AS chem Ch 1.2 Hess Law

AS Chemistry Unit1 Ch 2 Energetics

Topic 1.4 Energetics

1. demonstrate an understanding of the term enthalpy change, ΔH

2. construct simple enthalpy level diagrams showing the enthalpy change

3. recall the sign of ΔH for exothermic and endothermic reactions,

eg illustrated by the use of exo- and endothermic reactions in hot and cold packs

4. recall the definition of standard enthalpy changes of reaction, formation, combustion, neutralization and

atomization and use experimental data to calculate energy transferred in a reaction and hence the enthalpy

change of the reaction.

This will be limited to experiments where substances are mixed in an insulated container, and combustion

experiments

5. recall Hess’s Law and apply it to calculating enthalpy changes of reaction from data provided, selected from a

table of data or obtained from experiments and understand why standard data is necessary to carry out

calculations of this type

6. evaluate the results obtained from experiments using the expression:

energy transferred in joules = mass x specific heat capacity x temperature change

and comment on sources of error and assumptions made in the experiments.

The following types of experiments should be performed:

i experiments in which substances are mixed in an insulated container and the temperature rise measured

ii simple enthalpy of combustion experiments using, eg a series of alcohols in a spirit burner

iii plan and carry out an experiment where the enthalpy change cannot be measured directly, eg the enthalpy

change for the decomposition of calcium carbonate using the enthalpy changes of reaction of calcium

carbonate and calcium oxide with hydrochloric acid

7. demonstrate an understanding of the terms bond enthalpy and mean bond enthalpy, and use bond enthalpies in

Hess cycle calculations and recognize their limitations.

Understand that bond enthalpy data gives some indication about which bond will break first in a reaction, how

easy or difficult it is and therefore how rapidly a reaction will take place at room temperature

1


2.1 Thermochemistry

Thermodynamics - study of the exchange of heat, energy and work between a system and its surroundings

System: the reaction mixture, containing reactants, products and solvents

Surroundings: everything outside the system

For example, in the reaction of HCl and NaOH

System: acid, alkali, salt and water

Surroundings: test tubes, air in the room, the entire rest of the universe

First Law of thermodynamics (or Law of conservation of energy) states that energy may be converted from one form

to another, but it is never created nor destroyed.

In a chemical reaction, the most usual form of energy exchanged is heat energy, which causes temperature

changes. The relationship between energy and temperature is calculated by

energy transferred in joules = mass x specific heat capacity x temperature change

2.2 Enthalpy change

The enthalpy of a substance, sometimes called its heat content, is an indication of the total energy content of the

substance at constant pressure that can be converted into heat.

The absolute enthalpy (H) value of a substance depends on

i. potential energy inherent in the electrical and nuclear interactions of the constituent particles (chemical energy),

ii. kinetic energy possessed by the atoms and sub-atomic particles (kinetic energy)

The absolute enthalpy value (H) of a substance cannot be found directly since it is still not possible to measure all the

interaction between the sub-atomic particles in an atom. Instead, we can measure the enthalpy change (ΔH) of a

system through experiments.

In a chemical change. the enthalpy of the products is different from the enthalpy of the reactants

Enthalpy change, ΔH, is the heat transferred at constant pressure by a chemical reaction. The units for ΔH is kJ mol-1

Exothermic reaction Endothermic reaction

Definition A reaction in which heat is released

to the surroundings

A reaction in which heat is absorbed from

the surroundings

Enthalpy of products and

reactants

Hp < Hr

Enthalpy is converted into heat

energy and released to the

surroundings

Hp>Hr

Heat energy is converted into enthalpy

(energy being stored in the products as

chemical energy)

Enthalpy change Negative (loss in enthalpy) ΔH = - Positive (gain in enthalpy) ΔH = +

2

ΔH = Hp – Hr

Temperature of surrounding Rises (energy releases from reacting

chemicals to the solution)

Falls (the reacting chemicals absorb energy

from the solution)

Enthalpy level diagram

In terms of bond forming

and breaking

Energy absorbed from surroundings

for breaking bonds in the reactants is

smaller than the energy released to

the surrounding during bond

formation of products

Energy absorbed from surroundings for

breaking bonds in the reactants is greater

than the energy released to the surrounding

during bond formation of products

Examples (1) combustion

(2) precipitation

(3) displacement

(4) neutralization

(1) cracking

(2) melting of solid

(3) boiling and evaporation of liquid

2.3 Standard Enthalpy changes

Is there any difference between the enthalpy changes of the following two reactions?

CH4(g) + 1½O2(g) CO(g) + 2H2O(l) CH4(g) + 1½O2(g) CO(g) + 2H2O(g)

..............................................................................................................................................................................................

..............................................................................................................................................................................................

