TOPIC 5

TOPIC 5

ENERGETICS/THERMOCHEMISTRY

5.1

MEASURING

ENERGY CHANGES

By: Merinda Sautel Alameda Int’l Jr/Sr High School Lakewood, CO [email protected]

ESSENTIAL IDEA

The enthalpy changes from chemical reactions can be calculated from

their effect on the temperature of their surroundings.

NATURE OF SCIENCE (2.6)

Fundamental principal – conservation of energy is a fundamental principle of science.

NATURE OF SCIENCE (3.1)

Making careful observations – measurable energy

transfers between systems and surroundings.

INTERNATIONAL- MINDEDNESS

The SI unit of temperature is the Kelvin (K), but the Celsius scale (

C), which has

the same incremental scaling, is

commonly used in most countries. The exception if the USA which continues to use the Fahrenheit scale (

F) for all non-

scientific communication.

THEORY OF KNOWLEDGE

What criteria do we use in judging discrepancies between

experimental and theoretical

values? Which ways of knowing do we use when assessing

experimental limitations and

theoretical assumptions?

UNDERSTANDING/KEY IDEA 5.1.A

Heat is a form of energy.

ENERGY

• All chemical reactions are accompanied by energy changes.

• Energy is a measure of the ability to do work, that is to move an object against an opposing force.

• Examples: heat, light, sound, electricity and chemical energy which is the energy released or absorbed during chemical

reactions.

HEAT

• Heat is a form of energy which is

transferred as a result of a difference in temperature and produces an

increase in disorder of the behavior of particles.

• Heat increases the average kinetic energy of the molecules in a

disordered fashion.

Needed definitions

• Enthalpy

• Heat content of a substance (at constant pressure)

• Enthalpy is also the internal energy stored in the reactants.

• The absolute value for the enthalpy of reactants and products cannot be known, but the what

can be measured is the difference between the two.

• System – area of interest (Example: beaker and its contents)

• Surroundings – everything else in the

universe

• An “open” system can exchange matter and energy with the

surroundings.

• A “closed” system can only

exchange energy, not matter, with

the surroundings.

• The joule (J) is the SI unit for

energy and work.

UNDERSTANDING/KEY IDEA 5.1.B

Temperature is a measure of

the average kinetic energy of

the particles.

Temperature

• Def – measure of the average kinetic energy of the particles

• Temperature increase depends upon:

• Mass of the object

• Amount of heat added

• Nature of the substance

• Different substances need different

amounts of heat to increase the temp of

a unit mass by 1K or 1ºC.

UNDERSTANDING/KEY IDEA 5.1.C

Total energy is conserved in

chemical reactions.

UNDERSTANDING/KEY IDEA 5.1.D

Chemical reactions that involve transfer of heat

between the system and the

surroundings are described as

endothermic or exothermic.

EXOTHERMIC REACTIONS

A reaction which results in a transfer of energy from the system to the

surroundings.

• Heat is given off or produced.

• Products have less energy or heat content than the reactants.

• ΔH is negative.

• The bonds in the products are stronger than

the bonds in the reactants.

Exothermic reaction Enthalpy diagram

reactants

enthalpy

H

extent of reaction

ENDOTHERMIC REACTIONS

• A reaction which results in a transfer of energy from the surroundings to the system.

• Heat is absorbed.

• Reactants have less energy than the products.

• ΔH is positive.

• The bonds in the reactants are stronger

than those in the products.

Endothermic reaction Enthalpy diagram

products

enthalpy

H

extent of reaction

UNDERSTANDING/KEY IDEA 5.1.E

The enthalpy change (ΔH) for chemical reactions is

indicated in kJ mol ^-1

UNDERSTANDING/KEY IDEA 5.1.F

ΔH values are usually

expressed under standard conditions, given by ΔH ^◦ ,

including the standard states.

GUIDANCE

Standard state refers to the

normal, most pure stable state

of a substance measured at 100

kPa. Temperature is not part of

the definition of standard state,

but 298K is commonly given as

the temperature of interest.

GUIDANCE

Be familiar with the enthalpy

change of combustion (ΔH _{c ◦} )

and the enthalpy change of

formation (ΔH _{f ◦} ).

APPLICATION/SKILLS

Be able to calculate the heat change when the

temperature of a pure

substance is changed using q

= mcΔT.

