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Topic 5. Energetics/thermochemistry

5.1 Measuring energy changes

Important definitions

Temperature is the average kinetic energy of molecules (K)

Heat is the amount of energy exchanged due to a temperature difference between two substances (J)

Enthalpy (H) is the amount of stored energy or heat content of a substance. It cannot be measured directly, but we can measure enthalpy change.

 

Change in enthalpy (∆H) is the heat energy change per mole, calculated by the enthalpy of products minus the enthalpy of reactants. 

∆H = ∑Hₚ - ∑Hᵣ

The standard enthalpy change of a reaction (∆Hº) is the difference between the enthalpy of the products and the enthalpy of the reactants at 298K and 1.00 * 10⁵ Pa (1 mol/dm³ for concentration). All reactants and products are in standard states.
 

The specific heat capacity (c) is the property of a substance that gives the heat needed to increase the temperature of unit mass by 1ºK.

  • when c↑, temperature change ↓

  • when c↓, temperature change 

Energy is the measure of the ability to do work; the total energy conserved in chemical reactions.

Exothermic V.S. Endothermic

Exothermic: giving off heat, thus the products become more stable

  • ∆H is negative because energy is given out

  • the temperature of the surroundings increases

  • all combustion, neutralization, condensation reactions are exothermic

    • gas → liquid releases energy​

  • products are more stable, have less stored energy, have stronger bonds

    • low energy = more stable​

Endothermic: energy is absorbed, thus products are less stable

  • ∆H is positive

  • the temperature of the surroundings decreases

  • forming bonds takes energy

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From Ms. Fu's powerpoint

From Ms. Fu's powerpoint

Calculating enthalpy change (kJ/mol)

The first step is to find the heat energy change of the surroundings (Q) in kJ or J.

 

Q = mc∆T

Q = heat energy change (in kJ or J)

m = mass (in kg or g)

  • if needed, convert 1cm³ = 1g

  • water's density is 1g/mL

c = specific heat capacity (in  kJ/kgºK or  kJ/kJºC  or  J/gºK  or  J/gºC)

  • water: 4.18 J/gºK

∆T = temperature change (in ºK or ºC)

  • the units must match up! either use only the kilo- ones or the ones without kilo-

Second, find the enthalpy change, or the heat energy change per mole

∆H = - Q/n

∆H = enthalpy change (in kJ/mol)

  • ​negative if exothermic, positive if endothermic

Q = heat energy change (in kJ)

  • if you found Q in Joules from step 1, you MUST convert it to kJ.

n = number of moles in the system

There is a negative sign in this equation because the first step measures the heat energy change of the surroundings, but the second step concerns the reaction system

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Calorimetry (Practical)

Calorimetry is a technique used to measure the enthalpy change of a reaction

Bomb calorimetry measures ∆H at a constant volume by measuring the ∆T of a solution with the reasoning ∆Hreaction = -∆HH2O

If the solution is water:

  1. c(H2O) = 4.18 J/gºC

  2. measure the mass of water

  3. measure then calculate Tfinal - Tinitial = ∆T

    • extrapolate the line (check the figure below)​

  4. calculate Q of water using Q = mc∆T

  5. find amount (n) of chemicals in the reaction

  6. calculate ∆H using ∆H = - Q/n

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From Ms. Fu's powerpoint

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From Ms. Fu's powerpoint

Assumptions and sources of error:

  1. assume there is no heat loss or gain to/from the surroundings before affecting the T of the water

    • heat loss → lower recorded value of T​

    • heat gain → higher recorded value of T

  2. assume the solution has the same specific heat capacity as pure water

  3. reaction mechanisms (many reactions could be occurring at the same time)

  4. experiment not done in standard conditions

  5. for exothermic only: incomplete combustion

  6. water has a high specific heat capacity → possible error: it absorbs heat but does not necessarily convert it to heat energy

    • therefore it is not detected by the thermometer

5.2 Hess's law

Hess's law states that in a chemical reaction, the total change in chemical potential energy (enthalpy change) must be equal to the energy lost or gained by the reaction system due to the law of conservation of energy. 

  • in other words: ∆H for the overall reaction is the sum of the ∆H for each individual step

∆H = ∆H₁ + ∆H₂ + ∆H₃

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From Ms. Fu's powerpoint

The standard enthalpy of formation (∆Hºf) is the enthalpy change that results when one mole of a compound is formed from its elements in their standard states (298ºK and 1.00 * 10⁵ Pa)

  • the ∆Hºf of an element = 0;     for example, ∆Hº(H2) = 0

    • it does not take much energy to turn an element into an element​

  • found in Table 12 of the IB Chemistry Data Booklet

Example: the ∆Hºf  of CH4

IMG_3587.jpg

∆Hreaction = ∑∆Hf of products - ∑∆Hf of reactants

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From Ms. Fu's powerpoint

The standard enthalpy of combustion (∆Hºc) is the enthalpy change that results when one mole of a compound reacts with oxygen at 298ºK and 1.00 * 10⁵ Pa when all reactants and products are in their standard states.

  • always negative because it is exothermic (energy released)

  • some reactants/products do not have this value because they do not combust

  • found in Table 13 of the IB Chemistry Data Booklet

∆Hreaction = ∑∆Hc of reactants - ∑∆Hc of products

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From Ms. Fu's powerpoint

You can use both ∆Hºf and ∆Hºc to solve the same problem:

IMG_3588.jpg
5.3 Bond enthalpies

Average bond enthalpy is the energy needed to break one mole of a bond in a gaseous molecule averaged over similar compounds.

  • only for covalent bonds and gases

  • because it does not account for IMFs, it can be inaccurate

∆H = ∆Hbreak – ∆Hform

∆H = ∆Hreactants – ∆Hproducts

* in the data booklet, all the values are positive. determining the sign (+ or -) is our responsibility

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Breaking bonds is endothermic as it requires energy → reactants

Forming bonds is exothermic as it releases energy → products

A general rule of thumb: single bonds < double bonds < triple bonds

Ozone depletion

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Pearson textbook pg. 236

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