For basis of comparison, enthalpy changes are stated under standard conditions

(1) pressure: 1 atm

(2) temperature: 25oC, 298K

(3) substance involved must be in their normal physical states under standard conditions

eg: NaCl(s), H2(g), H2O(l), C(graphite)

(4) when solution is involved: concentration of solution: 1M

Standard enthalpy change of formation ( Δ H o f) ( 可正數可負數 )

The standard enthalpy change of formation of a compound is the enthalpy change when 1 mole of the compound is

formed from its element in their standard states under standard conditions (at 25oC, 1 atm).

According to the definition, standard enthalpy change of formation of an element, e.g. oxygen, is the enthalpy change

3

associated with the following change, O2(g) → O2(g). Actually, there is no change, the standard enthalpy change of

formation of an element is always zero .


Example

Explain the meaning ΔHof [Na2CO3(s)] = -1130.7 kJ mol-1.

..............................................................................................................................................................................................

..........................................................................................................................................................................................

Complete the table.

1. ΔHof [CH4(g)] = -74.8 kJ mol-1

2. ΔHof [H2O(l)] = -286 kJ mol-1

3. N2(g) N2(g)

4. ΔHof [NO(g)] = +90 kJ mol-1

5. ΔHo = 5 ×ΔHof [NO(g)] = + 450 kJ mol-1

6. ΔHof [KMnO4(s)] = -813 kJ mol-1

7. C(graphite) + 2H2(g) + ½O2(g) CH3OH(l) = -239 kJ mol-1

8. 2C(graphite) + 4H2(g) + O2(g) 2CH3OH(l) ΔHo =

9. ΔHof [S8(s)] = 0 kJ mol-1

10. ΔHof [S8(g)] = + 103kJ mol-1

11. ΔHof [S(g)] = + 279kJ mol-1

12. If ΔHof [H2O(l)] = -286 kJ mol-1, then ΔHo of 2H2O(l) 2H2(g) + O2(g) = kJ mol-1

Standard enthalpy of combustion ( Δ H o c) ( 必定負數 )

The standard enthalpy change of combustion is the enthalpy change when 1 mole of a substance is completely burnt in

oxygen under standard conditions (at 25oC, 1 atm).

e.g. S(s) + O2(g) → SO2(g) ΔHoc [S(s)] = -297 kJmol-1

Example

1. Explain the meaning ΔHoc [C2H4(g)] = -1411 kJ mol-1.

..............................................................................................................................................................................................

..........................................................................................................................................................................................

2. Which of the following represent the standard enthalpy change of combustion of graphite?

A. C(diamond) + O2(g) → CO2(g) [C(graphite)] = -393.5 kJmol-1

B. C(diamond) + O2(g) → CO2(g) [C(diamond)] = -393.5 kJmol-1

C. C(graphite) + O2(g) → ½CO (g) [C(graphite)] = -393.5 kJmol-1

D. C(graphite) + O2(g) → CO(g) [C(graphite)] = -393.5 kJmol-1

4

3. Given that the standard enthalpy change of combustion of ethanol, ethane and hydrogen are -1367.3 kJmol-1, -

1560kJmol-1 and – 285.5kJmol-1 respectively. Write thermochemical equations to represent these data.

..............................................................................................................................................................................................

..............................................................................................................................................................................................

........................................................................................................................................................................................


4. The molar enthalpy change of combustion of some alkanes is given below in kJ mol-1.

C3H8 C4H10 C5H12 C6H14

−¿2219 −¿2877 −¿3509 −¿4163

Another alkane was found to have an enthalpy change of combustion of -6125kJ mol-. The alkane is

A. C7H16 B. C8H18 C. C9H20 D. C10H22

Standard enthalpy of atomization ΔHoat ( 必定正數 : energy must be absorbed to pull the atoms apart)

Standard enthalpy of atomization is the enthalpy change when 1 mole of gaseous atoms is formed from an element in

its standard state.

i. C(graphite) → C(g) ΔHoat [C(graphite)]= +715 kJmol-1

ii. ½H2(g) → H(g) ΔHoat [½H2(g)]= +218 kJmol-1

1. ΔHoat [½O2(g)] = +249.2 kJ mol-1

2. ΔHoat [Na(s)] = +108 kJ mol-1

3. ½Br2(l) → Br(g)

4. ΔHo = 4× ΔHo

at [½H2(g)]= +856 kJmol

5. The enthalpy change of atomization of iodine is the value of ΔH for the process

A. I2(s) I2(g) B. I2(s) 2I(g) C. I2(g) 2I(g) D. ½I2(s) I(g)

Standard enthalpy of neutralization Δ H o neut ( 必定負數 )

It is the heat change when an acid and a base react to form 1 mole of water under standard condition.

e.g. HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) ΔHoneut = -57.1 kJmol-1

H+(aq) + OH-(aq) H2O(l) ΔHoneut = -57.1 kJmol-1

CH3COOH(aq) + NaOH(aq) → CH3COONa(aq) + H2O(l) ΔHoneut= -55.2 kJmol-1

Neutralization of a ___________ acid is found to be less exothermic than that of a ______________ acid because

energy is required to ___________ the molecules of ___________ acid first.