GUIDANCE

The specific heat capacity of

water is provided in the data

booklet in section 2.

APPLICATION/SKILLS

Be able to evaluate

calorimetry experiments for

the enthalpy of a reaction.

GUIDANCE

Be familiar with reactions occurring in aqueous

solutions and combustion

reactions.

GUIDANCE

Students can assume the specific heat and density of aqueous

solutions are equal to those of water, but be aware of this

limitation.

GUIDANCE

Heat losses to the environment and the heat capacity of the

calorimeter in experiments

should be considered, but the

use of a bomb calorimeter is not

required.

THERMOCHEMICAL

EQUATION EXAMPLES

Combustion of Methane

CH4(g) + 2O2(g) CO2(g) + 2H2O(l) ΔH = -890 kJ mol^-1

Photosynthesis

6CO_2(g) + 6H₂O_(l) C₆H₁₂O_6(aq) + 6O_2(g) ΔH = +2802.5 kJ mol^-1

Thermite Reaction

2Al(s) + Fe2O3(s) Al2O3(s) + 2Fe(s) ΔH = -841 kJ mol^-1

• You must give the “state” symbols such as (s), (g), (l), (aq) in

thermochemical equations

because energy changes depend

upon the state of the reactants and

the products.

• Standard enthalpy change of reaction ΔHº

• This is the heat or enthalpy change of a reaction when carried out at standard conditions.

• Temperature = 298K or 25ºC

• Pressure = 101.3 kPa (1atm)

• Solution concentration = 1mol dm

(1M)

• All substances are in their standard states (how

they are found in nature)

• The actual amount of heat absorbed or produced in a chemical reaction

depends upon several factors:

• Nature of the reactants and products

• The amount or concentration of the

reactants. (The greater the amount that reacts, the greater the heat change.)

• The states of the reactants and products.

Changing states involves an enthalpy change so this will affect the total heat change.

• The temperature of the reaction.

• All combustion reactions are exothermic processes.

• All neutralization reactions are

exothermic processes.

• Apply the relationship between temperature change, enthalpy

change and be able to classify the reaction as endothermic or

exothermic.

• This will be determined by

calorimetry.

Specific Heat Capacity

• The following relationship allows the heat change in a material to be

calculated from the temperature change.

q = mcΔT

heat = mass x specific heat x ΔT

• Specific heat capacity (c) is the heat needed to increase the

temperature of a unit mass (usually 1g) by 1K or 1ºC.

• Specific heat of water

c = 4.18 J K

g

• Be able to deduce from an enthalpy level diagram, the

relative stabilities of reactants and products and the sign of the

enthalpy change for the reaction.

• There is a natural direction for change.

• The direction of change is in the direction of lower stored energy.

• We generally expect a reaction to occur if ΔH is negative (exothermic), but some endothermic reactions do occur. These occur when the entropy (S) of the

system is large.

Enthalpy Diagram for Hydrogen Peroxide (H

O

)

ΔH1

ΔH₃

ΔH₂ Enthalpy

In this diagram, hydrogen peroxide in the middle is stable compared to H₂ and O₂, but unstable compared to the decomposition of water and oxygen on the bottom line.

• The heat produced when one mole of a substance is burned in excess oxygen is called the enthalpy of

combustion.

• ΔH

is exothermic and always

negative.

• Calculate the heat energy change

when the temperature of a pure

substance is changed.

How to calculate heat changes from temp changes.

• When heat released by an exothermic reaction is absorbed by water, the

temperature of the water increases.

• The heat produced by the reaction can be calculated if it is assumed that all the

heat is absorbed by the water.

heat change of rxn = -heat change of water = - m

x c

x ΔT

As the water has gained the heat produced by the reaction, the heat change of the reaction is negative when the temperature of the water increases.

Example Problem 1

• Calculate the enthalpy of combustion of ethanol from the following data.

Assume all heat from the reaction is

absorbed by the water. Compare your value with the IB Data booklet value

and suggest reasons for any

differences.

• Mass of water in Cu Calorimeter = 200.00g

• Temperature increase in water = 13.00ºC

• Mass of ethanol burned = 0.45g

The heat calculated above is the heat gained by the water which is also the heat lost by the combustion of the ethanol.

Since ΔHc is per mole of substance combusted, you have to find moles of ethanol burned and divide the heat it lost by the moles.

