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Lesson Exothermic and Endothermic Reactions and Calorimetry IB Chemistry Power Points Topic 05 Energetics www.pedagogics.ca
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Page 1: Lesson : Enthalpy and Calorimetry

Lesson Exothermic and

Endothermic Reactions

and Calorimetry

IB Chemistry Power Points

Topic 05Energetics

www.pedagogics.ca

Page 2: Lesson : Enthalpy and Calorimetry

Great thanks toJONATHAN HOPTON & KNOCKHARDY PUBLISHING

www.knockhardy.org.uk/sci.htm

Much taken from

ENTHALPYCHANGES

Page 3: Lesson : Enthalpy and Calorimetry

Background and ReviewFirst Law of Thermodynamics (Law of Energy Conservation)Energy can be neither created nor destroyed but it can be converted from one form to another

Energy Changes in Chemical Reactions

All chemical reactions are accompanied by some form of energy change

Exothermic Energy is given out

Endothermic Energy is absorbed

Examples Exothermic combustion reactionsneutralization (acid + base)

Endothermic photosynthesisthermal decomposition of calcium

carbonate

Activity : observing exothermic and endothermic reactions

Page 4: Lesson : Enthalpy and Calorimetry

The heat content of a chemical system is called enthalpy (represented by H).

Key Concept

We cannot measure enthalpy directly, only the change in enthalpy ∆H i.e. the amount of heat released or absorbed when a chemical reaction occurs at constant pressure.

∆H = H(products) – H(reactants)

∆Ho (the STANDARD enthalpy of reaction) is the value measured when temperature is 298 K and pressure is 101.3 kPa.

Page 5: Lesson : Enthalpy and Calorimetry

If ∆H is negative, H(products) < H(reactants)

There is an enthalpy decrease and heat is released to the surroundings.

Enthalpy Diagram -Exothermic Change

enthalpy

Page 6: Lesson : Enthalpy and Calorimetry

If ∆H is positive, H(products) > H(reactants)

There is an enthalpy increase and heat is absorbed from the surroundings.

Enthalpy Diagram - Endothermic Change

enthalpy

Page 7: Lesson : Enthalpy and Calorimetry

Exothermic reactions release heat

Example: Enthalpy change in a chemical reaction

N2(g) + 3H2(g) 2 NH3(g) ∆H = -92.4 kJ/mol

The coefficients in the balanced equation represent the number of moles of reactants and products.

Page 8: Lesson : Enthalpy and Calorimetry

N2(g) + 3H2(g) 2 NH3(g) ∆H = -92.4 kJ/mol

State symbols are ESSENTIAL as changes of state involve changes in thermal energy.

The enthalpy change is directly proportional to the number of moles of substance involved in the reaction.

For the above equation, 92.4 kJ is released- for each mole of N2 reacted- for every 3 moles of H2 reacted- for every 2 moles of NH3 produced.

Page 9: Lesson : Enthalpy and Calorimetry

The reverse reaction 2 NH3(g) N2(g) + 3H2(g)

∆H = +92.4 kJ/mol

Note: the enthalpy change can be read directly from the enthalpy profile diagram.

Page 10: Lesson : Enthalpy and Calorimetry

Thermochemical Standard ConditionsThe ∆H value for a given reaction will depend on reaction conditions.

Values for enthalpy changes are standardized :

for the standard enthalpy ∆Ho

- Temperature is 298 K- Pressure is 1 atmosphere- All solutions involved are 1 M concentration

Page 11: Lesson : Enthalpy and Calorimetry

Assigned : Review Exercise 1

The ∆H value for a given reaction will depend on reaction conditions.

Values for enthalpy changes are standardized :

for the standard enthalpy ∆Ho

- Temperature is 298 K- Pressure is 1 atmosphere- All solutions involved are 1 M concentration

Page 12: Lesson : Enthalpy and Calorimetry

Calorimetry – Part 1

Specific Heat CapacityThe specific heat capacity of a substance is a physical property. It is defined as the amount of heat (Joules) required to change the temperature (oC or K) of a unit mass (g or kg) of substance by ONE degree.

Specific heat capacity (SHC) is measured inJ g-1 K-1 or kJ kg-1 K-1 (chemistry)J kg-1 K-1 (physics)

Page 13: Lesson : Enthalpy and Calorimetry

Calorimetry – continued

Heat and temperature changeKnowing the SHC is useful in thermal chemistry. Heat added or lost can be determined by measuring temperature change of a known substance (water).Q = mc∆T

heat = mass x SHC x ∆Temp

Page 14: Lesson : Enthalpy and Calorimetry

Determining the Specific Heat Capacity of WaterRead the background information and lab activity instructions carefully

A kettle (and other electrical heating devices) have power ratings (given in Watts). This tells you the amount of heat energy supplied by the device. The watt is equivalent to 1 J per second.

(heat) Q = P x t

DON’T FORGET TO CONSIDER UNCERTAINTY

Page 15: Lesson : Enthalpy and Calorimetry

Determining the Specific Heat Capacity of WaterWhat you need to do (DCP)Plot data and choose a suitable range for analysis (look for constant increase in temperature)Organize the data from the range in a suitable table. You may choose to do further processing at this time.Plot the selected range of raw/processed data. Use the slope to determine the specific heat capacity of water.Present all calculations / data processing clearlyDON’T FORGET TO CONSIDER UNCERTAINTY

Page 16: Lesson : Enthalpy and Calorimetry

Determining the Specific Heat Capacity of WaterWhat you need to do (CE)Write a conclusion paragraph. Don’t forget to compare data to the accepted value. Mention any indication of presence (or absence) of random or systematic error. Is your result valid?