Standard enthalpy of reaction Δ H o rxn

Standard enthalpy change of reaction (ΔHorxn) is the energy transferred when the number of moles of reactants shown

in the equation react under standard condition..

e.g. 4NH3(g) + 3O2(g) 2N2(g) + 6H2O(l) ΔHorxn = -1260 kJmol-1

NH3(g) + 3/4O2(g) 1/2N2(g) +3/2H2O(l) ΔHorxn = kJmol-1

2N2(g) + 6H2O(l) 4NH3(g) + 3O2(g) ΔHorxn = kJmol-1

Class Practice

5

Write balanced equations for reactions that have the following enthalpy changes:

enthalpy of combustion of methane

enthalpy of formation of water

enthalpy of combustion of hydrogen

twice the enthalpy of formation of hydrogen sulfide

enthalpy of atomisation of chlorine


2.4 Hess Law

Not ΔHof of all compounds can be determined directly. This is because

(a) not all compounds can be synthesized directly from its constituent elements at standard state.

sucrose (C12H22O11) cannot be formed by direct combination of 12C(graphite), 11H2(g) and 5½O2(g)

(b) formation of side products

4Na(s) + O2(g) → 2Na2O(s), 2Na(s) + O2(g) → Na2O2(s)

It is difficult to determine the enthalpy change of C(s) + ½O2(g) CO(g) since some carbon dioxide is always

formed in the reaction.

For the ΔHof of the compound which cannot be determined directly, Hess’s Law can be used to determine its value.

Hess’s Law

The total enthalpy change for a reaction is independent of the route taken for the reaction.

Example

Copper(II) sulfate exists as blue hydrated crystals and white anhydrous crystals. The enthalpy changes of solution for

these two substances may be represented by the following simplified equations:

CuSO4.5H2O(s) + aq CuSO4(aq) ΔH1 = +11.5 kJ mol–1

blue

CuSO4(s) + aq CuSO4(aq) ΔH2 = –66.1 kJ mol–1

white

(a) (i) Fill in the box and add labelled arrows to complete the Hess cycle to enable you to calculate ΔHreaction. (3)

(ii) Calculate a value for the enthalpy change ΔHreaction. (2)

..............................................................................................................................................................................................

..............................................................................................................................................................................................

6

(b) Suggest why it is not possible to directly measure the enthalpy change for the conversion of the blue hydrated

copper(II) sulfate crystals into the white anhydrous crystals. (1)

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................


Applying Hess's Law using enthalpy of formation data

ΔH[reaction] =

∑ H f o[reactants]−∑ H f o [ products ]

Steps:

1. Write the balance equation interested at the top

2. Write no of mole of each constituent element of

the reactants and products at the bottom.

3. Complete the cycle by adding arrows, pointing

from the constituent elements to the reactants and

products.

4. Label the arrows with

ΔHof[chemical] × appropriate no of mole

Example

1. Given that the standard enthalpy change of formation of SO2(g) and SO3(g) is -297kJ mol-1 -441kJ mol-1

respectively. Find the standard enthalpy change of contact process.

2. Find the standard enthalpy of reaction for Fe2O3(s) + 3CO (g) 2Fe(s) + 3CO2 at 25oC by drawing a Hess cycle.

Given: ΔHof [Fe2O3(s)] = -822.2kJ mol-1, ΔHo

f [CO(g)]= -110.5kJ mol-1 and ΔHof [CO2(g)] = -393.5kJ mol-1

7

3. Hydrogen peroxide is a powerful oxidizing which decomposes slowly to form water and oxygen. Find the

standard enthalpy of this reaction. Given that ΔHof [H2O2(l)] = -187.8 kJ mol-1, ΔHo

f [H2O(l)]= -285.8 kJ mol-1.


Applying Hess's Law using enthalpy of combustion data

ΔH[reaction] =

∑ H c o [reactants ]−∑ H c o[ products ]

Steps:

1. Write the balance equation interested at the top

2. Write the combustion products of the reactants and

products at the bottom.

3. Complete the cycle by adding arrows, pointing

from the reactants and products to their combustion

products.

4. Label the arrows with

no of mole of oxygen needed for that combustion

ΔHoc[chemical] × appropriate no of mole

1. Use the standard enthalpy changes of combustion, ΔHc○, given in the table below to find the standard enthalpy

change of formation for ethanoic acid, CH3COOH, in kJ mol–1.