To find moles, divide mass by molar mass:

0.45g / 46.08 g mol-1 = .0098mol

ΔHc = -10868J / .0098mol = -1112883 J/mol = -1112.883 kJ/mol Using sig figs, the answer is limited by the mass = -1100 kJ/mol

• The IB Data booklet value is -1367 kJ/mol.

• What are some reasons for the difference in values?

• Not all the heat produced by the combustion is transferred to the water.

• Some is needed to heat the Cu calorimeter and some has passed to the surroundings.

• The combustion of ethanol is unlikely to be complete due to the limited oxygen available.

Heat of Combustion

Set up

Example Problem 2

• When 0.824 g of ethanol was burned, it produced a temperature rise of

15.7 K in 275g of water.

• a. Draw the calorimetry set up.

• b. Write the ΔHc reaction.

• c. Calculate the ΔH for the reaction.

• d. The Data Book value is -1367 kJ/mol.

Include three reasons for the discrepancy.

• It is important to record qualitative as well as quantitative data.

• When asked to evaluate experiments and suggest improvements, avoid giving trivial answers such as incorrect measurements.

• Incomplete combustion can be reduced by burning the fuel in oxygen and heat loss can be reduced by insulating the

apparatus.

• Design suitable experimental

procedures for measuring the heat

energy change of reactions.

ΔH of reactions in solution

• The enthalpy changes of reaction in solution can be calculated by carrying out the

reaction in an insulated system such as a polystyrene cup or Styrofoam calorimeter.

• The heat released or absorbed by the

reaction can be measured by the change in temp of the water.

• The largest source of error in this type of

experiment is heat loss to the environment.

• To compensate for the heat lost by the water in exothermic reactions to the

surroundings as the reaction proceeds, a plot of temperature vs time can be

drawn.

• We want the highest temperature that would have been reached had the

reaction been instantaneous.

• If we graph the temp vs time, we can extrapolate back to determine the

highest temperature.

T2 extrapolation at same rate of cooling

T₁

∆T ∆T for rxn = T2-T0

Temp

(K or ⁰C)

T₀

point where reactants are mixed

Compensating for Heat Loss

T₀ = initial temp of reactants

T₁ = highest temp actually reached

T₂ = temp that would have been reached if no heat were lost to the surroundings

• Calculate the enthalpy change for a

reaction using experimental data on

temperature changes, quantities of

reactants and mass of water.

Example problem 3

• 50.0 cm

of 1.00 mol/dm

hydrochloric acid was added to 50.0 cm

of 1.00 mol/dm

sodium hydroxide solution in a polystyrene beaker. The initial temperature of both

solutions was 16.7ºC. After stirring and accounting for the heat loss, the highest

temperature reached was 23.5ºC. Calculate

the enthalpy change for the reaction.

• Step 1 – write the equation

HCl

+ NaOH

NaCl

+ H

O

• Step 2 – calculate molar quantities

Anytime you are given a volume and a molarity, multiply them together (watch your units) to get moles. Convert cm

to dm

first.

50.0cm

x 1dm

/1000cm

= 0.0500 dm

Then multiply the volume by the molarity.

0.0500dm

x 1.00 mol/dm

= 0.0500 mol

• Step 3 – calculate the heat evolved

• Assumptions:

• The solution has the same density as water which is 1.00 g/cm³.

• The solution has the same specific heat as water which is 4.18 J/gK.

• Total volume of solution

50.0 cm³ + 50.0 cm³ = 100.0 cm³

• Using the density of water, calculate the mass of the solution

100.0 cm³ x 1.00 g/cm³ = 100.0 g

• Temperature change

23.5ºC – 16.7ºC = 6.8ºC = 6.8 K

• The heat evolved from the reaction = mcΔT.

q = (100.g)(4.18 J/gK)(6.8 K) = 2840 J = -28.4kJ

• The ΔH is per mole so

ΔH = -28.4kJ/.0500mol = -56.8 kJ/mol

• Evaluate the results of

experiments to determine enthalpy

changes.

Example Problem to evaluate the results of an experiment.

• The experiment involves the neutralization of

sodium hydroxide with sulfuric acid to form water and sodium sulfate.

• Assumptions:

• The NaOH and H

SO

are 1M (1mol dm

).

• The total volume is kept constant at 120 cm

with varying volumes of NaOH and H

SO

.

• No heat is lost from the system.

• All heat is transferred to the water.

