5.1 Measuring Energy Change

The Law of Conservation of Energy.

Certain laws are thought to be universally true.

One of these is the first law of thermodynamics which says that energy can be converted from one form to another but can never be created or destroyed.

{E=mc² explains how energy can be “stored” as mass but this only affects Physics.}

In a chemical reaction energy may be stored in bonds this is a form of _______________________ energy.

There is also ____________________ energy in the particles since they are all in constant

motion (even in a solid reactant).

The___________________ of a system is simply the Average Kinetic Energy of the molecules

(each of which will be at different energy due to collisions with neighbours).

A closed system is one in which no energy escapes so all energy changes are between the reactants are their immediate surroundings.

Why do we need to put this energy in?

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In the reaction in the energy profile above, more energy will be released than was put in.

• If we thought of the Activation Energy as E_in and the energy released as E_out, how could we work out ∆H?

∆H = ____________________________________________________________

• If you were holding a test-tube in which the reaction was happening, would your hand be heated or cooled?

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• vIs this reaction Exo-or Endothermic? How do you know?

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• Is ∆H positive or negative?

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In the reaction in this energy profile, less energy will be released than was put in.

• If we thought of the Activation Energy as E_in and the energy released as E_out, how could we work out ∆H?

∆H = ____________________________________________________________

• If you were holding a test-tube in which the reaction was happening, would your hand be heated or cooled?

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• Is this reaction Exo-or Endothermic? How do you know?

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• Is ∆H positive or negative?

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• Why would this reaction stop if we stopped heating it?

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• What is the purpose of the lid?

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• All three of these precautions reduce Systematic Errors- what does this phrase mean?

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• Another way to reduce systematic errors is to stir. Why is this necessary?

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{Assume that densities of solutions = 1.00 g cm^-3, specific heat capacity of solutions is 4.18 J g^-1 K^-1. Ignore the heat capacity of solids}.

1. 100 cm³ of 0.20 mol dm^-3 copper sulphate solution was put in a calorimeter and 2.0g of magnesium powder added. The temperature of the solution rose by 25.1°C. Work out which reagent was in excess and then calculate the enthalpy change for the reaction.

2. 25 cm³ of 2.0 mol dm^-3 nitric acid was added to 25 cm³ of 2.0 mol dm^-3 potassium hydroxide solution. The temperature rose by 13.7°C. Calculate the enthalpy of neutralisation for this reaction.

3. 50 cm³ of 2.0 mol dm^-3 hydrochloric acid was added to 50 cm³ of 2.0 mol dm^-3 ammonia solution. The temperature rose by 12.4°C. Calculate the enthalpy of neutralisation for this reaction.

4. 50 cm³ of 1.0 mol dm^-3 nitric acid was added to 20 cm³ of 1.0 mol dm^-3 barium hydroxide solution. The temperature rose by 7.9°C. Calculate enthalpy of neutralisation (per mole of nitric acid reacting).

5. 50 cm³ of 0.10 mol dm^-3 silver nitrate solution was put in a calorimeter and 0.2 g of zinc powder added. The temperature of the solution rose by 4.3°C. Work out which reagent was in excess and then calculate the enthalpy change for the reaction (per mole of zinc that reacts).

2 AgNO_3(aq) + Zn_(s) → 2 Ag_(s) + Zn(NO₃)_2(aq)

6. 3.53 g of sodium hydrogencarbonate was added to 30.0 cm³ of 2.0 mol dm^-3 hydrochloric acid. The temperature fell by 10.3 K. Work out which reagent was in excess and calculate the enthalpy change of reaction.

NaHCO_3(s) + HCl_(aq) → NaCl_(aq) + H₂O_(l) + CO_2(g)

7. In an experiment, 0.750 g of benzene (C6H6) were completely burned in air. The heat evolved raised the temperature of 200 g of water by 43.7°C. Use this data to calculate the enthalpy of combustion of benzene

8. a) Write an equation to represent the ∆H_c of butan-1-ol (C₄H₉OH_(l)).

b) A simple flame calorimeter was used to measure the ∆H_c of butan-1-ol. 0.600 g of butan-1ol was burned in a simple spirit burner under a container of water. There was 250 g of water in the container and its temperature rose by 19.4°C.