Substance C(s, graphite) H2(g) CH3COOH(l)

ΔHc○/ kJ mol–1 –394 –286 –870

2. Calculate the enthalpy change of formation of ethanol in kJ mol–1.

Substance C(s, graphite) H2(g) C2H5OH(l)

ΔHc○/ kJ mol–1 –394 –286 –1370

8

3. Hex-1-ene can be converted to hexane in the following reaction.C6H12(l) + H2(g) → C6H14(l)

(i) What catalyst is used in this reaction? .................................................................

(ii) Find the enthalpy change of this reaction ΔHreaction from the following data of enthalpy changes of combustion.

Substance Hex-1-ene, C6H12 Hydrogen, H2 Hexane, C6H14

ΔHc○/ kJ mol–1 –4003 –286 –4163


Class Practice/HW

1. Calculate the heat of formation of ethene given that its heat of combustion is -1305 kJ mol-1, the heat of formation

of CO2 is -393.4 kJ mol-1 and the heat of formation of water is -285.5 kJ mol-1. [-52.8kJ mol-1]

2. Given that 2CO(g) + 2NO(g) 2CO2(g) + N2(g) ΔH= -746.8kJ mol-1

Calculate the standard enthalpy change of formation of nitrogen monoxide if

ΔHof [CO(g)]= -111kJ mol-1 and ΔHo

f [CO2(g)] = -394kJ mol-1 [90.4kJ mol-1]

3. Calculate the standard enthalpy change of formation of ethane. Draw a Hess’s Law cycle

as part of your

answer. [-84kJ mol-1]

4. Given the following enthalpies:

ΔHof [MgCO3(s)]= -1113kJ mol-1

ΔHof [CO2(s)] = -394kJ mol-1

ΔHof [MgO(s)]= -602kJ mol-1

Calculate the enthalpy for the thermal decomposition of magnesium carbonate. [+117kJ mol-1]

5. The standard enthalpy change of formation of CO2(g) is –394 kJ mol–1 and that of H2O(l) is –286 kJ mol–1.

Draw a Hess cycle to calculate ΔHof [C5H12(g)] if ΔHc

○ [pentane(g)] = –3509 kJ mol–1. [-177kJ mol-1]

6. Methylhydrazine (CH3NHNH2) and dinitrogen tetraoxide (N2O4) are fuels of Apolo 11 which landed the first man

on moon on 1969. They are expected to react very exothermically according to the following reaction:

CH3NHNH2(l) + N2O4(l) CO2(g) + H2O(l) + N2(g)

(a) Balance the equation.

(b) The standard enthalpy changes of formation are: CH3NHNH2(l) = +53kJ mol-1 N2O4(l) = -20kJ mol-1

The standard enthalpy changes of combustion are: C(s, graphite) = -393 kJ mol-1 H2(g) = -286 kJ mol-1

Calculate the enthalpy change for the reaction between CH3NHNH2(l) and N2O4(l). [-5116kJ mol-1]

9

Substance C(s, graphite) H2(g) Ethane(g)

ΔHc○/ kJ mol–1 –393 –286 –1560

2.5 Experimental Determination of enthalpy changes

Simple calorimetric method is used to determine enthalpy change of reaction and enthalpy change of neutralization.

Basic steps of simple calorimetric methods :

(1) Mixing known mass of reactants in an insulated container called calorimeter**

(2) measure the initial temperature and final temperature

(3) calculate the heat energy transferred in joules = mass × specific heat capacity × temperature change

**Example of calorimeter is a “coffee-cup calorimeter”, which is two expanded polystyrene cup, one inside

the other, with a lid. It is an excellent insulator of heat and also has low specific heat capacity.


Example

In an experiment performed to measure the enthalpy change for the reaction Cu2+ (aq) + Zn (s) Cu (s) + Zn2+ (aq)

3g of zinc powder (an excess) was added to 30.0 cm3 of copper(II) sulfate solution of concentration 1.00 mol dm–3.

The temperature rise of the mixture was 47.6 K.

Assuming that the heat capacity of the solution is 4.2 J K–1 g–1, the enthalpy change for the reaction is given by

A ΔH = - (30 × 4.2 × 47.6)÷ 0.03 B ΔH = - (33 × 4.2 × 47.6)÷ 0.03

C ΔH = - (30 × 4.2 × 47.6)× 0.03 D ΔH = - (33 × 4.2 × 47.6)× 0.03

Jun 2011 Unit 1 Question 18c (Design experiment to measure the enthalpy change of the reaction)

CuSO4.5H2O(s) + aq CuSO4(aq) ΔH1 = +11.5 kJ mol–1

(i) Describe briefly the experimental procedure that you would use to obtain the data necessary to calculate ΔH1,

given a known mass of hydrated copper(II) sulfate crystals, CuSO4.5H2O(s).

