Hess's Law Lab | Lab Reports Chemistry

HESS’S LAW LAB

Pre lab assignment: You will need to complete the following parts prior to doing the

lab: Title, Purpose, and Storyboard of the procedures for each part, Blank Data

tables, and the Prelab questions (below).

Chemical Background Information:

Magnesium reacts with oxygen in air to for magnesium oxide, according to

equation 1.

Mg(s) + ½ O2 (g)  MgO(s) + heat Equation 1

In the reaction above, a great deal of heat and light are produced – the

temperature of the flame can reach as high as 2400 oC. The amount of heat energy

produced in this reaction cannot be measured directly using standard methods in the

high school lab. It is possible, however, to determine the amount of heat produced by

an indirect method, using Hess’s Law.

The heat or enthalpy change for a chemical reaction is called the H reaction. The

enthalpy change – defined as the difference in enthalpy between the products and

reactants – is equal to the amount of heat transferred at a constant pressure and does

not depend on how the transformation occurs. This definition of enthalpy makes it

possible to determine the heats of reaction for reactions that cannot be measured

directly. According to Hess’s Law, if the same overall reaction is achieved in a series of

steps, rather than in one step, the enthalpy change for the overall reaction is equal to

the sum of the enthalpy changes for each step in the reaction series. Consider the

following three reactions:

Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g) Equation

MgO(s) + 2HCl(aq)  MgCl2(aq) + H2O(l) Equation B

H2(g) + ½ O2(g)  H2O(l) ∆H = -285.8 kJ Equation C

It is possible to express the combustion of magnesium (Equation 1 above) as an

sum of Equations A, B, and C.

Purpose

The purpose of this experiment is to use Hess’s Law to determine the heat of

reaction for the combustion of magnesium (Equation 1).

Materials

Hydrochloric acid, HCl, 1 M, 60 mL Graduated Cylinder, 25- or 50-mL

Magnesium ribbon, Mg, 7-cm strip Metric ruler, marked in mm

Magnesium oxide, MgO, 0.04 g Scissors

Balance, centigram (0.01 g precision) Spatula

Calorimeter, small-scale Stirring rod

Digital thermometer or temperature sensor Wash bottle and water

Forceps Weighing dish

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HESS’S LAW LAB

HESS’S LAW LAB

1. Obtain a 7-cm strip of magnesium ribbon and cut it into two pieces of

unequal length, roughly 3- and 4-cm each. Note: Handle the magnesium

ribbon using the forceps.

2. Measure the exact length of each piece of magnesium ribbon to the

nearest 0.1 cm.

3. Multiply the length of each piece of Mg ribbon by the conversion factor

(g/cm) provided by your teacher to obtain the mass of each piece of Mg.

4. Mass a clean, dry calorimeter to the nearest 0.01 g.

5. Using a graduated cylinder, add 15 mL of 1 M hydrochloric acid to the

calorimeter and measure the combined mass of the calorimeter and

acid.

6. Using a digital thermometer or temperature sensor, measure the initial

temperature of the hydrochloric acid solution to the nearest 0.1°C.

7. Add the first (shorter) piece of magnesium ribbon to the acid and stir the

solution until the magnesium has dissolved and the temperature of the

solution remains constant.

8. Record the final temperature of the solution to the nearest 0.1°C.

9. Rinse the contents of the calorimeter down the drain with excess water.

10. Dry the calorimeter and mass it again to the nearest 0.01 g.

11. Repeat steps 5-9 using the second (larger) piece of magnesium

ribbon.

12. Mass a clean, dry calorimeter to the nearest 0.01 g.

13. Using a graduated cylinder, add 15 mL of 1 M HCl to the calorimeter

and measure the combined mass of the calorimeter and hydrochloric acid.

14. Mass a small weighing dish and add about 0.20 g of magnesium oxide.

Measure the exact mass of magnesium oxide and the dish to the nearest

0.01 g.

15. Using a digital thermometer or a temperature sensor, measure the

initial temperature of the hydrochloric acid solution to the nearest 0.1°C.

16. Using a stirring rod, add the magnesium oxide to the acid. Stir the

reaction mixture until the temperature remains constant for several five-

second intervals. Record the final temperature of the solution to the nearest

0.1°C.

17. Pour the reaction mixture down the drain with excess water. Rinse and

dry the calorimeter.

18. Repeat steps 12-16 using a second sample of magnesium oxide.