Lesson

Exothermic and

Endothermic Reactions

and Calorimetry

IB Chemistry Power Points

Topic 05

Energetics

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JONATHAN HOPTON & KNOCKHARDY PUBLISHING

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Much taken from

ENTHALPY

CHANGES

Background and Review

First 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 reactions

neutralization (acid + base)

Endothermic photosynthesis

thermal decomposition of calcium carbonate

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 100.0 kPa.

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

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

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.

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.

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.

Thermochemical Standard Conditions

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 100 kPa

- All solutions involved are 1 M concentration

Calorimetry – Part 1

Specific Heat Capacity

The 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 in

J g-1 K-1 or kJ kg-1 K-1 (chemistry)

J kg-1 K-1 (physics)

Calorimetry – continued

Heat and temperature change

Knowing 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

Calorimetry – Part 2

Applications

A 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.

Sample Problem 1

When 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.

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 HCl

both 26.7o

.

26.7o 33.5o

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.

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.

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.

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?

First step – calculate heat evolved using calorimetry

Last 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.

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?

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.