You should state the apparatus that you would use and any measurements that you would make.

You are not required to calculate the amounts of substances or to explain how you would use the data obtained. (4)

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

(ii) List three possible errors which do not relate to the quality of the apparatus or chemicals used or possible

mistakes in carrying out the procedure. (3)

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

10

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

Graphical method for obtaining Δ T

To be more accurate, instead of just measuring the initial and final temperature, we can measure the temperature

of the reactants along time.

A graph is then plotted and extrapolated (外推) to determine a more accurate value of ΔT.


Two possible forms of graph you encountered: heat of reaction / heat of neutralization + experiment steps

(1) Enthalpy change of reaction (eg.metal displacement;

CO32- + H+)

Take Zn and CuSO4(aq) as an example.

Known volume and molarity of CuSO4 is added to a

polystyrene cup.

The temperature of CuSO4(aq) was measured every

minute for 3 minutes.

After exactly 3.5 minutes, solid zinc was added and

the mixture was stirred.

The temperature of the mixture was then taken every

minute for another 8 minutes. Extrapolates the two

lines formed.

Measure the difference between these two lines

at the time when zinc is added .

(2) For neutralization titration,

Measure the temperature of alkali in the conical flask.

Add acid from burette with 2cm3 each time, stirring

and measuring the temperature after each addition.

Continue to add alkali for a few times after reaching

the maximum temperature. (end point)

Extrapolates the line of rising temperature and falling

temperature until they intersect. This is the maximum

temperature reached in the neutralization reaction.

11

The point at which the two extrapolated lines meet

corresponds to the volume of titre required for

neutralization.

Read off this volume from the graph.

Possible source of errors of calorimetric methods

(1) There was heat loss to the surroundings

(2) The specific heat capacities of the expanded polystyrene cup and the thermometer were not taken into account

(3) The specific heat capacity of the reaction mixture was not the same as that of water.

(4) The density of the reaction mixture was not the same as that of water.

(5) The experiment was not done under standard conditions.

Error (3)and Error (4) are not suitable for enthalpy change of combustion since water is used as the medium for

absorbing heat directly. So its density is really 1g/cm 3 and c = 4.18J/g/K.


Jan 2011 Unit 1 Question 19 ( Determination of enthalpy change of combustion)

The enthalpy change of combustion of ethanol was determined using the apparatus shown below. In the experiment,

the temperature increase of the water in the beaker is measured when a known mass of the ethanol is burned.

(a) The results of the experiment are summarised in the table below.

Mass of water in the beaker 250.00 g

Mass of spirit burner + contents (initial) 63.21 g

12

Mass of spirit burner + contents (final) 62.47 g

Temperature of water (initial) 21.0 °C

Temperature of water (final) 31.5 °C

(i) Calculate the heat energy produced by the combustion of the alcohol using the equation

heat energy produced (J) = mass of water × 4.18 × temperature change (1)

(ii) Calculate the number of moles of ethanol burned in this experiment (the formula of ethanol is C2H5OH). (3)

(iii) Use the equation below to calculate the enthalpy change of combustion of ethanol in kJ mol–1. Give the value an

appropriate sign. (2)

(b) The data book value for the enthalpy change of combustion of ethanol is –1370 kJ mol–1.

(i) Calculate the percentage error in the value calculated in (a)(iii) in comparison with the data book value. (1)

(ii) List three ways in which the design of the experiment causes the results to be so different from the data book

value. (You should be specific but detailed explanations are not required.) (3)

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................


Jan 2012 Unit 3 Question 3 ( Determination of enthalpy change of reaction)

Magnesium carbonate reacts with dilute nitric acid : MgCO3(s) + 2HNO3(aq) Mg(NO3)2(aq) + CO2(g) + H2O(l)

The enthalpy change for this reaction can be determined as follows:

Procedure

13

1. Weigh 3.50 g of finely powdered magnesium carbonate.

2. Transfer 50.0 cm3 of 2.00 mol dm–3 nitric acid into a polystyrene cup and record the temperature of the acid.

3. Add the magnesium carbonate to the nitric acid.

4. Stir the mixture and record the maximum temperature reached.

Results

Temperature of nitric acid before addition of magnesium carbonate 21.0 oC

Final temperature of solution 29.7 oC

(a) Explain why the magnesium carbonate used in this experiment should be finely powdered rather than in lumps. (1)

..............................................................................................................................................................................................

(b) (i) Calculate the number of moles of magnesium carbonate in 3.50 g. (1)

(ii) The volume of dilute nitric acid used contained 0.100 mol of HNO3. Suggest why this amount is suitable.(1)

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

(c) (i) Calculate the heat energy transferred, in joules, in this reaction between magnesium carbonate and nitric acid.