• Assume 120cm

of solution contains 120cm

of

water.

• Your task is to evaluate the following when the temperature changes are measured when

different volumes of NaOH and H

SO

are mixed:

• Determine the volumes of the solutions which produce the largest increase in temperature.

• Calculate the heat produced by the reaction when the maximum temperature was produced.

• Calculate the heat produced for one mole of NaOH.

• Calculate the %error and suggest a reason for the discrepancy between the experimental and literature values.

• The literature value at standard conditions is -57.5 kJ/mol^-1.

35 34 33 32 31 30 29 28 27 26 25

Temp ºC

□

• The highest increase in temp is produced at 80 cm

of NaOH which means that there is 120-80 = 40 cm

of H

SO

.

• The maximum temp where the lines cross on the graph was 33.5ºC.

q = -mcΔT

= -(120.0g)(4.18J/gºC)(33.5ºC-25.0ºC) = -4264J

• To calculate the heat produced per mole:

• Percent Error

%error = (-57.5-

53.3)/-57.5 x 100% = 7%

• There are uncertainties in temp, vol and

concentration measurements and it is assumed all heat is transferred to the water and no heat is lost from the system.

Another fine example

• 50.0 cm

of 0.200 mol dm

copper II sulfate was placed in a

polystyrene cup. After 2 min, 1.20 g of powdered zinc was added.

The temperature was taken every 30 seconds and the following

graph was obtained. Calculate the enthalpy change for the reaction

taking place.

T2 extrapolation at same rate of cooling

T₁

∆T ∆T for rxn = T2-T0

Temp (⁰C)

T₀

point where zinc was added

Temp Change for the Reaction of Copper II Sulfate and Zinc

T₀ = initial temp of reactants = 17⁰C

T₁ = highest temp actually reached = 26.5 ⁰C T₂ = temp that would have been reached if

no heat were lost to the surroundings = 27.4 ⁰C

• From the graph, we are able to extrapolate back and determine that T

would have been 27.4⁰C.

• That gives us a ∆T of

27.4⁰C - 17.0⁰C = 10.4⁰C

• Next we should recognize that we are mixing two given amounts which gives us a limiting reactant problem.

CuSO

+ Zn

Cu

+ ZnSO

0.0183mol Zn x 1/1 = 0.0183mol CuSO4 needed 0.0100mol CuSO4 x 1/1 = 0.0100mol Zn needed

We need 0.0183 mol CuSO4 and are only given 0.0100 mol so CuSO4 is the limiting reactant.

• We know that the reaction will stop

after the 0.0100 mol of CuSO

is used up so the heat evolved is per 0.0100 mol of CuSO

.

• q = mc∆T

= (50.0 cm3 x 1.00 g/cm3)(4.18J/g⁰C)(10.4⁰C) = 2170 J = 2.17 kJ

• ∆H = -2.17kJ / 0.0100 mol = -217 kJ mol

Comments on problem

• Why was there a wait time at the beginning?

• The highest temp was reached 5 min after the zinc had been added. What exactly was happening at this point?

• Copper II sulfate solution is blue. Zinc sulfate

solution is colorless. Zinc is a silver – grey metal and copper is a reddish brown metal. State what you would think you would observe as the reaction proceeds.

• What are three assumptions you have made in arriving at your answer?

• The literature value is -218 kJ/mol. Comment on

the validity of your assumptions you stated.

Final comments on calculating ∆H

• There are 4 types of problems that you should be familiar with.

• Combustion calorimetry

• Neutralization rxns

• Extrapolation for heat loss

• Finding answers from a graph when 2

solutions are mixed and volumes are

kept constant.

• Make sure you always record

qualitative as well as quantitative results.

• Be sure to state the assumptions you are making which are often that no heat is lost from the

system, all heat is transferred to

the water and that the mass of the

solution is assumed to be water so

you are using the specific heat of

water in your calculations.

Citations

International Baccalaureate Organization. Chemistry Guide, First assessment 2016. Updated 2015.

Brown, Catrin, and Mike Ford. Higher Level Chemistry. 2nd ed.

N.p.: Pearson Baccalaureate, 2014. Print.

ISBN 978 1 447 95975 5 eBook 978 1 447 95976 2

Most of the information found in this power point comes directly from this textbook.

The power point has been made to directly complement the

Higher Level Chemistry textbook by Brown and Ford and is used for direct instructional purposes only.