9. When 1.30 g of zinc reacts with 100 cm³ of 2.00 mol dm^-3 nitric acid, the temperature rises by 6.0oC. The equation for the reaction is:

Zn_(s) + 2 HNO_3(aq) → Zn(NO₃)_2(aq) + H_2(g)

a) Calculate which reagent is in excess.

b) Calculate the heat given out in the experiment.

c) Calculate the enthalpy change for the reaction.

10. When 0.40 g of calcium reacts with 100 cm³ of 2.00 mol dm^-3 hydrochloric acid, the temperature rises by 12.0^oC. The equation for the reaction is

Ca_(s) + 2 HCl_(aq) → CaCl_2(aq) + H_2(g)

a) Calculate the heat released in the reaction.

b) Calculate which reagent is in excess.

c) Calculate the enthalpy change for this reaction per mole of calcium reacting.

It is impossible for a reaction to be over immediately the two reactants meet.

Why?

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But as the reaction goes on energy is released and some will escape. Why might this be a problem to the final q=mc∆T calculation?

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It is impossible to stop all heat loss so we have to try to account for it. Imagine an experiment that has the following results. Plot them.

The last 7 points show the system cooling steadily once the reaction was over (or almost over).

We can assume that it was cooling at that rate all the time.

So we extrapolate the line backwards to account for the heat lost while the reaction was still going quickly during the first few points.

• Extrapolate your own line backwards and use it to estimate a better ∆T

• Why might is have been better to measure the temperature for 2 or 3 minutes before starting the reaction?

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• Learn the definition:

The standard enthalpy change of a substance is the enthalpy change when ____________ moles of a substance is formed from its _________________ ____________________________________________ with all species in their standard state, under standard conditions of __________kelvin and ___________kPa pressure.

ΔH^⊖_reaction=∑ΔH^⊖_f(products)− ∑ΔH^⊖_f(Reactants)

Imagine a simple reaction in which A and B form C:

A+B⇋C

That has standard enthalpies of formation:

· ΔH_f^o[A] = 433 KJ/mol

· ΔH_f^o[B] = -256 KJ/mol

· ΔH_f^o[C] = 523 KJ/mol

the equation for the standard enthalpy change of formation is as follows:

ΔH_reaction^o = ΔH_f^o[C] - (ΔH_f^o[A] + ΔH_f^o[B])

ΔH_reaction^o = (523 kJ/mol) - ((433 kJ/mol) + (-256 kJ/mol))

ΔH_reaction^o = 346 kJ/mol

• Using the values in the above table of standard enthalpies of formation, calculate the ΔH_reaction^o for the formation of NO₂(g).

• Using the values in the above table of standard enthalpies of formation, calculate the ΔH_reaction^o for the formation of N₂O₄(g)

• Using the values in the above table of standard enthalpies of formation, calculate the ΔH_reaction^o for

NO₂ + CO--> NO + CO₂

• Using the values in the table, calculate the ΔH_reaction^o for:

SO₂ + N₂O₄ --> NO₂ + NO + SO₃

• Using the values in the table, calculate the ΔH_reaction^o for:

2H₂O + F₂--> 2HF + H₂ + O₂

• Why is it safe to assume ΔH_f^o =0kJ/mol for F₂?

• Learn the definition:

The standard enthalpy change of a combustion is the enthalpy change when ____________ moles of a substance is burned in _________________ with all species in their standard state, under standard conditions of _____kelvin and ___________kPa pressure.

Combustion enthalpies can generally be found by careful Calorimetry but could always be found by calculation from Formation Enthalpies.

• Using the values in the table of standard enthalpies of formation, calculate the ΔH_combustion^o of C₂H₄ (g).

{Write a balanced equation for complete combustion first!}