(ii) Calculate the enthalpy change, ΔH, for the reaction of one mole of magnesium carbonate with nitric acid. Your

answer should be in units of kJ mol–1, expressed to three significant figures, and include a sign. (2)

(d) (i) The nitric acid for this experiment could be measured using either a pipette or a measuring cylinder. Give one

practical advantage of using each piece of apparatus. (2)

Pipette ..................................................................................................................................................................................

Measuring cylinder ..............................................................................................................................................................

(ii) The total error in measuring the mass of the magnesium carbonate was ±0.01 g.

Calculate the percentage error in the weighing. (1)


(e) State and explain the effect, if any, on the calculated enthalpy change, ΔH, if

(i) a copper calorimeter were used instead of the polystyrene cup. (2)

..............................................................................................................................................................................................

..............................................................................................................................................................................................

14

..............................................................................................................................................................................................

(ii) 3.50 g of damp magnesium carbonate were used. (2)

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

*(f) The experiment was repeated with a change in the procedure. The temperature of the dilute nitric acid was

measured every minute for three minutes. After exactly three and a half minutes, the magnesium carbonate was added

and the mixture was stirred. The temperature of the mixture was then taken every minute for another six minutes.

A graph of the temperature readings against time using this procedure is shown below.

*(i) Use this graph to calculate the maximum temperature change. Show your working on the graph. (2)

Maximum temperature change = ............................................................................o C

*(ii) Why does this method give a more accurate result than the original procedure? (1)

..............................................................................................................................................................................................

..............................................................................................................................................................................................


June 2011 Unit 3 Question 3 ( Graphical determination of enthalpy change of neutralization)

A titration was carried out in order to investigate the neutralization reaction between NaOH and HCl.

Procedure : 1. Using a pipette, transfer 50.0 cm3 of sodium hydroxide solution, concentration 1.00 mol dm–3, to a

polystyrene cup. Allow to stand for a few minutes.

15

2. Record the temperature of the solution.

3. From a burette, add 5.00 cm3 of dilute hydrochloric acid to the solution in the cup.

4. Stir the mixture with the thermometer and record the temperature.

5. Add successive 5.00 cm3 portions of hydrochloric acid, stirring the mixture and recording the

temperature after each addition.

6. Continue adding hydrochloric acid until a total of 50.00 cm3 of the acid has been added.

Results

Volume of HCl added / cm3 0.00 5.00 10.00 15.00 20.0

0

25.00 30.0

0

35.00 40.00 45.00 50.00

Temperature / °C 22.2 23.7 25.1 26.6 28.0 29.5 29.2 28.4 27.6 26.8 26.0

A graph of the temperature (y-axis) against the volume of hydrochloric acid added (x-axis) enables the maximum

temperature rise and the volume of acid required for neutralization to be determined.

From this information it is possible to calculate (a) the concentration of the hydrochloric acid (b) the enthalpy change

for the reaction.

(a) (i) Plot a graph of temperature against volume of acid added on the axes below. Draw two straight lines on your

graph and extrapolate the lines until they intersect.


(ii) Use the extrapolated lines on your graph to read off the maximum temperature reached in the neutralization. (2)

16

Maximum temperature ................................................oC

(iii) The point at which the two extrapolated lines meet corresponds to the volume of hydrochloric acid required for

neutralization. Read off this volume from your graph. (1)

Volume of hydrochloric acid ...................................................cm3

(iv) Calculate the number of moles of sodium hydroxide in 50.0 cm3 of a 1.00 mol dm–3 solution. (1)

(v) HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Use this equation and your answers to (iii) and (iv) to calculate the molarity of the hydrochloric acid (2)

(b) (i) Use your graph and answer to (a)(ii) to calculate the maximum temperature change, ΔT, for the reaction. (1)

ΔT = ................................................ °C

(ii) Use your value for the temperature rise, ΔT, to calculate the heat energy produced when 50.0 cm3 of sodium

hydroxide is exactly neutralized by the volume of hydrochloric acid you obtained in (a)(iii).

Use the expression energy = total mass of solution × specific heat capacity of solution × temperature rise

[Assume the specific heat capacity of the solution to be 4.2 J g–1 °C–1 and the density of the solution to be 1.0 g cm–3]

(2)

(iii) Use your answers to (a)(iv) and (b)(ii) to calculate the enthalpy change, in kJ mol–1, for this reaction.

Give your answer to two significant figures and include a sign. (3)ΔH = ................................................. kJ mol–1


Jan 2010 Unit 3 Question 3d ( Determination of enthalpy change of neutralization)

17

The standard enthalpy change for the reaction of calcium hydroxide with hydrochloric acid was found by reacting

0.0100 mol of solid calcium hydroxide with 50.0 cm3 of a 1.00 mol dm–3 solution of hydrochloric acid (an excess), in a

polystyrene cup. The temperature rose from 21.2 °C to 26.7 °C.

