Experiment 1 HESS\'S LAW PDF

Preview

Summary

LABORATORY REPORTPHYSICAL CHEMISTRY - CHM 271(EXPERIMENT 1: HESS LAW)PREPARED BY:STUDENT NUMBER:GROUP: A4AS1202_DATE OF EXPERIMENT: 10 MARCH 2020DATE OF SUBMISSION: 20 MAY 2020INSTRUCTOR:OBJECTIVE:To determine the heat of formation of magnesium oxide using Hess’s law.INTRODUCTION:The magnitude of th...

Description

LABORATORY REPORT

PHYSICAL CHEMISTRY - CHM 271 (EXPERIMENT 1: HESS LAW)

PREPARED BY: STUDENT NUMBER: GROUP:

A4AS1202_19

DATE OF EXPERIMENT:

10 MARCH 2020

DATE OF SUBMISSION:

20 MAY 2020

INSTRUCTOR:

OBJECTIVE: To determine the heat of formation of magnesium oxide using Hess’s law. INTRODUCTION: The magnitude of the enthalpy change of a chemical reaction depends only on the difference in the enthalpy content of the products and the reactants, and does not depend on how the reaction is completed. The heat liberated or absorbed during a chemical reaction is independent of the route by which the chemical change occurs, provided the initial and final conditions are the same as stated in Hess’s law. This is another way of expressing the law of conservation energy, which states that energy can neither be created nor destroyed. Hess’s law is useful in calculating those enthalpy changes which cannot be measured from experiments. The standard enthalpy change of the formation of a substance is the enthalpy change when 1 mole of the substance is formed from its elements under standard conditions. The general equation of formation of magnesium oxide is: Reaction 1:

Mg

+

½ O2



MgO

ΔH1 = a kJ mol-1

The standard enthalpy of formation of MgO cannot directly measure from experiment but can be determined by applying Hess’s law. The reaction of magnesium crystal and magnesium oxide crystals with hydrochloric acid respectively will give a certain enthalpy value. The reactions are as below: Reaction 2:

Mg

+

2HCl

 MgCl2

+

H2

ΔH2 = b kJ mol-1

Reaction 3:

MgO

+

2HCl

 MgCl2

+

H 2O

ΔH3 = c kJ mol-1

Reaction 4:

H2

+

½ O2 

H 2O

ΔH4 = d kJ mol-1

We can determine the value for ΔH2 and ΔH3 by using the following Formula 1 and Formula 2. Formula 1

ΔQ

=

ma × c × ΔT

=

(ma × c × ΔT) J

=

[(ma × c × ΔT) ÷ 1000] kJ

Where,

ma c ΔT

Formula 2

= =

mass of acid specific heat capacity = 4.2 J g-1 °c-1

= change in temperature

ΔH (kJ mol-1) = ΔQ/n where n is number of mole of Mg or MgO

The number of mole of Mg or MgO can be calculated by using: Formula 3

n = mass/Mr

Where mass is the mass of Mg or MgO and Mr is the corresponding relative molar mass. Hence, ΔH1 value can be determined by using Formula 4 which is derived by using Hess’s law. Formula 4 ΔH1 = ΔH2 + (-ΔH3) + ΔH4

APPARATUS 

Calorimeter (polystyrene cup or plastic with a lid with a hole to insert thermometer)



Thermometer (110°c)



Stirrer



Burette



Retort stand



Beaker



Electronic balance

CHEMICALS 

2.0 M Hydrochloric acid (HCL)



Magnesium powder



Magnesium oxide powder

PROSEDURE 1. A burette was filled with 2.0 M hydrochloric acid. 2. 30 mL 2.0 M of hydrochloric acid was measured and poured into the calorimeter. Thermometer was inserted. The temperature after a few minutes was recorded. 3. 0.60 g of magnesium powder (limiting reactant) was weighed. The magnesium powder was added into the calorimeter containing hydrochloric acid. 4. The mixture was gently stirred and the highest temperature reached was recorded. 5. Step (1) until step (4) was repeated using 1.00 g magnesium oxide powder instead of magnesium powder. 6. All readings was recorded in Table 2.1

RESULTS

Mass, m (g) Initial temperature (°c) (HCl) Final temperature (°c) (HCl + Mg or MgO) Temperature change (ΔT)

Magnesium (Mg) 0.6039 25.7

Magnesium oxide (MgO) 1.0090 25.2

94.7

50.7

69

25.5

Volume of acid is used = 30 mL Mass of acid, ma = 30 g (assume density of acid is 1 g mL-1)

QUESTIONS: 1. Calculate ΔH2 and ΔH3 ΔH2 ΔQ = - [(30 g) × (4.2 J g-1 °c-1) × (69°c)] = -8694 J = -8.70 kJ n = (0.6039 g) / (24 g mol-1) = 0.02516 mol ΔH2 = -ΔQ / n = (-8.70 kJ) / (0.02516 mol) = -345.79 kJ mol-1 ΔH3 ΔQ = - [(30 g) × (4.2 J g-1 °c-1) × (25.5 °c)] = -3213 J = -3.21 kJ n = (1.0090 g)/ (40 g mol-1) = 0.02523 mol ΔH3 = (-3.21 kJ) / (0.02523 mol) =-127.23 kJ mol-1

2. State whether the reaction is endothermic or exothermic



The reaction is exothermic

3. Calculate ΔH1, the standard enthalpy of formation of magnesium oxide. Given that the enthalpy of formation of water (value of ΔH4) is -286 kJ mol -1

ΔH1 = ΔH2 + (-ΔH3) + ΔH4 = -345.79 + (127.23) + (-286) = -504.56 kJ mol-1

DISCUSSION The purpose of this experiment was to find the enthalpy of formation of magnesium oxide, MgO using Hess’s Law. Through Hess’s Law, we could not determine the enthalpy directl y but can be determine throughout addition of another enthalpy in the reaction. In reaction 2, the enthalpy change was -345.79 kJ mol-1. 345.79 kJ mol-1 was released to surrounding. For reaction 3, the enthalpy change that we managed to obtain was -127.23 kJ mol-1. In this reaction, the reaction was reversed so the sign for the enthalpy also change. To obtain the enthalpy change of magnesium oxide, ΔH1 add all the enthalpy of the reaction together. The enthalpy changes for ΔH1 are -504.56 kJ mol-1. This experiment was conducted using simple calorimeter. Simple calorimeter was used to measure enthalpy change for reactions that take place in aqueous solution. There is a possibility to the loss of heat to the surrounding during the reaction take place. Thus the right value of MgO could not be obtained. This is because value that we obtain through this experiment was less than the actual value of MgO.

CONCLUSION During the experiment, HCl were added to produce an exothermic reaction which results in an increase of heat. The rise in temperature was measures with the calorimeter and the enthalpy change was calculated through Hess’s Law. We discovered that the change in enthalpy is the same weather the reaction takes places in one steps or series of steps. This known fact is called the Hess’s law which can be proven by the concept of additional of heats of reaction....