization reactions. Reactions(2) could be called a solution reaction, and we will call the enthalpy change

for this process the heat of solution. Note also that if we add reactions (2)

and (3), we get reaction (1). Thus, reaction (1) and the sum of reaction (2)

and (3) represent two different pathways by which we can get from the

same initial state to the same final state.

Experimental ProcedurePreparing the calorimeters. Prepare two calorimeters, ach like that shown in

figure. Using a pen, label the calorimeters as calorimeters 1 and 2.

In each experiment, you will be measuring initial temperatures of two

solutions and the final temperature of the mixed solutions. You must be

careful to use the same thermometer to measure the initial and final

temperatures in an experiment because inexpensive lab thermometers often

give readings that do not perfectly agree. If you are using a thermometer

with 1 C intervals, estimate the temperature reading to the nearest 1/s of adegree.

1. The heat of solution and neutralization of 10 M H2SO4 and 1 M NaOH. Place 50 ml of 1 M NaOH and 45 ml of deionized or distilled

water in calorimeter 1. Measure out exactly 5 ml of 10 M H2SO4 in

a 5 or 10 ml graduated cylinder. Measure the temperature of the 10

M H2SO4 and record the reading. Remove the thermometer and

readjust the volume to exactly 5 ml, if necessary, using a polyethylene transfer pipet (or a medicine dropper). Then rinse and

dry the thermometer and insert it into calorimeter 1, which containsthe NaOH solution. Stir the solution gently; read and record the

temperature. Then add the 5 ml of 10 M H2SO4. Stir for 30 s (or

until the reading is steady or only slowly decreasing) and record the

temperature (to the nearest 0,2 C if you are using a thermometer with

1 C intervals). Neutralize the solution with 5 g of sodium bicarbonate

and discard the solution down the drain. Rinse and dry the cups and

carry out a duplicate set of measurements. Then rinse and dry the

cups in preparation for part 2.

2. The Heat of solution of 10 M H2SO4. Measure 90 ml of deionized or distilled water in a graduated cylinder and put into calorimeter 1.



Measure out 10 ml of 10 M H2SO4 in a 10 ml graduated cylinder.

Measure and record the temperature of the 10 M H2SO4. Remove

the thermometer and readjust the volume to exactly 10 ml, if

necessary. Then rinse and dry the thermometer, place it in

calorimeter 1 (which contains the water), and gently stir. Read andrecord the initial temperature of the water as soon as the reading is

steady. Then pour the 10 M H2SO4 into the calorimeter 1. Continue

to stir for 30 s (or until the reading is steady), then read and record

the final temperature of the solution. Neutralize the solution with 10

g of sodium bicarbonate and discard the solution down the drain.Rinse and dry the calorimeter cups and carry out a set of duplicate

measurements. Rinse and dry the cups to prepare for part 3.

3. The heat of neutralization of 1 M H2SO4 and 1 M NaOH. Place 50

ml of 1 M NaOH in calorimeter 1 and 50 ml of 1 M H2SO4 in

calorimeter 2. Put the thermometer in calorimeter 1 and stir gently

until the temperature reading is steady, read and record thetemperature of the sulfuric acid solution. Then pour the H2SO4

solution in calorimeter 2 quickly and completely into calorimeter 1.

Stir for 30 s (or until the reading is steady) and record the final

temperature after stirring. Neutralize the solution with 5 g of sodium

bicarbonate and discard the solution down the drain. Rinse and drythe cups. Carry out a duplicate set of measurements.

Calculations. For each reaction in parts 1, 2 and 3, determine the

temperature change for the solution in calorimeter 1 and the added

reactant solution by subtracting the initial temperature from the final

temperature.

The enthalpies of the several common reactions or physical

processes can be measured quite accurately with the Styrofoam-cup

calorimeter. Try measuring:

a) Mg(s), Zn(s),

Compare your results to those found from Hess’s law and tabulated

enthalpies of formation.

Does stirring the water in your calorimeter change the temperature?

Try to duplicate Joule’s original experiment to determine themechanical equivalent of heat.

enthalpy changes in chemical reactions

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