(i) Calculate the energy, in joules, transferred in the reaction. Use the expression

Energy transferred = mass × specific heat capacity × temperature change

[Assume density of solution = 1.0 g cm–3, specific heat capacity of solution = 4.18 J g–1°C –1] (1)

(ii) Calculate the standard enthalpy change, ΔH, for the reaction. Include a sign and units in your answer. (2)

(iii) Calculate the percentage error in the temperature change caused by an uncertainty of 0.1°C in each

thermometer reading. (1)

(iv) The experiment was repeated using 50.0 cm3 of a 1.00 mol dm–3 solution of nitric acid instead of the

hydrochloric acid. Explain why the temperature change was the same in both experiments. (1)

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

(v) The experiment was repeated again using 25 cm3 of 2.00 mol dm−3 hydrochloric acid. Predict the temperature

change in this experiment. (1)

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................

(vi) Which of the experiments in (iv) and (v) gave the least error in the temperature change? Justify your answer.

(1)

..............................................................................................................................................................................................

..............................................................................................................................................................................................

..............................................................................................................................................................................................


2.6 Bond enthalpy

18

Bond enthalpy of a particular bond in a specific compound is the energy required to break one mole of this covalent

bond in gaseous state.

H2(g) 2H(g) ΔH = +436kJ mol-1 E(H – H) = +436kJ mol-1

HCl(g) H(g) + Cl(g) ΔH = + 432 kJ mol-1 E(H – Cl) = +432kJ mol-1

Example

Explain and write equation: bond enthalpy of chlorine is +242kJ mol-1

..............................................................................................................................................................................................

..............................................................................................................................................................................................

Example

Find the C – H bond enthalpy in methane if CH4(g) C(g) + 4H(g) ΔH = 1663.5kJ mol-1

This process involves the breaking of four C – H bond. So the C – H bond enthalpy in methane is 1663.4 / 4 = +415kJ

mol -1 . Symbol: E( C – H) = +415kJ mol -1

Mean bond enthalpy

The same type of bond in different compound can have slightly different bond enthalpies depending on the

environment of the bond within a particular molecule.

For example:

(1) Consider CH4 and CH3Cl. The Cl atom in CH3Cl affects the environment of the C – H bonds. Thus, C – H bonds

have different strengths and different bond enthalpies value in different compound.

(2) The C – H bond enthalpy in methane will be slightly different from the value of the C – H bond in ethanol.

The mean bond enthalpy of a particular kind of bond is the average value of the bond enthalpy for that bond in a wide

range of different compounds.

Example

1. If the mean C - H bond enthalpy is +x, which of the

following represents a process with an enthalpy

change of +4x?

A C(g) + 4H(g) CH4(g)

B CH4(g) C(g) + 4H(g)

C CH4(g) C(s, graphite) + 2H2(g)

D C(s, graphite) + 2H2(g) CH4(g)

2. Which equation represents the reaction for which the

enthalpy change, ΔH, is the mean bond enthalpy of

the C–H bond?

A ¼CH4(g) → ¼C(g) + H(g)

B CH4(g) → C(s) + 2H2(g)

C CH4(g) → C(g) + 4H(g)

D CH4(g) → C(g) + 2H2(g)

3. For which of the following reactions is the enthalpy

change equal to the bond enthalpy of H–I?

A HI(g) ½H2(g) + ½I2(s)

B HI(g) ½H2(g) + ½I2(g)

C HI(g) H(g) + I(g)

D HI(g) H+(g) + I–(g)

4. The enthalpy change for the reaction

CH4(g) C(g) + 4H(g)

is +1648 kJ mol–1. Hence the mean bond enthalpy for the

C–H bond is

A +329.6 kJ mol–1 C +1648 kJ mol–1

B +412.0 kJ mol–1 D +6592 kJ mol–1

5. This question is about some standard enthalpy changes, ΔH○

19

A enthalpy of reaction B enthalpy of combustion C mean bond enthalpy D bond enthalpy

(a) Which enthalpy change is represented by p?

CH4 (g) → CH3(g) + H(g) ΔH ○ = p

(b) Which enthalpy change is represented by q?

CH4(g) → C(g) + 4H(g) ΔH ○ = 4q

Bond enthalpies can be used in the calculation in Hess Cycle.

Given that E (C – H) = 413 kJ mol-1, E (C – O) = 358 kJ mol-1, E (H – O) = 463 kJ mol-1.