Brainstorm a list of procedural/measurement weaknesses or limitations.- do you have evidence (data) that any of these significantly affected the result (as indicated by the data)? How could you improve the investigation?

Be prepared to discuss in class

Page 17: Lesson : Enthalpy and Calorimetry

Determining Specific Heat Capacity of Water

Sample results1300 W kettle supplies 1300 J of heat each second.

Slope tells us temperature increases 0.2215 oC each second.

1300 J heat gives ∆T of 0.2215 oC

A ∆T of 1oC would therefore require 5870 J of heat.

Page 18: Lesson : Enthalpy and Calorimetry

Determining Specific Heat Capacity of Water

Sample results (continued)Mass of water in kettle was 1405 g. To change the temperature of this water by 1 degree, 5870 J of heat were required.

The amount of heat required to change 1 g of water is therefore 4.18 J. This is the SHC of water (very close to the literature value)

SHC of H2O is 4.18 J g-1 K-1

Page 19: Lesson : Enthalpy and Calorimetry

For exampleWhen 3 g of sodium carbonate are added to 50 cm3 of 1.0 M HCl, the temperature rises from 22 °C to 28.5 °C. Calculate the heat required for this temperature change.

Page 20: Lesson : Enthalpy and Calorimetry

Calorimetry – Part 2

ApplicationsA calorimeter is used to measure the heat absorbed or released in a chemical (or other) process by measuring the temperature change of an insulated mass of water.

Page 21: Lesson : Enthalpy and Calorimetry

Sample Problem 1When 3 g of sodium carbonate are added to 50 cm3 of 1.0 M HCl, the temperature rises from 22.0 °C to 28.5 °C. Calculate the heat required for this temperature change.

Page 22: Lesson : Enthalpy and Calorimetry

Sample problem 2: 50.0 cm3 of a 1.00 mol dm-3 HCl solution is mixed with 25.0 cm3 of 2.00 mol dm-3 NaOH. A neutralization reaction occurs. The initial temperature of both solutions was 26.7oC. After stirring and accounting for heat loss, the highest temperature reached was 33.5 oC. Calculate the enthalpy change for this reaction.

NaOH HClboth 26.7o

.

26.7o 33.5o

Page 23: Lesson : Enthalpy and Calorimetry

After writing a balanced equation, the molar quantities and limiting reactant needs to be determined.

Note that in this example there is exactly the right amount of each reactant. If one reactant is present in excess, the heat evolved will associated with the mole amount of limiting reactant.

Sample problem 2: 50.0 cm3 of a 1.00 mol dm-3 HCl solution is mixed with 25.0 cm3 of 2.00 mol dm-3 NaOH. A neutralization reaction occurs. The initial temperature of both solutions was 26.7oC. After stirring and accounting for heat loss, the highest temperature reached was 33.5 oC. Calculate the enthalpy change for this reaction.

Page 24: Lesson : Enthalpy and Calorimetry

Next step – determine how much heat was released.

There are some assumptions in this calculation- Density of reaction mixture (to determine mass)- SHC of reaction mixture (to calculate Q)

Sample problem 2: 50.0 cm3 of a 1.00 mol dm-3 HCl solution is mixed with 25.0 cm3 of 2.00 mol dm-3 NaOH. A neutralization reaction occurs. The initial temperature of both solutions was 26.7oC. After stirring and accounting for heat loss, the highest temperature reached was 33.5 oC. Calculate the enthalpy change for this reaction.

Page 25: Lesson : Enthalpy and Calorimetry

Final step – calculate ΔH for the reaction

Sample problem 2: 50.0 cm3 of a 1.00 mol dm-3 HCl solution is mixed with 25.0 cm3 of 2.00 mol dm-3 NaOH. A neutralization reaction occurs. The initial temperature of both solutions was 26.7oC. After stirring and accounting for heat loss, the highest temperature reached was 33.5 oC. Calculate the enthalpy change for this reaction.

Page 26: Lesson : Enthalpy and Calorimetry

Sample problem 3: to determine the enthalpy of combustion for ethanol (see reaction), a calorimeter setup (below) was used. The burner was lit and allowed to heat the water for 60 s. The change in mass of the burner was 0.518 g and the temperature increase was measured to be 9.90 oC.

What is the big assumption made with this type of experiment?

Page 27: Lesson : Enthalpy and Calorimetry

First step – calculate heat evolved using calorimetryLast step – determine ΔH for the reaction

Sample problem 3: to determine the enthalpy of combustion for ethanol (see reaction), a calorimeter setup (below) was used. The burner was lit and allowed to heat the water for 60 s. The change in mass of the burner was 0.518 g and the temperature increase was measured to be 9.90 oC.

Page 28: Lesson : Enthalpy and Calorimetry

Sample problem 4: 100.0 cm3 of 0.100 mol dm-3 copper II sulphate solution is placed in a styrofoam cup. 1.30 g of powdered zinc is added and a single replacement reaction occurs. The temperature of the solution over time is shown in the graph below. Determine the enthalpy value for this reaction.

First step Make sure you understand the graph.

Extrapolate to determine the change in temperature.

The extrapolation is necessary to compensate for heat loss while the reaction is occurring. Why would powdered zinc be used?

Page 29: Lesson : Enthalpy and Calorimetry

Determine the limiting reactant

Calculate Q

Calculate the enthalpy for the reaction.

Review Exercise 2

Sample problem 4: 100.0 cm3 of 0.100 mol dm-3 copper II sulphate solution is placed in a styrofoam cup. 1.30 g of powdered zinc is added and a single replacement reaction occurs. The temperature of the solution over time is shown in the graph below. Determine the enthalpy value for this reaction.