ΔHoat [C(graphite)]= +716.7 kJmol-1 ΔHo

at [½H2(g)]= +218 kJmol-1 ΔHoat [½O2(g)] = +249.2 kJ mol-1

Find the enthalpy change of formation of methanol by drawing a Hess Cycle. [-222 kJ mol1-]

Use of bond enthalpy

(1) Compare the strength of different bonds. It indicates the covalent bond strength. Bonds with small bond enthalpies

are weak and break first. In contrast, strong bonds have high bond enthalpies. Triple bond has a high enthalpy.

E(H – H)= +436kJ mol-1 , E(Cl – Cl) = +242kJ mol-1

. Thus Cl – Cl bond is a weaker bond and will break first in the

reaction.

(2) Indicate how fast a reaction might be. Low bond enthalpies indicate that a reaction will take place quickly.

(3) Estimate the enthalpy change in chemical reactions involving covalent molecules.

Example

Calculate the standard enthalpy change of combustion of ethanol according to the data given.

(1) The standard enthalpy change of formation of CO2(g) is –394 kJ mol–1 and that of H2O(l) is –286 kJ mol–1.

Draw a Hess cycle to calculate ΔHoc [C2H5OH(l)] if ΔHf

○ [C2H5OH(l)] = –278kJ mol–1.

20

(2) 0.388g of ethanol is burnt in a spirit burner. 100cm3 of water is used to absorb the heat released. The temperature

change of the water is + 24.6oC.

(3) E (C-C) = 347 kJ mol–1., E(C-H) = 413 kJ mol–1., E(C-O) = 336 kJ mol–1., E (C=O) = 805 kJ mol–1.,

E (H-O) = 464 kJ mol–1., E (O=O) = 498 kJ mol–1.

Summary

Enthalpy change of a reaction can be determined by

(1) Using standard enthalpy change data and calculate through Hess Cycle

(2) Doing experiment ( Energy transfer = mcΔT)

(3) Using mean bond enthalpies (∑ bond breaking−∑ bond forming¿¿

Value of (2.experiment value) different from (1.value in data booklet):

Mainly due to experimental error such as

a) heat loss to the surroundings

b) specific heat capacity of apparatus not considered

c) experiment not done under standard conditions

Value of (3. by mean bond enthalpies) different form (2. experiment value) :

a) Mean bond enthalpies values are used for calculation. It is only an average value from wide range of compounds.

There are variations in the strength of one particular bond in different molecules.

b) Bond enthalpies values only apply to substances in the gaseous state. Not all the reactants or products of the

reaction are in the gaseous state.

21

HW

1. Calculate the heat of formation of ethene given that its heat of combustion is -1305 kJ mol-1, the heat of formation

of CO2 is -393.4 kJ mol-1 and the heat of formation of water is -285.5 kJ mol-1. [-52.8kJ mol-1]

2. Given that 2CO(g) + 2NO(g) 2CO2(g) + N2(g) ΔH= -746.8kJ mol-1

Calculate the standard enthalpy change of formation of nitrogen monoxide if

ΔHof [CO(g)]= -111kJ mol-1 and ΔHo

f [CO2(g)] = -394kJ mol-1 [90.4kJ mol-1]

3. Calculate the standard enthalpy change of formation of ethane. Draw a Hess’s Law cycle

as part of your

answer. [-84kJ mol-1]

4. Given the following enthalpies:

ΔHof [MgCO3(g)]= -1113kJ mol-1

ΔHof [CO2(g)] = -394kJ mol-1

ΔHof [MgO(g)]= -602kJ mol-1

Calculate the enthalpy for the thermal decomposition of magnesium carbonate. [+117kJ mol-1]

5. The standard enthalpy change of formation of CO2(g) is –394 kJ mol–1 and that of H2O(l) is –286 kJ mol–1.

Draw a Hess cycle to calculate ΔHof [C5H12(g)] if ΔHc

○ [pentane(g)] = –3509 kJ mol–1. [-177kJ mol-1]

6. Methylhydrazine (CH3NHNH2) and dinitrogen tetraoxide (N2O4) are fuels of Apolo 11 which landed the first man

on moon on 1969. They are expected to react very exothermically according to the following reaction:

CH3NHNH2(l) + N2O4(l) CO2(g) + H2O(l) + N2(g)

(a) Balance the equation.

(b) The standard enthalpy changes of formation are: CH3NHNH2(l) = +53kJ mol-1 N2O4(l) = -20kJ mol-1

The standard enthalpy changes of combustion are: C(s, graphite) = -393 kJ mol-1 H2(g) = -286 kJ mol-1

Calculate the enthalpy change for the reaction between CH3NHNH2(l) and N2O4(l). [-5116kJ mol-1]

22

Substance C(s, graphite) H2(g) Ethane(g)

ΔHc○/ kJ mol–1 –393 –286 –1560

AS chem Ch 1.2 Hess Law

Documents

Transcript of AS chem Ch 1.2 Hess Law